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On adding chlorine water, bromine and Iodine are displaced from their solutions by chlorine. Bromine is more electronegative than iodide but less 6than chlorine. On adding Bromine water, iodine is displaced from its solution but not chlorine. Table showing the displacement of the halogens (V) means there is displacement (x ) means there is no displacement Chemical /ionic equations With Fluorine F2(g) + 2NaCl-(aq) -> 2NaF(aq) + Cl2(aq) F2(g) + 2Cl-(aq) -> 2F-(aq) + Cl2(aq) F2(g) + 2NaBr-(aq) -> 2NaF(aq) + Br2(aq) F2(g) + 2Br-(aq) -> 2F-(aq) + Br2(aq) F2(g) + 2NaI-(aq) -> 2NaF(aq) + I2(aq) F2(g) + 2I-(aq) -> 2F-(aq) + I2(aq) With chlorine Cl2(g) + 2NaCl-(aq) -> 2NaCl(aq) + Br2(aq) Cl2(g) + 2Br-(aq) -> 2Cl-(aq) + Br2(aq) Halogen ion in solution Halogen F- Cl- Br- I- F2 X Cl2 X X Br2 X X X I2 X X X X
46 Cl2(g) + 2NaI-(aq) -> 2NaCl(aq) + I2(aq) Cl2(g) + 2I-(aq) -> 2Cl-(aq) + I2(aq) With Bromine Br2(g) + 2NaI-(aq) -> 2NaBr(aq) + I2(aq) Br2(g) + 2I-(aq) -> 2Br-(aq) + I2(aq) Uses of halogens (i) Florine – manufacture of P.T.F.E (Poly tetra fluoroethene) synthetic fiber. - Reduce tooth decay when added in small amounts/quantities in tooth paste. NB –large small quantities of fluorine /fluoride ions in water cause browning of teeth/flourosis. - Hydrogen fluoride is used to engrave words /pictures in glass.
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- Reduce tooth decay when added in small amounts/quantities in tooth paste. NB –large small quantities of fluorine /fluoride ions in water cause browning of teeth/flourosis. - Hydrogen fluoride is used to engrave words /pictures in glass. (ii) Bromine - Silver bromide is used to make light sensitive photographic paper/films. (iii) Iodide – Iodine dissolved in alcohol is used as medicine to kill bacteria in skin cuts. It is called tincture of iodine. The table below to show some compounds of halogens. (i) Below is the table showing the bond energy of four halogens. Bond Bond energy k J mole-1 Cl-Cl 242 Br-Br 193 I-I 151 I. What do you understand by the term “bond energy” Element Halogen H Na Mg Al Si C P F HF NaF MgH2 AlF3 SiF4 CF4 PF3 Cl HCl NaCl MgCl AlCl3 SiCl3 CCl4 PCl3 Br HBr NaBr MgBr2 AlBr3 SiBr4 CBr4 PBr3 I Hl Nal Mgl2 All3 SiI4 Cl2 PBr3
47 Bond energy is the energy required to break/ form one mole of chemical bond II. Explain the trend in bond Energy of the halogens above: -Decrease down the group from chlorine to Iodine -Atomic radius increase down the group decreasing the energy required to break the covalent bonds between the larger atom with reduced effective nuclear @ charge an outer energy level that take part in bonding. (c)Group VIII elements: Noble gases Group VIII elements are called Noble gases. They are all non metals. Noble gases occupy about 1.0% of the atmosphere as colourless gaseous mixture. Argon is the most abundant with 0.9%. They exists as monatomic molecules with very weak van-der-waals /intermolecular forces holding the molecules. They include: All noble gas atoms have a stable duplet(two electrons in the 1st energy level) or octet(eight electrons in other outer energy level)in the outer energy level. They therefore do not acquire/gain extra electron in the outer energy level or donate/lose. They therefore are therefore zerovalent . The number of energy levels increases down the group from Helium to Randon.
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They therefore do not acquire/gain extra electron in the outer energy level or donate/lose. They therefore are therefore zerovalent . The number of energy levels increases down the group from Helium to Randon. The more the number of energy levels the bigger/larger the atomic size/radius. e.g. The atomic size/radius of Argon is bigger/larger than that of Neon because Argon has more/3 energy levels than Neon (2 energy levels). Atomic radius noble gases increase down the group as the number of energy levels increases. The effective nuclear attraction on the outer electrons thus decrease down the group. The noble gases are generally unreactive because the outer energy level has the stable octet/duplet. The stable octet/duplet in noble gas atoms lead to a Element Symbol Atomic number Electron structure State at room temperature Helium He 2 2: Colourless gas Neon Ne 10 2:8 Colourless gas Argon Ar 18 2:8:8 Colourless gas Krypton Kr 36 2:8:18:8 Colourless gas Xenon Xe 54 2:8:18:18:8 Colourless gas Radon Rn 86 2:8:18:32:18:8 Radioctive
48 comparatively very high 1st ionization energy. This is because losing /donating an electron from the stable atom require a lot of energy to lose/donate and make it unstable. As atomic radius increase down the group and the 1st ionization energy decrease, very electronegative elements like Oxygen and Fluorine are able to react and bond with lower members of the noble gases.e.g Xenon reacts with Fluorine to form a covalent compound XeF6.This is because the outer electrons/energy level if Xenon is far from the nucleus and thus experience less effective nuclear attraction. Noble gases have low melting and boiling points. This is because they exist as monatomic molecules joined by very weak intermolecular/van-der-waals forces that require very little energy to weaken and form liquid and break to form a gas. The intermolecular/van-der-waals forces increase down the group as the atomic radius/size increase from Helium to Radon. The melting and boiling points thus increase also down the group. Noble gases are insoluble in water and are poor conductors of electricity.
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The intermolecular/van-der-waals forces increase down the group as the atomic radius/size increase from Helium to Radon. The melting and boiling points thus increase also down the group. Noble gases are insoluble in water and are poor conductors of electricity. Element Formula of molecule Electrical conductivity Solubility in water Atomic radius(nM) 1st ionization energy Melting point(0C) Boiling point(0C) Helium He Poor Insoluble 0.128 2372 -270 -269 Neon Ne Poor Insoluble 0.160 2080 -249 -246 Argon Ar Poor Insoluble 0.192 1520 -189 -186 Krypton Kr Poor Insoluble 0.197 1350 -157 -152 Xenon Xe Poor Insoluble 0.217 1170 -112 -108 Radon Rn Poor Insoluble 0.221 1134 -104 -93 Uses of noble gases Argon is used in light bulbs to provide an inert environment to prevent oxidation of the bulb filament Argon is used in arch welding as an insulator. Neon is used in street and advertisement light Helium is mixed with Oxygen during deep sea diving and mountaineering. Helium is used in weather balloon for meteorological research instead of Hydrogen because it is unreactive/inert.Hydrogen when impure can ignite with an explosion. Helium is used in making thermometers for measuring very low temperatures. 49 C. PERIODICITY OF ACROSS THE PERIOD. (See Chemical bonding and Structure)
PERIODICITY OF CHEMICAL FAMILES (Patterns down the group) The number of valence electrons and the number of occupied energy levels in an atom of an element determine the position of an element in the periodic table. i.e The number of occupied energy levels determine the Period and the valence electrons determine the Group. Elements in the same group have similar physical and chemical properties. The trends in physical and chemical properties of elements in the same group vary down the group. Elements in the same group thus constitute a chemical family. (a) Group I elements: Alkali metals Group I elements are called Alkali metals except Hydrogen which is a non metal.
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The trends in physical and chemical properties of elements in the same group vary down the group. Elements in the same group thus constitute a chemical family. (a) Group I elements: Alkali metals Group I elements are called Alkali metals except Hydrogen which is a non metal. The alkali metals include: Element Symbol Atomic number Electron structure Oxidation state Valency Lithium Li 3 2:1 Li+ 1 Sodium Na 11 2:8:1 Na+ 1 Potassium K 19 2:8:8:1 K+ 1 Rubidium Rb 37 2:8:18:8:1 Rb+ 1 Caesium Cs 55 2:8:18:18:8:1 Cs+ 1 Francium Fr 87 2:8:18:32:18:8:1 Fr+ 1 All alkali metals atom has one electron in the outer energy level. They therefore are monovalent. They donate /lose the outer electron to have oxidation state M+ The number of energy levels increases down the group from Lithium to Francium. The more the number of energy levels the bigger/larger the atomic size. e.g. The atomic size of Potassium is bigger/larger than that of sodium because Potassium has more/4 energy levels than sodium (3 energy levels). Atomic and ionic radius The distance between the centre of the nucleus of an atom and the outermost energy level occupied by electron/s is called atomic radius. Atomic radius is measured in nanometers(n).The higher /bigger the atomic radius the bigger /larger the atomic size. 2 The distance between the centre of the nucleus of an ion and the outermost energy level occupied by electron/s is called ionic radius. Ionic radius is also measured in nanometers(n).The higher /bigger the ionic radius the bigger /larger the size of the ion. Atomic radius and ionic radius depend on the number of energy levels occupied by electrons. The more the number of energy levels the bigger/larger the atomic /ionic radius. e.g. The atomic radius of Francium is bigger/larger than that of sodium because Francium has more/7 energy levels than sodium (3 energy levels). Atomic radius and ionic radius of alkali metals increase down the group as the number of energy levels increases. The atomic radius of alkali metals is bigger than the ionic radius.
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The atomic radius of Francium is bigger/larger than that of sodium because Francium has more/7 energy levels than sodium (3 energy levels). Atomic radius and ionic radius of alkali metals increase down the group as the number of energy levels increases. The atomic radius of alkali metals is bigger than the ionic radius. This is because alkali metals react by losing/donating the outer electron and hence lose the outer energy level. Table showing the atomic and ionic radius of some alkali metals Element Symbol Atomic number Atomic radius(nM) Ionic radius(nM) Lithium Li 3 0.133 0.060 Sodium Na 11 0.157 0.095 Potassium K 19 0.203 0.133 The atomic radius of sodium is 0.157nM .The ionic radius of Na+ is 0.095nM. This is because sodium reacts by donating/losing the outer electrons and hence the outer energy level. The remaining electrons/energy levels experience more effective / greater nuclear attraction/pull towards the nucleus reducing the atomic radius. Electropositivity The ease of donating/losing electrons is called electropositivity. All alkali metals are electropositive. Electropositivity increase as atomic radius increase. This is because the effective nuclear attraction on outer electrons decreases with increase in atomic radius. The outer electrons experience less nuclear attraction and can be lost/ donated easily/with ease. Francium is the most electropositive element in the periodic table because it has the highest/biggest atomic radius. Ionization energy
3 The minimum amount of energy required to remove an electron from an atom of element in its gaseous state is called 1st ionization energy. The SI unit of ionization energy is kilojoules per mole/kJmole-1 .Ionization energy depend on atomic radius. The higher the atomic radius, the less effective the nuclear attraction on outer electrons/energy level and thus the lower the ionization energy. For alkali metals the 1st ionization energy decrease down the group as the atomic radius increase and the effective nuclear attraction on outer energy level electrons decrease. e.g. The 1st ionization energy of sodium is 496 kJmole-1 while that of potassium is 419 kJmole-1 .This is because atomic radius increase and thus effective nuclear attraction on outer energy level electrons decrease down the group from sodium to Potassium.
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For alkali metals the 1st ionization energy decrease down the group as the atomic radius increase and the effective nuclear attraction on outer energy level electrons decrease. e.g. The 1st ionization energy of sodium is 496 kJmole-1 while that of potassium is 419 kJmole-1 .This is because atomic radius increase and thus effective nuclear attraction on outer energy level electrons decrease down the group from sodium to Potassium. It requires therefore less energy to donate/lose outer electrons in Potassium than in sodium. Physical properties Soft/Easy to cut: Alkali metals are soft and easy to cut with a knife. The softness and ease of cutting increase down the group from Lithium to Francium. This is because an increase in atomic radius, decreases the strength of metallic bond and the packing of the metallic structure Appearance: Alkali metals have a shiny grey metallic luster when freshly cut. The surface rapidly/quickly tarnishes on exposure to air. This is because the metal surface rapidly/quickly reacts with elements of air/oxygen. Melting and boiling points: Alkali metals have a relatively low melting/boiling point than common metals like Iron. This is because alkali metals use only one delocalized electron to form a weak metallic bond/structure. Electrical/thermal conductivity: Alkali metals are good thermal and electrical conductors. Metals conduct using the outer mobile delocalized electrons. The delocalized electrons move randomly within the metallic structure. Summary of some physical properties of the 1st three alkali metals Alkali metal Appearance Ease of cutting Melting point (oC) Boiling point (oC) Conductivity 1st ionization energy Lithium Silvery white Not easy 180 1330 Good 520 Sodium Shiny grey Easy 98 890 Good 496 Potassium Shiny grey Very easy 64 774 Good 419
4 Chemical properties (i)Reaction with air/oxygen On exposure to air, alkali metals reacts with the elements in the air. Example On exposure to air, Sodium first reacts with Oxygen to form sodium oxide. 4Na(s) + O2(g) -> 2Na2O(s) The sodium oxide formed further reacts with water/moisture in the air to form sodium hydroxide solution.
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Summary of some physical properties of the 1st three alkali metals Alkali metal Appearance Ease of cutting Melting point (oC) Boiling point (oC) Conductivity 1st ionization energy Lithium Silvery white Not easy 180 1330 Good 520 Sodium Shiny grey Easy 98 890 Good 496 Potassium Shiny grey Very easy 64 774 Good 419
4 Chemical properties (i)Reaction with air/oxygen On exposure to air, alkali metals reacts with the elements in the air. Example On exposure to air, Sodium first reacts with Oxygen to form sodium oxide. 4Na(s) + O2(g) -> 2Na2O(s) The sodium oxide formed further reacts with water/moisture in the air to form sodium hydroxide solution. Na2O(s) + H2O(l) -> 2NaOH(aq) Sodium hydroxide solution reacts with carbon(IV)oxide in the air to form sodium carbonate. 2NaOH(aq) + CO2(g) -> Na2CO3(g) + H2O(l) (ii)Burning in air/oxygen Lithium burns in air with a crimson/deep red flame to form Lithium oxide 4Li (s) + O2(g) -> 2Li2O(s) Sodium burns in air with a yellow flame to form sodium oxide 4Na (s) + O2(g) -> 2Na2O(s) Sodium burns in oxygen with a yellow flame to form sodium peroxide 2Na (s) + O2(g) -> Na2O2 (s) Potassium burns in air with a lilac/purple flame to form potassium oxide 4K (s) + O2(g) -> 2K2O (s) (iii) Reaction with water: Experiment Measure 500 cm3 of water into a beaker. Put three drops of phenolphthalein indicator. Put about 0.5g of Lithium metal into the beaker. Determine the pH of final product Repeat the experiment using about 0.1 g of Sodium and Potassium.
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Put three drops of phenolphthalein indicator. Put about 0.5g of Lithium metal into the beaker. Determine the pH of final product Repeat the experiment using about 0.1 g of Sodium and Potassium. Caution: Keep a distance Observations
5 Alkali metal Observations Comparative speed/rate of the reaction Lithium -Metal floats in water -rapid effervescence/fizzing/bubbling -colourless gas produced (that extinguishes burning splint with explosion /“pop” sound) -resulting solution turn phenolphthalein indicator pink -pH of solution = 12/13/14 Moderately vigorous Sodium -Metal floats in water -very rapid effervescence /fizzing /bubbling -colourless gas produced (that extinguishes burning splint with explosion /“pop” sound) -resulting solution turn phenolphthalein indicator pink -pH of solution = 12/13/14 Very vigorous Potassium -Metal floats in water -explosive effervescence /fizzing /bubbling -colourless gas produced (that extinguishes burning splint with explosion /“pop” sound) -resulting solution turn phenolphthalein indicator pink -pH of solution = 12/13/14 Explosive/burst into flames Explanation Alkali metals are less dense than water. They therefore float in water.They react with water to form a strongly alkaline solution of their hydroxides and producing hydrogen gas. The rate of this reaction increase down the group. i.e. Potassium is more reactive than sodium .Sodium is more reactive than Lithium. The reactivity increases as electropositivity increases of the alkali increases. This is because as the atomic radius increases , the ease of donating/losing outer electron increase during chemical reactions.
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Potassium is more reactive than sodium .Sodium is more reactive than Lithium. The reactivity increases as electropositivity increases of the alkali increases. This is because as the atomic radius increases , the ease of donating/losing outer electron increase during chemical reactions. 6 Chemical equations 2Li(s) + 2H2O(l) -> 2LiOH(aq) + H2(g) 2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g) 2K(s) + 2H2O(l) -> 2KOH(aq) + H2(g) 2Rb(s) + 2H2O(l) -> 2RbOH(aq) + H2(g) 2Cs(s) + 2H2O(l) -> 2CsOH(aq) + H2(g) 2Fr(s) + 2H2O(l) -> 2FrOH(aq) + H2(g) Reactivity increase down the group (iv) Reaction with chlorine: Experiment Cut about 0.5g of sodium into a deflagrating spoon with a lid cover. Introduce it on a Bunsen flame until it catches fire. Quickly and carefully lower it into a gas jar containing dry chlorine to cover the gas jar. Repeat with about 0.5g of Lithium. Caution: This experiment should be done in fume chamber because chlorine is poisonous /toxic. Observation Sodium metal continues to burn with a yellow flame forming white solid/fumes. Lithium metal continues to burn with a crimson flame forming white solid / fumes. Alkali metal react with chlorine gas to form the corresponding metal chlorides. The reactivity increase as electropositivity increase down the group from Lithium to Francium.The ease of donating/losing the outer electrons increase as the atomic radius increase and the outer electron is less attracted to the nucleus.
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Lithium metal continues to burn with a crimson flame forming white solid / fumes. Alkali metal react with chlorine gas to form the corresponding metal chlorides. The reactivity increase as electropositivity increase down the group from Lithium to Francium.The ease of donating/losing the outer electrons increase as the atomic radius increase and the outer electron is less attracted to the nucleus. Chemical equations 2Li(s) + Cl2(g) -> 2LiCl(s) 2Na(s) + Cl2(g) -> 2NaCl(s) 2K(s) + Cl2(g) -> 2KCl(s) 2Rb(s) + Cl2(g) -> 2RbCl(s) 2Cs(s) + Cl2(g) -> 2CsCl(s) 2Fr(s) + Cl2(g) -> 2FrCl(s) Reactivity increase down the group The table below shows some compounds of the 1st three alkali metals
7 Some uses of alkali metals include: (i)Sodium is used in making sodium cyanide for extracting gold from gold ore. (ii)Sodium chloride is used in seasoning food. (iii)Molten mixture of sodium and potassium is used as coolant in nuclear reactors. (iv)Sodium is used in making sodium hydroxide used in making soapy and soapless detergents. (v)Sodium is used as a reducing agent for the extraction of titanium from Titanium(IV)chloride. (vi)Lithium is used in making special high strength glasses (vii)Lithium compounds are used to make dry cells in mobile phones and computer laptops. Group II elements: Alkaline earth metals Group II elements are called Alkaline earth metals .
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(v)Sodium is used as a reducing agent for the extraction of titanium from Titanium(IV)chloride. (vi)Lithium is used in making special high strength glasses (vii)Lithium compounds are used to make dry cells in mobile phones and computer laptops. Group II elements: Alkaline earth metals Group II elements are called Alkaline earth metals . The alkaline earth metals include: Lithium sodium Potassium Hydroxide LiOH NaOH KOH Oxide Li2O Na2O K2O Sulphide Li2S Na2S K2S Chloride LiCl NaCl KCl Carbonate Li2CO3 Na2CO3 K2CO3 Nitrate(V) LiNO3 NaNO3 KNO3 Nitrate(III) - NaNO2 KNO2 Sulphate(VI) Li2SO4 Na2SO4 K2SO4 Sulphate(IV) - Na2SO3 K2SO3 Hydrogen carbonate - NaHCO3 KHCO3 Hydrogen sulphate(VI) - NaHSO4 KHSO4 Hydrogen sulphate(IV) - NaHSO3 KHSO3 Phosphate - Na3PO4 K3PO4 Manganate(VI) - NaMnO4 KMnO4 Dichromate(VI) - Na2Cr2O7 K2Cr2O7 Chromate(VI) - Na2CrO4 K2CrO4
8 Element Symbol Atomic number Electron structure Oxidation state Valency Beryllium Be 4 2:2 Be2+ 2 Magnesium Mg 12 2:8:2 Mg2+ 2 Calcium Ca 20 2:8:8:2 Ca2+ 2 Strontium Sr 38 2:8:18:8:2 Sr2+ 2 Barium Ba 56 2:8:18:18:8:2 Ba2+ 2 Radium Ra 88 2:8:18:32:18:8:2 Ra2+ 2 All alkaline earth metal atoms have two electrons in the outer energy level. They therefore are divalent. They donate /lose the two outer electrons to have oxidation state M2+ The number of energy levels increases down the group from Beryllium to Radium. The more the number of energy levels the bigger/larger the atomic size. e.g.
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They donate /lose the two outer electrons to have oxidation state M2+ The number of energy levels increases down the group from Beryllium to Radium. The more the number of energy levels the bigger/larger the atomic size. e.g. The atomic size/radius of Calcium is bigger/larger than that of Magnesium because Calcium has more/4 energy levels than Magnesium (3 energy levels). Atomic radius and ionic radius of alkaline earth metals increase down the group as the number of energy levels increases. The atomic radius of alkaline earth metals is bigger than the ionic radius. This is because they react by losing/donating the two outer electrons and hence lose the outer energy level. Table showing the atomic and ionic radius of the 1st three alkaline earth metals Element Symbol Atomic number Atomic radius(nM) Ionic radius(nM) Beryllium Be 4 0.089 0.031 Magnesium Mg 12 0.136 0.065 Calcium Ca 20 0.174 0.099 The atomic radius of Magnesium is 0.136nM .The ionic radius of Mg2+ is 0.065nM. This is because Magnesium reacts by donating/losing the two outer electrons and hence the outer energy level. The remaining electrons/energy levels experience more effective / greater nuclear attraction/pull towards the nucleus reducing the atomic radius. 9 Electropositivity All alkaline earth metals are also electropositive like alkali metals. The electropositivity increase with increase in atomic radius/size. Calcium is more electropositive than Magnesium. This is because the effective nuclear attraction on outer electrons decreases with increase in atomic radius. The two outer electrons in calcium experience less nuclear attraction and can be lost/ donated easily/with ease because of the higher/bigger atomic radius. Ionization energy For alkaline earth metals the 1st ionization energy decrease down the group as the atomic radius increase and the effective nuclear attraction on outer energy level electrons decrease. e.g. The 1st ionization energy of Magnesium is 900 kJmole-1 while that of Calcium is 590 kJmole-1 .This is because atomic radius increase and thus effective nuclear attraction on outer energy level electrons decrease down the group from magnesium to calcium. It requires therefore less energy to donate/lose outer electron in calcium than in magnesium.
