1. Topic 6: Acids and Bases (including writing chemical & ionic equations)

4. Acids solutions used in the laboratory may be concentrated or dilute. Concentrated acids contain a large amount of pure acid dissolved in water, while dilute acids contain small amounts of acid in water. Two bottles of HCl have concentrations of 5M & 0.5M. Which is more concentrated? Explain The one with 5 mol / dm3 5 M 0.5 M

8. Example Write the ionic equation for the reaction between magnesium and hydrochloric acid. Step 1: Write the balanced chemical equation (including the state symbols) Mg(s) + 2HCl(aq)  MgCl 2 (aq) + H 2 (g) Step 2: Split aqueous compounds into their ions. (Only compounds that are aqueous are split into ions) LHS RHS Mg (s) + 2H + (aq) + 2Cl - (aq) -> Mg 2 + (aq) + 2Cl - (aq) + H 2 (g)

10. Question: Write the ionic equation for reaction between Copper (II) Sulphate solution and sodium hydroxide solution. Above reaction is an example of insoluble salt preparation (Precipitation also occurs between aqueous solutions of 2 soluble salts )

15. Other acids behave in a similar way. Pure sulphuric acid and nitric acid (without water) are liquids which consist of simple molecules. In water, the acid ionises and hydrogen ions, H + (aq) are produced. So what is an Acid?? An acid i s a substance that produces hydrogen ions, H + (aq) , in water/aqueous solution . Thus, acids have acidic properties only when dissolved in water or when in aqueous solution , where hydrogen ions, H + (aq) are present. The properties and reactions of acids are due to the presence of the hydrogen ions! Lets see what happened in the reaction between Magnesium and HCI again!

16. Exercise 1. Citric acid, a white solid, can be dissolved in two solvents, water and propanone to form solutions. The solution in water turns Universal Indicator paper an orange-red colour. The solution in propanone has no effect on Universal Indicator paper. Explain the above differences . When citric acid dissolves in water, the acid ionises to produce hydrogen ions, H + (aq). Due to the presence of the hydrogen ions, citric acid in water has acidic properties, thus it turns Universal Indicator paper orange-red. When citric acid dissolves in propanone, the acid does not ionise but remains as covalent molecules. As hydrogen ions are not present, citric acid in propanone does not have acidic properties, thus it has no effect on Universal Indicator paper.

18. (4) Basicity of Acids The basicity of an acid refers to the maximum number of hydrogen ions produced by a molecule of the acid in aqueous solution . Acid Reaction with water Basicity Hydrochloric acid HCl(aq)  H + (aq) + Cl - (aq) Monobasic Nitric acid HNO 3 (aq)  H + (aq) + NO 3 - (aq) Monobasic Ethanoic acid CH 3 COOH(aq)  H + (aq) + CH 3 COO - (aq) Monobasic Sulphuric acid H 2 SO 4 (aq)  2H + (aq) + SO 4 2- (aq) Dibasic Phosphoric(V) acid H 3 PO 4 (aq)  3H + (aq) + PO 4 3- (aq) Tribasic

30. a) Ca(OH) 2 (aq) + 2HNO 3 (aq)  Ca(NO 3 ) 2 (aq) + 2H 2 O(l) calcium nitrate H + (aq) + OH - (aq)  H 2 O(l) b) (NH 4 ) 2 SO 4 (aq) + 2KOH(aq)  K 2 SO 4 (aq) + 2H 2 O(l) + 2NH 3 (g) potassium sulphate NH 4 + (aq) + OH - (aq)  H 2 O(l) + NH 3 (g) c) Fe 2 (SO 4 ) 3 (aq) + 6NaOH(aq)  2Fe(OH) 3 (s) + 3Na 2 SO 4 (aq) iron(III) hydroxide Fe 3+ (aq) + 3OH - (aq)  Fe(OH) 3 (s) d) Pb(NO 3 ) 2 (aq) + 2NH 4 OH(aq)  Pb(OH) 2 (s) + 2NH 4 NO 3 (aq) lead(II) hydroxide Pb 2+ (aq) + 2OH - (aq)  Pb(OH) 2 (s)

31. (3) Alkalis & Hydroxide ions So what is an alkali? An alkali i s a substance that produces hydroxide ions, OH - (aq) , in water/aqueous solution . The properties and reactions of alkalis are due to the presence of the hydroxide ions. Thus, alkalis have alkaline properties only when dissolved in water or when in aqueous solution , where hydroxide ions, OH - (aq) are present.

33. (5) Uses of Alkalis 1. An alkali, calcium hydroxide (slaked lime) , or the base calcium oxide (quicklime) are added by farmers to neutralise excess acids in soils that have become too acidic from extensive use of chemical fertilizers (e.g. ammonium sulphate) and from acid rain. 2. Mild alkalis are used in medicine such as antacid to relieve the pain of indigestion as caused by excess acid in the stomach . Antacid contains a base such as magnesium hydroxide or aluminium hydroxide that neutralises the excess acid.

