ACIDS AND ALKALIS
Acids dissolve in water and the resulting solution contains hydrogen ions, H+(aq). Alkalis dissolve in water and the resulting solution contains hydroxide ions, OH-(aq) When a substance dissolves in water it forms an aqueous solution which may be acidic, alkaline or neutral. Water itself is neutral. Some metal oxides and hydroxides (e.g. the oxides and hydroxides of sodium, potassium and, to some extent, calcium) dissolve in water to produce alkaline solutions. Soluble oxides of non-metals (e.g. carbon dioxide, sulphur dioxide and nitrogen dioxide) produce acidic solutions. (a) pH scale
The pH scale is used to show how acidic or alkaline a solution is. 0 7 14 increasing neutral increasing acidity alkalinity Indicators can be used to show whether a solution is acidic, alkaline or neutral by the way their colours change. Most indicators show one colour at a lower pH and a second colour at a higher pH. e.g. methyl orange phenolphthalein red in acid colourless in acid yellow in alkali pink in alkali
Universal indicator (which is a mixture of several indicators) shows a range of colours (see below) and gives a more precise pH value than single indicators . Colour pH ACIDIC Red 0-2 Orange 3-4 Yellow 5-6 Green 7 NEUTRAL Blue 8-9 Navy blue Purple 10-12 13-14 ALKALINE
Universal indicator should not be used in titrations (see below).
(b) Acids The common reactions of dilute acids are: to change the colours of indicators, e.g. turn blue litmus red; with moderately reactive metals they form hydrogen and a solution of the salt. e.g. Mg(s) + H2SO4(aq) MgSO4(aq) + H2(g) (iii) with bases such as oxides and hydroxides they form salts and water only. (neutralisation) e.g. CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l) NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l) (iv) with carbonates and hydrogencarbonates they form carbon dioxide, water and a salt. (this is used as a test for carbonates) e.g. ZnCO3(s) + 2HNO3(aq) Zn(NO3)2(aq) + H2O(l) + CO2(g) NaHCO3(s) HCl(aq) NaCl(aq) + H2O(l) + CO2(g) In these reactions the acid produces H+(aq) ions in water. An acid is a proton donor. A base is a substance which neutralises an acid. A base is a proton acceptor. (c) Alkalis All metal oxides and hydroxides are known as bases. Most metal oxides and hydroxides are insoluble in water but can neutralise acids to produce salts and water. Alkalis are soluble bases which produce OH (aq) ions in water, e.g. sodium hydroxide, potassium hydroxide and ammonia solution. Alkalis: (i) turn litmus paper blue. (ii) neutralise acids; in general acid + alkali salt + water The reaction of any acid with any alkali (i.e. neutralisation) can be described by the following ionic equation: H+(aq) + OH (aq) H2O(l) from acid from alkali (iii) react with metal salts in solution to form precipitates of the metal hydroxide. This can be used to identify metal ions in solution, e.g. Cu2+(aq) + blue solution 2OH(aq) colourless(aq) Cu(OH)2(s) blue precipitate (i) (ii)
(d) Strength of acids and alkalis Acids and alkalis are classified by the extent of their ionisation in water. A strong acid or alkali is one that is 100% ionised in water. Examples of strong acids are hydrochloric acid, sulphuric acid and nitric acid. Examples of strong alkalis are sodium hydroxide and potassium hydroxide. A weak acid or alkali is only partly ionised in water. Examples of weak acids are ethanoic acid, citric acid and carbonic acid. Examples of weak alkalis are sodium hydrogencarbonate and ammonia solutions. Candidates should be able to describe how to distinguish between strong and weak acids of the same concentration by using the pH scale or the rate of reaction with metals. The pH scale helps to distinguish strong and weak acids and alkalis. strong acids have pH of about 0 - 3 weak acids have pH of about 4 - 6 neutral solutions or pure water have pH of 7 weak alkalis have pH of about 8 - 10 strong alkalis have pH of about 11 - 14 As strong acids have a much higher amount of H+(aq) ions than weak acids of the same concentration. They react much faster with metals such as magnesium. Hence the rate of production of hydrogen can also be used to distinguish between the two types of acid. (e) Salts
There are several general methods of producing salts: reaction of a metal with an acid; reaction of an insoluble base with an acid; reaction of a soluble base with an acid; by mixing two solutions to form an insoluble salt (precipitation) by direct combination of the elements. Candidates should be able to give practical details of salt preparations based on each of these general methods.
Salts are produced in neutralisation reactions.
The particular salt made depends on the acid used and the metal ion in the alkali or base or carbonate. neutralising hydrochloric acid (HCl) produces chlorides, neutralising nitric acid (HNO3) produces nitrates, neutralising sulphuric acid (H2SO4) produces sulphates.
