Group 2 Elements Revision Guide (AQA)

2.2 Group 2 Melting points Melting points decrease down the group. The metallic Atomic radius bonding weakens as the atomic size...

2.2 Group 2 Melting points Melting points decrease down the group. The metallic Atomic radius bonding weakens as the atomic size increases. The Atomic radius increases down the group. distance between the positive ions and delocalized As one goes down the group, the atoms have more electrons increases. Therefore the electrostatic shells of electrons making the atom bigger. attractive forces between the positive ions and the delocalized electrons weaken. 1st Ionisation Energy The outermost electrons are held more weakly because they are successively further from the nucleus in additional shells. In addition, the outer shell electrons become more shielded from the attraction of the nucleus by the repulsive force of inner shell electrons. Group 2 reactions Reactivity of group 2 metals increases down the group. Reactions with oxygen. Magnesium will also react slowly in oxygen without a flame. The group 2 metals will burn in oxygen. Magnesium ribbon will often have a thin layer of magnesium oxide Mg burns with a bright white flame. The on it formed by the reaction with oxygen in the air. MgO appears as a white powder. 2 Mg + O2  2 MgO 2 Mg + O2  2 MgO The magnesium oxide needs to be removed by emery paper before doing reactions with magnesium ribbon. MgO is a white solid with a high melting If testing for reaction rates with Mg and acid, an un-cleaned Mg point due to its ionic bonding. ribbon would give a false result because both the Mg and MgO would react but at different rates. Mg + 2 HCl  MgCl2 + H2 MgO + 2 HCl  MgCl2 + H2O Reactions with water. Make sure you learn the difference between the reaction of magnesium with steam and that with warm water. Magnesium reacts in steam to produce Mg will also react with warm water, giving a different magnesium oxide and hydrogen. The Mg magnesium hydroxide product. would burn with a bright white flame. The Mg + 2 H2O  Mg(OH)2 + H2 MgO appears as a white powder. This is a much slower reaction than the reaction with Mg(s) + H2O(g)  MgO(s) + H2(g) steam and there is no flame. The other group 2 metals will react with cold water with One would observe: increasing vigour down the group to form hydroxides. fizzing, (more vigorous down group) Ca + 2 H2O (l) Ca(OH)2 (aq) + H2 (g) the metal dissolving, (faster down group) Sr + 2 H2O (l) Sr(OH)2 (aq) + H2 (g) the solution heating up (more down group) Ba + 2 H2O (l) Ba(OH)2 (aq) + H2 (g) with calcium a white precipitate appearing The hydroxides produced make the water alkaline (less precipitate forms down group with (if they are soluble in water). other metals) N Goalby chemrevise.org 1 Using magnesium to extract titanium Titanium cannot be extracted with carbon Titanium is a very useful metal because it is abundant, because titanium carbide (TiC) it is formed has a low density and is corrosion resistant – it is used for rather than titanium. making strong, light alloys for use in aircraft. Titanium cannot be extracted by electrolysis because it has to be very pure. Titanium is extracted by reaction with a more reactive metal (e.g. Magnesium). Steps in extracting titanium TiO2 + 2 Cl2 + 2 C TiCl4 + 2 CO 1. TiO2 (solid) is converted to TiCl4 (liquid) at 900C: 2. The TiCl4 is purified by fractional distillation in an argon atmosphere. TiCl4 + 2Mg Ti + 2 MgCl2 3. The Ti is extracted by Mg in an argon atmosphere at 500C Titanium is expensive because: TiO2 is converted to TiCl4 as it can be purified by fractional distillation, 1. The expensive cost of the magnesium TiCl4 being molecular (liquid at 2. This is a batch process which makes it expensive because the room temperature) rather than process is slower (having to fill up and empty reactors takes ionic like TiO2 (solid at room time) and requires more labour and the energy is lost when the temperature). reactor is cooled down after stopping 3. The process is also expensive due to the argon, and the need to remove moisture (because TiCl4 is susceptible to hydrolysis). 4. High temperatures required in both steps This all makes titanium expensive even though it is a relatively abundant metal. It is only therefore used to a limited amount even though it has useful properties. Calcium oxide can be used to remove SO2 from the waste gases from furnaces (e.g. coal fired power stations) by flue gas desulfurisation. The The calcium sulfite which is gases pass through a scrubber containing basic calcium oxide which formed can be used to make reacts with the acidic sulfur dioxide in a neutralisation reaction. calcium sulfate for plasterboard. SO2 + CaO  CaSO3 calcium sulfite N Goalby chemrevise.org 2 Solubility of Hydroxides Group II hydroxides become more soluble down the group. All Group II hydroxides when not soluble appear as white precipitates. Magnesium hydroxide is classed as insoluble in water. Calcium hydroxide is classed as partially soluble in water and will appear as a white Simplest ionic equation for