This document provides information about acids and bases, including their properties and reactions. It defines acids as substances that produce hydrogen ions in aqueous solution, and bases as metal oxides or hydroxides. Strong acids are fully ionized in water, while weak acids are only partially ionized. The strength of an acid does not relate to its concentration. Common uses of acids include battery electrolytes, rust removal, and food preservation.

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

2. Learning Objectives Define acids and alkalis in terms of the ions they produce in aqueous solution and their effects on indicators Describe the characteristic properties of acids in reactions with metals, bases, and carbonates State some uses of acids including sulphuric acid Describe the characteristic properties of bases in reactions with acids, metal ions and ammonium compounds Describe the difference between strong and weak acid and alkalis in terms of the extent of ionisation Describe the pH scale as a measure of relative acidity and alkalinity

3. Acids (1) Introduction Some common acids in our daily lives! Ethanoic acid, CH 3 COOH – found in vinegar, tomato juice Citric acid, C 6 H 8 O 7 – found in citrus fruits, e.g. oranges, lemons, grapefruit Lactic acid – found in sour milk, yoghurt Tartaric acid – found in grapes Tannic acid – found in tea What about those commonly found in our lab? Hydrochloric acid, HCl Sulphuric acid, H 2 SO 4 Nitric acid, HNO 3 Which of these acids is used in car batteries? Sulphuric Acid!

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

5. (2) Properties Of Dilute Acids 1. Most acids have a sour taste . 2. Acids are hazardous and dilute acids are irritants . 3. Acids change the colour of indicators, e.g. acids turn blue litmus red . 4. Most dilute acids react with many metals to produce hydrogen gas. e.g. Mg(s) + 2HCl(aq)  MgCl 2 (aq) + H 2 (g) We can show that hydrogen is produced by testing the gas with a lighted/burning splint . Hydrogen extinguishes a lighted splint with a ‘pop’ sound .

6. 5. Acids react with most bases (metal oxides and hydroxides) to form salt and water e.g. FeO(s) + H 2 SO 4 (aq)  FeSO 4 (aq) + H 2 O(l) 6. Acids react with carbonates (and hydrogencarbonates) to produce carbon dioxide gas. e.g. CaCO 3 (s) + 2HCl(aq)  CaCl 2 (aq) + H 2 O(l) + CO 2 (g) e.g. NaHCO 3 (s)/(aq) + HCl(aq)  NaCl(aq) + H 2 O(l) + CO 2 (g) We can show that carbon dioxide is produced by bubbling the gas through limewater . The formation of a white precipitate with limewater indicates the presence

7. Writing Ionic Equations What is an ionic equation?? An ionic equation is an equation involving ions in aqueous solution; only the ions that are involved in the reaction are included. In ionic equations: 1. Formulae of ions that are involved in the reaction are included, while formulae of ions that are not involved in the reaction are omitted. 2. Formulae of solids, liquids and gases are written in full.

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 )

12. Exercise Briefly describe what you would observe when the following substances are mixed. a) A piece of iron and lemon juice. Effervescence of a colourless and odourless gas (hydrogen gas). b) Baking soda (sodium hydrogencarbonate) and vinegar. Effervescence of a colourless and odourless gas (carbon dioxide). 2. Write balanced chemical and ionic equations (include state symbols) for the reactions between the following: a) zinc and dilute sulphuric acid b) aqueous sodium hydroxide and dilute hydrochloric acid c) copper(II) oxide and dilute nitric acid d) potassium carbonate solution and dilute sulphuric acid

13. a) Zn(s) + H 2 SO 4 (aq)  ZnSO 4 (aq) + H 2 (g) zinc sulphate Zn(s) + 2H + (aq)  Zn 2+ (aq) + H 2 (g) b) CuO(s) + 2HNO 3 (aq)  Cu(NO 3 ) 2 (aq) + H 2 O(l) calcium nitrate CuO(s) + 2H + (aq)  Cu 2+ (aq) + H 2 O(l) c) NaOH(aq) + HCl(aq)  NaCl(aq) + H 2 O(l) sodium chloride H + (aq) + OH - (aq)  H 2 O(l) d) K 2 CO 3 (aq) + H 2 SO 4 (aq)  K 2 SO 4 (aq) + H 2 O(l) + CO 2 (g) potassium sulfate CO 3 2- (aq) + 2H + (aq)  H 2 O(l) + CO 2 (g)

14. (3) Acids and Hydrogen ions Take a look at the below experiments Experiment A Experiment B 1. Litmus paper changes colour (to red ) only in experiment B 2. Bubbles of hydrogen gas are produced only in experiment B WHY? Pure acids (without water) consist of simple covalent molecules . However, in the presence of water , the acid ionises to form ions. e.g. when hydrogen chloride gas is dissolved in water, ionisation occurs and hydrogen ions, H + (aq) are produced: HCl(g)  H + (aq) + Cl - (aq)

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.