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e.g. The 1st ionization energy of Magnesium is 900 kJmole-1 while that of Calcium is 590 kJmole-1 .This is because atomic radius increase and thus effective nuclear attraction on outer energy level electrons decrease down the group from magnesium to calcium. It requires therefore less energy to donate/lose outer electron in calcium than in magnesium. The minimum amount of energy required to remove a second electron from an ion of an element in its gaseous state is called the 2nd ionization energy. The 2nd ionization energy is always higher /bigger than the 1st ionization energy. This because once an electron is donated /lost form an atom, the overall effective nuclear attraction on the remaining electrons/energy level increase. Removing a second electron from the ion require therefore more energy than the first electron. The atomic radius of alkali metals is higher/bigger than that of alkaline earth metals.This is because across/along the period from left to right there is an increase in nuclear charge from additional number of protons and still additional number of electrons entering the same energy level. Increase in nuclear charge increases the effective nuclear attraction on the outer energy level which pulls it closer to the nucleus. e.g. Atomic radius of Sodium (0.157nM) is higher than that of Magnesium (0.137nM). This is because Magnesium has more effective nuclear attraction on the outer energy level than Sodium hence pulls outer energy level more nearer to its nucleus. Physical properties Soft/Easy to cut: Alkaline earth metals are not soft and easy to cut with a knife like alkali metals. This is because of the decrease in atomic radius of
10 corresponding alkaline earth metal, increases the strength of metallic bond and the packing of the metallic structure. Alkaline earth metals are (i)ductile(able to form wire/thin long rods) (ii)malleable(able to be hammered into sheet/long thin plates) (iii)have high tensile strength(able to be coiled without breaking/ not brittle/withstand stress) Appearance: Alkali earth metals have a shiny grey metallic luster when their surface is freshly polished /scrubbed. The surface slowly tarnishes on exposure to air. This is because the metal surface slowly undergoes oxidation to form an oxide. This oxide layer should be removed before using the alkaline earth metals.
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The surface slowly tarnishes on exposure to air. This is because the metal surface slowly undergoes oxidation to form an oxide. This oxide layer should be removed before using the alkaline earth metals. Melting and boiling points: Alkaline earth metals have a relatively high melting/ boiling point than alkali metals. This is because alkali metals use only one delocalized electron to form a weaker metallic bond/structure. Alkaline earth metals use two delocalized electrons to form a stronger metallic bond /structure. The melting and boiling points decrease down the group as the atomic radius/size increase reducing the strength of metallic bond and packing of the metallic structure. e.g. Beryllium has a melting point of 1280oC. Magnesium has a melting point of 650oC.Beryllium has a smaller atomic radius/size than magnesium .The strength of metallic bond and packing of the metallic structure is thus stronger in beryllium. Electrical/thermal conductivity: Alkaline earth metals are good thermal and electrical conductors. The two delocalized valence electrons move randomly within the metallic structure. Electrical conductivity increase down the group as the atomic radius/size increase making the delocalized outer electrons less attracted to nucleus. Alkaline earth metals are better thermal and electrical conductors than alkali metals because they have more/two outer delocalized electrons.e.g. Magnesium is a better conductor than sodium because it has more/two delocalized electrons than sodium. The more delocalized electrons the better the electrical conductor. Calcium is a better conductor than magnesium. Calcium has bigger/larger atomic radius than magnesium because the delocalized electrons are less attracted to the nucleus of calcium and thus more free /mobile and thus better the electrical conductor Summary of some physical properties of the 1st three alkaline earth metals
11 Alkaline earth metal Appearance Ease of cutting Melting point (oC) Boiling point (oC) Conduct- ivity 1st ionization energy 2nd ionization energy Beryllium Shiny grey Not easy 1280 3450 Good 900 1800 Magnesium Shiny grey Not Easy 650 1110 Good 736 1450 calcium Shiny grey Not easy 850 1140 Good 590 970 Chemical properties (i)Reaction with air/oxygen On exposure to air, the surface of alkaline earth metals is slowly oxidized to its oxide on prolonged exposure to air.
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The more delocalized electrons the better the electrical conductor. Calcium is a better conductor than magnesium. Calcium has bigger/larger atomic radius than magnesium because the delocalized electrons are less attracted to the nucleus of calcium and thus more free /mobile and thus better the electrical conductor Summary of some physical properties of the 1st three alkaline earth metals
11 Alkaline earth metal Appearance Ease of cutting Melting point (oC) Boiling point (oC) Conduct- ivity 1st ionization energy 2nd ionization energy Beryllium Shiny grey Not easy 1280 3450 Good 900 1800 Magnesium Shiny grey Not Easy 650 1110 Good 736 1450 calcium Shiny grey Not easy 850 1140 Good 590 970 Chemical properties (i)Reaction with air/oxygen On exposure to air, the surface of alkaline earth metals is slowly oxidized to its oxide on prolonged exposure to air. Example On exposure to air, the surface of magnesium ribbon is oxidized to form a thin film of Magnesium oxide . 2Mg(s) + O2(g) -> 2MgO(s) (ii)Burning in air/oxygen Experiment Hold a about 2cm length of Magnesium ribbon on a Bunsen flame. Stop heating when it catches fire/start burning. Caution: Do not look directly at the flame Put the products of burning into 100cm3 beaker. Add about 5cm3 of distilled water. Swirl. Test the mixture using litmus papers. Repeat with Calcium Observations -Magnesium burns with a bright blindening flame -White solid /ash produced -Solid dissolves in water to form a colourless solution -Blue litmus paper remain blue -Red litmus paper turns blue -colourless gas with pungent smell of urine Explanation Magnesium burns in air with a bright blindening flame to form a mixture of Magnesium oxide and Magnesium nitride. 2Mg (s) + O2(g) -> 2MgO(s) 3Mg (s) + N2 (g) -> Mg3N2 (s) Magnesium oxide dissolves in water to form magnesium hydroxide. 12 MgO(s) + H2O (l) -> Mg(OH)2(aq) Magnesium nitride dissolves in water to form magnesium hydroxide and produce ammonia gas.
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Repeat with Calcium Observations -Magnesium burns with a bright blindening flame -White solid /ash produced -Solid dissolves in water to form a colourless solution -Blue litmus paper remain blue -Red litmus paper turns blue -colourless gas with pungent smell of urine Explanation Magnesium burns in air with a bright blindening flame to form a mixture of Magnesium oxide and Magnesium nitride. 2Mg (s) + O2(g) -> 2MgO(s) 3Mg (s) + N2 (g) -> Mg3N2 (s) Magnesium oxide dissolves in water to form magnesium hydroxide. 12 MgO(s) + H2O (l) -> Mg(OH)2(aq) Magnesium nitride dissolves in water to form magnesium hydroxide and produce ammonia gas. Mg3N2 (s) + 6H2O(l) -> 3Mg(OH)2(aq) + 2NH3 (g) Magnesium hydroxide and ammonia are weakly alkaline with pH 8/9/10/11 and turns red litmus paper blue. Calcium burns in air with faint orange/red flame to form a mixture of both Calcium oxide and calcium nitride. 2Ca (s) + O2(g) -> 2CaO(s) 3Ca (s) + N2 (g) -> Ca3N2 (s) Calcium oxide dissolves in water to form calcium hydroxide. CaO(s) + H2O(l) -> Ca(OH)2(aq) Calcium nitride dissolves in water to form calcium hydroxide and produce ammonia gas. Ca3N2 (s) + 6H2O(l) -> 3Ca(OH)2(aq) + 2NH3 (g) Calcium hydroxide is also weakly alkaline solution with pH 8/9/10/11 and turns red litmus paper blue. (iii)Reaction with water Experiment Measure 50 cm3 of distilled water into a beaker. Scrub/polish with sand paper 1cm length of Magnesium ribbon Place it in the water. Test the product-mixture with blue and red litmus papers. Repeat with Calcium metal. Observations -Surface of magnesium covered by bubbles of colourless gas. -Colourless solution formed. -Effervescence/bubbles/fizzing takes place in Calcium.
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Observations -Surface of magnesium covered by bubbles of colourless gas. -Colourless solution formed. -Effervescence/bubbles/fizzing takes place in Calcium. -Red litmus paper turns blue. -Blue litmus paper remains blue. Explanations Magnesium slowly reacts with cold water to form Magnesium hydroxide and bubbles of Hydrogen gas that stick on the surface of the ribbon. Mg(s) + 2H2O (l) -> Mg(OH)2(aq) + H2 (g)
13 Calcium moderately reacts with cold water to form Calcium hydroxide and produce a steady stream of Hydrogen gas. Ca(s) + 2H2O (l) -> Ca(OH)2(aq) + H2 (g) (iv)Reaction with water vapour/steam Experiment Put some cotton wool soaked in water/wet sand in a long boiling tube. Coil a well polished magnesium ribbon into the boiling tube. Ensure the coil touches the side of the boiling tube. Heat the cotton wool/sand slightly then strongly heat the Magnesium ribbon . Set up of apparatus Observations -Magnesium glows red hot then burns with a blindening flame. -Magnesium continues to glow/burning even without more heating. -White solid/residue. -colourless gas collected over water. Explanation On heating wet sand, steam is generated which drives out the air that would otherwise react with /oxidize the ribbon. Magnesium burns in steam/water vapour generating enough heat that ensures the reaction goes to completion even without further heating. White Magnesium oxide is formed and hydrogen gas is evolved. To prevent suck back, the delivery tube should be removed from the water before heating is stopped at the end of the experiment. 14 Mg(s) + H2O (l) -> MgO(s) + H2 (g) (v)Reaction with chlorine gas. Experiment Lower slowly a burning magnesium ribbon/shavings into a gas jar containing Chlorine gas. Repeat with a hot piece of calcium metal. Observation -Magnesium continues to burn in chlorine with a bright blindening flame. -Calcium continues to burn for a short time. -White solid formed . -Pale green colour of chlorine fades. Explanation Magnesium continues to burn in chlorine gas forming white magnesium oxide solid. Mg(s) + Cl2 (g) -> MgCl2 (s) Calcium burns slightly in chlorine gas to form white calcium oxide solid.
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-Pale green colour of chlorine fades. Explanation Magnesium continues to burn in chlorine gas forming white magnesium oxide solid. Mg(s) + Cl2 (g) -> MgCl2 (s) Calcium burns slightly in chlorine gas to form white calcium oxide solid. Calcium oxide formed coat unreacted Calcium stopping further reaction Ca(s) + Cl2 (g) -> CaCl2 (s) (v)Reaction with dilute acids. Experiment Place about 4.0cm3 of 0.1M dilute sulphuric(VI)acid into a test tube. Add about 1.0cm length of magnesium ribbon into the test tube. Cover the mouth of the test tube using a thumb. Release the gas and test the gas using a burning splint. Repeat with about 4.0cm3 of 0.1M dilute hydrochloric/nitric(V) acid. Repeat with 0.1g of Calcium in a beaker with all the above acid Caution: Keep distance when using calcium Observation -Effervescence/fizzing/bubbles with dilute sulphuric(VI) and nitric(V) acids -Little Effervescence/fizzing/bubbles with calcium and dilute sulphuric(VI) acid. -Colourless gas produced that extinguishes a burning splint with an explosion/ “pop” sound. -No gas is produced with Nitric(V)acid. -Colourless solution is formed. Explanation Dilute acids react with alkaline earth metals to form a salt and produce hydrogen gas. 15 Nitric(V)acid is a strong oxidizing agent. It quickly oxidizes the hydrogen produced to water. Calcium is very reactive with dilute acids and thus a very small piece of very dilute acid should be used.
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15 Nitric(V)acid is a strong oxidizing agent. It quickly oxidizes the hydrogen produced to water. Calcium is very reactive with dilute acids and thus a very small piece of very dilute acid should be used. Chemical equations Mg(s) + H2SO4 (aq) -> MgSO4(aq) + H2 (g) Mg(s) + 2HNO3 (aq) -> Mg(NO3)2(aq) + H2 (g) Mg(s) + 2HCl (aq) -> MgCl2(aq) + H2 (g) Ca(s) + H2SO4 (aq) -> CaSO4(s) + H2 (g) (insoluble CaSO4(s) coat/cover Ca(s)) Ca(s) + 2HNO3 (aq) -> Ca(NO3)2(aq) + H2 (g) Ca(s) + 2HCl (aq) -> CaCl2(aq) + H2 (g) Ba(s) + H2SO4 (aq) -> BaSO4(s) + H2 (g) (insoluble BaSO4(s) coat/cover Ba(s)) Ba(s) + 2HNO3 (aq) -> Ba(NO3)2(aq) + H2 (g) Ba(s) + 2HCl (aq) -> BaCl2(aq) + H2 (g) The table below shows some compounds of some alkaline earth metals Some uses of alkaline earth metals include: (i)Magnesium hydroxide is a non-toxic/poisonous mild base used as an anti acid medicine to relieve stomach acidity. (ii)Making duralumin. Duralumin is an alloy of Magnesium and aluminium used for making aeroplane bodies because it is light.
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Chemical equations Mg(s) + H2SO4 (aq) -> MgSO4(aq) + H2 (g) Mg(s) + 2HNO3 (aq) -> Mg(NO3)2(aq) + H2 (g) Mg(s) + 2HCl (aq) -> MgCl2(aq) + H2 (g) Ca(s) + H2SO4 (aq) -> CaSO4(s) + H2 (g) (insoluble CaSO4(s) coat/cover Ca(s)) Ca(s) + 2HNO3 (aq) -> Ca(NO3)2(aq) + H2 (g) Ca(s) + 2HCl (aq) -> CaCl2(aq) + H2 (g) Ba(s) + H2SO4 (aq) -> BaSO4(s) + H2 (g) (insoluble BaSO4(s) coat/cover Ba(s)) Ba(s) + 2HNO3 (aq) -> Ba(NO3)2(aq) + H2 (g) Ba(s) + 2HCl (aq) -> BaCl2(aq) + H2 (g) The table below shows some compounds of some alkaline earth metals Some uses of alkaline earth metals include: (i)Magnesium hydroxide is a non-toxic/poisonous mild base used as an anti acid medicine to relieve stomach acidity. (ii)Making duralumin. Duralumin is an alloy of Magnesium and aluminium used for making aeroplane bodies because it is light. Beryllium Magnesium Calcium Barium Hydroxide Be(OH)2 Mg(OH)2 Ca(OH)2 Ba(OH)2 Oxide BeO MgO CaO BaO Sulphide - MgS CaS BaS Chloride BeCl2 MgCl2 CaCl2 BaCl2 Carbonate BeCO3 MgCO3 CaCO3 BaCO3 Nitrate(V) Be(NO3)2 Mg(NO3)2 Ca(NO3)2 Ba(NO3)2 Sulphate(VI) BeSO4 MgSO4 CaSO4 BaSO4 Sulphate(IV) - - CaSO3 BaSO3 Hydrogen carbonate - Mg(HCO3)2 Ca(HCO3)2 - Hydrogen sulphate(VI) - Mg(HSO4)2 Ca(HSO4)2 -
16 (iii) Making plaster of Paris-Calcium sulphate(VI) is used in hospitals to set a fractures bone.
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(ii)Making duralumin. Duralumin is an alloy of Magnesium and aluminium used for making aeroplane bodies because it is light. Beryllium Magnesium Calcium Barium Hydroxide Be(OH)2 Mg(OH)2 Ca(OH)2 Ba(OH)2 Oxide BeO MgO CaO BaO Sulphide - MgS CaS BaS Chloride BeCl2 MgCl2 CaCl2 BaCl2 Carbonate BeCO3 MgCO3 CaCO3 BaCO3 Nitrate(V) Be(NO3)2 Mg(NO3)2 Ca(NO3)2 Ba(NO3)2 Sulphate(VI) BeSO4 MgSO4 CaSO4 BaSO4 Sulphate(IV) - - CaSO3 BaSO3 Hydrogen carbonate - Mg(HCO3)2 Ca(HCO3)2 - Hydrogen sulphate(VI) - Mg(HSO4)2 Ca(HSO4)2 -
16 (iii) Making plaster of Paris-Calcium sulphate(VI) is used in hospitals to set a fractures bone. (iii)Making cement-Calcium carbonate is mixed with clay and sand then heated to form cement for construction/building. (iv)Raise soil pH-Quicklime/calcium oxide is added to acidic soils to neutralize and raise the soil pH in agricultural farms. (v)As nitrogenous fertilizer-Calcium nitrate(V) is used as an agricultural fertilizer because plants require calcium for proper growth. (vi)In the blast furnace-Limestone is added to the blast furnace to produce more reducing agent and remove slag in the blast furnace for extraction of Iron. (c)Group VII elements: Halogens Group VII elements are called Halogens. They are all non metals. They include: Element Symbol Atomic number Electronicc configuration Charge of ion Valency State at Room Temperature Fluorine Chlorine Bromine Iodine Astatine F Cl Br I At 9 17 35 53 85 2:7 2:8:7 2:8:18:7 2:8:18:18:7 2:8:18:32:18:7 F- Cl- Br- I- At- 1 1 1 1 1 Pale yellow gas Pale green gas Red liquid Grey Solid Radioactive All halogen atoms have seven electrons in the outer energy level. They acquire/gain one electron in the outer energy level to be stable.
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They are all non metals. They include: Element Symbol Atomic number Electronicc configuration Charge of ion Valency State at Room Temperature Fluorine Chlorine Bromine Iodine Astatine F Cl Br I At 9 17 35 53 85 2:7 2:8:7 2:8:18:7 2:8:18:18:7 2:8:18:32:18:7 F- Cl- Br- I- At- 1 1 1 1 1 Pale yellow gas Pale green gas Red liquid Grey Solid Radioactive All halogen atoms have seven electrons in the outer energy level. They acquire/gain one electron in the outer energy level to be stable. They therefore are therefore monovalent .They exist in oxidation state X- The number of energy levels increases down the group from Fluorine to Astatine. The more the number of energy levels the bigger/larger the atomic size. e.g. The atomic size/radius of Chlorine is bigger/larger than that of Fluorine because Chlorine has more/3 energy levels than Fluorine (2 energy levels). 17 Atomic radius and ionic radius of Halogens increase down the group as the number of energy levels increases. The atomic radius of Halogens is smaller than the ionic radius. This is because they react by gaining/acquiring extra one electron in the outer energy level. The effective nuclear attraction on the more/extra electrons decreases. The incoming extra electron is also repelled causing the outer energy level to expand to reduce the repulsion and accommodate more electrons. Table showing the atomic and ionic radius of four Halogens Element Symbol Atomic number Atomic radius(nM) Ionic radius(nM) Fluorine F 9 0.064 0.136 Chlorine Cl 17 0.099 0.181 Bromine Br 35 0.114 0.195 Iodine I 53 0.133 0.216 The atomic radius of Chlorine is 0.099nM .The ionic radius of Cl- is 0.181nM. This is because Chlorine atom/molecule reacts by gaining/acquiring extra one electrons. The more/extra electrons/energy level experience less effective nuclear attraction /pull towards the nucleus .The outer enegy level expand/increase to reduce the repulsion of the existing and incoming gained /acquired electrons.
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Table showing the atomic and ionic radius of four Halogens Element Symbol Atomic number Atomic radius(nM) Ionic radius(nM) Fluorine F 9 0.064 0.136 Chlorine Cl 17 0.099 0.181 Bromine Br 35 0.114 0.195 Iodine I 53 0.133 0.216 The atomic radius of Chlorine is 0.099nM .The ionic radius of Cl- is 0.181nM. This is because Chlorine atom/molecule reacts by gaining/acquiring extra one electrons. The more/extra electrons/energy level experience less effective nuclear attraction /pull towards the nucleus .The outer enegy level expand/increase to reduce the repulsion of the existing and incoming gained /acquired electrons. Electronegativity The ease of gaining/acquiring extra electrons is called electronegativity. All halogens are electronegative. Electronegativity decreases as atomic radius increase. This is because the effective nuclear attraction on outer electrons decreases with increase in atomic radius. The outer electrons experience less nuclear attraction and thus ease of gaining/acquiring extra electrons decrease. It is measured using Pauling’s scale. Where Fluorine with Pauling scale 4.0 is the most electronegative element and thus the highest tendency to acquire/gain extra electron. Table showing the electronegativity of the halogens. Halogen F Cl Br I At Electronegativity (Pauling scale) 4.0 3.0 2.8 2.5 2.2
18 The electronegativity of the halogens decrease down the group from fluorine to Astatine. This is because atomic radius increases down the group and thus decrease electron – attracting power down the group from fluorine to astatine. Fluorine is the most electronegative element in the periodic table because it has the small atomic radius. Electron affinity The minimum amount of energy required to gain/acquire an extra electron by an atom of element in its gaseous state is called 1st electron affinity. The SI unit of electron affinity is kilojoules per mole/kJmole-1 . Electron affinity depend on atomic radius. The higher the atomic radius, the less effective the nuclear attraction on outer energy level electrons and thus the lower the electron affinity.
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The SI unit of electron affinity is kilojoules per mole/kJmole-1 . Electron affinity depend on atomic radius. The higher the atomic radius, the less effective the nuclear attraction on outer energy level electrons and thus the lower the electron affinity. For halogens the 1st electron affinity decrease down the group as the atomic radius increase and the effective nuclear attraction on outer energy level electrons decrease. Due to its small size/atomic radius Fluorine shows exceptionally low electron affinity. This is because a lot of energy is required to overcome the high repulsion of the existing and incoming electrons. Table showing the election affinity of halogens for the process X + e -> X- The higher the electron affinity the more stable theion.i.e Cl- is a more stable ion than Br- because it has a more negative / exothermic electron affinity than Br- Electron affinity is different from: (i) Ionization energy. Ionization energy is the energy required to lose/donate an electron in an atom of an element in its gaseous state while electron affinity is the energy required to gain/acquire extra electron by an atom of an element in its gaseous state. (ii) Electronegativity. -Electron affinity is the energy required to gain an electron in an atom of an element in gaseous state. It involves the process: X(g) + e -> X-(g) Electronegativity is the ease/tendency of gaining/ acquiring electrons by an element during chemical reactions. Halogen F Cl Br I Electron affinity kJmole-1 -333 -364 -342 -295
19 It does not involve use of energy but theoretical arbitrary Pauling’ scale of measurements. Physical properties State at room temperature Fluorine and Chlorine are gases, Bromine is a liquid and Iodine is a solid. Astatine is radioactive . All halogens exist as diatomic molecules bonded by strong covalent bond. Each molecule is joined to the other by weak intermolecular forces/ Van-der-waals forces. Melting/Boiling point The strength of intermolecular/Van-der-waals forces of attraction increase with increase in molecular size/atomic radius. Iodine has therefore the largest atomic radius and thus strongest intermolecular forces to make it a solid. Iodine sublimes when heated to form (caution: highly toxic/poisonous) purple vapour.
