1. What are acids and bases?

2. Monoprotic and diprotic acids Many acids are called monoprotic acids. This means that they only donate one mole of protons per mole of acid; e.g. HCl, HNO 3, CH 3 COOH. Triprotic acids donate three moles of protons per mole of acid; e.g. H 3 PO 4. Some acids are diprotic acids. This means that they can donate two moles of protons per mole of acid; e.g. H 2 SO 4, HOOCCOOH. H 2 SO 4(aq) + 2H 2 O (l)  2H 3 O + (aq) + SO 4 2- (aq) HCl (aq) + H 2 O (l) H 3 O + (aq) + Cl - (aq)

3. Conjugate acids and bases A conjugate acid is a species formed from a Brønsted-Lowry base by the addition of a proton. A conjugate base is a species formed from a Brønsted-Lowry acid by the loss of a proton. NH 3(aq) + H 2 O (l) NH 4 + (aq) + OH – (aq) In the example above, the conjugate acid of ammonia (NH 3 ) is NH 4 + and the conjugate base of H 2 O is OH –. NH 3 accepts a proton, so it is a Brønsted-Lowry base. H 2 O donates a proton, so it is a Brønsted-Lowry acid.

4. Interpreting equations

18. What is pH? The acidity of an aqueous solution depends on the number of H + (H 3 O + ) ions in solution. The pH is defined as: The pH scale is a logarithmic scale with base 10. This means that each value is 10 times the value below it. For example, pH 5 is 10 times more acidic than pH 6. pH values are usually given to 2 decimal places. pH = – log 10 [H + ] where [H + ] is the concentration of H + in mol dm –3.

22. Strong and weak acids The strength of an acid refers to its pH, not its concentration. A strong acid has a low pH (usually 0 or 1). This means that the concentration of H + is high. This is because the acid is fully dissociated into its ions. A weak acid has a higher pH (but still less than 7). This means that the concentration of H + is lower than for a strong acid. This is because the acid is not fully dissociated into its ions. CH 3 COOH (aq) + H 2 O (l) CH 3 COO – (aq) + H 3 O + (aq) HCl (aq) + H 2 O (l) H 3 O + (aq) + Cl – (aq)

25. Calculating the pH of a strong acid

26. Calculations of strong acid pH

27. Calculating [H + ] from the pH Example: Calculate the concentration of H + (aq) ions given that the pH of a solution of hydrochloric acid is 2.50. Make sure you know which buttons to push on your calculator. pH = –log 10 [H + ] The equation for pH can be rearranged so that H + concentration (in mol dm –3 ) can be found. [H + ] = 10 –pH therefore [H + ] = 3.16 × 10 –3 mol dm –3 [H + ] = 10 –2.5 [H + ] = 10 –pH

28. Calculations of [H + ] from pH

29. Weak acids and K a Weak acids are only partially dissociated into their ions when dissolved in water. A weak acid, HA, dissociates in water according to the following equilibrium: The expression for the equilibrium constant, K a, is as follows: HA (aq) + H 2 O (l) H 3 O + (aq) + A – (aq) Or more simply: HA (aq) H + (aq) + A – (aq) K a = [H + (aq) ][A – (aq) ] [HA (aq) ] The units for K a are mol dm –3.

30. K a and pK a The extent to which an acid dissociates is shown by the value of K a. The larger the value of K a, the stronger the acid. The lower the value of pK a, the larger the value of K a, and the stronger the acid. pK a = –log 10 K a The values of K a span a wide range. To make them easier to interpret, a new term pK a is used. In general: if K a is greater than 1, the acid is strong. if K a is less than 1, the acid is weak

31. Calculating the pH of a weak acid 1

32. Calculating the pH of a weak acid 2

33. Calculations of weak acid pH

34. The ionic product for water, K w Water acts as both a Brønsted-Lowry acid and a Brønsted-Lowry base in this dissociation. [H 2 O] is not included in this expression because it is essentially constant. A constant called the ionic product for water, K w, is defined as: Water exists, not just as molecules, but in the equilibrium: 2H 2 O (l) H 3 O + (aq) + OH – (aq) Or, more simply: H 2 O (l) H + (aq) + OH – (aq) K w = [H + ][OH – ] The units for K w are mol 2 dm –6

35. The pH of water

36. Calculating the pH of a strong base

37. Calculations of strong base pH

40. pH calculations summary

41. pH calculations

81. Titration In a titration, a solution of known concentration is usually titrated with a solution of unknown concentration. An acid and a base can be titrated to find the concentration of one of them. If a pH meter is used, a pH curve can be plotted. This is a graph showing how the pH of a solution changes as base (or acid) is added. During an acid-base titration, an indicator is often used to show when the solution has been neutralized. This point is called the equivalence point.

82. pH curves

83. Titration calculations

84. colourlesspink yellowblue redyellow purple redyellow redyellow Selecting an appropriate indicator Around the equivalence point of a titration, the pH changes very rapidly. Indicators change colour over a narrow pH range approximately centred around the pK a of the indicator. An indicator will be appropriate for a titration if the pH range of the indicator falls within the rapid pH change for that titration. bromophenol blue methyl orange methyl red phenolphthalein bromothymol blue thymol blue IndicatorColour in acidpH rangeColour in alkali 1.2–2.8 3.1–4.4 4.4–6.2 6.0–7.6 8.3–10.0 3.0–4.6

90. Indicators and pH curves

94. The half equivalence point

97. Titration curves: summary

98. Strong acid–strong base titrations

99. Strong acid–strong base calculations

101. Buffer solutions A buffer solution is one that is able to oppose changes in pH when small quantities of acid or base are added. A basic buffer or alkaline buffer consists of a weak base and its conjugate acid (usually in the form of one of its salts). E.g. a mixture of NH 3(aq) and NH 4 Cl (aq) is a basic buffer. An acidic buffer consists of a weak acid and its conjugate base (usually in the form of one of its salts). E.g. a mixture of CH 3 COOH (aq) and CH 3 COONa (aq) is an acidic buffer.

102. How do buffers work?

103. Buffers in the body It is essential that the pH of the blood remains constant for the chemical reactions that occur in the body. In the blood, the following equilibrium is set up: If H + ions are added, the equilibrium shifts to the right to remove them. If OH – ions are added, they react with the H + ions present. The equilibrium shifts to the left, releasing more H + ions. HCO 3 – + H + CO 2 + H 2 O

117. Calculating the pH of a buffer solution

118. Calculations: buffer pH

119. Making buffer from acid and base When a small quantity of a strong base is added to a weak acid, a buffer solution is formed. Some of the weak acid reacts with the strong base to form the salt of the weak acid. For example: CH 3 COOH (aq) + NaOH (aq) CH 3 COONa (aq) + H 2 O (l) If the acid is in excess, the solution contains some of the salt of the weak acid, along with some of the weak acid, so the solution is a buffer.

120. Calculating the buffer pH

121. Weak acid–strong base calculation

122. Changes in the pH of a buffer When acid or alkali is added to a buffer solution, the equilibrium shifts to oppose the change. The pH does not remain unchanged, but the buffer action ensures that only very small changes occur. If acid is added, the equilibrium shifts to the left, increasing [CH 3 COOH] and reducing [CH 3 COO – ]. If base is added, the equilibrium shifts to the right, reducing [CH 3 COOH] and increasing [CH 3 COO – ]. These facts can be used to calculate the change in pH. CH 3 COOH + H 2 O CH 3 COO – + H 3 O +

123. Calculating changes in pH

124. Calculations: changes in pH