17. 2. A drink tablet contains a mixture of a solid acid (often citric acid) and solid sodium hydrogencarbonate. When the tablet is added to water, the drink fizzes. Explain why. When the drink tablet is added to water, the acid ionises to produce hydrogen ions, H + (aq). The hydrogen ions react with the sodium hydrogencarbonate to produce carbon dioxide gas, which causes the fizz.

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

19. (5) Strength of Acids The strength of an acid depends on the degree of ionisation dissociation of the acid in aqueous solution . A strong acid is completely ionised in water/aqueous solution. In a solution of strong acid, all acid molecules are ionised to form ions in water. E.g. a solution of hydrochloric acid only contains hydrogen ions and chloride ions and no HCl molecules, as all the HCl molecules have ionised in water. HCl(aq)  H + (aq) + Cl - (aq) H 2 SO 4 (aq)  2H + (aq) + SO 4 2- (aq) HNO 3 (aq)  H + (aq) + NO 3 - (aq)

20. A weak acid is partially ionised in water/aqueous solution. In a solution of weak acid, most of the acid molecules remain as covalent molecules in water, only few acid molecules are ionised to form hydrogen ions. E.g. in a solution of 1M ethanoic acid , only 4 out of 1000 CH 3 COOH molecules ionise to produce hydrogen ions, the remaining 99% CH 3 COOH molecules remain unchanged in water. Examples of weak acids and equations for their ionisation in aqueous solution: CH 3 COOH(aq)  H + (aq) + CH 3 COO - (aq) H 2 CO 3 (aq)  2H + (aq) + CO 3 2- (aq) A solution of strong acid, e.g. hydrochloric acid, has three kinds of particles: hydrogen ions , chloride ions and water molecules . A solution of weak acid, e.g. ethanoic acid, has four kinds of particles: hydrogen ions , ethanoate ions , ethanoic acid molecules and water molecules .

21. Take a look at the difference between strong and weak acids below: Strong acids react more vigorously than weak acids. e.g. if strips of magnesium ribbon are added to solutions of hydrochloric acid and ethanoic acid (in separate experiments), effervescence of hydrogen gas occurs more rapidly with hydrochloric acid than with ethanoic acid.

22. Question : Does a strong acid means it is concentrated? Note that the strength of acids does not relate to their concentration! The terms “strong” and “weak” refer to the extent of ionisation of the acid in water. The terms “concentrated” and “dilute” refer to the amount of solute in the solution. a) 10.0 moldm -3 solution of hydrochloric acid = concentrated solution of a strong acid b) 0.100 moldm -3 solution of hydrochloric acid = dilute solution of a strong acid c) 10.0 moldm -3 solution of ethanoic acid = concentrated solution of a weak acid d) 0.100 moldm -3 solution of ethanoic acid = dilute solution of a weak acid

23. Question: a) Is H 2 SO 4 (l) a good conduct of electricity? b) Is H 2 SO 4 (aq) a good conduct of electricity? a) No. Pure sulphuric acid (no water) consist of small covalent molecules, hence are poor electrical conductors. b) Yes! (acids in water act as strong electrolyte ) Why? H 2 SO 4 (aq)  2H + (aq) + SO 4 2- (aq) Acids in water contain moving ions that could carry the electric current!

24. Question: How do strength and concentration affect conduction of electricity? Explain why dilute sulphuric acid conducts electricity better than concentrated sulphuric acid. In dilute sulphuric acid, the acid ionizes completely to produce a high concentration of ions which move freely to conduct electricity. In concentrated sulphuric acid, due to small amount of water present, the acid is only partially ionized (although it is still a strong acid), thus there are much less number of free-moving ions present to conduct electricity.

25. (6) Uses of Acids 1. Acids are used to make many useful products. The most important industrial acid is sulphuric acid , which is mainly used in the manufacture of agricultural fertilisers . e.g. sulphuric acid reacts with ammonia in making the fertiliser ammonium sulphate. Sulphuric acid is also used in the manufacture of detergents, paints, dyes, artificial fibres and plastics . Dilute sulphuric acid is used in batteries for road vehicles such as cars and buses. 2. Acids such as hydrochloric acid and sulphuric acid are also used in the removal of rust from iron and steel objects. Acids remove rust from such objects by reacting and ‘dissolving’ the iron(III) oxide in the rust .