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Melting/Boiling point The strength of intermolecular/Van-der-waals forces of attraction increase with increase in molecular size/atomic radius. Iodine has therefore the largest atomic radius and thus strongest intermolecular forces to make it a solid. Iodine sublimes when heated to form (caution: highly toxic/poisonous) purple vapour. This is because Iodine molecules are held together by weak van-derwaals/intermolecular forces which require little heat energy to break. Electrical conductivity All Halogens are poor conductors of electricity because they have no free delocalized electrons. Solubility in polar and non-polar solvents All halogens are soluble in water(polar solvent). When a boiling tube containing either chlorine gas or bromine vapour is separately inverted in a beaker containing distilled water and tetrachloromethane (non-polar solvent), the level of solution in boiling tube rises in both water and tetrachloromethane. This is because halogen are soluble in both polar and non-polar solvents. Solubility of halogens in water/polar solvents decrease down the group. Solubility of halogens in non-polar solvent increase down the group. The level of water in chlorine is higher than in bromine and the level of tetrachloromethane in chlorine is lower than in bromine. Caution: Tetrachloromethane , Bromine vapour and Chlorine gas are all highly toxic/poisonous. 20 Table showing the physical properties of Halogens Halogen Formula of molecule Electrical conductivity Solubility in water Melting point(oC) Boiling point(oC) Fluorine F2 Poor Insoluble/soluble in tetrachloromethane -238 -188 Chlorine Cl2 Poor Insoluble/soluble in tetrachloromethane -101 -35 Bromine Br2 Poor Insoluble/soluble in tetrachloromethane 7 59 Iodine I2 Poor Insoluble/soluble in tetrachloromethane 114 sublimes Chemical properties (i)Displacement Experiment Place separately in test tubes about 5cm3 of sodium chloride, Sodium bromide and Sodium iodide solutions.
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The level of water in chlorine is higher than in bromine and the level of tetrachloromethane in chlorine is lower than in bromine. Caution: Tetrachloromethane , Bromine vapour and Chlorine gas are all highly toxic/poisonous. 20 Table showing the physical properties of Halogens Halogen Formula of molecule Electrical conductivity Solubility in water Melting point(oC) Boiling point(oC) Fluorine F2 Poor Insoluble/soluble in tetrachloromethane -238 -188 Chlorine Cl2 Poor Insoluble/soluble in tetrachloromethane -101 -35 Bromine Br2 Poor Insoluble/soluble in tetrachloromethane 7 59 Iodine I2 Poor Insoluble/soluble in tetrachloromethane 114 sublimes Chemical properties (i)Displacement Experiment Place separately in test tubes about 5cm3 of sodium chloride, Sodium bromide and Sodium iodide solutions. Add 5 drops of chlorine water to each test tube: Repeat with 5 drops of bromine water instead of chlorine water Observation Using Chlorine water -Yellow colour of chlorine water fades in all test tubes except with sodium chloride. -Coloured Solution formed. Using Bromine water Yellow colour of bromine water fades in test tubes containing sodium iodide. -Coloured Solution formed. Explanation The halogens displace each other from their solution. The more electronegative displace the less electronegative from their solution. Chlorine is more electronegative than bromine and iodine. On adding chlorine water, bromine and Iodine are displaced from their solutions by chlorine. Bromine is more electronegative than iodide but less 6than chlorine. 21 On adding Bromine water, iodine is displaced from its solution but not chlorine.
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On adding chlorine water, bromine and Iodine are displaced from their solutions by chlorine. Bromine is more electronegative than iodide but less 6than chlorine. 21 On adding Bromine water, iodine is displaced from its solution but not chlorine. Table showing the displacement of the halogens (V) means there is displacement (x ) means there is no displacement Chemical /ionic equations With Fluorine F2(g) + 2NaCl-(aq) -> 2NaF(aq) + Cl2(aq) F2(g) + 2Cl-(aq) -> 2F-(aq) + Cl2(aq) F2(g) + 2NaBr-(aq) -> 2NaF(aq) + Br2(aq) F2(g) + 2Br-(aq) -> 2F-(aq) + Br2(aq) F2(g) + 2NaI-(aq) -> 2NaF(aq) + I2(aq) F2(g) + 2I-(aq) -> 2F-(aq) + I2(aq) With chlorine Cl2(g) + 2NaCl-(aq) -> 2NaCl(aq) + Br2(aq) Cl2(g) + 2Br-(aq) -> 2Cl-(aq) + Br2(aq) Cl2(g) + 2NaI-(aq) -> 2NaCl(aq) + I2(aq) Cl2(g) + 2I-(aq) -> 2Cl-(aq) + I2(aq) With Bromine Br2(g) + 2NaI-(aq) -> 2NaBr(aq) + I2(aq) Br2(g) + 2I-(aq) -> 2Br-(aq) + I2(aq) Halogen ion in solution Halogen F- Cl- Br- I- F2 X Cl2 X X Br2 X X X I2 X X X X
22 Uses of halogens (i) Florine – manufacture of P.T.F.E (Poly tetra fluoroethene) synthetic fiber. - Reduce tooth decay when added in small amounts/quantities in tooth paste. NB –large small quantities of fluorine /fluoride ions in water cause browning of teeth/flourosis. - Hydrogen fluoride is used to engrave words /pictures in glass.
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- Reduce tooth decay when added in small amounts/quantities in tooth paste. NB –large small quantities of fluorine /fluoride ions in water cause browning of teeth/flourosis. - Hydrogen fluoride is used to engrave words /pictures in glass. (ii) Bromine - Silver bromide is used to make light sensitive photographic paper/films. (iii) Iodide – Iodine dissolved in alcohol is used as medicine to kill bacteria in skin cuts. It is called tincture of iodine. The table below to show some compounds of halogens. (i) Below is the table showing the bond energy of four halogens. Bond Bond energy k J mole-1 Cl-Cl 242 Br-Br 193 I-I 151 I. What do you understand by the term “bond energy” Bond energy is the energy required to break/ form one mole of chemical bond II. Explain the trend in bond Energy of the halogens above: -Decrease down the group from chlorine to Iodine Element Halogen H Na Mg Al Si C P F HF NaF MgF2 AlF3 SiF4 CF4 PF3 Cl HCl NaCl MgCl2 AlCl3 SiCl 4 CCl4 PCl3 Br HBr NaBr MgBr2 AlBr3 SiBr4 CBr4 PBr3 I Hl Nal Mgl2 All3 SiI4 C l 4 PBr3
23 -Atomic radius increase down the group decreasing the energy required to break the covalent bonds between the larger atom with reduced effective nuclear @ charge an outer energy level that take part in bonding. (c)Group VIII elements: Noble gases Group VIII elements are called Noble gases. They are all non metals. Noble gases occupy about 1.0% of the atmosphere as colourless gaseous mixture. Argon is the most abundant with 0.9%. They exists as monatomic molecules with very weak van-der-waals /intermolecular forces holding the molecules. They include: All noble gas atoms have a stable duplet(two electrons in the 1st energy level) or octet(eight electrons in other outer energy level)in the outer energy level. They therefore do not acquire/gain extra electron in the outer energy level or donate/lose. They therefore are therefore zerovalent . The number of energy levels increases down the group from Helium to Randon.
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They therefore do not acquire/gain extra electron in the outer energy level or donate/lose. They therefore are therefore zerovalent . The number of energy levels increases down the group from Helium to Randon. The more the number of energy levels the bigger/larger the atomic size/radius. e.g. The atomic size/radius of Argon is bigger/larger than that of Neon because Argon has more/3 energy levels than Neon (2 energy levels). Atomic radius noble gases increase down the group as the number of energy levels increases. The effective nuclear attraction on the outer electrons thus decrease down the group. The noble gases are generally unreactive because the outer energy level has the stable octet/duplet. The stable octet/duplet in noble gas atoms lead to a comparatively very high 1st ionization energy. This is because losing /donating an electron from the stable atom require a lot of energy to lose/donate and make it unstable. Element Symbol Atomic number Electron structure State at room temperature Helium He 2 2: Colourless gas Neon Ne 10 2:8 Colourless gas Argon Ar 18 2:8:8 Colourless gas Krypton Kr 36 2:8:18:8 Colourless gas Xenon Xe 54 2:8:18:18:8 Colourless gas Radon Rn 86 2:8:18:32:18:8 Radioctive
24 As atomic radius increase down the group and the 1st ionization energy decrease, very electronegative elements like Oxygen and Fluorine are able to react and bond with lower members of the noble gases.e.g Xenon reacts with Fluorine to form a covalent compound XeF6.This is because the outer electrons/energy level if Xenon is far from the nucleus and thus experience less effective nuclear attraction. Noble gases have low melting and boiling points. This is because they exist as monatomic molecules joined by very weak intermolecular/van-der-waals forces that require very little energy to weaken and form liquid and break to form a gas. The intermolecular/van-der-waals forces increase down the group as the atomic radius/size increase from Helium to Radon. The melting and boiling points thus increase also down the group. Noble gases are insoluble in water and are poor conductors of electricity.
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The intermolecular/van-der-waals forces increase down the group as the atomic radius/size increase from Helium to Radon. The melting and boiling points thus increase also down the group. Noble gases are insoluble in water and are poor conductors of electricity. Element Formula of molecule Electrical conductivity Solubility in water Atomic radius(nM) 1st ionization energy Melting point(0C) Boiling point(0C) Helium He Poor Insoluble 0.128 2372 -270 -269 Neon Ne Poor Insoluble 0.160 2080 -249 -246 Argon Ar Poor Insoluble 0.192 1520 -189 -186 Krypton Kr Poor Insoluble 0.197 1350 -157 -152 Xenon Xe Poor Insoluble 0.217 1170 -112 -108 Radon Rn Poor Insoluble 0.221 1134 -104 -93 Uses of noble gases Argon is used in light bulbs to provide an inert environment to prevent oxidation of the bulb filament Argon is used in arch welding as an insulator. Neon is used in street and advertisement light Helium is mixed with Oxygen during deep sea diving and mountaineering. Helium is used in weather balloon for meteorological research instead of Hydrogen because it is unreactive/inert. Hydrogen when impure can ignite with an explosion. Helium is used in making thermometers for measuring very low temperatures. STRUCTURE AND BONDING IONIC (ELECTROVALENT) BONDING Noble gases like neon or argon have eight electrons in their outer shells (or two in the case of helium). These noble gas structures are thought of as being in some way a "desirable" thing for an atom to have. When other atoms react, they try to organise electrons such that their outer shells are either completely full or completely empty. Chemical reactions occur so that atoms attain inert gas configuration by either losing valency electrons as in the case of metals, or gaining electrons as in the case of non metals. Ionic bonding in sodium chloride Sodium (2,8,1) has 1 electron more than a stable noble gas structure (2,8). If it gave away that electron it would become more stable. Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8).
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Ionic bonding in sodium chloride Sodium (2,8,1) has 1 electron more than a stable noble gas structure (2,8). If it gave away that electron it would become more stable. Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8). If it could gain an electron from somewhere it too would become more stable. If a sodium atom gives an electron to a chlorine atom, both become more stable. The sodium has lost an electron, so it no longer has equal numbers of electrons and protons. Because it has one more proton than electron, it has a charge of 1+. If electrons are lost from an atom, positive ions are formed. Positive ions are sometimes called cations because they move to the cathode during electrolysis. The chlorine has gained an electron, so it now has one more electron than proton. It therefore has a charge of 1-. If electrons are gained by an atom, negative ions are formed. A negative ion is sometimes called an anion since it drifts to the anode during electrolysis. The nature of ionic bond The sodium ions and chloride ions are held together by the strong electrostatic attractions between the positive and negative charges. You need one sodium atom to provide the extra electron for one chlorine atom, so they combine together 1:1. The formula is therefore NaCl. Properties of ionic compounds All compounds with ionic bonding produce giant ionic structures. Consist of oppositely charged ions arranged in an ionic lattice, the ions are held together by strong ionic bonds. e.g. NaCl is composed of Na+ ions and Cl- ions. These bonds are hard to break, therefore ionic substances have very high melting and boiling points. All exist as solids. They conduct electricity when molten, because the ions are free to move, but do not conduct when solid. They conduct electricity in the aqueous state because the ions are free to move. Most ionic substances are soluble in water because the polar water molecules can accommodate the charged ions. COVALENT BONDING - SINGLE BONDS As well as achieving noble gas structures by transferring electrons from one atom to another as in ionic bonding, it is also possible for atoms to reach these stable structures by sharing electrons to give covalent bonds.
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They conduct electricity in the aqueous state because the ions are free to move. Most ionic substances are soluble in water because the polar water molecules can accommodate the charged ions. COVALENT BONDING - SINGLE BONDS As well as achieving noble gas structures by transferring electrons from one atom to another as in ionic bonding, it is also possible for atoms to reach these stable structures by sharing electrons to give covalent bonds. Depending on the number of electron pairs shared between atoms which participate in bonding, covalent bonds are classified as follows: Some simple covalent molecules Chlorine For example, two chlorine atoms could both achieve stable structures by sharing their single unpaired electron as in the diagram. The fact that one chlorine has been drawn with electrons marked as crosses and the other as dots is simply to show where all the electrons come from. In reality there is no difference between them. The two chlorine atoms are said to be joined by a covalent bond. The reason that the two chlorine atoms stick together is that the shared pair of electrons is attracted to the nucleus of both chlorine atoms. Hydrogen Hydrogen atoms only need two electrons in their outer level to reach the noble gas structure of helium. Once again, the covalent bond holds the two atoms together because the pair of electrons is attracted to both nuclei. This is another single bond. Hydrogen chloride
The hydrogen has a helium structure, and the chlorine an argon structure. Water Oxygen atom has six electrons in the outer shell, while each of the two hydrogen atoms has one each. After bonding, oxygen has 8 electrons while each hydrogen atom has two as shown by the molecule. NITROGEN GAS Each nitrogen atom has five electrons in the outer shell. Each needs 3 electrons to complete the outer shell. In the formation of the molecule, each nitrogen atom contributes three electrons and a triple bond is formed Characteristics of Covalent Compounds 1) Covalent compounds consist of molecules and not ions. The molecules do not have any electric charge on them. The molecules are held together by weak forces called Van der Waal's forces. 2) Covalent compounds are gases, volatile liquids or soft solids. As there are weak, Van der Waal's forces between the molecules, they are not held in rigid position. The state depends on the bond energy. If the bond energy is very low, they stay as gases, if it is appreciable they are volatile liquids.
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As there are weak, Van der Waal's forces between the molecules, they are not held in rigid position. The state depends on the bond energy. If the bond energy is very low, they stay as gases, if it is appreciable they are volatile liquids. If very high, they exist as soft solids. 3) Covalent compounds generally have low melting and boiling points. As Van der Waal's forces are weak, a very small amount of energy is required to break the bond between the molecules corresponding to low melting point and boiling point. 4) Covalent compounds dissolve in organic solvents. As they do not contain ions, solvation does not take place when water is added to the compound. Hence they do not dissolve in water. 5) Covalent compounds are bad conductors of electricity. They do not contain ions in the fused state, nor do ions migrate on application of an electric potential. Hence, there is no conduction of current. 6) Covalent compounds are less dense when compared to water. Very weak Van der Waal's forces hold the molecules together, hence there are large inter molecular spaces. Consequently less number of molecules per unit volume, which means mass per unit volume is also less. Hence they have a low density. Exceptions Diamond and graphite, the allotropes of carbon have high melting point. Hydrogen chloride in the aqueous state conducts electricity. The giant covalent structure of diamond Carbon has an electronic arrangement of 2, 4. In diamond, each carbon shares electrons with four other carbon atoms - forming four single bonds. In the diagram some carbon atoms only seem to be forming two bonds (or even one bond), but that's not really the case. We are only showing a small bit of the whole structure. This is a giant covalent structure - it continues on and on in three dimensions. It is not a molecule, because the number of atoms joined up in a real diamond is completely variable - depending on the size of the crystal. The physical properties of diamond Diamond Has a very high melting point (almost 4000°C). Very strong carbon-carbon covalent bonds have to be broken throughout the structure before melting occurs. Is very hard. This is again due to the need to break very strong covalent bonds operating in 3dimensions. Doesnಬt conduct electricity. All the electrons are held tightly between the atoms, and aren't free to move.
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This is again due to the need to break very strong covalent bonds operating in 3dimensions. Doesnಬt conduct electricity. All the electrons are held tightly between the atoms, and aren't free to move. Is insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently bound carbon atoms. The giant covalent structure of graphite Graphite has a layer structure which is quite difficult to draw convincingly in three dimensions. The diagram below shows the arrangement of the atoms in each layer, and the way the layers are spaced. The bonding in graphite Each carbon atom uses three of its electrons to form simple bonds to its three close neighbours. That leaves a fourth electron in the bonding level. These "spare" electrons in each carbon atom become delocalised over the whole of the sheet of atoms in one layer. They are no longer associated directly with any particular atom or pair of atoms, but are free to wander throughout the whole sheet. The important thing is that the delocalised electrons are free to move anywhere within the sheet - each electron is no longer fixed to a particular carbon atom. There is, however, no direct contact between the delocalised electrons in one sheet and those in the neighbouring sheets. The atoms within a sheet are held together by strong covalent bonds - stronger, in fact, than in diamond because of the additional bonding caused by the delocalised electrons. So what holds the sheets together? In graphite you have the ultimate example of van der Waals dispersion forces. As the delocalised electrons move around in the sheet, very large temporary dipoles can be set up which will induce opposite dipoles in the sheets above and below - and so on throughout the whole graphite crystal. The physical properties of graphite Graphite Has a high melting point, similar to that of diamond. In order to melt graphite, it isn't enough to loosen one sheet from another. You have to break the covalent bonding throughout the whole structure. Has a soft, slippery feel, and is used in pencils and as a dry lubricant for things like locks. You can think of graphite rather like a pack of cards - each card is strong, but the cards will slide over each other, or even fall off the pack altogether. When you use a pencil, sheets are rubbed off and stick to the paper. Has a lower density than diamond.
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You can think of graphite rather like a pack of cards - each card is strong, but the cards will slide over each other, or even fall off the pack altogether. When you use a pencil, sheets are rubbed off and stick to the paper. Has a lower density than diamond. This is because of the relatively large amount of space that is "wasted" between the sheets. Is insoluble in water and organic solvents - for the same reason that diamond is insoluble. Attractions between solvent molecules and carbon atoms will never be strong enough to overcome the strong covalent bonds in graphite. Conducts electricity. The delocalised electrons are free to move throughout the sheets. If a piece of graphite is connected into a circuit, electrons can fall off one end of the sheet and be replaced with new ones at the other end. The structure of silicon dioxide, SiO2 Silicon dioxide is also known as silicon (IV) oxide. The giant covalent structure of silicon dioxide There are three different crystal forms of silicon dioxide. The easiest one to remember and draw is based on the diamond structure. Crystalline silicon has the same structure as diamond. To turn it into silicon dioxide, all you need to do is to modify the silicon structure by including some oxygen atoms. Notice that each silicon atom is bridged to its neighbours by an oxygen atom. Don't forget that this is just a tiny part of a giant structure extending on all 3 dimensions. The physical properties of silicon dioxide Silicon dioxide Has a high melting point - varying depending on what the particular structure is (remember that the structure given is only one of three possible structures), but around 1700°C. Very strong silicon-oxygen covalent bonds have to be broken throughout the structure before melting occurs. Is hard. This is due to the need to break the very strong covalent bonds. Doesnಬt conduct electricity. There aren't any delocalised electrons. All the electrons are held tightly between the atoms, and aren't free to move. Is insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and the silicon or oxygen atoms which could overcome the covalent bonds in the giant structure. Uses of Silica i) Quartz glass is used for manufacturing optical instruments. ii) Colored quartz is used for manufacturing gems. iii) Sand is used in manufacture of glass, porcelain, sand paper and mortar etc.
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Uses of Silica i) Quartz glass is used for manufacturing optical instruments. ii) Colored quartz is used for manufacturing gems. iii) Sand is used in manufacture of glass, porcelain, sand paper and mortar etc. iv) Sand stone is used as a building material. CO-ORDINATE (DATIVE COVALENT) BONDING Co-ordinate (dative covalent) bonding A covalent bond is formed by two atoms sharing a pair of electrons. The atoms are held together because the electron pair is attracted by both of the nuclei. In the formation of a simple covalent bond, each atom supplies one electron to the bond - but that doesn't have to be the case. A co-ordinate bond (also called a
dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom. The reaction between ammonia and hydrogen chloride If these colourless gases are allowed to mix, a thick white smoke of solid ammonium chloride is formed. Ammonium ions, NH4+, are formed by the transfer of a hydrogen ion from the hydrogen chloride to the lone pair of electrons on the ammonia molecule. When the ammonium ion, NH4+, is formed, the fourth hydrogen is attached by a dative covalent bond, because only the hydrogen's nucleus is transferred from the chlorine to the nitrogen. The hydrogen's electron is left behind on the chlorine to form a negative chloride ion. Once the ammonium ion has been formed it is impossible to tell any difference between the dative covalent and the ordinary covalent bonds. Although the electrons are shown differently in the diagram, there is no difference between them in reality. INTERMOLECULAR BONDING - VAN DER WAALS FORCES (a) VAN DER WAALS FORCES Intermolecular attractions are attractions between one molecule and a neighbouring molecule. The forces of attraction which hold an individual molecule together (for example, the covalent bonds) are known as
intramolecular attractions. All molecules experience intermolecular attractions, although in some cases those attractions are very weak. Even in a gas like hydrogen, H2, if you slow the molecules down by cooling the gas, the attractions are large enough for the molecules to stick together eventually to form a liquid and then a solid.