34. 3. Alkalis are used in toothpastes to neutralise acids on our teeth produced by bacteria when they feed on sugars in our food. If the acid is not destroyed, it corrodes the teeth causing tooth decay . Toothpastes usually contain magnesium hydroxide which neutralises the acids in the mouth. 4. Alkalis are used to remove grease . Mild/weak alkalis are contained in soaps and detergents . Aqueous ammonia , a mild/weak alkali, is used in window cleaners to remove grease and dirt from glass. Sodium hydroxide , a powerful/strong alkali, is used in floor and oven cleaners .

39. (2) Measurement of pH There are other indicators that not only indicate whether a solution is acidic or alkaline, but also how acidic or alkaline it is, by indicating the pH value of the solution, e.g. Universal Indicator . The pH value of a solution can be measured by using Universal Indicator, a pH meter or by a pH sensor connected to a computer. 1. Universal Indicator (or pH indicator) is a mixture of indicators which gives different colours at different pH values . It is used in the form of a solution or a paper (pH paper). The pH of a solution is measured by adding a few drops of Universal Indicator to the solution or by dipping a piece of Universal Indicator paper in the solution. The pH of the solution is then found by comparing the colour obtained with a colour chart .

41. (3) Importance of pH 1. pH and the body Substances in the body have different pH values. Acidic conditions in the stomach (pH ~ 1.5) and alkaline conditions in the small intestine (pH ~ 8.4) are needed for good digestion. Slightly acidic condition in the blood (pH ~ 6.5) that goes to the heart and lungs is due to carbon dioxide present in the blood. 2. pH and food preservation Many fresh food quickly go bad mainly due to microorganisms , such as bacteria , present in the food. Acids can be used to preserve foods because microorganisms do not grow well in solutions of low pH .

46. 1. Acidic Oxides a. Acidic oxides react with water to produce acids SO 3 (g) + H 2 O(l)  H 2 SO 4 (aq) SO 2 (g) + H 2 O(l)  H 2 SO 3 (aq) CO 2 (g) + H 2 O(l)  H 2 CO 3 (aq) P 4 O 10 (s) + 6H 2 O(l)  4H 3 PO 4 (aq) b. Acidic oxides react with alkalis to produce salt and water . CO 2 (g) + 2NaOH(aq)  Na 2 CO 3 (aq) + H 2 O(l) SO 3 (g) + Ca(OH) 2 (aq)  CaSO 4 (s) + H 2 O(l) (2) Reactions of Oxides

47. 2. Basic Oxides a. Basic oxides react with water to produce alkalis . (most basic oxides are insoluble in water) K 2 O(s) + H 2 O(l)  2KOH(aq) b. Basic oxides react with acids to produce salt and water . CaO(s) + 2HNO 3 (aq)  Ca(NO 3 ) 2 (aq) + H 2 O(l) 3. Neutral Oxides Neutral oxides do not react with either acids or bases, thus they do not form salts. 4. Amphoteric Oxides Amphoteric oxides behave as an acidic oxide or as a basic oxide , hence they react both with acids and with alkalis to form salts . Al 2 O 3 (s) + 6HCl(aq)  2AlCl 3 (aq) + 3H 2 O(l) acts as base Al 2 O 3 (s) + NaOH(aq)  sodium aluminate + H 2 O(l) acts as acid NaAl(OH) 4 (aq )

48. Question: ‘ Gallium oxide reacts with hydrochloric acid and sodium hydroxide to form salts.’ What can you deduce from this statement? Gallium oxide is an amphoteric oxide (and gallium is a metal)!

49. (3) Sulphur Dioxide & Its Uses Sulphur dioxide is formed when sulphur burns in air. As it is an acidic oxide , it reacts with sodium hydroxide solution to form the salt, sodium sulphite : SO 2 (g) + 2NaOH(aq)  Na 2 SO 3 (aq) + H 2 O(l) sodium sulphite 1. Sulphur dioxide and sodium sulphite has many uses: Sulphur dioxide is most importantly used in the manufacture of sulphuric acid . 2. Sulphur dioxide and sodium sulphite are used as food preservatives , where they kill bacteria that make foods and drinks go bad. They are used in fruits, dried fruits, some meats and vegetables, sauces, fruit j juices, soft drinks and wines.

50. 3. Sulphur dioxide and sulphites are also bleaches , where uses include: Bleaching delicate materials such as wool and silk . Giving foods such as flour and some cheeses a white appearance. Paper making, where wood is converted into wood pulp and bleached to make white paper .