26. 3. Some acids are used in food preservation . Examples are: Ethanoic acid (in vinegar) – used to preserve vegetables Benzoic acid – used in fruit juices, jams and oyster sauce Citric acid – used in foods as both a preservative and flavouring

27. Bases & Alkalis (1 ) Introduction Bases are the oxides or hydroxides of metals , e.g. copper(II) oxide, sodium hydroxide. Soluble bases are called alkalis , thus an alkali is a base which is soluble in water. Thus, all alkalis are bases , but not all bases are alkalis . Alkalis, like acids, are common in our daily lives. For example, they are found in soaps, detergents, window cleaners, floor cleaners, oven cleaners, toothpastes and medicine for indigestion. Alkalis are also used in the school laboratory. Three such common alkalis are: Sodium hydroxide, NaOH Aqueous ammonia/Ammonia solution, NH 3 (aq) Calcium hydroxide (Limewater), Ca(OH) 2 Insoluble Bases Alkalis Bases

28. (2) Properties Of Dilute Alkalis 1. Dilute solutions of alkalis have a ‘slippery’ feel 2. Alkalis are hazardous and like dilute acids, dilute solutions of sodium hydroxide and potassium hydroxide are irritants . 3. Concentrated solutions of sodium hydroxide and potassium hydroxide are corrosive and burn skin. 4. Alkalis change the colour of indicators, e.g. alkalis turn red litmus blue . 5. Alkalis react with acids to form a salt and water , in a reaction called neutralisation . 6. Alkalis react with (solid) ammonium salts on warming to produce ammonia gas. NH 4 Cl(s) + NaOH(aq)  NaCl(aq) + H 2 O(l) + NH 3 (g)

29. We can show that ammonia is produced by testing the gas evolved with damp red litmus paper . The damp red litmus paper would turn blue , indicating the presence of ammonia. 7. Alkalis react with most aqueous metal salts to form insoluble precipitates of metal hydroxides . CuSO 4 (aq) + 2NaOH(aq)  Cu(OH) 2 (s) + Na 2 SO 4 (aq) Exercise: Write balanced chemical and ionic equations (include state symbols) for the reactions between the following: a) aqueous calcium hydroxide and dilute nitric acid b) ammonium sulphate solution and aqueous potassium hydroxide c) iron(III) sulphate solution and aqueous sodium hydroxide d) lead(II) nitrate solution and aqueous ammonia Blue PPT

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.

32. (4) Strength of Alkalis Like for acids, the strength of an alkali depends on the degree of ionisation /dissociation of the alkali in aqueous solution . A strong alkali is completely ionised in water/aqueous solution. Examples of strong alkalis and equations for their ionisation in aqueous solution: NaOH(aq)  Na + (aq) + OH - (aq) KOH(aq)  K + (aq) + OH - (aq) A weak alkali is partially ionised in water/aqueous solution. The most common example of a weak alkali is aqueous ammonia ammonia solution . When ammonia gas dissolves in water, most of the ammonia molecules remain as covalent molecules , only a small amount of ammonia molecules are ionised to form hydroxide ions:

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 .

35. (7) Applications of Neutralisation Neutralisation is the reaction between an acid and a base to form a salt and water only. Neutralisation is important as it has many applications in daily life. 1. Controlling the pH of soil – as in point 1. under ‘ Uses of Alkalis ’ section. However, if soil is too alkaline , compost which consists of rotting plant material such as leaves & vegetables can be added. The plants decompose to give off carbon dioxide which dissolves in water to form carbonic acid that reduces alkalinity of the soil.

36. 2. Treatment of indigestion – as in point 2. under ‘ Uses of Alkalis ’ section. 3. In toothpastes – as in point 3. under ‘ Uses of Alkalis ’ section 4. Treatment of insect stings Bee stings are acidic and can be neutralised by applying baking soda (solution or paste) which contains alkaline sodium hydrogencarbonate to the affected area. Wasp stings are alkaline and can be neutralised by applying vinegar or lemon juice to the affected area. Bee Sting – Acidic! Wasp Sting – Alkaline! 5. Treatment of industrial wastewater Wastewater from many industrial processes can be highly acidic or alkaline , e.g. wastewater from dyeing factories can have a pH of 10 to 12. In the past, such wastewater went straight into rivers or the sea where it caused environmental damages such as killing of marine life and corrosion of metal pipes

37. Indicators & pH (1) Indicators and pH scale The pH value of a solution indicates the degree of acidity or alkalinity of the substance. The pH scale is numbered between 0 and 14 . On this scale, a pH of 7 is neutral (this is the pH of pure water), a solution with pH less than 7 is acidic, and a solution with pH greater than 7 is alkaline. 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 the pH scale Acidity increases Neutral Alkalinity increases For an acidic solution: the smaller the pH , the more acidic the solution is, the more hydrogen ions it contains. For an alkaline solution: the greater the pH , the more alkaline the solution is, the more hydroxide ions it contains.

