32 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 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.


(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. 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
33 (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: 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.


Group II elements: Alkaline earth metals Group II elements are called Alkaline earth metals . The alkaline earth metals include: 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. 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
34 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. Electropositivity All alkaline earth metals are also electropositive like alkali metals. The electropositivity increase with increase in atomic radius/size.


The remaining electrons/energy levels experience more effective / greater nuclear attraction/pull towards the nucleus reducing the atomic radius. 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. 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
35 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.


Atomic radius of Sodium (0.157nM) is higher than that of Magnesium (0.137nM). This is because Magnesium has more effective nuclear attraction on
35 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 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. 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.


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. 36 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 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.


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. 37 2Mg (s) + O2(g) -> 2MgO(s) 3Mg (s) + N2 (g) -> Mg3N2 (s) Magnesium oxide dissolves in water to form magnesium hydroxide. 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.


Repeat with Calcium metal. 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)
38 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. Mg(s) + H2O (l) -> MgO(s) + H2 (g)
39 (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.


-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. 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.


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
40 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. (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.


(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. 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 -
41 (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. 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.


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). 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
42 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.


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 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. For halogens the 1st electron affinity decrease down the group as the
43 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.


(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. 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. Halogen F Cl Br I Electron affinity kJmole-1 -333 -364 -342 -295
44 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.


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. 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
45 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. On adding Bromine water, iodine is displaced from its solution but not chlorine.


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.


- 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.


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.


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.


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.


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.


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.


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.


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.


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.


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 .


(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.


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.


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.


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.


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.


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.


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.


-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.


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.


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.


(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.


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.


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.


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.


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.


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.


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.


- 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.


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.


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).


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.


 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.


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.


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.


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.


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.