52. Questions: The following are names of some salts: iron(III) nitrate, copper(II) chloride, ammonium sulphate. Name the possible reactants that can react together to produce each of the given salts. For iron(III) nitrate: iron / iron(III) oxide / iron(III) hydroxide / iron(III) carbonate with dilute nitric acid. For copper(II) chloride: copper(II) oxide / copper(II) hydroxide / copper(II) carbonate with dilute hydrochloric acid. For ammonium sulphate: aqueous ammonia / ammonium carbonate with dilute sulphuric acid

53. (2) Solubility of Salts (Solubility Rules!) Salts Solubility (in water) Na + , K + , NH 4 + salts (sodium, potassium & ammonium salts) All soluble NO 3 - (nitrate salts) All soluble Cl - (chloride salts) All soluble except AgCl, PbCl 2 I - (iodide salts) All soluble except AgI, PbI 2 SO 4 2- (sulphate salts) All soluble except BaSO 4 , PbSO 4 , CaSO 4 CO 3 2- (carbonate salts) All insoluble except Na 2 CO 3 , K 2 CO 3 , (NH 4 ) 2 CO 3

54. (3) Preparation of Salts The method used to make/prepare a salt in the laboratory depends on the solubility of the salt in water . There are 3 steps in writing out the preparation of a salt. Step 1: Check solubility of the salt to be prepared Step 2: Check solubility of the parent acid and parent base to be used Step 3: Check solubility of parent carbonate and/or parent oxide

60. Chemical equation: Zn(s) + H 2 SO 4 (aq)  ZnSO 4 (aq) + H 2 (g) / ZnO(s) + H 2 SO 4 (aq)  ZnSO 4 (aq) + H 2 O(l) / ZnCO 3 (s) + H 2 SO 4 (aq)  ZnSO 4 (aq) + H 2 O(l) + CO 2 (g) 4. Filter the mixture to remove excess zinc powder / zinc oxide / zinc carbonate as the residue, and to obtain zinc sulphate solution as the filtrate. 5. Pour the zinc sulphate solution into an evaporating dish, and heat to evaporate some of the water to obtain a hot saturated solution. 6. Allow the hot saturated solution to cool, where crystals of zinc sulphate would form.

62. Example 3: Preparation of Na + /K + Salts Let’s say we want to prepare the salt crystals of sodium sulphate: 1. Sodium sulphate is a soluble salt (soluble in water), hence it is prepared by the reaction between dilute sulphuric acid and dilute sodium hydroxide . In this case, the titration method is required. Chemical equation: 2NaOH(aq) + H 2 SO 4 (aq)  Na 2 SO 4 (aq) + 2H 2 O(l) 2. Using a pipette, place a fixed volume, e.g. 25.0 cm 3 , of dilute sulphuric acid in a conical flask. Add a few drops of a suitable indicator to the acid. 3. Fill up a burette with dilute sodium hydroxide. Add the alkali gradually from the burette to the acid in the conical flask until the end-point is reached where the indicator changes colour.

63. 4. Measure this volume of alkali added from the burette to the acid. 5. Repeat the experiment with 25.0 cm 3 of dilute sulphuric acid placed in a conical flask, but without the indicator added. 6. From the burette, as before, add the same volume of dilute sodium hydroxide as measured previously to the acid. 7. A solution of sodium sulphate is thus produced without any excess acid or alkali present. 8. Pour the sodium sulphate chloride solution into an evaporating dish, and heat to evaporate the solution to dryness to obtain crystals of sodium sulphate. Note: To obtain crystals of certain salts, the solution of the salt has to be evaporated to dryness. Examples of such salts are sodium chloride and sodium sulphate. Note: If potassium chloride or some other salts are to be prepared, then step 8. would have to be replaced by steps 5. to 8. as in Example 2 above.

65. Exercise: For each of the following given salts, a. Write down the suitable starting materials/reactants for its preparation in the lab. b. Write the chemical equation (with state symbols) for each reaction used in making the salt. 1. Copper(II) chloride 2. Lead(II) chloride 3. Magnesium carbonate 4. Potassium nitrate 1. Copper(II) oxide / Copper(II) hydroxide / Copper(II) carbonate and dilute hydrochloric acid. CuO(s) + 2HCl(aq)  CuCl 2 (aq) + H 2 O(l) / Cu(OH) 2 (s) + 2HCl(aq)  CuCl 2 (aq) + 2H 2 O(l) / CuCO 3 (s) + 2HCl(aq)  CuCl 2 (aq) + H 2 O(l) + CO 2 (g)

69. Food preservatives Sodium nitrate Sodium sulphate Sodium citrate Photography Silver salts Sodium chloride in industry Chemicals Medical Uses Plaster of Paris (calcium sulphate) Food Flavouring Sodium chloride monosodium glutamate (MSG) Fertilizers in agriculture Ammonium Sulphate/Nitrate