38. An indicator is a substance that has different colours in acidic and alkaline solutions. Simple indicators tell whether a solution is acidic or alkaline , e.g. litmus, methyl orange, phenolphthalein. The colours of these three indicators in solutions of acids and alkalis are as shown: In addition, different indicators change colours over different ranges of pH values: Indicator Colour in Acids Colour in Alkalis Litmus Red Blue Methyl orange Red Yellow Phenolphthalein Colourless Pink Common Indicators Colour of Indicator at different pH values Litmus red blue 0 5.0 8.0 14 Methyl orange red yellow 0 3.1 4.4 14 Phenolphthalein colourless pink 0 8.3 10 14

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 .

40. 2. A pH meter is an electrical method of measuring the pH of a solution. It consists of a pH electrode or probe connected to a meter . The probe is dipped into the solution and the meter shows the pH value either on a scale or digitally. It is much more reliable and accurate than the Universal Indicator, though the latter is often more convenient. 3. A pH sensor connected to a computer through an interface can be used to measure the pH of a solution. The pH reading is displayed on the computer screen. The pH values of some common substances are shown below:

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 .

42. 3. pH in the garden The pH of soil is important for optimal plant growth. Most plants grow best when the soil has a pH of about 5.6 , that is, soil that is weakly acidic. However, some plants grow well in more acidic soil or more alkaline soil. Examples of pH of soil in which plants grow best: a. Orchids – more acidic soil (pH 4 – 5) b. Azaleas, strawberries, apples, potatoes – weakly acidic soil (pH 5 – 6) c. Common vegetables, e.g. beans, lettuce, onion – about neutral soil (pH 6 – 8) d. Water lilies – more alkaline soil (pH 8 – 9)

43. 4. pH and hair Our normal hair is weakly acidic with a pH of about 5 where hair is smooth, strong and healthy. In hair perming , hair is treated with alkaline perm solutions which make hair dull, weaker, easier to break and even damaged. Rinsing the hair with water or weakly acidic solutions reverses these changes and returns the hair to its normal state. In cleaning our hair, most shampoos used are alkaline so as to dissolve grease present in the hair, but alkalis can damage hair. Rinsing the hair with hair conditioners which contain weak natural acids (such as citric acid or tartaric acid) neutralise any excess alkali left in the hair after shampooing and restores the pH of hair to its normal value.

44. Oxides (1) Types of Oxides An Oxide is a compound of oxygen and another element. Most oxides can be classified into four types: acidic , basic , neutral and amphoteric . Acidic and neutral oxides are oxides of non-metals . Basic and amphoteric oxides are oxides of metals . Iron Ores Copper (II) Oxide Lead (II) Oxide

45. Oxides Metallic Oxides Non-metallic oxides Basic oxides Amphoteric oxides Neutral oxides Acidic oxides Li 2 O Al 2 O 3 H 2 O SO 2 Na 2 O ZnO CO SO 3 MgO PbO NO NO 2 CaO CO 2 FeO P 4 O 10 CuO SiO 2 etc…..

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 .

51. Salts (1 ) Introduction A salt is produced when an acid reacts with a base . The salt consists of two parts; one part comes from the base, the other from the base part NaOH + HCl  NaCl + H 2 O acid part Thus, a salt is a substance formed when the hydrogen ions of an acid are completely or partly replaced by a metal ion or an ammonium ion . Sodium Chloride “ Common Salt”

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

55. Is Salt Soluble? Precipitation Method React: 2 soluble salts, or Soluble salt + acid E.g. NaCI + Ag NO 3  AgCI + Na NO 3 React Acid + *Metal, or -Acid + insoluble base, or -Acid + insoluble carbonate e.g. CuO + H 2 SO 4  CuSO 4 + H 2 O Titration Method React: Acid + Alkalis Using a Suitable Indicator E.g. KCl + H 2 O Filter off salt prepared (residue) Wash and Dry salt Add **excess metal/base/carbonate to hot acid Filter off excess metal/base/carbonate (residue) Salt solution remains behind (filtrate) Do crystallisation if salts crystals are desired Filter off salt crystals prepared Dry the salt crystals NO YES, and is Na + /K + ,/NH 4 + salts YES Given Salt to be prepared

