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Chemguide: Support for CIE A level Chemistry Learning outcome 7.2 Equilibria The Bronsted-Lowry theory of acids and bases I am treating statements 7.2.1 to 7.2.8 together because a lot of what they contain is fairly trivial. Statements 7.2.1 and 7.2.2 Read the statements - there is nothing to add! Statement 7.2.3 This statement is about the Bronsted-Lowry theory of acids and bases. Before you go on, you should find and read the statement in your copy of the syllabus. You will find more than you need on the page about theories of acids and bases. Read the bit about the Arrhenius theory as background, although it isn't mentioned in the syllabus. Then concentrate on the section about the Bronsted-Lowry theory. Read down to, but not including, the section on conjugate pairs. You won't need the rest of this page until the second half of the course. Statement 7.2.4 This statement is an introduction to strong and weak acids and bases. Before you go on, you should find and read the statement in your copy of the syllabus. You need to start by reading bits of the page about strong and weak acids. This page covers work that you will need in the second half of the course as well as the first. You should ignore all the numerical parts of the page. Note that when the syllabus talks about acids in solution, they use the symbol H+(aq) and not the version H3O+. That is what you should use in an exam unless a question specifically asks about the H3O+ ion. Read down as far as the first paragraph under the heading "Strong acids and pH". For now, all you need to know is that strong acids of the sort of concentrations that we normally use in the lab have pHs of around 0 - 1. For example, 1 mol dm-3 hydrochloric acid has a pH of 0. (More concentrated solutions of hydrochloric acid will have slightly negative pHs.) Ignore the sections headed "Defining pH" and "Working out the pH of a strong acid", and go on to the section on weak acids. In the weak acid section, just read the introductory bit about what a weak acid is. Weak acids of the sort of concentrations we use in the lab, will typically have pHs in the 2 - 6 range. For example, 1 mol dm-3 ethanoic acid has a pH of 2.4. Now do much the same thing on the page about strong and weak bases. In the strong base section, just read the introductory bit about what a strong base is, and then jump down to the similar introductory bit about weak bases. Ignore everything else on the page. Typically, a strong base like sodium hydroxide solution will have a pH of 14 for a 1 mol dm-3 solution. Ammonia solution, a weak base, has a pH of 11.6 for a solution of the same concentration. Statement 7.2.5 Just read the statement. There isn't a lot to add to it. As you have seen above, solutions of strong acids have lower pHs than weak acids, and solutions of strong bases have higher pHs than weak bases. A solution with a pH less than 7 is described as acidic; one with a pH above 7 is described as alkaline. An alkali is simply a soluble base. Not all bases dissolve in water. Fot example, most metal oxides and hydroxides are insoluble in water, and so aren't alkalis. But the oxide or hydroxide ions present in the solid do react with acids by accepting hydrogen ions, and so they are definitely bases - just not soluble ones. Statement 7.2.6 This statement is about the practical differences between strong and weak acids. Reactions with reactive metals A reactive metal like magnesium reacts more slowly with a weak acid than it does with a strong acid, but both produce hydrogen gas and a salt. With hydrochloric acid, a strong acid, magnesium produces magnesium chloride and hydrogen. Mg(s) + 2HCl(aq) With ethanoic acid, a weak acid, magnesium produces magnesium ethanoate and hydrogen, but this is a slower reaction. Mg(s) + 2CH3COOH(aq) | |
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Note: If you haven't come across salts of organic acids before, it is normal to write the metal after the residue from the acid after it has lost its hydrogen. | |
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The reason for the difference is that the reaction is actually between the metal and hydrogen ions in solution. Mg(s) + 2H+(aq) In the weak acid, at any one time there are fewer hydrogen ions in solution. However, the ionisation of the weak acid is a reversible reaction, and as hydrogen ions are removed by reaction with the metal, more acid splits up to replace them, So the reaction will continue until it is complete - just more slowly. Testing with a pH meter or universal indicator For acids of the same concentration, a strong acid has a lower pH than a weak one. This has already been discussed in statement 7.2.4. Conductivity The electrical conductivity of a solution depends on the number and nature of the ions present. That means that the conductivity of a strong acid will be higher than a weak one of the same concentration. A 1 mol dm-3 solution of hydrochloric acid, for example, contains a far greater concentration of ions than a solution of ethanoic acid of the same concentration. Its conductivity will also be far greater. Statements 7.2.7 and 7.2.8 These statements are about simple neutralisation reactions. The statements tell you all you really need to know. If you react sodium hydroxide solution and hydrochloric acid together in the right proportions, they neutralise each other. NaOH(aq) + HCl(aq) You get a salt (in this case sodium chloride) formed as well as water. But if you rewrite this as an ionic equation, you find that the sodium ions and chloride ions are spectator ions, and that the underlying reaction is OH-(aq) + H+(aq) As I have written it, it matches the order of substances in the full equation, but it is more usually written H+(aq) + OH-(aq) This is true for all reactions involving solutions of soluble hydroxides and solutions of acids. The spectator ions will of course be different in each case, and so the salt formed will be different. Potassium hydroxide solution and nitric acid will give a solution of potassium nitrate; sodium hydroxide and sulfuric acid will give a solution of sodium sulfate; and so on. But the ionic equation for the reaction is identical in each case: H+(aq) + OH-(aq)
© Jim Clark 2020 |
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MgCl2(aq) + H2(g)