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The forces of attraction which hold an individual molecule together (for example, the covalent bonds) are known as
intramolecular attractions. All molecules experience intermolecular attractions, although in some cases those attractions are very weak. Even in a gas like hydrogen, H2, if you slow the molecules down by cooling the gas, the attractions are large enough for the molecules to stick together eventually to form a liquid and then a solid. In hydrogen's case the attractions are so weak that the molecules have to be cooled to (-252°C) before the attractions are enough to condense the hydrogen as a liquid. Helium's intermolecular attractions are even weaker - the molecules won't stick together to form a liquid until the temperature drops to (-269°C). HYDROGEN BONDING Polar molecules, such as water molecules, have a weak, partial negative charge at one region of the molecule (the oxygen atom in water) and a partial positive charge elsewhere (the hydrogen atoms in water). Thus when water molecules are close together, their positive and negative regions are attracted to the oppositely-charged regions of nearby molecules. The force of attraction, shown here as a dotted line, is called a hydrogen bond. Each water molecule is hydrogen bonded to four others. The hydrogen bonds that form between water molecules account for some of the essential — and unique — properties of water. The attraction created by hydrogen bonds keeps water liquid over a wider range of temperature than is found for any other molecule its size. The energy required to break multiple hydrogen bonds causes water to have a high heat of vaporization; that is; a large amount of energy is needed to convert liquid water, where the molecules are attracted through their hydrogen bonds, to water vapor, where they are not. Liquid Water and Hydrogen Bonding Why water is a liquid? In many ways, water is a miracle liquid. Since the hydrogen and oxygen atoms in the molecule carry opposite (though partial) charges, nearby water molecules are attracted to each other like tiny little magnets. Hydrogen bonding makes water molecules "stick" together. This makes water have high melting and boiling points compared to other covalent compounds such as ammonia (NH3) which have similar molecular mass but are gases Ice and Hydrogen Bonding The structure that forms in the solid ice crystal actually has large holes in it. Therefore, in a given volume of ice, there are fewer water molecules than in the same volume of liquid water.
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Hydrogen bonding makes water molecules "stick" together. This makes water have high melting and boiling points compared to other covalent compounds such as ammonia (NH3) which have similar molecular mass but are gases Ice and Hydrogen Bonding The structure that forms in the solid ice crystal actually has large holes in it. Therefore, in a given volume of ice, there are fewer water molecules than in the same volume of liquid water. In other words, ice is less dense than liquid water and will float on the surface of the liquid. Surface Tension and hydrogen bonding As we just discussed, neighboring water molecules are attracted to one another. Molecules at the surface of liquid water have fewer neighbors and, as a result, have a greater attraction to the few water molecules that are nearby. This enhanced attraction is called surface tension. It makes the surface of the liquid slightly more difficult to break through than the interior. Water as a Solvent The partial charge that develops across the water molecule helps make it an excellent solvent. Water dissolves many substances by surrounding charged particles and "pulling" them into solution. For example, common table salt, sodium chloride, is an ionic substance that contains alternating sodium and chlorine ions. When table salt is added to water, the partial charges on the water molecule are attracted to the Na+ and Cl- ions. Why does ethanol have a higher boiling point than methoxymethane? Ethanol, CH3CH2-O-H, and methoxymethane, CH3-O-CH3, both have the same molecular formula, C2H6O. They have the same number of electrons, and a similar length to the molecule. The van der Waals attractions (both dispersion forces and dipole-dipole attractions) in each will be much the same. However, ethanol has a hydrogen atom attached directly to oxygen - and that oxygen still has exactly the same two lone pairs as in a water molecule. Hydrogen bonding can occur between ethanol molecules, although not as effectively as in water. The hydrogen bonding is limited by the fact that there is only one hydrogen in each ethanol molecule with sufficient + charge. In methoxymethane, the lone pairs on the oxygen are still there, but the hydrogens aren't sufficiently + for hydrogen bonds to form. Except in some rather unusual cases, the hydrogen atom has to be attached directly to the very electronegative element for hydrogen bonding to occur.
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The hydrogen bonding is limited by the fact that there is only one hydrogen in each ethanol molecule with sufficient + charge. In methoxymethane, the lone pairs on the oxygen are still there, but the hydrogens aren't sufficiently + for hydrogen bonds to form. Except in some rather unusual cases, the hydrogen atom has to be attached directly to the very electronegative element for hydrogen bonding to occur. The boiling points of ethanol and methoxymethane show the dramatic effect that the hydrogen bonding has on the stickiness of the ethanol molecules:
ethanol (with hydrogen bonding) 78.5°C methoxymethane (without hydrogen bonding) -24.8°C The hydrogen bonding in the ethanol has lifted its boiling point about 100°C.It is important to realise that hydrogen bonding exists in addition to van der Waals attractions. For example, all the following molecules contain the same number of electrons, and the first two are much the same length. The higher boiling point of the butan-1-ol is due to the additional hydrogen bonding. 4. BONDING IN METALS Bonding in metals Metal atoms have relatively few electrons in their outer shells. When they are packed together, each metal atom loses its outer electrons into a ಫseaಬ of free electrons (or mobile electrons). Having lost electrons, the atoms are no longer electrically neutral. They become positive ions because they have lost electrons but the number of protons in the nucleus has remained unchanged. Therefore the structure of a metal is made up of positive ions packed together. These ions are surrounded by electrons, which can move freely between the ions. An ion is a charged particle made from an atom by the loss or gain of electrons. Metal atoms most easily lose electrons, so they become positive ions. In doing so they achieve a more stable electron arrangement, usually that of the nearest noble gas. These free electrons are delocalized (not restricted to orbiting one positive ion) and form a kind of electrostatic ಫglueಬ holding the structure together. In an electrical circuit, metals can conduct electricity because the mobile electrons can move through the structure carrying charge. His type of bonding (called metallic boding) is present in alloys as well. Alloys, for example solder and brass, will conduct electricity. The physical properties of metals: This strong bonding generally results in dense, strong materials with high melting and boiling points.
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His type of bonding (called metallic boding) is present in alloys as well. Alloys, for example solder and brass, will conduct electricity. The physical properties of metals: This strong bonding generally results in dense, strong materials with high melting and boiling points. Usually a relatively large amount of energy is needed to melt or boil metals. a. Metals are good conductors of electricity because these 'free' electrons carry the charge of an electric current when a potential difference (voltage!) is applied across a piece of metal. b. Metals are also good conductors of heat. This is also due to the free moving electrons. Nonmetallic solids conduct heat energy by hotter more strongly vibrating atoms, knocking against cooler less strongly vibrating atoms to pass the particle kinetic energy on. In metals, as well as this effect, the 'hot' high kinetic energy electrons move around freely to transfer the particle kinetic energy more efficiently to 'cooler' atoms. c. Typical metals also have a silvery surface but remember this may be easily tarnished by corrosive oxidation in air and water. d. Unlike ionic solids, metals are very malleable, they can be readily bent, pressed or hammered into shape. INTRODUCTION TO SALTS 1.(a) A salt is an ionic compound formed when the cation from a base combine with the anion derived from an acid. A salt is therefore formed when the hydrogen ions in an acid are replaced wholly/fully or partially/partly ,directly or indirectly by a metal or ammonium radical. (b) The number of ionizable/replaceable hydrogen in an acid is called basicity of an acid. Some acids are therefore: (i)monobasic acids generally denoted HX e.g. HCl, HNO3,HCOOH,CH3COOH. (ii)dibasic acids ; generally denoted H2X e.g. H2SO4, H2SO3, H2CO3,HOOCOOH. (iii)tribasic acids ; generally denoted H3X e.g. H3PO4. (c) Some salts are normal salts while other are acid salts. (i)A normal salt is formed when all the ionizable /replaceable hydrogen in an acid is replaced by a metal or metallic /ammonium radical. (ii)An acid salt is formed when part/portion the ionizable /replaceable hydrogen in an acid is replaced by a metal or metallic /ammonium radical.
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(c) Some salts are normal salts while other are acid salts. (i)A normal salt is formed when all the ionizable /replaceable hydrogen in an acid is replaced by a metal or metallic /ammonium radical. (ii)An acid salt is formed when part/portion the ionizable /replaceable hydrogen in an acid is replaced by a metal or metallic /ammonium radical. Table showing normal and acid salts derived from common acids Acid name Chemical formula Basicity Normal salt Acid salt Hydrochloric acid HCl Monobasic Chloride(Cl-) None Nitric(V)acid HNO3 Monobasic Nitrate(V)(NO3-) None Nitric(III)acid HNO2 Monobasic Nitrate(III)(NO2-) None Sulphuric(VI)acid H2SO4 Dibasic Sulphate(VI) (SO42-) Hydrogen sulphate(VI) (HSO4-) Sulphuric(IV)acid H2SO3 Dibasic Sulphate(IV) (SO32-) Hydrogen sulphate(IV) (HSO3-)
Carbonic(IV)acid H2CO3 Dibasic Carbonate(IV)(CO32-) Hydrogen carbonate(IV) (HCO3-) Phosphoric(V) acid H3PO4 Tribasic Phosphate(V)(PO43-) Dihydrogen phosphate(V) (H2PO42-) Hydrogen diphosphate(V) (HP2O42-) The table below show shows some examples of salts.
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(i)A normal salt is formed when all the ionizable /replaceable hydrogen in an acid is replaced by a metal or metallic /ammonium radical. (ii)An acid salt is formed when part/portion the ionizable /replaceable hydrogen in an acid is replaced by a metal or metallic /ammonium radical. Table showing normal and acid salts derived from common acids Acid name Chemical formula Basicity Normal salt Acid salt Hydrochloric acid HCl Monobasic Chloride(Cl-) None Nitric(V)acid HNO3 Monobasic Nitrate(V)(NO3-) None Nitric(III)acid HNO2 Monobasic Nitrate(III)(NO2-) None Sulphuric(VI)acid H2SO4 Dibasic Sulphate(VI) (SO42-) Hydrogen sulphate(VI) (HSO4-) Sulphuric(IV)acid H2SO3 Dibasic Sulphate(IV) (SO32-) Hydrogen sulphate(IV) (HSO3-)
Carbonic(IV)acid H2CO3 Dibasic Carbonate(IV)(CO32-) Hydrogen carbonate(IV) (HCO3-) Phosphoric(V) acid H3PO4 Tribasic Phosphate(V)(PO43-) Dihydrogen phosphate(V) (H2PO42-) Hydrogen diphosphate(V) (HP2O42-) The table below show shows some examples of salts. Base/alkali Cation Acid Anion Salt Chemical name of salts NaOH Na+ HCl Cl NaCl Sodium(I)chloride Mg(OH)2 Mg2+ H2SO4 SO42 MgSO4 Mg(HSO4)2 Magnesium sulphate(VI) Magnesium hydrogen sulphate(VI) Pb(OH)2 Pb2+ HNO3 NO3 Pb(NO3)2 Lead(II)nitrate(V) Ba(OH)2 Ba2+ HNO3 NO3 Ba(NO3)2 Barium(II)nitrate(V) Ca(OH)2 Ba2+ H2SO4 SO42 MgSO4 Calcium sulphate(VI) NH4OH NH4+ H3PO4 PO43 (NH4 )3PO4 (NH4 )2HPO4 NH4 H2PO4 Ammonium phosphate(V) Diammonium phosphate(V) Ammonium diphosphate(V) KOH K+ H3PO4 PO43 K3PO4 Potassium phosphate(V) Al(OH)3 Al3+ H2SO4 SO42 Al2(SO4)2 Aluminium(III)sulphate(VI) Fe(OH)2 Fe2+ H2SO4 SO42 FeSO4 Iron(II)sulphate(VI)
Fe(OH)3 Fe3+ H2SO4 SO42 Fe2(SO4)2 Iron(III)sulphate(VI) (d) Some salts undergo hygroscopy, deliquescence and efflorescence.
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(ii)An acid salt is formed when part/portion the ionizable /replaceable hydrogen in an acid is replaced by a metal or metallic /ammonium radical. Table showing normal and acid salts derived from common acids Acid name Chemical formula Basicity Normal salt Acid salt Hydrochloric acid HCl Monobasic Chloride(Cl-) None Nitric(V)acid HNO3 Monobasic Nitrate(V)(NO3-) None Nitric(III)acid HNO2 Monobasic Nitrate(III)(NO2-) None Sulphuric(VI)acid H2SO4 Dibasic Sulphate(VI) (SO42-) Hydrogen sulphate(VI) (HSO4-) Sulphuric(IV)acid H2SO3 Dibasic Sulphate(IV) (SO32-) Hydrogen sulphate(IV) (HSO3-)
Carbonic(IV)acid H2CO3 Dibasic Carbonate(IV)(CO32-) Hydrogen carbonate(IV) (HCO3-) Phosphoric(V) acid H3PO4 Tribasic Phosphate(V)(PO43-) Dihydrogen phosphate(V) (H2PO42-) Hydrogen diphosphate(V) (HP2O42-) The table below show shows some examples of salts. Base/alkali Cation Acid Anion Salt Chemical name of salts NaOH Na+ HCl Cl NaCl Sodium(I)chloride Mg(OH)2 Mg2+ H2SO4 SO42 MgSO4 Mg(HSO4)2 Magnesium sulphate(VI) Magnesium hydrogen sulphate(VI) Pb(OH)2 Pb2+ HNO3 NO3 Pb(NO3)2 Lead(II)nitrate(V) Ba(OH)2 Ba2+ HNO3 NO3 Ba(NO3)2 Barium(II)nitrate(V) Ca(OH)2 Ba2+ H2SO4 SO42 MgSO4 Calcium sulphate(VI) NH4OH NH4+ H3PO4 PO43 (NH4 )3PO4 (NH4 )2HPO4 NH4 H2PO4 Ammonium phosphate(V) Diammonium phosphate(V) Ammonium diphosphate(V) KOH K+ H3PO4 PO43 K3PO4 Potassium phosphate(V) Al(OH)3 Al3+ H2SO4 SO42 Al2(SO4)2 Aluminium(III)sulphate(VI) Fe(OH)2 Fe2+ H2SO4 SO42 FeSO4 Iron(II)sulphate(VI)
Fe(OH)3 Fe3+ H2SO4 SO42 Fe2(SO4)2 Iron(III)sulphate(VI) (d) Some salts undergo hygroscopy, deliquescence and efflorescence. (i) Hygroscopic salts /compounds are those that absorb water from the atmosphere but do not form a solution.
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Table showing normal and acid salts derived from common acids Acid name Chemical formula Basicity Normal salt Acid salt Hydrochloric acid HCl Monobasic Chloride(Cl-) None Nitric(V)acid HNO3 Monobasic Nitrate(V)(NO3-) None Nitric(III)acid HNO2 Monobasic Nitrate(III)(NO2-) None Sulphuric(VI)acid H2SO4 Dibasic Sulphate(VI) (SO42-) Hydrogen sulphate(VI) (HSO4-) Sulphuric(IV)acid H2SO3 Dibasic Sulphate(IV) (SO32-) Hydrogen sulphate(IV) (HSO3-)
Carbonic(IV)acid H2CO3 Dibasic Carbonate(IV)(CO32-) Hydrogen carbonate(IV) (HCO3-) Phosphoric(V) acid H3PO4 Tribasic Phosphate(V)(PO43-) Dihydrogen phosphate(V) (H2PO42-) Hydrogen diphosphate(V) (HP2O42-) The table below show shows some examples of salts. Base/alkali Cation Acid Anion Salt Chemical name of salts NaOH Na+ HCl Cl NaCl Sodium(I)chloride Mg(OH)2 Mg2+ H2SO4 SO42 MgSO4 Mg(HSO4)2 Magnesium sulphate(VI) Magnesium hydrogen sulphate(VI) Pb(OH)2 Pb2+ HNO3 NO3 Pb(NO3)2 Lead(II)nitrate(V) Ba(OH)2 Ba2+ HNO3 NO3 Ba(NO3)2 Barium(II)nitrate(V) Ca(OH)2 Ba2+ H2SO4 SO42 MgSO4 Calcium sulphate(VI) NH4OH NH4+ H3PO4 PO43 (NH4 )3PO4 (NH4 )2HPO4 NH4 H2PO4 Ammonium phosphate(V) Diammonium phosphate(V) Ammonium diphosphate(V) KOH K+ H3PO4 PO43 K3PO4 Potassium phosphate(V) Al(OH)3 Al3+ H2SO4 SO42 Al2(SO4)2 Aluminium(III)sulphate(VI) Fe(OH)2 Fe2+ H2SO4 SO42 FeSO4 Iron(II)sulphate(VI)
Fe(OH)3 Fe3+ H2SO4 SO42 Fe2(SO4)2 Iron(III)sulphate(VI) (d) Some salts undergo hygroscopy, deliquescence and efflorescence. (i) Hygroscopic salts /compounds are those that absorb water from the atmosphere but do not form a solution. Some salts which are hygroscopic include anhydrous copper(II)sulphate(VI), anhydrous cobalt(II)chloride, potassium nitrate(V) common table salt.
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Base/alkali Cation Acid Anion Salt Chemical name of salts NaOH Na+ HCl Cl NaCl Sodium(I)chloride Mg(OH)2 Mg2+ H2SO4 SO42 MgSO4 Mg(HSO4)2 Magnesium sulphate(VI) Magnesium hydrogen sulphate(VI) Pb(OH)2 Pb2+ HNO3 NO3 Pb(NO3)2 Lead(II)nitrate(V) Ba(OH)2 Ba2+ HNO3 NO3 Ba(NO3)2 Barium(II)nitrate(V) Ca(OH)2 Ba2+ H2SO4 SO42 MgSO4 Calcium sulphate(VI) NH4OH NH4+ H3PO4 PO43 (NH4 )3PO4 (NH4 )2HPO4 NH4 H2PO4 Ammonium phosphate(V) Diammonium phosphate(V) Ammonium diphosphate(V) KOH K+ H3PO4 PO43 K3PO4 Potassium phosphate(V) Al(OH)3 Al3+ H2SO4 SO42 Al2(SO4)2 Aluminium(III)sulphate(VI) Fe(OH)2 Fe2+ H2SO4 SO42 FeSO4 Iron(II)sulphate(VI)
Fe(OH)3 Fe3+ H2SO4 SO42 Fe2(SO4)2 Iron(III)sulphate(VI) (d) Some salts undergo hygroscopy, deliquescence and efflorescence. (i) Hygroscopic salts /compounds are those that absorb water from the atmosphere but do not form a solution. Some salts which are hygroscopic include anhydrous copper(II)sulphate(VI), anhydrous cobalt(II)chloride, potassium nitrate(V) common table salt. (ii)Deliquescent salts /compounds are those that absorb water from the atmosphere and form a solution. Some salts which are deliquescent include: Sodium nitrate(V),Calcium chloride, Sodium hydroxide, Iron(II)chloride, Magnesium chloride. (iii)Efflorescent salts/compounds are those that lose their water of crystallization to the atmosphere.
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(ii)Deliquescent salts /compounds are those that absorb water from the atmosphere and form a solution. Some salts which are deliquescent include: Sodium nitrate(V),Calcium chloride, Sodium hydroxide, Iron(II)chloride, Magnesium chloride. (iii)Efflorescent salts/compounds are those that lose their water of crystallization to the atmosphere. Some salts which effloresces include: sodium carbonate decahydrate, Iron(II)sulphate(VI)heptahydrate, sodium sulphate (VI)decahydrate. (e)Some salts contain water of crystallization.They are hydrated.Others do not contain water of crystallization. They are anhydrous. Table showing some hydrated salts. Name of hydrated salt Chemical formula Copper(II)sulphate(VI)pentahydrate CuSO4.5H2O Aluminium(III)sulphate(VI)hexahydrate Al2 (SO4) 3.6H2O Zinc(II)sulphate(VI)heptahydrate ZnSO4.7H2O Iron(II)sulphate(VI)heptahydrate FeSO4.7H2O Calcium(II)sulphate(VI)heptahydrate CaSO4.7H2O Magnesium(II)sulphate(VI)heptahydrate MgSO4.7H2O Sodium sulphate(VI)decahydrate Na2SO4.10H2O Sodium carbonate(IV)decahydrate Na2CO3.10H2O Potassium carbonate(IV)decahydrate K2CO3.10H2O Potassium sulphate(VI)decahydrate K2SO4.10H2O (f)Some salts exist as a simple salt while some as complex salts. Below are some complex salts.
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Table showing some hydrated salts. Name of hydrated salt Chemical formula Copper(II)sulphate(VI)pentahydrate CuSO4.5H2O Aluminium(III)sulphate(VI)hexahydrate Al2 (SO4) 3.6H2O Zinc(II)sulphate(VI)heptahydrate ZnSO4.7H2O Iron(II)sulphate(VI)heptahydrate FeSO4.7H2O Calcium(II)sulphate(VI)heptahydrate CaSO4.7H2O Magnesium(II)sulphate(VI)heptahydrate MgSO4.7H2O Sodium sulphate(VI)decahydrate Na2SO4.10H2O Sodium carbonate(IV)decahydrate Na2CO3.10H2O Potassium carbonate(IV)decahydrate K2CO3.10H2O Potassium sulphate(VI)decahydrate K2SO4.10H2O (f)Some salts exist as a simple salt while some as complex salts. Below are some complex salts. Table of some complex salts Name of complex salt Chemical formula Colour of the complex salt Tetraamminecopper(II)sulphate(VI) Cu(NH3) 4 SO4 H2O Royal/deep blue solution
Tetraamminezinc(II)nitrate(V) Zn(NH3) 4 (NO3 )2 Colourless solution Tetraamminecopper(II) nitrate(V) Cu(NH3) 4 (NO3 )2 Royal/deep blue solution Tetraamminezinc(II)sulphate(VI) Zn(NH3) 4 SO4 Colourless solution (g)Some salts exist as two salts in one. They are called double salts.
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Below are some complex salts. Table of some complex salts Name of complex salt Chemical formula Colour of the complex salt Tetraamminecopper(II)sulphate(VI) Cu(NH3) 4 SO4 H2O Royal/deep blue solution
Tetraamminezinc(II)nitrate(V) Zn(NH3) 4 (NO3 )2 Colourless solution Tetraamminecopper(II) nitrate(V) Cu(NH3) 4 (NO3 )2 Royal/deep blue solution Tetraamminezinc(II)sulphate(VI) Zn(NH3) 4 SO4 Colourless solution (g)Some salts exist as two salts in one. They are called double salts. Table of some double salts Name of double salts Chemical formula Trona(sodium sesquicarbonate) Na2CO3 NaHCO3.2H2O Ammonium iron(II)sulphate(VI) FeSO4(NH4) 2SO4.2H2O Ammonium aluminium(III)sulphate(VI) Al2(SO4) 3(NH4) 2SO4.H2O (h)Some salts dissolve in water to form a solution. They are said to be soluble. Others do not dissolve in water. They form a suspension/precipitate in water.