56. *Metals such as Zn and Mg react moderately with dilute acids, thus they are suitable for this method. metals that are too unreactive (e.g. Cu) would not react with dilute acids at all, metals that are too reactive (e.g. Na, K) would react too vigorously and explosively with dilute acids. * *Excess amount of metal/base/carbonate must be used to use up all the acid , where reaction is then complete; otherwise, any excess acid remaining after the reaction would contaminate the salt solution produced in the reaction. We will look at some examples on the preparation of soluble/insoluble salts

57. Example 1: Preparation of Insoluble Salts Let’s say we want to prepare the salt barium sulphate: 1. Barium sulphate is an insoluble salt (insoluble in water), hence it is prepared by the precipitation method. 2. Mix aqueous barium(II) nitrate and aqueous sodium sulphate for reaction. / Mix aqueous barium(II) nitrate and dilute sulphuric acid for reaction. A white precipitate of barium sulphate would be produced. Chemical equation: Ba(NO 3 ) 2 (aq) + Na 2 SO 4 (aq)  BaSO 4 (s) + 2NaNO 3 (aq) / Ba(NO 3 ) 2 (aq) + H 2 SO 4 (aq)  BaSO 4 (s) + 2HNO 3 (aq) Barium Sulphate (white ppt!)

58. 3. Filter the mixture to obtain lead(II) iodide (the yellow precipitate) as the residue. 4. Wash the lead(II) iodide (the residue) with distilled water. 5. Dry the lead(II) iodide between pieces of filter paper.

59. Example 2: Preparation of soluble Salts Let’s say we want to prepare the salt crystals of zinc sulphate: 1. Zinc sulphate is a soluble salt (soluble in water), hence it is prepared by the reaction between dilute sulphuric acid and zinc / zinc oxide / zinc carbonate . 2. Add zinc powder / zinc oxide / zinc carbonate with stirring, to hot dilute sulphuric acid for reaction. 3. Continue adding until some of the zinc powder / zinc oxide / zinc carbonate no longer reacts with the acid. The zinc powder / zinc oxide / zinc carbonate is now in excess, and the acid is used up.

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.

61. 7. Filter this mixture to obtain zinc sulphate crystals as the residue. 8. Dry the zinc sulphate crystals between pieces of filter paper.

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)

66. 2. Aqueous lead(II) nitrate and aqueous sodium chloride / aqueous potassium chloride / dilute hydrochloric acid. Pb(NO 3 ) 2 (aq) + 2NaCl(aq)  PbCl 2 (s) + 2NaNO 3 (aq) / Pb(NO 3 ) 2 (aq) + 2HCI(aq)  PbCl 2 (s) + 2HNO 3 (aq) 3. Aqueous magnesium nitrate and aqueous sodium carbonate /aqueous potassium carbonate. Mg(NO 3 ) 2 (aq) + Na 2 CO 3 (aq)  MgCO 3 (s) +2NaNO 3 (aq) 4. Potassium oxide / Aqueous potassium hydroxide / Aqueous potassium carbonate and dilute nitric acid. K 2 O(s) + 2HNO 3 (aq)  2KNO 3 (aq) + H 2 O(l) / KOH(aq) + HNO 3 (aq)  KNO 3 (aq) + H 2 O(l) / K 2 CO 3 (aq) + 2HNO 3 (aq)  2KNO 3 (aq) + H 2 O(l) + CO 2 (g)

67. Exercise: A student attempted to prepare lead(II) sulphate in the laboratory by reacting lead(II) oxide and dilute sulphuric acid. Would he succeed in the preparation? Explain why. No, he would not succeed. This is because once the product, lead(II) sulphate is formed during the reaction, being an insoluble salt (insoluble in water), it would coat onto the unreacted lead(II) oxide, thus inhibiting further reaction between lead(II) oxide and dilute sulphuric acid. [in fact, since lead(II) sulphate is an insoluble salt, the precipitation method should be used => e.g. using aqueous lead(II) nitrate and aqueous sodium sulphate as reactants.]

68. (4) Uses of Salts Salts are important to our bodies. Salt is lost from our bodies through sweat and urine , where it must be replaced through food and drink . Some uses of salts in society: 1. Salts such as ammonium sulphate and ammonium nitrate are used as fertilisers in agriculture. 2. Salts such as sodium sulphite , sodium nitrite and sodium citrate are used as food preservatives . 3. Salts such as sodium chloride and monosodium glutamate (MSG) are used as food flavourings . 4. The salt calcium sulphate has medicinal uses, e.g. as ‘ plaster of Paris’ . 5. The salt sodium chloride is used in many chemicals in the industry . 6. Salts such as silver salts are used in photography

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