formation of Mg(OH)2(s) precipitate It is used in agriculture to neutralise Mg2+(aq) + 2OH-(aq)  Mg(OH)2(s). acidic soils. A suspension of magnesium hydroxide in water will A suspension of calcium hydroxide in water will appear slightly alkaline (pH 9) so some hydroxide ions appear more alkaline (pH 11) than magnesium must therefore have been produced by a very slight hydroxide as it is more soluble so there will be dissolving. more hydroxide ions present in solution. Magnesium hydroxide is used in medicine (in suspension An aqueous solution of calcium hydroxide is as milk of magnesia) to neutralise excess acid in the called lime water and can be used a test for stomach and to treat constipation. carbon dioxide. The limewater turns cloudy as white calcium carbonate is produced. Mg(OH)2 + 2HCl  MgCl2 + 2H2O Ca(OH)2 (aq) + CO2 (g)  CaCO3 (s) + H2O(l) It is safe to use because it is so weakly alkaline. It is preferable to using calcium carbonate as it will not Barium hydroxide would easily dissolve in produce carbon dioxide gas. water. The hydroxide ions present would make the solution strongly alkaline. Solubility of sulfates Ba(OH)2 (s) + aq  Ba2+ (aq) + 2OH-(aq) Group II sulfates become less soluble down the group. BaSO4 is the least soluble. An equation for the formation of the precipitate can be written as a full equation or simplest ionic equation. Full equation : SrCl2(aq) + Na2SO4 (aq)  2NaCl (aq) + SrSO4 (s) Ionic equation: Sr2+ (aq) + SO42-(aq)  SrSO4 (s). BaSO4 is used in medicine as a ‘Barium meal’ given to patients who need x-rays of their intestines. The barium absorbs the x-rays and so the gut shows up on the x-ray image. Even though barium compounds are toxic, it is safe to use here because barium sulfate’s low solubility means it is not absorbed into the blood. If barium metal is reacted with sulfuric acid it will only react slowly, as the insoluble barium sulfate produced will cover the surface of the metal and act as a barrier to further attack. Ba + H2SO4  BaSO4 + H2 The same effect happens to a lesser extent with metals going up the group as the solubility of the sulfates increases. The same effect does not happen with other acids like hydrochloric or nitric as they form soluble group 2 salts. Testing for presence of a sulfate ion BaCl2 solution acidified with hydrochloric acid is used as a reagent to Other anions should give a test for sulfate ions. negative result which is no If acidified barium chloride is added to a solution that contains sulfate ions a precipitate forming. white precipitate of barium sulfate forms. Simplest ionic equation Ba2+ (aq) + SO42-(aq)  BaSO4 (s). The hydrochloric acid is needed to react with carbonate impurities that are often found in salts which would form a white barium carbonate precipitate and so give a false result. You could not use sulfuric acid because it contains sulfate ions and so would give a false positive result. 2HCl + Na2CO3  2NaCl + H2O + CO2 Fizzing due to CO2 would be observed if a carbonate was present. N Goalby chemrevise.org 3  More on Insoluble salts and Precipitation reactions Insoluble salts can be made by mixing appropriate solutions of ions so that a precipitate is formed barium nitrate (aq) + sodium sulfate (aq)  barium sulfate (s) + sodium nitrate (aq) These are called precipitation reactions. A precipitate is a solid. There are some common rules for solubility of salts. No syllabus requires these to be learnt but a good chemist does know them. Soluble salts Insoluble salts All sodium, potassium and ammonium salts All nitrates Most chlorides, bromides, iodides Silver, lead chlorides, bromides iodides Most sulfates Lead, strontium and barium sulfates Sodium, potassium and ammonium Most other carbonates carbonates Sodium, potassium and ammonium Most other hydroxides hydroxides When making an insoluble salt, normally the salt would be removed by filtration, washed with distilled water to remove soluble impurities and then dried on filter paper. Filtration Filter paper Buchner residue funnel Filter Filter paper funnel Buchner flask (has Air outlet to thicker glass walls water pump than a normal flask filtrate to cope with the vacuum ) This is gravitational filtration. This is vacuum filtration. The apparatus is Use if small amounts of solid connected to a water pump which will are formed. produce a vacuum. Use if larger amounts of solid are formed. For both types of filtration apparatus AQA expect filter paper to be drawn on the diagram Writing ionic equations for precipitation reactions We usually write ionic equations to show precipitation Spectator ions are ions that are reactions. Ionic equations only show the ions that are not changing state reacting and leave out spectator ions. not changing oxidation number Take full equation Ba(NO3)2 (aq) + Na2SO4 (aq)  BaSO4 (s) + 2 NaNO3 (aq) Separate aqueous Ba2+(aq) + 2 NO3-(aq) + 2 Na+ (aq)+ SO42-(aq)  BaSO4(s) + 2 Na+(aq)+ 2 NO3- (aq) solutions into ions. Cancel out spectator ions leaving Ba2+ (aq) + SO42-(aq)  BaSO4 (s). the simplest ionic equation. N Goalby chemrevise.org 4