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They are said to be soluble. Others do not dissolve in water. They form a suspension/precipitate in water. Table of solubility of salts Soluble salts Insoluble salts All nitrate(V)salts All sulphate(VI)/SO42- salts except Barium(II) sulphate(VI)/BaSO4 Calcium(II) sulphate(VI)/CaSO4 Lead(II) sulphate(VI)/PbSO4 All sulphate(IV)/SO32- salts except Barium(II) sulphate(IV)/BaSO3 Calcium(II) sulphate(IV)/CaSO3 Lead(II) sulphate(IV)/PbSO3 All chlorides/Cl- except Silver chloride/AgCl Lead(II)chloride/PbCl2(dissolves in hot water) All phosphate(V)/PO43- All sodium,potassium and ammonium salts All hydrogen carbonates/HCO3- All hydrogen sulphate(VI)/ HSO4- Sodium carbonate/Na2CO3, potassium carbonate/ K2CO3, ammonium carbonate (NH4) 2CO3 except All carbonates All alkalis(KOH,NaOH, NH4OH) except All bases 13 Salts can be prepared in a school laboratory by a method that uses its solubility in water. (a) Soluble salts may be prepared by using any of the following methods: (i)Direct displacement/reaction of a metal with an acid. By reacting a metal higher in the reactivity series than hydrogen with a dilute acid,a salt is formed and hydrogen gas is evolved. Excess of the metal must be used to ensure all the acid has reacted. When effervescence/bubbling /fizzing has stopped ,excess metal is filtered. The filtrate is heated to concentrate then allowed to crystallize. Washing with distilled water then drying between filter papers produces a sample crystal of the salt. i.e.
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The filtrate is heated to concentrate then allowed to crystallize. Washing with distilled water then drying between filter papers produces a sample crystal of the salt. i.e. M(s) + H2X -> MX(aq) + H2(g) Examples Mg(s) + H2SO4(aq) -> MgSO4 (aq) + H2(g) Zn(s) + H2SO4(aq) -> ZnSO4 (aq) + H2(g) Pb(s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + H2(g) Ca(s) + 2HNO3(aq) -> Ca(NO3) 2(aq) + H2(g) Mg(s) + 2HNO3(aq) -> Mg(NO3) 2(aq) + H2(g) Mg(s) + 2HCl(aq) -> MgCl 2(aq) + H2(g) Zn(s) + 2HCl(aq) -> ZnCl 2(aq) + H2(g) (ii)Reaction of an insoluble base with an acid By adding an insoluble base (oxide/hydroxide )to a dilute acid until no more dissolves, in the acid,a salt and water are formed. Excess of the base is filtered off. The filtrate is heated to concentrate ,allowed to crystallize then washed with distilled water before drying between filter papers e.g.
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M(s) + H2X -> MX(aq) + H2(g) Examples Mg(s) + H2SO4(aq) -> MgSO4 (aq) + H2(g) Zn(s) + H2SO4(aq) -> ZnSO4 (aq) + H2(g) Pb(s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + H2(g) Ca(s) + 2HNO3(aq) -> Ca(NO3) 2(aq) + H2(g) Mg(s) + 2HNO3(aq) -> Mg(NO3) 2(aq) + H2(g) Mg(s) + 2HCl(aq) -> MgCl 2(aq) + H2(g) Zn(s) + 2HCl(aq) -> ZnCl 2(aq) + H2(g) (ii)Reaction of an insoluble base with an acid By adding an insoluble base (oxide/hydroxide )to a dilute acid until no more dissolves, in the acid,a salt and water are formed. Excess of the base is filtered off. The filtrate is heated to concentrate ,allowed to crystallize then washed with distilled water before drying between filter papers e.g. PbO(s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + H2O (l) Pb(OH)2(s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + 2H2O (l) CaO (s) + 2HNO3(aq) -> Ca(NO3) 2(aq) + H2O (l) MgO (s) + 2HNO3(aq) -> Mg(NO3) 2(aq) + H2O (l) MgO (s) + 2HCl(aq) -> MgCl 2(aq) + H2O (l) ZnO (s) + 2HCl(aq) -> ZnCl 2(aq) + H2O (l) Zn(OH)2(s) + 2HNO3(aq) -> Zn(NO3) 2(aq) + 2H2O (l) CuO (s) + 2HCl(aq) -> CuCl 2(aq) + H2O (l) CuO (s) + H2SO4(aq) -> CuSO4(aq) + H2O (l) Ag2O(s) + 2HNO3(aq) -> 2AgNO3(aq) + H2O (l) Na2O(s) + 2HNO3(aq) -> 2NaNO3(aq) + H2O (l) (iii)reaction of insoluble /soluble carbonate /hydrogen carbonate with an acid.
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Excess of the base is filtered off. The filtrate is heated to concentrate ,allowed to crystallize then washed with distilled water before drying between filter papers e.g. PbO(s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + H2O (l) Pb(OH)2(s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + 2H2O (l) CaO (s) + 2HNO3(aq) -> Ca(NO3) 2(aq) + H2O (l) MgO (s) + 2HNO3(aq) -> Mg(NO3) 2(aq) + H2O (l) MgO (s) + 2HCl(aq) -> MgCl 2(aq) + H2O (l) ZnO (s) + 2HCl(aq) -> ZnCl 2(aq) + H2O (l) Zn(OH)2(s) + 2HNO3(aq) -> Zn(NO3) 2(aq) + 2H2O (l) CuO (s) + 2HCl(aq) -> CuCl 2(aq) + H2O (l) CuO (s) + H2SO4(aq) -> CuSO4(aq) + H2O (l) Ag2O(s) + 2HNO3(aq) -> 2AgNO3(aq) + H2O (l) Na2O(s) + 2HNO3(aq) -> 2NaNO3(aq) + H2O (l) (iii)reaction of insoluble /soluble carbonate /hydrogen carbonate with an acid. By adding an excess of a soluble /insoluble carbonate or hydrogen carbonate to adilute acid, effervescence /fizzing/bubbling out of carbon(IV)oxide gas shows the
reaction is taking place. When effervescence /fizzing/bubbling out of the gas is over, excess of the insoluble carbonate is filtered off. The filtrate is heated to concentrate ,allowed to crystallize then washed with distilled water before drying between filter paper papers e.g.
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By adding an excess of a soluble /insoluble carbonate or hydrogen carbonate to adilute acid, effervescence /fizzing/bubbling out of carbon(IV)oxide gas shows the
reaction is taking place. When effervescence /fizzing/bubbling out of the gas is over, excess of the insoluble carbonate is filtered off. The filtrate is heated to concentrate ,allowed to crystallize then washed with distilled water before drying between filter paper papers e.g. PbCO3 (s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + H2O (l)+ CO2(g) ZnCO3 (s) + 2HNO3(aq) -> Zn(NO3) 2(aq) + H2O (l)+ CO2(g) CaCO3 (s) + 2HNO3(aq) -> Ca(NO3) 2(aq) + H2O (l)+ CO2(g) MgCO3 (s) + H2SO4(aq) -> MgSO4(aq) + H2O (l)+ CO2(g) Cu CO3 (s) + H2SO4(aq) -> CuSO4(aq) + H2O (l) + CO2(g) Ag2CO3 (s) + 2HNO3(aq) -> 2AgNO3(aq) + H2O (l) + CO2(g) Na2CO3 (s) + 2HNO3(aq) -> 2NaNO3(aq) + H2O (l) + CO2(g) K2CO3 (s) + 2HCl(aq) -> 2KCl(aq) + H2O (l) + CO2(g) NaHCO3 (s) + HNO3(aq) -> NaNO3(aq) + H2O (l) + CO2(g) KHCO3 (s) + HCl(aq) -> KCl(aq) + H2O (l) + CO2(g) (iv)neutralization/reaction of soluble base/alkali with dilute acid By adding an acid to a burette into a known volume of an alkali with 2-3 drops of an indicator, the colour of the indicator changes when the acid has completely reacted with an alkali at the end point.
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When effervescence /fizzing/bubbling out of the gas is over, excess of the insoluble carbonate is filtered off. The filtrate is heated to concentrate ,allowed to crystallize then washed with distilled water before drying between filter paper papers e.g. PbCO3 (s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + H2O (l)+ CO2(g) ZnCO3 (s) + 2HNO3(aq) -> Zn(NO3) 2(aq) + H2O (l)+ CO2(g) CaCO3 (s) + 2HNO3(aq) -> Ca(NO3) 2(aq) + H2O (l)+ CO2(g) MgCO3 (s) + H2SO4(aq) -> MgSO4(aq) + H2O (l)+ CO2(g) Cu CO3 (s) + H2SO4(aq) -> CuSO4(aq) + H2O (l) + CO2(g) Ag2CO3 (s) + 2HNO3(aq) -> 2AgNO3(aq) + H2O (l) + CO2(g) Na2CO3 (s) + 2HNO3(aq) -> 2NaNO3(aq) + H2O (l) + CO2(g) K2CO3 (s) + 2HCl(aq) -> 2KCl(aq) + H2O (l) + CO2(g) NaHCO3 (s) + HNO3(aq) -> NaNO3(aq) + H2O (l) + CO2(g) KHCO3 (s) + HCl(aq) -> KCl(aq) + H2O (l) + CO2(g) (iv)neutralization/reaction of soluble base/alkali with dilute acid By adding an acid to a burette into a known volume of an alkali with 2-3 drops of an indicator, the colour of the indicator changes when the acid has completely reacted with an alkali at the end point. The procedure is then repeated without the indicator .The solution mixture is then heated to concentrate , allowed to crystallize ,washed with distilled water before drying with filter papers. e.g.
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PbCO3 (s) + 2HNO3(aq) -> Pb(NO3) 2(aq) + H2O (l)+ CO2(g) ZnCO3 (s) + 2HNO3(aq) -> Zn(NO3) 2(aq) + H2O (l)+ CO2(g) CaCO3 (s) + 2HNO3(aq) -> Ca(NO3) 2(aq) + H2O (l)+ CO2(g) MgCO3 (s) + H2SO4(aq) -> MgSO4(aq) + H2O (l)+ CO2(g) Cu CO3 (s) + H2SO4(aq) -> CuSO4(aq) + H2O (l) + CO2(g) Ag2CO3 (s) + 2HNO3(aq) -> 2AgNO3(aq) + H2O (l) + CO2(g) Na2CO3 (s) + 2HNO3(aq) -> 2NaNO3(aq) + H2O (l) + CO2(g) K2CO3 (s) + 2HCl(aq) -> 2KCl(aq) + H2O (l) + CO2(g) NaHCO3 (s) + HNO3(aq) -> NaNO3(aq) + H2O (l) + CO2(g) KHCO3 (s) + HCl(aq) -> KCl(aq) + H2O (l) + CO2(g) (iv)neutralization/reaction of soluble base/alkali with dilute acid By adding an acid to a burette into a known volume of an alkali with 2-3 drops of an indicator, the colour of the indicator changes when the acid has completely reacted with an alkali at the end point. The procedure is then repeated without the indicator .The solution mixture is then heated to concentrate , allowed to crystallize ,washed with distilled water before drying with filter papers. e.g. NaOH (aq) + HNO3(aq) -> NaNO3(aq) + H2O (l) KOH (aq) + HNO3(aq) -> KNO3(aq) + H2O (l) KOH (aq) + HCl(aq) -> KCl(aq) + H2O (l) 2KOH (aq) + H2SO4(aq) -> K2SO4(aq) + 2H2O (l) 2 NH4OH (aq) + H2SO4(aq) -> (NH4)2SO4(aq) + 2H2O (l) NH4OH (aq) + HNO3(aq) -> NH4NO3(aq) + H2O (l) (iv)direct synthesis/combination.
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The procedure is then repeated without the indicator .The solution mixture is then heated to concentrate , allowed to crystallize ,washed with distilled water before drying with filter papers. e.g. NaOH (aq) + HNO3(aq) -> NaNO3(aq) + H2O (l) KOH (aq) + HNO3(aq) -> KNO3(aq) + H2O (l) KOH (aq) + HCl(aq) -> KCl(aq) + H2O (l) 2KOH (aq) + H2SO4(aq) -> K2SO4(aq) + 2H2O (l) 2 NH4OH (aq) + H2SO4(aq) -> (NH4)2SO4(aq) + 2H2O (l) NH4OH (aq) + HNO3(aq) -> NH4NO3(aq) + H2O (l) (iv)direct synthesis/combination. When a metal burn in a gas jar containing a non metal , the two directly combine to form a salt. e.g. 2Na(s) + Cl2(g) -> 2NaCl(s) 2K(s) + Cl2(g) -> 2KCl(s) Mg(s) + Cl2(g) -> Mg Cl2 (s) Ca(s) + Cl2(g) -> Ca Cl2 (s) Some salts once formed undergo sublimation and hydrolysis. Care should be taken to avoid water/moisture into the reaction flask during their preparation.Such salts include aluminium(III)chloride(AlCl3) and iron (III)chloride(FeCl3)
1. Heated aluminium foil reacts with chlorine to form aluminium(III)chloride that sublimes away from the source of heating then deposited as solid again 2Al(s) + 3Cl2(g) -> 2AlCl3 (s/g) Once formed aluminium(III)chloride hydrolyses/reacts with water vapour / moisture present to form aluminium hydroxide solution and highly acidic fumes of hydrogen chloride gas. AlCl3(s)+ 3H2 O(g) -> Al(OH)3 (aq) + 3HCl(g) 2.
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Care should be taken to avoid water/moisture into the reaction flask during their preparation.Such salts include aluminium(III)chloride(AlCl3) and iron (III)chloride(FeCl3)
1. Heated aluminium foil reacts with chlorine to form aluminium(III)chloride that sublimes away from the source of heating then deposited as solid again 2Al(s) + 3Cl2(g) -> 2AlCl3 (s/g) Once formed aluminium(III)chloride hydrolyses/reacts with water vapour / moisture present to form aluminium hydroxide solution and highly acidic fumes of hydrogen chloride gas. AlCl3(s)+ 3H2 O(g) -> Al(OH)3 (aq) + 3HCl(g) 2. Heated iron filings reacts with chlorine to form iron(III)chloride that sublimes away from the source of heating then deposited as solid again 2Fe(s) + 3Cl2(g) -> 2FeCl3 (s/g) Once formed , aluminium(III)chloride hydrolyses/reacts with water vapour / moisture present to form aluminium hydroxide solution and highly acidic fumes of hydrogen chloride gas. FeCl3(s)+ 3H2 O(g) -> Fe(OH)3 (aq) + 3HCl(g) (b)Insoluble salts can be prepared by reacting two suitable soluble salts to form one soluble and one insoluble. This is called double decomposition or precipitation. The mixture is filtered and the residue is washed with distilled water then dried.
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FeCl3(s)+ 3H2 O(g) -> Fe(OH)3 (aq) + 3HCl(g) (b)Insoluble salts can be prepared by reacting two suitable soluble salts to form one soluble and one insoluble. This is called double decomposition or precipitation. The mixture is filtered and the residue is washed with distilled water then dried. CuSO4(aq) + Na2CO3 (aq) -> CuCO3 (s) + Na2 SO4(aq) BaCl2(aq) + K2SO4 (aq) -> BaSO4 (s) + 2KCl (aq) Pb(NO3)2(aq) + K2SO4 (aq) -> PbSO4 (s) + 2KNO3 (aq) 2AgNO3(aq) + MgCl2 (aq) -> 2AgCl(s) + Mg(NO3)2 (aq) Pb(NO3)2(aq) + (NH4) 2SO4 (aq) -> PbSO4 (s) + 2NH4NO 3(aq) BaCl2(aq) + K2SO3 (aq) -> BaSO3 (s) + 2KCl (aq) 14. Salts may lose their water of crystallization , decompose ,melt or sublime on heating on a Bunsen burner flame. The following shows the behavior of some salts on heating gently /or strongly in a laboratory school burner: (a)effect of heat on chlorides All chlorides have very high melting and boiling points and therefore are not affected by laboratory heating except ammonium chloride. Ammonium chloride sublimes on gentle heating. It dissociate into the constituent ammonia and hydrogen chloride gases on strong heating. NH4Cl(s) NH4Cl(g) NH3(g) + HCl(g) (sublimation) (dissociation) (b)effect of heat on nitrate(V) (i) Potassium nitrate(V)/KNO3 and sodium nitrate(V)/NaNO3 decompose on heating to form Potassium nitrate(III)/KNO2 and sodium nitrate(III)/NaNO2 and producing Oxygen gas in each case.
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Ammonium chloride sublimes on gentle heating. It dissociate into the constituent ammonia and hydrogen chloride gases on strong heating. NH4Cl(s) NH4Cl(g) NH3(g) + HCl(g) (sublimation) (dissociation) (b)effect of heat on nitrate(V) (i) Potassium nitrate(V)/KNO3 and sodium nitrate(V)/NaNO3 decompose on heating to form Potassium nitrate(III)/KNO2 and sodium nitrate(III)/NaNO2 and producing Oxygen gas in each case. 2KNO3 (s) -> 2KNO2(s) + O2(g) 2NaNO3 (s) -> 2NaNO2(s) + O2(g) (ii)Heavy metal nitrates(V) salts decompose on heating to form the oxide and a mixture of brown acidic nitrogen(IV)oxide and oxygen gases. e.g. 2Ca(NO3)2 (s) -> 2CaO(s) + 4NO2(g) + O2(g) 2Mg(NO3)2(s) -> 2MgO(s) + 4NO2(g) + O2(g) 2Zn(NO3)2(s) -> 2ZnO(s) + 4NO2(g) + O2(g) 2Pb(NO3)2(s) -> 2PbO(s) + 4NO2(g) + O2(g) 2Cu(NO3)2(s) -> 2CuO(s) + 4NO2(g) + O2(g) 2Fe(NO3)2(s) -> 2FeO(s) + 4NO2(g) + O2(g) (iii)Silver(I)nitrate(V) and mercury(II) nitrate(V) are lowest in the reactivity series. They decompose on heating to form the metal(silver and mercury)and the Nitrogen(IV)oxide and oxygen gas. i.e.
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2Ca(NO3)2 (s) -> 2CaO(s) + 4NO2(g) + O2(g) 2Mg(NO3)2(s) -> 2MgO(s) + 4NO2(g) + O2(g) 2Zn(NO3)2(s) -> 2ZnO(s) + 4NO2(g) + O2(g) 2Pb(NO3)2(s) -> 2PbO(s) + 4NO2(g) + O2(g) 2Cu(NO3)2(s) -> 2CuO(s) + 4NO2(g) + O2(g) 2Fe(NO3)2(s) -> 2FeO(s) + 4NO2(g) + O2(g) (iii)Silver(I)nitrate(V) and mercury(II) nitrate(V) are lowest in the reactivity series. They decompose on heating to form the metal(silver and mercury)and the Nitrogen(IV)oxide and oxygen gas. i.e. 2AgNO3(s) -> 2Ag (s) + 2NO2(g) + O2(g) 2Hg(NO3)2 (s) -> 2Hg (s) + 4NO2(g) + O2(g) (iv)Ammonium nitrate(V) and Ammonium nitrate(III) decompose on heating to Nitrogen(I)oxide(relights/rekindles glowing splint) and nitrogen gas respectively.Water is also formed.i.e. NH4NO3(s) -> N2O (g) + H2O(l) NH4NO2(s) -> N2 (g) + H2O(l) (c) effect of heat on nitrate(V) Only Iron(II)sulphate(VI), Iron(III)sulphate(VI) and copper(II)sulphate(VI) decompose on heating. They form the oxide, and produce highly acidic fumes of acidic sulphur(IV)oxide gas.
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2AgNO3(s) -> 2Ag (s) + 2NO2(g) + O2(g) 2Hg(NO3)2 (s) -> 2Hg (s) + 4NO2(g) + O2(g) (iv)Ammonium nitrate(V) and Ammonium nitrate(III) decompose on heating to Nitrogen(I)oxide(relights/rekindles glowing splint) and nitrogen gas respectively.Water is also formed.i.e. NH4NO3(s) -> N2O (g) + H2O(l) NH4NO2(s) -> N2 (g) + H2O(l) (c) effect of heat on nitrate(V) Only Iron(II)sulphate(VI), Iron(III)sulphate(VI) and copper(II)sulphate(VI) decompose on heating. They form the oxide, and produce highly acidic fumes of acidic sulphur(IV)oxide gas. 2FeSO4 (s) -> Fe2O3(s) + SO3(g) + SO2(g) Fe2(SO4) 3(s) -> Fe2O3(s) + SO3(g) CuSO4 (s) -> CuO(s) + SO3(g)
(d) effect of heat on carbonates(IV) and hydrogen carbonate(IV). (i)Sodium carbonate(IV)and potassium carbonate(IV)do not decompose on heating. (ii)Heavy metal nitrate(IV)salts decompose on heating to form the oxide and produce carbon(IV)oxide gas. Carbon (IV)oxide gas forms a white precipitate when bubbled in lime water. The white precipitate dissolves if the gas is in excess. e.g. CuCO3 (s) -> CuO(s) + CO2(g) CaCO3 (s) -> CaO(s) + CO2(g) PbCO3 (s) -> PbO(s) + CO2(g) FeCO3 (s) -> FeO(s) + CO2(g) ZnCO3 (s) -> ZnO(s) + CO2(g) (iii)Sodium hydrogen carbonate(IV) and Potassium hydrogen carbonate(IV)decompose on heating to give the corresponding carbonate (IV) and form water and carbon(IV)oxide gas. i.e.