Solubilities of salts The method used to prepare a particular salt depends on the solubilities of the reactants and products. SOLUBLE All sodium, potassium and ammonium salts All nitrates Most chlorides, bromides and iodides Silver and lead chlorides, bromides and iodides Lead sulphate, barium sulphate, calcium sulphate (slightly) Most other carbonates Most other hydroxides INSOLUBLE
Most sulphates Sodium, potassium and ammonium carbonates Sodium and potassium hydroxides Ammonia solution (‘ammonium hydroxide’) Calcium hydroxide is slightly soluble
There are four general methods of producing salts. You should be able to give practical details of salt preparations based on each of these four general methods. Soluble salts method 1
reaction of a dilute acid with excess insoluble metal; reaction of a dilute acid with excess insoluble base (metal oxide or hydroxide); reaction of a dilute acid with excess insoluble carbonate; method 2 reaction of a soluble base with an acid (titration);
Insoluble salts method 3
precipitation, i.e. by mixing two solutions to form an insoluble salt.
Anhydrous salts method 4 by direct combination of the elements. Method 1 starting with insoluble metal or insoluble carbonate e.g. Mg(s) or + H2SO4(aq) 2HNO3(aq) MgSO4(aq) + H2(g) H2O(l) +
ZnCO3(s) + CO2(g)
Place some acid in a beaker. Continue adding insoluble metal (or carbonate) to the acid until the acid is all used up, i.e. effervescence of hydrogen (or carbon dioxide) stops. Filter off the unreacted metal (or carbonate). Evaporate some water from the solution to form a saturated solution. Leave to cool to allow crystals to form. Filter off the crystals, wash with a little cold distilled water. Dry the crystals using filter paper. Method 1 starting with insoluble base e.g. CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l) Place some acid in a beaker and warm it. Continue adding insoluble base to the acid until the acid is all used up and no more base reacts. (You could test the solution by removing a drop on a glass rod and checking with litmus that it was no longer acidic). then evaporate the solution and continue as above Method 2 Titration starting with soluble hydroxide or soluble carbonate NaOH(aq) K2CO3(s) + + HCl(aq) 2HCl(aq) NaCl(aq) 2KCl(aq) + + H2O(l) H2O(l) + CO2(g)
Rinse and fill a burette with dilute acid. Record the volume of acid present. Rinse a pipette and use it to put 25 cm3 of alkali into a conical flask. Add a few drops of indicator, phenolphthalein or methyl orange. Add acid from the burette until the indicator just changes colour. Record the volume of acid used. Repeat the titration using the same volumes of acid and alkali but with no indicator. Evaporate some water from the solution to form a saturated solution. Leave to cool to allow crystals to form. Filter off the crystals, wash with a little cold distilled water. Dry the crystals using filter paper.
Method 3 Precipitation starting with soluble reagents. e.g. AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq) Mix two solutions, one containing the required positive ion and one containing the required negative ion. Filter off the precipitate. Wash with distilled water. Dry by gentle warming.
Method 4 Direct synthesis of anhydrous salts from elements Some anhydrous salts are made by reacting the elements directly: (a) Anhydrous aluminium chloride 2Al (s) + 3Cl2 (g) 2Al Cl3 (s) Pass dry chlorine over heated aluminium. The aluminium chloride sublimes from the surface of the metal to a cool place further along the apparatus. This method will only work for salts which sublime, otherwise a coating of ionic salt forms on the surface of the metal and prevents further reaction. (b) Zinc sulphide Heat a mixture of powdered zinc and powdered sulphur. A very exothermic reaction occurs forming a white powder. Zn (s) + S (s) ZnS (s) (f) Precipitation Reactions These occur when solutions containing ions are mixed; positive ions from one solution combine with negative ions from the other to form an insoluble product which appears as a precipitate. e.g. AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)
This may be represented as an ionic equation, leaving out the spectator ions. i.e. leaving out the sodium and nitrate ions the equation becomes Ag+(aq) + Cl(aq) AgCl(s)
Precipitation reactions can be used: (i) to produce new materials such as salts as shown above. (ii) to remove unwanted ions from solution e.g. the softening of water by the addition of sodium carbonate to water containing Ca2+(aq) or Mg2+(aq) ions e.g. Ca2+(aq) + CO32(aq) CaCO3(s)
(iii) to identify some ions in solution: Chloride ions can be identified by adding silver nitrate solution and dilute nitric acid. If a chloride is present a white precipitate of silver chloride forms. (The nitric acid prevents any other precipitates forming.) Ag+(aq) + Cl(aq) AgCl(s)
Sulphate ions can be identified by adding barium chloride (or nitrate) solution and dilute hydrochloric (or nitric) acid. If a sulphate is present a white precipitate of barium sulphate forms. (The acid prevents any other precipitates forming.) Ba2+(aq) + SO42(aq) BaSO4(s)
Metal ions can be identified in solution by the colour of the hydroxide precipitate formed when aqueous sodium hydroxide is added:
iron (II) ions, Fe2+ give a green, gelatinous (jelly-like) precipitate of iron(II) hydroxide Fe2+(aq) + 2 OH(aq) Fe(OH)2(s)
iron (III) ions, Fe3+ give a brown, gelatinous precipitate of iron(III) hydroxide Fe3+(aq) + 3 OH(aq) Fe(OH)3(s)
copper (II) ions, Cu2+ give a pale blue, gelatinous precipitate of copper(II) hydroxide Cu2+(aq) + 2OH(aq) Cu(OH)2(s)
The same reactions also occur if ammonia solution is used, except that when copper compounds are used, the initial pale blue precipitate of copper(II) hydroxide dissolves when more (excess) ammonia solution is added and a deep blue solution is formed. Common mistakes Students often fail to understand the difference between the words (strong / weak) and (concentrated / dilute) Students fail to put in enough detail when answering questions regarding salt preparation.