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e.g. CuCO3 (s) -> CuO(s) + CO2(g) CaCO3 (s) -> CaO(s) + CO2(g) PbCO3 (s) -> PbO(s) + CO2(g) FeCO3 (s) -> FeO(s) + CO2(g) ZnCO3 (s) -> ZnO(s) + CO2(g) (iii)Sodium hydrogen carbonate(IV) and Potassium hydrogen carbonate(IV)decompose on heating to give the corresponding carbonate (IV) and form water and carbon(IV)oxide gas. i.e. 2NaHCO 3(s) -> Na2CO3(s) + CO2(g) + H2O(l) 2KHCO 3(s) -> K2CO3(s) + CO2(g) + H2O(l) (iii) Calcium hydrogen carbonate (IV) and Magnesium hydrogen carbonate(IV) decompose on heating to give the corresponding carbonate (IV) and form water and carbon(IV)oxide gas. i. e. Ca(HCO3) 2(aq) -> CaCO3(s) + CO2(g) + H2O(l) Mg(HCO3) 2(aq) -> MgCO3(s) + CO2(g) + H2O(l)
1 INTRODUCTION TO ELECTROLYSIS (ELECTROLYTIC CELL) 1.Electrolysis is defined simply as the decomposition of a compound by an electric current/electricity. A compound that is decomposed by an electric current is called an electrolyte. Some electrolytes are weak while others are strong. 2.Strong electrolytes are those that are fully ionized/dissociated into (many) ions. Common strong electrolytes include: (i)all mineral acids (ii)all strong alkalis/sodium hydroxide/potassium hydroxide. (iii)all soluble salts 3.Weak electrolytes are those that are partially/partly ionized/dissociated into (few) ions. Common weak electrolytes include: (i)all organic acids (ii)all bases except sodium hydroxide/potassium hydroxide. (iii)Water 4. A compound that is not decomposed by an electric current is called nonelectrolyte.
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Common weak electrolytes include: (i)all organic acids (ii)all bases except sodium hydroxide/potassium hydroxide. (iii)Water 4. A compound that is not decomposed by an electric current is called nonelectrolyte. Non-electrolytes are those compounds /substances that exist as molecules and thus cannot ionize/dissociate into(any) ions . Common non-electrolytes include: (i) most organic solvents (e.g. petrol/paraffin/benzene/methylbenzene/ethanol) (ii)all hydrocarbons(alkanes /alkenes/alkynes) (iii)Chemicals of life(e.g. proteins, carbohydrates, lipids, starch, sugar) 5. An electrolytes in solid state have fused /joined ions and therefore do not conduct electricity but the ions (cations and anions) are free and mobile in molten and aqueous (solution, dissolved in water) state. 6.During electrolysis, the free ions are attracted to the electrodes. An electrode is a rod through which current enter and leave the electrolyte during electrolysis. 2 An electrode that does not influence/alter the products of electrolysis is called an inert electrode. Common inert electrodes include: (i)Platinum (ii)Carbon graphite Platinum is not usually used in a school laboratory because it is very expensive. Carbon graphite is easily/readily and cheaply available (from used dry cells). 7.The positive electrode is called Anode.The anode is the electrode through which current enter the electrolyte/electrons leave the electrolyte 8.The negative electrode is called Cathode. The cathode is the electrode through which current leave the electrolyte / electrons enter the electrolyte 9. During the electrolysis, free anions are attracted to the anode where they lose /donate electrons to form neutral atoms/molecules. i.e. M(l) -> M+(l) + e (for cations from molten electrolytes) M(s) -> M+(aq) + e (for cations from electrolytes in aqueous state / solution / dissolved in water) The neutral atoms /molecules form the products of electrolysis at the anode. This is called discharge at anode 10. During electrolysis, free cations are attracted to the cathode where they gain /accept/acquire electrons to form neutral atoms/molecules.
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M(l) -> M+(l) + e (for cations from molten electrolytes) M(s) -> M+(aq) + e (for cations from electrolytes in aqueous state / solution / dissolved in water) The neutral atoms /molecules form the products of electrolysis at the anode. This is called discharge at anode 10. During electrolysis, free cations are attracted to the cathode where they gain /accept/acquire electrons to form neutral atoms/molecules. X+ (aq) + 2e -> X(s) (for cations from electrolytes in aqueous state / solution / dissolved in water) 2X+ (l) + 2e -> X (l) (for cations from molten electrolytes) The neutral atoms /molecules form the products of electrolysis at the cathode. This is called discharge at cathode. 11. The below set up shows an electrolytic cell. 3 BatteryAnode(+)Cathode(-)ElectrolyteSimple set up of electrolytic cellGaseous product at anodeGaseous product at cathode 12. For a compound /salt containing only two ion/binary salt the products of electrolysis in an electrolytic cell can be determined as in the below examples: a)To determine the products of electrolysis of molten Lead(II)chloride (i)Decomposition of electrolyte into free ions; PbCl2 (l) -> Pb 2+(l) + 2Cl-(l) (Compound decomposed into free cation and anion in liquid state) (ii)At the cathode/negative electrode(-); Pb 2+(l) + 2e -> Pb (l) (Cation / Pb 2+ gains / accepts / acquires electrons to form free atom) (iii)At the anode/positive electrode(+); 2Cl-(l) -> Cl2 (g) + 2e (Anion / Cl- donate/lose electrons to form free atom then a gas molecule) (iv)Products of electrolysis therefore are; I.At the cathode grey beads /solid lead metal. II.At the anode pale green chlorine gas.
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3 BatteryAnode(+)Cathode(-)ElectrolyteSimple set up of electrolytic cellGaseous product at anodeGaseous product at cathode 12. For a compound /salt containing only two ion/binary salt the products of electrolysis in an electrolytic cell can be determined as in the below examples: a)To determine the products of electrolysis of molten Lead(II)chloride (i)Decomposition of electrolyte into free ions; PbCl2 (l) -> Pb 2+(l) + 2Cl-(l) (Compound decomposed into free cation and anion in liquid state) (ii)At the cathode/negative electrode(-); Pb 2+(l) + 2e -> Pb (l) (Cation / Pb 2+ gains / accepts / acquires electrons to form free atom) (iii)At the anode/positive electrode(+); 2Cl-(l) -> Cl2 (g) + 2e (Anion / Cl- donate/lose electrons to form free atom then a gas molecule) (iv)Products of electrolysis therefore are; I.At the cathode grey beads /solid lead metal. II.At the anode pale green chlorine gas. b)To determine the products of electrolysis of molten Zinc bromide
4 (i)Decomposition of electrolyte into free ions; ZnBr2 (l) -> Zn 2+(l) + 2Br-(l) (Compound decomposed into free cation and anion in liquid state) (ii)At the cathode/negative electrode(-); Zn 2+(l) + 2e -> Zn(l) (Cation / Zn2+ gains / accepts / acquires electrons to form free atom) (iii)At the anode/positive electrode(+); 2Br-(l) -> Br2 (g) + 2e (Anion / Br- donate/lose electrons to form free atom then a liquid molecule which change to gas on heating) (iv)Products of electrolysis therefore are; I.At the cathode grey beads /solid Zinc metal. II.At the anode red bromine liquid / red/brown bromine gas.
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II.At the anode pale green chlorine gas. b)To determine the products of electrolysis of molten Zinc bromide
4 (i)Decomposition of electrolyte into free ions; ZnBr2 (l) -> Zn 2+(l) + 2Br-(l) (Compound decomposed into free cation and anion in liquid state) (ii)At the cathode/negative electrode(-); Zn 2+(l) + 2e -> Zn(l) (Cation / Zn2+ gains / accepts / acquires electrons to form free atom) (iii)At the anode/positive electrode(+); 2Br-(l) -> Br2 (g) + 2e (Anion / Br- donate/lose electrons to form free atom then a liquid molecule which change to gas on heating) (iv)Products of electrolysis therefore are; I.At the cathode grey beads /solid Zinc metal. II.At the anode red bromine liquid / red/brown bromine gas. c)To determine the products of electrolysis of molten sodium chloride (i)Decomposition of electrolyte into free ions; NaCl (l) -> Na +(l) + Cl-(l) (Compound decomposed into free cation and anion in liquid state) (ii)At the cathode/negative electrode(-); 2Na+(l) + 2e -> Na (l) (Cation / Na+ gains / accepts / acquires electrons to form free atom) (iii)At the anode/positive electrode(+); 2Cl-(l) -> Cl2 (g) + 2e (Anion / Cl- donate/lose electrons to form free atom then a gas molecule) (iv)Products of electrolysis therefore are; I.At the cathode grey beads /solid sodium metal. II.At the anode pale green chlorine gas.
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II.At the anode red bromine liquid / red/brown bromine gas. c)To determine the products of electrolysis of molten sodium chloride (i)Decomposition of electrolyte into free ions; NaCl (l) -> Na +(l) + Cl-(l) (Compound decomposed into free cation and anion in liquid state) (ii)At the cathode/negative electrode(-); 2Na+(l) + 2e -> Na (l) (Cation / Na+ gains / accepts / acquires electrons to form free atom) (iii)At the anode/positive electrode(+); 2Cl-(l) -> Cl2 (g) + 2e (Anion / Cl- donate/lose electrons to form free atom then a gas molecule) (iv)Products of electrolysis therefore are; I.At the cathode grey beads /solid sodium metal. II.At the anode pale green chlorine gas. d)To determine the products of electrolysis of molten Aluminium (III)oxide (i)Decomposition of electrolyte into free ions; Al2O3 (l) -> 2Al 3+(l) + 3O2-(l) (Compound decomposed into free cation and anion in liquid state)
5 (ii)At the cathode/negative electrode(-); 4Al 3+ (l) + 12e -> 4Al (l) (Cation / Al 3+ gains / accepts / acquires electrons to form free atom) (iii)At the anode/positive electrode(+); 6O2-(l) -> 3O2 (g) + 12e (Anion /6O2- donate/lose 12 electrons to form free atom then three gas molecule) (iv)Products of electrolysis therefore are; I.At the cathode grey beads /solid aluminium metal. II.At the anode colourless gas that relights/rekindles glowing splint. 13.In industries electrolysis has the following uses/applications: (a)Extraction of reactive metals from their ores. Potassium, sodium ,magnesium, and aluminium are extracted from their ores using electrolytic methods. (b)Purifying copper after exraction from copper pyrites ores. Copper obtained from copper pyrites ores is not pure.
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Potassium, sodium ,magnesium, and aluminium are extracted from their ores using electrolytic methods. (b)Purifying copper after exraction from copper pyrites ores. Copper obtained from copper pyrites ores is not pure. After extraction, the copper is refined by electrolysing copper(II)sulphate(VI) solution using the impure copper as anode and a thin strip of pure copper as cathode. Electrode ionization take place there: (i)At the cathode; Cu2+ (aq) + 2e -> Cu(s) (Pure copper deposits on the strip (ii)At the anode; Cu(s) ->Cu2+ (aq) + 2e (impure copper erodes/dissolves) (c)Electroplating The label EPNS(Electro Plated Nickel Silver) on some steel/metallic utensils mean they are plated/coated with silver and/or Nickel to improve their appearance(add their aesthetic value)and prevent/slow corrosion(rusting of iron). Electroplating is the process of coating a metal with another metal using an electric current. During electroplating, the cathode is made of the metal to be coated/impure. Example: During the electroplating of a spoon with silver (i)the spoon/impure is placed as the cathode(negative terminal of battery) (ii)the pure silver is placed as the anode(positive terminal of battery) (iii)the pure silver erodes/ionizes/dissociates to release electrons: Ag(s) ->Ag+ (aq) + e (impure silver erodes/dissolves)
6 (iv) silver (Ag+)ions from electrolyte gain electrons to form pure silver deposits / coat /cover the spoon/impure Ag+ (aq) + e ->Ag(s) (pure silver deposits /coat/cover on spoon)
CARBON AND ITS COMPOUNDS Carbon is an element in Group IV(Group 4)of the Periodic table .It has atomic number 6 and electronic configuration 2:4 and thus has four valence electrons (tetravalent).It does not easily ionize but forms strong covalent bonds with other elements including itself. (a)Occurrence Carbon mainly naturally occurs as: (i)allotropes of carbon i.e graphite, diamond and fullerenes. (ii)amorphous carbon in coal, peat ,charcoal and coke.
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Example: During the electroplating of a spoon with silver (i)the spoon/impure is placed as the cathode(negative terminal of battery) (ii)the pure silver is placed as the anode(positive terminal of battery) (iii)the pure silver erodes/ionizes/dissociates to release electrons: Ag(s) ->Ag+ (aq) + e (impure silver erodes/dissolves)
6 (iv) silver (Ag+)ions from electrolyte gain electrons to form pure silver deposits / coat /cover the spoon/impure Ag+ (aq) + e ->Ag(s) (pure silver deposits /coat/cover on spoon)
CARBON AND ITS COMPOUNDS Carbon is an element in Group IV(Group 4)of the Periodic table .It has atomic number 6 and electronic configuration 2:4 and thus has four valence electrons (tetravalent).It does not easily ionize but forms strong covalent bonds with other elements including itself. (a)Occurrence Carbon mainly naturally occurs as: (i)allotropes of carbon i.e graphite, diamond and fullerenes. (ii)amorphous carbon in coal, peat ,charcoal and coke. (iii)carbon(IV)oxide gas accounting 0.03% by volume of normal air in the atmosphere. (b)Allotropes of Carbon Carbon naturally occur in two main crystalline allotropic forms, carbon-graphite and carbon-diamond Carbon-diamond Carbon-graphite Shiny crystalline solid Black/dull crystalline solid Has a very high melting/boiling point because it has a very closely packed giant tetrahedral structure joined by strong covalent bonds Has a high melting/boiling point because it has a very closely packed giant hexagonal planar structure joined by strong covalent bonds Has very high density(Hardest known Soft
2 natural substance) Abrassive Slippery Poor electrical conductor because it has no free delocalized electrons Good electrical conductor because it has free 4th valency delocalized electrons Is used in making Jewels, drilling and cutting metals Used in making Lead-pencils,electrodes in batteries and as a lubricant Has giant tetrahedral structure Has giant hexagonal planar structure c)Properties of Carbon (i)Physical properties of carbon Carbon occur widely and naturally as a black solid It is insoluble in water but soluble in carbon disulphide and organic solvents.
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(ii)amorphous carbon in coal, peat ,charcoal and coke. (iii)carbon(IV)oxide gas accounting 0.03% by volume of normal air in the atmosphere. (b)Allotropes of Carbon Carbon naturally occur in two main crystalline allotropic forms, carbon-graphite and carbon-diamond Carbon-diamond Carbon-graphite Shiny crystalline solid Black/dull crystalline solid Has a very high melting/boiling point because it has a very closely packed giant tetrahedral structure joined by strong covalent bonds Has a high melting/boiling point because it has a very closely packed giant hexagonal planar structure joined by strong covalent bonds Has very high density(Hardest known Soft
2 natural substance) Abrassive Slippery Poor electrical conductor because it has no free delocalized electrons Good electrical conductor because it has free 4th valency delocalized electrons Is used in making Jewels, drilling and cutting metals Used in making Lead-pencils,electrodes in batteries and as a lubricant Has giant tetrahedral structure Has giant hexagonal planar structure c)Properties of Carbon (i)Physical properties of carbon Carbon occur widely and naturally as a black solid It is insoluble in water but soluble in carbon disulphide and organic solvents. It is a poor electrical and thermal conductor. (ii)Chemical properties of carbon I. Burning Experiment Introduce a small piece of charcoal on a Bunsen flame then lower it into a gas jar containing Oxygen gas. Put three drops of water. Swirl. Test the solution with blue and red litmus papers. Observation -Carbon chars then burns with a blue flame -Colourless and odourless gas produced -Solution formed turn blue litmus paper faint red. Red litmus paper remains red. Explanation Carbon burns in air and faster in Oxygen with a blue non-sooty/non-smoky flame forming Carbon (IV) oxide gas. Carbon burns in limited supply of air with a blue non-sooty/non-smoky flame forming Carbon (IV) oxide gas. Carbon (IV) oxide gas dissolve in water to form weak acidic solution of Carbonic (IV)acid.
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Explanation Carbon burns in air and faster in Oxygen with a blue non-sooty/non-smoky flame forming Carbon (IV) oxide gas. Carbon burns in limited supply of air with a blue non-sooty/non-smoky flame forming Carbon (IV) oxide gas. Carbon (IV) oxide gas dissolve in water to form weak acidic solution of Carbonic (IV)acid. Chemical Equation C(s) + O2(g) -> CO2(g) (in excess air) 2C(s) + O2(g) -> 2CO(g) (in limited air) CO2(g) + H2O (l) -> H2CO3 (aq) (very weak acid) II. Reducing agent Experiment Mix thoroughly equal amounts of powdered charcoal and copper (II)oxide into a crucible. Heat strongly. Observation Colour change from black to brown Explanation
3 Carbon is a reducing agent. For ages it has been used to reducing metal oxide ores to metal, itself oxidized to carbon(IV)oxide gas.
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Heat strongly. Observation Colour change from black to brown Explanation
3 Carbon is a reducing agent. For ages it has been used to reducing metal oxide ores to metal, itself oxidized to carbon(IV)oxide gas. Carbon reduces black copper(II)oxide to brown copper metal Chemical Equation 2CuO(s) + C(s) -> 2Cu(s) + CO2(g) (black) (brown) 2PbO(s) + C(s) -> 2Pb(s) + CO2(g) (brown when hot/ (grey) yellow when cool) 2ZnO(s) + C(s) -> 2Zn(s) + CO2(g) (yellow when hot/ (grey) white when cool) Fe2O3(s) + 3C(s) -> 2Fe(s) + 3CO2(g) (brown when hot/cool (grey) Fe3O4 (s) + 4C(s) -> 3Fe(s) + 4CO2(g) (brown when hot/cool (grey)
4 B: COMPOUNDS OF CARBON The following are the main compounds of Carbon (i)Carbon(IV)Oxide(CO2) (ii)Carbon(II)Oxide(CO) (iii)Carbonate(IV) (CO32-)and hydrogen carbonate(IV(HCO3-) (iv)Sodium carbonate(Na2CO3) (i) Carbon(IV)Oxide (CO2) (a)Occurrence Carbon(IV)oxide is found: -in the air /atmosphere as 0.03% by volume. -a solid carbon(IV)oxide mineral in Esageri near Eldame Ravine and Kerita near Limuru in Kenya. (b)School Laboratory preparation In the school laboratory carbon(IV)oxide can be prepared in the school laboratory from the reaction of marble chips(CaCO3)or sodium hydrogen carbonate(NaHCO3) with dilute hydrochloric acid. 5 (c)Properties of carbon(IV)oxide gas(Questions) 1.Write the equation for the reaction for the school laboratory preparation of carbon (IV)oxide gas. Any carbonate reacted with dilute hydrochloric acid should be able to generate carbon (IV)oxide gas.
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(b)School Laboratory preparation In the school laboratory carbon(IV)oxide can be prepared in the school laboratory from the reaction of marble chips(CaCO3)or sodium hydrogen carbonate(NaHCO3) with dilute hydrochloric acid. 5 (c)Properties of carbon(IV)oxide gas(Questions) 1.Write the equation for the reaction for the school laboratory preparation of carbon (IV)oxide gas. Any carbonate reacted with dilute hydrochloric acid should be able to generate carbon (IV)oxide gas. Chemical equations CaCO3(s) + 2HCl(aq) -> CaCO3 (aq) + H2O(l) + CO2 (g) ZnCO3(s) + 2HCl(aq) -> ZnCO3 (aq) + H2O(l) + CO2 (g) MgCO3(s) + 2HCl(aq) -> MgCO3 (aq) + H2O(l) + CO2 (g) CuCO3(s) + 2HCl(aq) -> CuCO3 (aq) + H2O(l) + CO2 (g) NaHCO3(s) + HCl(aq) -> Na2CO3 (aq) + H2O(l) + CO2 (g) KHCO3(s) + HCl(aq) -> K2CO3 (aq) + H2O(l) + CO2 (g) 2.What method of gas collection is used in preparation of Carbon(IV)oxide gas. Explain. Downward delivery /upward displacement of air/over mercury Carbon(IV)oxide gas is about 1½ times denser than air. 3.What is the purpose of : (a)water? To absorb the more volatile hydrogen chloride fumes produced during the vigorous reaction. (b)sodium hydrogen carbonate? To absorb the more volatile hydrogen chloride fumes produced during the vigorous reaction and by reacting with the acid to produce more carbon (IV)oxide gas . 6 Chemical equation NaHCO3(s) + HCl(aq) -> Na2CO3 (aq) + H2O(l) + CO2 (g) (c)concentrated sulphuric(VI)acid? To dry the gas/as a drying agent 4.Describe the smell of carbon(IV)oxide gas Colourless and odourless 5. Effect on lime water.
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6 Chemical equation NaHCO3(s) + HCl(aq) -> Na2CO3 (aq) + H2O(l) + CO2 (g) (c)concentrated sulphuric(VI)acid? To dry the gas/as a drying agent 4.Describe the smell of carbon(IV)oxide gas Colourless and odourless 5. Effect on lime water. Experiment Bubbled carbon(IV)oxide gas into a test tube containing lime water for about three minutes Observation White precipitate is formed. White precipitate dissolved when excess carbon(IV)oxide gas is bubbled . Explanation Carbon(IV)oxide gas reacts with lime water(Ca(OH)2) to form an insoluble white precipitate of calcium carbonate. Calcium carbonate reacts with more Carbon(IV) oxide gas to form soluble Calcium hydrogen carbonate. Chemical equation Ca(OH)2(aq) + CO2 (g) -> CaCO3 (s) + H2O(l) CaCO3 (aq) + H2O(l) + CO2 (g) -> Ca(HCO3) 2 (aq) 6. Effects on burning Magnesium ribbon Experiment Lower a piece of burning magnesium ribbon into a gas jar containing carbon (IV)oxide gas. Observation The ribbon continues to burn with difficulty White ash/solid is formed. Black speck/solid/particles formed on the side of gas jar. Explanation Carbon(IV)oxide gas does not support combustion/burning.Magnesium burn to produce/release enough heat energy to decompose Carbon(IV) oxide gas to carbon and oxygen.Magnesium continues to burn in Oxygen forming white Magnesium Oxide solid/ash.Black speck/particle of carbon/charcoal residue forms on the sides of reaction flask. During the reaction Carbon(IV) oxide is reduced(Oxidizing agent)to carbon while Magnesium is Oxidized to Magnesium Oxide. Chemical equation 2Mg(s) + CO2 (g) -> C (s) + 2MgO(l)
7 7. Dry and wet litmus papers were separately put in a gas jar containing dry carbon (IV)oxide gas. State and explain the observations made.
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Chemical equation 2Mg(s) + CO2 (g) -> C (s) + 2MgO(l)
7 7. Dry and wet litmus papers were separately put in a gas jar containing dry carbon (IV)oxide gas. State and explain the observations made. Observation Blue dry litmus paper remain blue Red dry litmus paper remain Red Blue wet/damp/moist litmus paper turn red Red wet/damp/moist litmus paper remain red Explanation Dry Carbon (IV) oxide gas is a molecular compound that does not dissociate/ionize to release H+ and thus has no effect on litmus papers. Wet/damp/moist litmus papers contains water that dissolves/react with dry carbon (IV) oxide gas to form the weak solution of carbonic (IV) acid(H2CO3). Carbonic (IV) acid dissociate/ionizes to a few /little free H+ and CO32-. The few H+ (aq) ions are responsible for turning blue litmus paper to faint red showing the gas is very weakly acidic. Chemical equation H2CO3(aq) -> 2H+ (aq) + CO32-(aq) 8. Explain why Carbon (IV)oxide cannot be prepared from the reaction of: (i) marble chips with dilute sulphuric(VI)acid. Explanation Reaction forms insoluble calcium sulphate(VI)that cover/coat unreacted marble chips stopping further reaction Chemical equation CaCO3(s) + H2SO4 (aq) -> CaSO4 (s) + H2O(l) + CO2 (g) PbCO3(s) + H2SO4 (aq) -> PbSO4 (s) + H2O(l) + CO2 (g) BaCO3(s) + H2SO4 (aq) -> BaSO4 (s) + H2O(l) + CO2 (g) (ii) Lead(II)carbonate with dilute Hydrochloric acid. Reaction forms insoluble Lead(II)Chloride that cover/coat unreacted Lead(II) carbonate stopping further reaction unless the reaction mixture is heated. Lead(II)Chloride is soluble in hot water. Chemical equation PbCO3(s) + 2HCl (aq) -> PbCl2 (s) + H2O(l) + CO2 (g) 9.
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Reaction forms insoluble Lead(II)Chloride that cover/coat unreacted Lead(II) carbonate stopping further reaction unless the reaction mixture is heated. Lead(II)Chloride is soluble in hot water. Chemical equation PbCO3(s) + 2HCl (aq) -> PbCl2 (s) + H2O(l) + CO2 (g) 9. Describe the test for the presence of Carbon (IV)oxide. Using burning splint Lower a burning splint into a gas jar suspected to contain Carbon (IV)oxide gas.The burning splint is extinguished. Using Lime water. Bubble the gas suspected to be Carbon (IV)oxide gas.A white precipitate that dissolve in excess bubbling is formed. 8 Chemical equation Ca(OH)2(aq) + CO2 (g) -> CaCO3 (s) + H2O(l) CaCO3 (aq) + H2O(l) + CO2 (g) -> Ca(HCO3) 2 (aq) 10.State three main uses of Carbon (IV)oxide gas (i)In the Solvay process for the manufacture of soda ash/sodium carbonate (ii)In preservation of aerated drinks (iii)As fire extinguisher because it does not support combustion and is denser than air. (iv)In manufacture of Baking powder. (ii) Carbon(II)Oxide (CO) (a)Occurrence Carbon(II)oxide is found is found from incomplete combustion of fuels like petrol charcoal, liquefied Petroleum Gas/LPG. (b)School Laboratory preparation In the school laboratory carbon(II)oxide can be prepared from dehydration of methanoic acid/Formic acid(HCOOH) or Ethan-1,2-dioic acid/Oxalic acid(HOOCCOOH) using concentrated sulphuric(VI) acid. Heating is necessary.
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(ii) Carbon(II)Oxide (CO) (a)Occurrence Carbon(II)oxide is found is found from incomplete combustion of fuels like petrol charcoal, liquefied Petroleum Gas/LPG. (b)School Laboratory preparation In the school laboratory carbon(II)oxide can be prepared from dehydration of methanoic acid/Formic acid(HCOOH) or Ethan-1,2-dioic acid/Oxalic acid(HOOCCOOH) using concentrated sulphuric(VI) acid. Heating is necessary. METHOD 1:Preparation of Carbon (IV)Oxide from dehydration of Oxalic/ethan-1,2-dioic acid METHOD 2:Preparation of Carbon (IV)Oxide from dehydration of Formic/Methanoic acid
9 (c)Properties of Carbon (II)Oxide(Questions) 1.Write the equation for the reaction for the preparation of carbon(II)oxide using; (i)Method 1; Chemical equation HOOCCOOH(s) –Conc.H2SO4--> CO(g) + CO2 (g) + H2O(l) H2C2O4(s) –Conc.H2SO4--> CO(g) + CO2 (g) + H2O(l) (ii)Method 2; Chemical equation HCOOH(s) –Conc.H2SO4--> CO(g) + H2O(l) H2CO2(s) –Conc.H2SO4--> CO(g) + H2O(l) 2.What method of gas collection is used during the preparation of carbon (II) oxide. Over water because the gas is insoluble in water. Downward delivery because the gas is 1 ½ times denser than air . 3.What is the purpose of : (i) Potassium hydroxide/sodium hydroxide in Method 1 To absorb/ remove carbon (II) oxide produced during the reaction. 2KOH (aq) + CO2 (g) -> K2CO3 (s) + H2O(l) 2NaOH (aq) + CO2 (g) -> Na2CO3 (s) + H2O(l) (ii) Concentrated sulphuric(VI)acid in Method 1 and 2.
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Downward delivery because the gas is 1 ½ times denser than air . 3.What is the purpose of : (i) Potassium hydroxide/sodium hydroxide in Method 1 To absorb/ remove carbon (II) oxide produced during the reaction. 2KOH (aq) + CO2 (g) -> K2CO3 (s) + H2O(l) 2NaOH (aq) + CO2 (g) -> Na2CO3 (s) + H2O(l) (ii) Concentrated sulphuric(VI)acid in Method 1 and 2. Dehydrating agent –removes the element of water (Hydrogen and Oxygen in ratio 2:1) present in both methanoic and ethan-1,2-dioic acid. 10 4. Describe the smell of carbon(II)oxide. Colourless and odourless. 5. State and explain the observation made when carbon(IV)oxide is bubbled in lime water for a long time. No white precipitate is formed. 6. Dry and wet/moist/damp litmus papers were separately put in a gas jar containing dry carbon(IV)oxide gas. State and explain the observations made. Observation -blue dry litmus paper remains blue -red dry litmus paper remains red - wet/moist/damp blue litmus paper remains blue - wet/moist/damp red litmus paper remains red Explanation Carbon(II)oxide gas is a molecular compound that does not dissociate /ionize to release H+ ions and thus has no effect on litmus papers. Carbon(II)oxide gas is therefore a neutral gas. 7. Carbon (II)oxide gas was ignited at the end of a generator as below. (i)State the observations made in flame K. Gas burns with a blue flame (ii)Write the equation for the reaction taking place at flame K. 2CO(g) + O2 (g) -> 2CO2 (g) 8. Carbon(II)oxide is a reducing agent. Explain Experiment Pass carbon(II)oxide through glass tube containing copper (II)oxide. Ignite any excess poisonous carbon(II)oxide. Observation Colour change from black to brown. Excess carbon (II)oxide burn with a blue flame. Flame K Dry carbon(II)oxide
11 Explanation Carbon is a reducing agent.
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Observation Colour change from black to brown. Excess carbon (II)oxide burn with a blue flame. Flame K Dry carbon(II)oxide
11 Explanation Carbon is a reducing agent. It is used to reduce metal oxide ores to metal, itself oxidized to carbon(IV)oxide gas. Carbon(II)Oxide reduces black copper(II)oxide to brown copper metal Chemical Equation CuO(s) + CO(g) -> Cu(s) + CO2(g) (black) (brown) PbO(s) + CO(g) -> Pb(s) + CO2(g) (brown when hot/ (grey) yellow when cool) ZnO(s) + CO(g) -> Zn(s) + CO2(g) (yellow when hot/ (grey) white when cool) Fe2O3(s) + 3CO(s) -> 2Fe(s) + 3CO2(g) (brown when hot/cool (grey) Fe3O4 (s) + 4CO(g) -> 3Fe(s) + 4CO2(g) (brown when hot/cool (grey) These reaction are used during the extraction of many metals from their ore. 9. Carbon (II) oxide is a pollutant. Explain. Carbon(II)oxide is highly poisonous/toxic.It preferentially combine with haemoglobin to form stable carboxyhaemoglobin in the blood instead of oxyhaemoglobin.This reduces the free haemoglobin in the blood causing nausea , coma then death. 10.The diagram below show a burning charcoal stove/burner/jiko. Use it to answer the questions that follow. 12 Explain the changes that take place in the burner Explanation Charcoal stove has air holes through which air enters. Air oxidizes carbon to carbon(IV)oxide gas at region I. This reaction is exothermic(-∆H) producing more heat. Chemical equation C(s) + O2(g) -> CO2(g) Carbon(IV)oxide gas formed rises up to meet more charcoal which reduces it to Carbon(II)oxide gas. Chemical equation 2CO2 (g) + O2(g) -> 2CO (g) At the top of burner in region II, Carbon (II)oxide gas is further oxidized to Carbon(IV)oxide gas if there is plenty of air but escape if the air is limited poisoning the living things around.
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This reaction is exothermic(-∆H) producing more heat. Chemical equation C(s) + O2(g) -> CO2(g) Carbon(IV)oxide gas formed rises up to meet more charcoal which reduces it to Carbon(II)oxide gas. Chemical equation 2CO2 (g) + O2(g) -> 2CO (g) At the top of burner in region II, Carbon (II)oxide gas is further oxidized to Carbon(IV)oxide gas if there is plenty of air but escape if the air is limited poisoning the living things around. Chemical equation 2CO (g) + O2(g) -> 2CO2 (g) (excess air) 11.Describe the test for the presence of carbon(II)oxide gas. Experiment Burn/Ignite the pure sample of the gas. Pass/Bubble the products into lime water/Calcium hydroxide . Observation Colourless gas burns with a blue flame. A white precipitate is formed that dissolve on further bubbling of the products. Chemical equation 2CO (g) + O2(g) -> 2CO2 (g) (gas burns with blue flame) Chemical equation Ca(OH) 2 (aq) + CO2 (g) -> CaCO3 (s) + H2O(l) Chemical equation CO2 (g) + CaCO3 (s) + H2O(l) -> Ca(HCO3) 2 (aq) 12. State the main uses of carbon (II)oxide gas. (i) As a fuel /water gas (ii)As a reducing agent in the blast furnace for extracting iron from iron ore(Magnetite/Haematite) (iii)As a reducing agent in extraction of Zinc from Zinc ore/Zinc blende (iv) As a reducing agent in extraction of Lead from Lead ore/Galena (v) As a reducing agent in extraction of Copper from Copper iron sulphide/Copper pyrites. 13 (iii)Carbonate(IV) (CO32-)and hydrogen carbonate(IV(HCO3-) 1.Carbonate (IV) (CO32-) are normal salts derived from carbonic(IV)acid (H2CO3) and hydrogen carbonate (IV) (HCO3-) are acid salts derived from carbonic(IV)acid.
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State the main uses of carbon (II)oxide gas. (i) As a fuel /water gas (ii)As a reducing agent in the blast furnace for extracting iron from iron ore(Magnetite/Haematite) (iii)As a reducing agent in extraction of Zinc from Zinc ore/Zinc blende (iv) As a reducing agent in extraction of Lead from Lead ore/Galena (v) As a reducing agent in extraction of Copper from Copper iron sulphide/Copper pyrites. 13 (iii)Carbonate(IV) (CO32-)and hydrogen carbonate(IV(HCO3-) 1.Carbonate (IV) (CO32-) are normal salts derived from carbonic(IV)acid (H2CO3) and hydrogen carbonate (IV) (HCO3-) are acid salts derived from carbonic(IV)acid. Carbonic(IV)acid(H2CO3) is formed when carbon(IV)oxide gas is bubbled in water. It is a dibasic acid with two ionizable hydrogens. H2CO3(aq) ->2H+(aq) + CO32-(aq) H2CO3(aq) -> H+(aq) + HCO3 - (aq) 2.Carbonate (IV) (CO32-) are insoluble in water except Na2CO3 , K2CO3 and (NH4)2CO3 3.Hydrogen carbonate (IV) (HCO3-) are soluble in water. Only five hydrogen carbonates exist. Na HCO3 , KHCO3 ,NH4HCO3 Ca(HCO3)2 and Mg(HCO3)2 Ca(HCO3)2 and Mg(HCO3)2 exist only in aqueous solutions. 3.The following experiments show the effect of heat on Carbonate (IV) (CO32-) and Hydrogen carbonate (IV) (HCO3-) salts: Experiment In a clean dry test tube place separately about 1.0 of the following: Zinc(II)carbonate(IV), sodium hydrogen carbonate(IV), sodium carbonate(IV), Potassium carbonate(IV) ammonium carbonate(IV), potassium hydrogen carbonate(IV), Lead(II)carbonate(IV), Iron(II)carbonate(IV), and copper(II)carbonate(IV). Heat each portion gently the strongly. Test any gases produced with lime water.
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3.The following experiments show the effect of heat on Carbonate (IV) (CO32-) and Hydrogen carbonate (IV) (HCO3-) salts: Experiment In a clean dry test tube place separately about 1.0 of the following: Zinc(II)carbonate(IV), sodium hydrogen carbonate(IV), sodium carbonate(IV), Potassium carbonate(IV) ammonium carbonate(IV), potassium hydrogen carbonate(IV), Lead(II)carbonate(IV), Iron(II)carbonate(IV), and copper(II)carbonate(IV). Heat each portion gently the strongly. Test any gases produced with lime water. Observation (i)Colorless droplets form on the cooler parts of test tube in case of sodium carbonate(IV) and Potassium carbonate(IV). (ii)White residue/solid left in case of sodium hydrogen carbonate(IV), sodium carbonate(IV), Potassium carbonate(IV) and potassium hydrogen carbonate(IV). (iii)Colour changes from blue/green to black in case of copper(II)carbonate(IV). (iv) Colour changes from green to brown/yellow in case of Iron (II)carbonate(IV). (v) Colour changes from white when cool to yellow when hot in case of Zinc (II) carbonate(IV). (vi) Colour changes from yellow when cool to brown when hot in case of Lead (II) carbonate(IV). (vii)Colourless gas produced that forms a white precipitate with lime water in all cases. Explanation 1. Sodium carbonate(IV) and Potassium carbonate(IV) exist as hydrated salts with 10 molecules of water of crystallization that condenses and collects on cooler parts of test tube as a colourless liquid. Chemical equation Na2CO3 .10H2O(s) -> Na2CO3 (s) + 10H2O(l)
14 K2CO3 .10H2O(s) -> K2CO3 (s) + 10H2O(l) 2. Carbonate (IV) (CO32-) and Hydrogen carbonate (IV) (HCO3-) salts decompose on heating except Sodium carbonate(IV) and Potassium carbonate(IV). (a) Sodium hydrogen carbonate(IV) and Potassium hydrogen carbonate(IV) decompose on heating to form sodium carbonate(IV) and Potassium carbonate(IV).Water and carbon(IV)oxide gas are also produced.
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Chemical equation Na2CO3 .10H2O(s) -> Na2CO3 (s) + 10H2O(l)
14 K2CO3 .10H2O(s) -> K2CO3 (s) + 10H2O(l) 2. Carbonate (IV) (CO32-) and Hydrogen carbonate (IV) (HCO3-) salts decompose on heating except Sodium carbonate(IV) and Potassium carbonate(IV). (a) Sodium hydrogen carbonate(IV) and Potassium hydrogen carbonate(IV) decompose on heating to form sodium carbonate(IV) and Potassium carbonate(IV).Water and carbon(IV)oxide gas are also produced. Chemical equation 2NaHCO3 (s) -> Na2CO3 (s) + H2O(l) + CO2 (g) (white) (white) 2KHCO3 (s) -> K2CO3 (s) + H2O(l) + CO2 (g) (white) (white) (b) Calcium hydrogen carbonate(IV) and Magnesium hydrogen carbonate(IV) decompose on heating to form insoluble Calcium carbonate(IV) and Magnesium carbonate(IV).Water and carbon(IV)oxide gas are also produced. Chemical equation Ca(HCO3)2 (aq) -> CaCO3 (s) + H2O(l) + CO2 (g) (Colourless solution) (white) Mg(HCO3)2 (aq) -> MgCO3 (s) + H2O(l) + CO2 (g) (Colourless solution) (white) (c) Ammonium hydrogen carbonate(IV) decompose on heating to form ammonium carbonate(IV) .Water and carbon(IV)oxide gas are also produced. Chemical equation 2NH4HCO3 (s) -> (NH4)2CO3 (s) + H2O(l) + CO2 (g) (white) (white) (d)All other carbonates decompose on heating to form the metal oxide and produce carbon(IV)oxide gas e.g.
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Chemical equation 2NaHCO3 (s) -> Na2CO3 (s) + H2O(l) + CO2 (g) (white) (white) 2KHCO3 (s) -> K2CO3 (s) + H2O(l) + CO2 (g) (white) (white) (b) Calcium hydrogen carbonate(IV) and Magnesium hydrogen carbonate(IV) decompose on heating to form insoluble Calcium carbonate(IV) and Magnesium carbonate(IV).Water and carbon(IV)oxide gas are also produced. Chemical equation Ca(HCO3)2 (aq) -> CaCO3 (s) + H2O(l) + CO2 (g) (Colourless solution) (white) Mg(HCO3)2 (aq) -> MgCO3 (s) + H2O(l) + CO2 (g) (Colourless solution) (white) (c) Ammonium hydrogen carbonate(IV) decompose on heating to form ammonium carbonate(IV) .Water and carbon(IV)oxide gas are also produced. Chemical equation 2NH4HCO3 (s) -> (NH4)2CO3 (s) + H2O(l) + CO2 (g) (white) (white) (d)All other carbonates decompose on heating to form the metal oxide and produce carbon(IV)oxide gas e.g. Chemical equation MgCO3 (s) -> MgO (s) + CO2 (g) (white solid) (white solid) Chemical equation BaCO3 (s) -> BaO (s) + CO2 (g) (white solid) (white solid) Chemical equation CaCO3 (s) -> CaO (s) + CO2 (g) (white solid) (white solid) Chemical equation CuCO3 (s) -> CuO (s) + CO2 (g) (blue/green solid) (black solid) Chemical equation ZnCO3 (s) -> ZnO (s) + CO2 (g) (white solid) (white solid when cool/
15 Yellow solid when hot) Chemical equation PbCO3 (s) -> PbO (s) + CO2 (g) (white solid) (yellow solid when cool/ brown solid when hot) 4.The following experiments show the presence of Carbonate (IV) (CO32-) and Hydrogen carbonate (IV) (HCO3-) ions in sample of a salt: (a)Using Lead(II) nitrate(V) I.
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Chemical equation Ca(HCO3)2 (aq) -> CaCO3 (s) + H2O(l) + CO2 (g) (Colourless solution) (white) Mg(HCO3)2 (aq) -> MgCO3 (s) + H2O(l) + CO2 (g) (Colourless solution) (white) (c) Ammonium hydrogen carbonate(IV) decompose on heating to form ammonium carbonate(IV) .Water and carbon(IV)oxide gas are also produced. Chemical equation 2NH4HCO3 (s) -> (NH4)2CO3 (s) + H2O(l) + CO2 (g) (white) (white) (d)All other carbonates decompose on heating to form the metal oxide and produce carbon(IV)oxide gas e.g. Chemical equation MgCO3 (s) -> MgO (s) + CO2 (g) (white solid) (white solid) Chemical equation BaCO3 (s) -> BaO (s) + CO2 (g) (white solid) (white solid) Chemical equation CaCO3 (s) -> CaO (s) + CO2 (g) (white solid) (white solid) Chemical equation CuCO3 (s) -> CuO (s) + CO2 (g) (blue/green solid) (black solid) Chemical equation ZnCO3 (s) -> ZnO (s) + CO2 (g) (white solid) (white solid when cool/
15 Yellow solid when hot) Chemical equation PbCO3 (s) -> PbO (s) + CO2 (g) (white solid) (yellow solid when cool/ brown solid when hot) 4.The following experiments show the presence of Carbonate (IV) (CO32-) and Hydrogen carbonate (IV) (HCO3-) ions in sample of a salt: (a)Using Lead(II) nitrate(V) I. Using a portion of salt solution in a test tube .add four drops of Lead(II)nitrate(V)solution.Preserve. Observation inference White precipitate/ppt CO32- ,SO32- ,SO42- ,Cl- - II. To the preserved solution ,add six drops of dilutte nitric(V)acid. Preserve.
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Observation inference White precipitate/ppt CO32- ,SO32- ,SO42- ,Cl- - II. To the preserved solution ,add six drops of dilutte nitric(V)acid. Preserve. Observation inference White precipitate/ppt persists White precipitate/ppt dissolves SO42- ,Cl- CO32- ,SO32- II. To the preserved sample( that forms a precipitate ),heat to boil. Observation inference White precipitate/ppt persists White precipitate/ppt dissolves SO42- Cl- II. To the preserved sample( that do not form a precipitate ),add three drops of acidified potassium manganate(VII)/lime water Observation inference Effervescence/bubbles/fizzing colourless gas produced Acidified KMnO4 decolorized/no white precipitate on lime water Effervescence/bubbles/fizzing colourless gas produced Acidified KMnO4 not decolorized/ white precipitate on lime water SO32- CO32-
16 Experiments/Observations: (b)Using Barium(II)nitrate(V)/ Barium(II)chloride I. To about 5cm3 of a salt solution in a test tube add four drops of Barium(II) nitrate (V) / Barium(II)chloride. Preserve. Observation Inference White precipitate/ppt SO42- , SO32- , CO32- ions II. To the preserved sample in (I) above, add six drops of 2M nitric(V) acid . Preserve. Observation 1 Observation Inference White precipitate/ppt persists SO42- , ions Observation 2 Observation Inference White precipitate/ppt dissolves SO32- , CO32- , ions III.To the preserved sample observation 2 in (II) above, add 4 drops of acidified potassium manganate(VII) /dichromate(VI).