Practice questions 1. The following data sheet should help you answer the questions that follow.
SOLUBILITY OF SALTS AND HYDROXIDES IN WATER AT ROOM TEMPERATURE SOLUBLE INSOLUBLE
All sodium, potassium and ammonium salts All nitrates Most chlorides, bromides and iodides Most sulphates Sodium potassium and ammonium carbonates. Sodium and potassium hydroxides. Ammonia solution (‘ammonium hydroxide’) Calcium hydroxide is slightly soluble a) b) c)
Silver and lead chlorides, bromides and iodides lead sulphate, barium sulphate, (calcium sulphate is slightly soluble) Most other carbonates. Most other hydroxides
d) e) f) g)
Using the data sheet provided, list the names of four salts included in the table which are insoluble in water: (1 mark) Choose one of these insoluble salts to be used as your example. Give its chemical formula: (1 mark) To make it you will need to add together two solutions one containing the positive ion and the other containing the negative ion. In the example of an insoluble salt you have chosen in b) write the formula of the positive ion it contains: (1 mark) Using the data sheet, name a compound containing this positive ion which is soluble in water: (1 mark) In the example of an insoluble salt you have chosen in b) write the formula of the negative ion it contains: (1 mark) Using the data sheet, name a compound containing this negative ion which is soluble in water: (1 mark) If you wished to have a pure dry sample of this salt; (i) What would you do to obtain the salt from the mixture in the beaker? (ii) Why would the salt collected not yet be pure ? (iii) What would you do to make it pure ? (iv) How would you dry the sample ? (4 marks)
Ethanoic (acetic) acid is present in commercially available vinegar. It is a weak acid. Sulphuric acid is used in car batteries. It is a strong acid. a) Give an example of an indicator that you could use to distinguish between a sample of dilute ethanoic acid and a sample of dilute sulphuric acid. You should include any observations that you might expect to be able to make. (3 marks) b) Write the formula for each of the acids mentioned. (2 marks) c) Define the following terms (i) acid (ii) strong acid (iii) weak acid (2 marks) (iv) How could you neutralise any of the acids mentioned? (1 mark)
Model answers 1. a) Silver chloride, lead chloride, silver bromide, lead bromide, silver iodide, lead iodide, lead sulphate, barium sulphate, calcium sulphate (4 for 1 mark) AgCl (1 mark) Ag+ (1 mark) Silver nitrate (1 mark) Cl- (1 mark) Sodium Chloride (1 mark) (i) Filter the suspension (1) (ii) Solid particles still wet with sodium nitrate solution (1) (iii) Wash the solid with distilled water (1) (iv) Either place in a warm oven or Leave overnight (1) Add a few drops of universal indicator to samples of each of the acids.(1) Ethanoic acid would turn the indicator an orange colour.(1) Sulphuric acid would turn the indicator a red colour.(1) Ethanoic acid CH3COOH Sulphuric Acid H2SO4 (i) An acid is a proton donor.(1) or acids produce H+ ions in solution (ii) A strong acid is fully ionised (dissociated) in solution.(1) (iii) A weak acid is partially ionised (dissociated) in solution.(1) (iv) Any one of the following (1 mark) add an alkali (or base) add a metal oxide add a metal carbonate add a metal hydroxide.
b) c) d) e) f) g)