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To the preserved sample in (I) above, add six drops of 2M nitric(V) acid . Preserve. Observation 1 Observation Inference White precipitate/ppt persists SO42- , ions Observation 2 Observation Inference White precipitate/ppt dissolves SO32- , CO32- , ions III.To the preserved sample observation 2 in (II) above, add 4 drops of acidified potassium manganate(VII) /dichromate(VI). Observation 1 Observation Inference (i)acidified potassium manganate(VII)decolorized (ii)Orange colour of acidified potassium dichromate(VI) turns to green SO32- ions Observation 2 Observation Inference (i)acidified potassium manganate(VII) not decolorized (ii)Orange colour of acidified potassium dichromate(VI) does not turns to green CO32- ions Explanations Using Lead(II)nitrate(V)
17 (i)Lead(II)nitrate(V) solution reacts with chlorides(Cl-), Sulphate (VI) salts (SO42- ), Sulphate (IV)salts (SO32-) and carbonates(CO32-) to form the insoluble white precipitate of Lead(II)chloride, Lead(II)sulphate(VI), Lead(II) sulphate (IV) and Lead(II)carbonate(IV). Chemical/ionic equation: Pb2+(aq) + Cl- (aq) -> PbCl2(s) Pb2+(aq) + SO42+ (aq) -> PbSO4 (s) Pb2+(aq) + SO32+ (aq) -> PbSO3 (s) Pb2+(aq) + CO32+ (aq) -> PbCO3 (s) (ii)When the insoluble precipitates are acidified with nitric(V) acid, - Lead(II)chloride and Lead(II)sulphate(VI) do not react with the acid and thus their white precipitates remain/ persists. - Lead(II) sulphate (IV) and Lead(II)carbonate(IV) reacts with the acid to form soluble Lead(II) nitrate (V) and produce/effervesces/fizzes/bubbles out sulphur(IV)oxide and carbon(IV)oxide gases respectively. .
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Chemical/ionic equation: Pb2+(aq) + Cl- (aq) -> PbCl2(s) Pb2+(aq) + SO42+ (aq) -> PbSO4 (s) Pb2+(aq) + SO32+ (aq) -> PbSO3 (s) Pb2+(aq) + CO32+ (aq) -> PbCO3 (s) (ii)When the insoluble precipitates are acidified with nitric(V) acid, - Lead(II)chloride and Lead(II)sulphate(VI) do not react with the acid and thus their white precipitates remain/ persists. - Lead(II) sulphate (IV) and Lead(II)carbonate(IV) reacts with the acid to form soluble Lead(II) nitrate (V) and produce/effervesces/fizzes/bubbles out sulphur(IV)oxide and carbon(IV)oxide gases respectively. . Chemical/ionic equation: PbSO3 (s) + 2H+(aq) -> H2 O (l) + Pb2+(aq) + SO2 (g) PbCO3 (s) + 2H+(aq) -> H2 O (l) + Pb2+(aq) + CO2 (g) (iii)When Lead(II)chloride and Lead(II)sulphate(VI) are heated/warmed; - Lead(II)chloride dissolves in hot water/on boiling(recrystallizes on cooling) - Lead(II)sulphate(VI) do not dissolve in hot water thus its white precipitate persists/remains on heating/boiling. (iv)When sulphur(IV)oxide and carbon(IV)oxide gases are produced; - sulphur(IV)oxide will decolorize acidified potassium manganate(VII) and / or Orange colour of acidified potassium dichromate(VI) will turns to green. Carbon(IV)oxide will not.
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Chemical/ionic equation: PbSO3 (s) + 2H+(aq) -> H2 O (l) + Pb2+(aq) + SO2 (g) PbCO3 (s) + 2H+(aq) -> H2 O (l) + Pb2+(aq) + CO2 (g) (iii)When Lead(II)chloride and Lead(II)sulphate(VI) are heated/warmed; - Lead(II)chloride dissolves in hot water/on boiling(recrystallizes on cooling) - Lead(II)sulphate(VI) do not dissolve in hot water thus its white precipitate persists/remains on heating/boiling. (iv)When sulphur(IV)oxide and carbon(IV)oxide gases are produced; - sulphur(IV)oxide will decolorize acidified potassium manganate(VII) and / or Orange colour of acidified potassium dichromate(VI) will turns to green. Carbon(IV)oxide will not. Chemical equation: 5SO32-(aq) + 2MnO4- (aq) +6H+(aq) -> 5SO42-(aq) + 2Mn2+(aq) + 3H2O(l) (purple) (colourless) 3SO32-(aq) + Cr2O72-(aq) +8H+(aq) -> 3SO42-(aq) + 2Cr3+(aq) + 4H2O(l) (Orange) (green) - Carbon(IV)oxide forms an insoluble white precipitate of calcium carbonate if three drops of lime water are added into the reaction test tube when effervescence is taking place. Sulphur(IV)oxide will not. Chemical equation: Ca(OH)2(aq) + CO2 (g) -> CaCO3(s) + H2O(l) These tests should be done immediately after acidifying to ensure the gases produced react with the oxidizing agents/lime water.
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Chemical equation: 5SO32-(aq) + 2MnO4- (aq) +6H+(aq) -> 5SO42-(aq) + 2Mn2+(aq) + 3H2O(l) (purple) (colourless) 3SO32-(aq) + Cr2O72-(aq) +8H+(aq) -> 3SO42-(aq) + 2Cr3+(aq) + 4H2O(l) (Orange) (green) - Carbon(IV)oxide forms an insoluble white precipitate of calcium carbonate if three drops of lime water are added into the reaction test tube when effervescence is taking place. Sulphur(IV)oxide will not. Chemical equation: Ca(OH)2(aq) + CO2 (g) -> CaCO3(s) + H2O(l) These tests should be done immediately after acidifying to ensure the gases produced react with the oxidizing agents/lime water. 18 Using Barium(II)nitrate(V)/ Barium(II)Chloride (i)Barium(II)nitrate(V) and/ or Barium(II)chloride solution reacts with Sulphate (VI) salts (SO42- ), Sulphate (IV)salts (SO32-) and carbonates(CO32-) to form the insoluble white precipitate of Barium(II)sulphate(VI), Barium(II) sulphate (IV) and Barium(II)carbonate(IV). Chemical/ionic equation: Ba2+(aq) + SO42+ (aq) -> BaSO4 (s) Ba2+(aq) + SO32+ (aq) -> BaSO3 (s) Ba2+(aq) + CO32+ (aq) -> BaCO3 (s) (ii)When the insoluble precipitates are acidified with nitric(V) acid, - Barium (II)sulphate(VI) do not react with the acid and thus its white precipitates remain/ persists. - Barium(II) sulphate (IV) and Barium(II)carbonate(IV) reacts with the acid to form soluble Barium(II) nitrate (V) and produce /effervesces /fizzes/ bubbles out sulphur(IV)oxide and carbon(IV)oxide gases respectively. .
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Chemical/ionic equation: Ba2+(aq) + SO42+ (aq) -> BaSO4 (s) Ba2+(aq) + SO32+ (aq) -> BaSO3 (s) Ba2+(aq) + CO32+ (aq) -> BaCO3 (s) (ii)When the insoluble precipitates are acidified with nitric(V) acid, - Barium (II)sulphate(VI) do not react with the acid and thus its white precipitates remain/ persists. - Barium(II) sulphate (IV) and Barium(II)carbonate(IV) reacts with the acid to form soluble Barium(II) nitrate (V) and produce /effervesces /fizzes/ bubbles out sulphur(IV)oxide and carbon(IV)oxide gases respectively. . Chemical/ionic equation: BaSO3 (s) + 2H+(aq) -> H2 O (l) + Ba2+(aq) + SO2 (g) BaCO3 (s) + 2H+(aq) -> H2 O (l) + Ba2+(aq) + CO2 (g) (iii) When sulphur(IV)oxide and carbon(IV)oxide gases are produced; - sulphur(IV)oxide will decolorize acidified potassium manganate(VII) and / or Orange colour of acidified potassium dichromate(VI) will turns to green. Carbon(IV)oxide will not. Chemical equation: 5SO32-(aq) + 2MnO4- (aq) +6H+(aq) -> 5SO42-(aq) + 2Mn2+(aq) + 3H2O(l) (purple) (colourless) 3SO32-(aq) + Cr2O72-(aq) +8H+(aq) -> 3SO42-(aq) + 2Cr3+(aq) + 4H2O(l) (Orange) (green) - Carbon(IV)oxide forms an insoluble white precipitate of calcium carbonate if three drops of lime water are added into the reaction test tube when effervescence is taking place. Sulphur(IV)oxide will not.
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Carbon(IV)oxide will not. Chemical equation: 5SO32-(aq) + 2MnO4- (aq) +6H+(aq) -> 5SO42-(aq) + 2Mn2+(aq) + 3H2O(l) (purple) (colourless) 3SO32-(aq) + Cr2O72-(aq) +8H+(aq) -> 3SO42-(aq) + 2Cr3+(aq) + 4H2O(l) (Orange) (green) - Carbon(IV)oxide forms an insoluble white precipitate of calcium carbonate if three drops of lime water are added into the reaction test tube when effervescence is taking place. Sulphur(IV)oxide will not. Chemical equation: Ca(OH)2(aq) + CO2 (g) -> CaCO3(s) + H2O(l)
19 These tests should be done immediately after acidifying to ensure the gases produced react with the oxidizing agents/lime water. (iii) Sodium carbonate(IV) (Na2CO3) (a)Extraction of sodium carbonate from soda ash Sodium carbonate naturally occurs in Lake Magadi in Kenya as Trona.trona is the double salt ; sodium sesquicarbonate. NaHCO3 .Na2CO3 .H2O.It is formed from the volcanic activity that takes place in Lake Naivasha, Nakuru ,Bogoria and Elementeita .All these lakes drain into Lake Magadi through underground rivers. Lake Magadi has no outlet. Solubility of Trona decrease with increase in temperature.High temperature during the day causes trona to naturally crystallize .It is mechanically scooped/dredged/dug and put in a furnace. Inside the furnace, trona decompose into soda ash/sodium carbonate.
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Lake Magadi has no outlet. Solubility of Trona decrease with increase in temperature.High temperature during the day causes trona to naturally crystallize .It is mechanically scooped/dredged/dug and put in a furnace. Inside the furnace, trona decompose into soda ash/sodium carbonate. Chemical equation 2NaHCO3 .Na2CO3 .H2O (s) -> 3Na2CO3 (s) + 5H2O(l) + CO2 (g) (trona) (soda ash) Soda ash is then bagged and sold as Magadi soda.It is mainly used: (i)in making glass to lower the melting point of raw materials (sand/SiO2 from 1650oC and CaO from 2500oC to around 1500oC) (ii)in softening hard water (iii)in the manufacture of soapless detergents. (iv)Swimming pool “pH increaser” Sodium chloride is also found dissolved in the lake. Solubility of sodium chloride decrease with decreases in temperature/ sodium chloride has lower solubility at lower temperatures. When temperatures decrease at night it crystallize out .The crystals are then mechanically dug/dredged /scooped then packed for sale as animal/cattle feeds and seasoning food. Summary flow diagram showing the extraction of Soda ash from Trona Sodium chloride and Trona dissolved in the sea Natural fractional crystallization Crystals of Trona (Day time) Dredging /scooping/ digging Crushing Furnace (Heating) Carbon(IV) oxide
20 b)The Solvay process for industrial manufacture of sodium carbonate(IV) (i)Raw materials. -Brine /Concentrated Sodium chloride from salty seas/lakes. -Ammonia gas from Haber. -Limestone /Calcium carbonate from chalk /limestone rich rocks. -Water from rivers/lakes. (ii)Chemical processes Ammonia gas is passed up to meet a downward flow of sodium chloride solution / brine to form ammoniated brine/ammoniacal brine mixture in the ammoniated brine chamber The ammoniated brine mixture is then pumped up, atop the carbonator/ solvay tower. In the carbonator/ solvay tower, ammoniated brine/ammoniacal brine mixture slowly trickle down to meet an upward flow of carbon(IV)oxide gas.
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-Water from rivers/lakes. (ii)Chemical processes Ammonia gas is passed up to meet a downward flow of sodium chloride solution / brine to form ammoniated brine/ammoniacal brine mixture in the ammoniated brine chamber The ammoniated brine mixture is then pumped up, atop the carbonator/ solvay tower. In the carbonator/ solvay tower, ammoniated brine/ammoniacal brine mixture slowly trickle down to meet an upward flow of carbon(IV)oxide gas. The carbonator is shelved /packed with quartz/broken glass to (i) reduce the rate of flow of ammoniated brine/ammoniacal brine mixture. (ii)increase surface area of the liquid mixture to ensure a lot of ammoniated brine/ammoniacal brine mixture react with carbon(IV)oxide gas. Insoluble sodium hydrogen carbonate and soluble ammonium chloride are formed from the reaction. Chemical equation CO2(g) + H2O(l) + NaCl (aq) + NH3(g) -> NaHCO3(s) + NH4Cl(aq)
21 The products are then filtered. Insoluble sodium hydrogen carbonate forms the residue while soluble ammonium chloride forms the filtrate. Sodium hydrogen carbonate itself can be used: (i) as baking powder and preservation of some soft drinks. (ii) as a buffer agent and antacid in animal feeds to improve fibre digestion. (iii) making dry chemical fire extinguishers. In the Solvay process Sodium hydrogen carbonate is then heated to form Sodium carbonate/soda ash, water and carbon (IV) oxide gas. Chemical equation 2NaHCO3 (s) -> Na2CO3(s) + CO2(g) + H2O(l) Sodium carbonate is stored ready for use in: (i) during making glass/lowering the melting point of mixture of sand/SiO2 from 1650oC and CaO from 2500oC to around 1500oC (ii) in softening hard water (iii) in the manufacture of soapless detergents. (iv) swimming pool “pH increaser”. Water and carbon(IV)oxide gas are recycled back to the ammoniated brine/ammoniacal brine chamber. More carbon(IV)oxide is produced in the kiln/furnace. Limestone is heated to decompose into Calcium oxide and carbon(IV)oxide.
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Water and carbon(IV)oxide gas are recycled back to the ammoniated brine/ammoniacal brine chamber. More carbon(IV)oxide is produced in the kiln/furnace. Limestone is heated to decompose into Calcium oxide and carbon(IV)oxide. Chemical equation CaCO3 (s) -> CaO(s) + CO2(g) Carbon(IV)oxide is recycled to the carbonator/solvay tower. Carbon (IV)oxide is added water in the slaker to form Calcium hydroxide. This process is called slaking. Chemical equation CaO(s) + H2O (l) -> Ca(OH)2 (aq) Calcium hydroxide is mixed with ammonium chloride from the carbonator/solvay tower in the ammonia regeneration chamber to form Calcium chloride , water and more ammonia gas. Chemical equation Ca(OH)2 (aq) +2NH4Cl (aq) -> CaCl2(s) + 2NH3(g) + H2O(l) NH3(g) and H2O(l) are recycled. Calcium chloride may be used: (i)as drying agent in the school laboratory during gas preparation (except ammonia gas)
22 (ii)to lower the melting point of solid sodium chloride / rock salt salts during the Downs process for industrial extraction of sodium metal. Detailed Summary flow diagram of Solvay Process
23 AmmoniatedbrineBrineAmmonia regenerationchamberSolvay Tower/CarbonatorKiln/FurnaceAmmoniumchlorideCarbon(IV)OxideHaberprocessSlakerCalcium hydroxideCalcium oxideWaterSodium hydrogen CarbonateRoasterSodiumcarbonateCalcium chlorideBrine saturated with ammoniaCoke & Limestone
24 Practice 1. The diagram below shows part of the Solvay process used in manufacturing sodium carbonate. Use it to answer the questions that follow. (a)Explain how Sodium Chloride required for this process is obtained from the sea. Sea water is pumped /scooped into shallow pods. Evaporation of most of the water takes place leaving a very concentrated solution. (b)(i) Name process: I. Filtration II.
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Evaporation of most of the water takes place leaving a very concentrated solution. (b)(i) Name process: I. Filtration II. Decomposition (ii) Write the equation for the reaction in process: Process I Chemical equation CO2(g) + H2O(l) + NaCl (aq) + NH3(g) -> NaHCO3(s) + NH4Cl(aq) Process II Chemical equation 2NaHCO3 (s) -> Na2CO3(s) + CO2(g) + H2O(l) (c)(i) Name two substances recycled in the solvay process Ammonia gas , Carbon(IV)Oxide and Water. (ii)Which is the by-product of this process? Calcium(II)Chloride /CaCl2 (iii)State two uses that the by-product can be used for: 1. As a drying agent in the school laboratory preparation of gases. 2. In the Downs cell/process for extraction of Sodium to lower the melting point of rock salt. Carbon (IV)oxide Ammonium chloride Process I Saturated sodium chloride solution Ammonia Sodium hydrogen carbonate Process II Sodium carbonate
25 (iv)Write the chemical equation for the formation of the byproducts in the Solvay process. Chemical equation Ca(OH)2 (aq) +2NH4Cl (aq) -> CaCl2(s) + 2NH3(g) + H2O(l) (d)In an experiment to determine the % purity of Sodium carbonate produced in the Solvay process ,2.15g of the sample reacted with exactly 40.0cm3 of 0.5M Sulphuric(VI)acid. (i)Calculate the number of moles of sodium carbonate that reacted. Chemical equation Na2CO3 (aq) +H2SO4 (aq) -> Na2SO4 (aq)+ CO2(g) + H2O(l) Mole ratio Na2CO3 :H2SO4 => 1:1 Moles H2SO4 = Molarity x Volume => 0.5 x 40.0 = 0.02 Moles 1000 1000 Moles of Na2CO3 = 0.02 Moles (ii)Determine the % of sodium carbonate in the sample.
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Chemical equation Ca(OH)2 (aq) +2NH4Cl (aq) -> CaCl2(s) + 2NH3(g) + H2O(l) (d)In an experiment to determine the % purity of Sodium carbonate produced in the Solvay process ,2.15g of the sample reacted with exactly 40.0cm3 of 0.5M Sulphuric(VI)acid. (i)Calculate the number of moles of sodium carbonate that reacted. Chemical equation Na2CO3 (aq) +H2SO4 (aq) -> Na2SO4 (aq)+ CO2(g) + H2O(l) Mole ratio Na2CO3 :H2SO4 => 1:1 Moles H2SO4 = Molarity x Volume => 0.5 x 40.0 = 0.02 Moles 1000 1000 Moles of Na2CO3 = 0.02 Moles (ii)Determine the % of sodium carbonate in the sample. Molar mass of Na2CO3 = 106g Mass of Na2CO3 = moles x Molar mass => 0.02 x 106 = 2.12 g % of Na2CO3 = ( 2.12 g x 100) = 98.6047% 2.15 (e) State two uses of soda ash. (i) during making glass/lowering the melting point of mixture of sand/SiO2 from 1650oC and CaO from 2500oC to around 1500oC (ii) in softening hard water (iii) in the manufacture of soapless detergents. (iv) swimming pool “pH increaser”. (f)The diagram below shows a simple ammonia soda tower used in manufacturing sodium carbonate .Use it to answer the questions that follow: Raw materialExcess Carbon(IV)oxide Metal plates Substance
26 (i)Name the raw materials needed in the above process -Ammonia -Water -Carbon(IV)oxide -Limestone -Brine/ Concentrated sodium chloride (ii)Identify substance A Ammonium chloride /NH4Cl (iii) Write the equation for the reaction taking place in: I.Tower.
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(i) during making glass/lowering the melting point of mixture of sand/SiO2 from 1650oC and CaO from 2500oC to around 1500oC (ii) in softening hard water (iii) in the manufacture of soapless detergents. (iv) swimming pool “pH increaser”. (f)The diagram below shows a simple ammonia soda tower used in manufacturing sodium carbonate .Use it to answer the questions that follow: Raw materialExcess Carbon(IV)oxide Metal plates Substance
26 (i)Name the raw materials needed in the above process -Ammonia -Water -Carbon(IV)oxide -Limestone -Brine/ Concentrated sodium chloride (ii)Identify substance A Ammonium chloride /NH4Cl (iii) Write the equation for the reaction taking place in: I.Tower. Chemical equation CO2(g) + NaCl (aq) + H2O(l) + NH3(g) -> NaHCO3(s) + NH4Cl(aq) II. Production of excess carbon (IV)oxide. Chemical equation CaCO3 (s) -> CaO(s) + CO2(g) III. The regeneration of ammonia Chemical equation Ca(OH)2 (aq) +2NH4Cl (aq) -> CaCl2(s) + 2NH3(g) + H2O(l) (iv)Give a reason for having the circular metal plates in the tower. -To slow the downward flow of brine. -To increase the rate of dissolving of ammonia. -To increase the surface area for dissolution (v)Name the gases recycled in the process illustrated above. Ammonia gas , Carbon(IV)Oxide and Water. 2. Describe how you would differentiate between carbon (IV)oxide and carbon(II)oxide using chemical method. Method I -Bubble both gases in lime water/Ca(OH)2 -white precipitate is formed if the gas is carbon (IV) oxide - No white precipitate is formed if the gas is carbon (II) oxide Method II
27 -ignite both gases - Carbon (IV) oxide does not burn/ignite - Carbon (II) oxide burn with a blue non-sooty flame. Method III -Lower a burning splint into a gas containing each gas separately.
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Describe how you would differentiate between carbon (IV)oxide and carbon(II)oxide using chemical method. Method I -Bubble both gases in lime water/Ca(OH)2 -white precipitate is formed if the gas is carbon (IV) oxide - No white precipitate is formed if the gas is carbon (II) oxide Method II
27 -ignite both gases - Carbon (IV) oxide does not burn/ignite - Carbon (II) oxide burn with a blue non-sooty flame. Method III -Lower a burning splint into a gas containing each gas separately. -burning splint is extinguished if the gas is carbon (IV) oxide -burning splint is not extinguished if the gas is carbon (II) oxide. 3.Using Magnesium sulphate(VI)solution ,describe how you can differentiate between a solution of sodium carbonate from a solution of sodium hydrogen carbonate -Add Magnesium sulphate(VI) solution to separate portions of a solution of sodium carbonate and sodium hydrogen carbonate in separate test tubes -White precipitate is formed in test tube containing sodium carbonate -No white precipitate is formed in test tube containing sodium hydrogen carbonate. Chemical equation Na2CO3 (aq) +MgSO4 (aq) -> Na2SO4 (aq) + MgCO3(s) (white ppt) Ionic equation CO32- (aq) + Mg2+ (aq) -> MgCO3(s) (white ppt) Chemical equation 2NaHCO3 (aq) +MgSO4 (aq) -> Na2SO4 (aq) + Mg(HCO3)2 (aq) (colourless solution) 4. The diagram below shows a common charcoal burner .Assume the burning take place in a room with sufficient supply of air. (a)Explain what happens around:
28 (i)Layer A Sufficient/excess air /oxygen enter through the air holes into the burner .It reacts with/oxidizes Carbon to carbon(IV)oxide Chemical equation C(s) + O2(g) -> CO2 (g) (ii)Layer B Hot carbon(IV)oxide rises up and is reduced by more carbon/charcoal to carbon (II)oxide. Chemical equation C(s) + CO2(g) -> 2CO (g) (ii)Layer C Hot carbon(II)oxide rises up and burns with a blue flame to be oxidized by the excess air to form carbon(IV)oxide.
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