1. Acids and Bases

2. Brønsted-Lowry Theory Brønsted-Lowry describes reactions of acids as involving the donation of a hydrogen ion (H + ) Brønsted-Lowry describes reactions of acids as involving the donation of a hydrogen ion (H + ) A hydrogen ion is a hydrogen that has lost its only electron. A hydrogen ion is a hydrogen that has lost its only electron. In most cases a hydrogen ion is a proton. In most cases a hydrogen ion is a proton. Chemists often use the terms hydrogen ion and protons interchangeably. Chemists often use the terms hydrogen ion and protons interchangeably.

3. According to Brønsted-Lowry theory, a substance behaves as an acid when it donates a proton to a base. According to Brønsted-Lowry theory, a substance behaves as an acid when it donates a proton to a base. In other words it donates a H + to a base. In other words it donates a H + to a base. A substance behaves as a base when it accepts a proton from an acid. A substance behaves as a base when it accepts a proton from an acid. Brønsted-Lowry Theory

4. HENCE: Acids are proton donors Acids are proton donors Bases are protons acceptors Bases are protons acceptors

5. As protons are exchanged from an acid to a base, this definition explains why acids and bases react together. As protons are exchanged from an acid to a base, this definition explains why acids and bases react together. For example: For example: Hydrochloric acid is very soluble in water. The molecules ionise in water. Hydrochloric acid is very soluble in water. The molecules ionise in water. HCl(g) + H 2 O(l) ―› H 3 O + (aq) + Cl - (aq) In an aqueous solution of hydrogen chloride, nearly all the hydrogen chloride is present as ions, virtually no molecules of HCl remain. This ionised solution is the hydrochloric acid we use regularly. In an aqueous solution of hydrogen chloride, nearly all the hydrogen chloride is present as ions, virtually no molecules of HCl remain. This ionised solution is the hydrochloric acid we use regularly. Brønsted-Lowry Theory

6. Hydrogen Chloride In the reaction each hydrogen chloride molecule donated a proton to the water molecule. In the reaction each hydrogen chloride molecule donated a proton to the water molecule. According to the Brønsted-Lowry theory is it an acid or a base? According to the Brønsted-Lowry theory is it an acid or a base? The water molecule has accepted a proton. The water molecule has accepted a proton. So is water an acid or a base? So is water an acid or a base?

7. Conjugate Pairs HCl and Cl - can be formed from each other by the loss or gain of a single H+ (proton). HCl and Cl - can be formed from each other by the loss or gain of a single H+ (proton). These are called conjugate acid/base pairs. These are called conjugate acid/base pairs. Similarly H 3 O + and H 2 O are also a conjugate pair. Similarly H 3 O + and H 2 O are also a conjugate pair. A conjugate pair is two species which differ by a proton. A conjugate pair is two species which differ by a proton.

8. Conjugate Pairs What are the conjugate pairs in this reaction. What are the conjugate pairs in this reaction. NH 3 (aq) + H 2 O(l) ―› NH 4 + (aq) + OH - (aq) Base Acid

9. Some Common acids and bases

10. Amphiprotic Substances Some substances can be acids or bases depending on what they react with. Some substances can be acids or bases depending on what they react with. They can donate or accept protons. They can donate or accept protons. These substances are said to be amphiprotic. These substances are said to be amphiprotic. Can you name any amphiprotic substances? Can you name any amphiprotic substances?

11. Water Water is an amphiprotic substance. Water is an amphiprotic substance. It can be an acid and a base. It can be an acid and a base. If the solute is a stronger acid than water, then water will act as a base. If the solute is a stronger acid than water, then water will act as a base. If the solute is a stronger base than water, then water will act as an acid. If the solute is a stronger base than water, then water will act as an acid.

12. Amphiprotic Substances

13. Acid and Base Strength Different acid solutions of the same concentration do not have the same pH. Different acid solutions of the same concentration do not have the same pH. Some acids donate a proton more readily than others. Some acids donate a proton more readily than others. The strength of an acid or a base is its ability to donate or accept an proton. The strength of an acid or a base is its ability to donate or accept an proton. We generally use an acids tendency to donate a proton to water or a base’s tendency to accept a proton from water, as a measure of its strength. We generally use an acids tendency to donate a proton to water or a base’s tendency to accept a proton from water, as a measure of its strength.

14. Strong Acids Acids that ionise completely in solution are called strong acids. Acids that ionise completely in solution are called strong acids. Solutions of strong acids would contain ions, with virtually no unreacted acid molecules remaining. Solutions of strong acids would contain ions, with virtually no unreacted acid molecules remaining. HCl(g) + H 2 O(l) ―› H 3 O + (aq) + Cl - (aq) H 2 SO 4 (aq) + H 2 O(l) ―› H 3 O + (aq) + HSO 4 - (aq) HNO 3 (aq) + H 2 O(l) ―› H 3 O + (aq) + NO 3 - (aq)

15. Weak Acids In water it ionises to produce ethanoate ions and hydronium ions. In water it ionises to produce ethanoate ions and hydronium ions. CH 3 COOH(l) + H 2 O(l) CH 3 COO - (aq) + H 3 O + (aq) However only a small proportion (less than 1%) of the ethanoic acid molecules actually ionise. However only a small proportion (less than 1%) of the ethanoic acid molecules actually ionise. So in water more is present as CH 3 COOH than CH 3 COO - So in water more is present as CH 3 COOH than CH 3 COO - We use a reversible arrow to represent a weak acid We use a reversible arrow to represent a weak acid

16. Strong and weak bases A strong base dissociates completely in water, all of the compound is now in the form of ions. Hydroxide ions are a strong base A strong base dissociates completely in water, all of the compound is now in the form of ions. Hydroxide ions are a strong base A weak base does not dissociate completely in water. Ammonia is a weak base. A weak base does not dissociate completely in water. Ammonia is a weak base. We also represent weak bases by the reversible arrows. We also represent weak bases by the reversible arrows.

17. Polyprotic acids Acids that are capable of donating more than one proton are polyprotic. Acids that are capable of donating more than one proton are polyprotic. Monoprotic acids can donate only one proton Monoprotic acids can donate only one proton These include HCl, HF, HNO 3, CH 3 COOH These include HCl, HF, HNO 3, CH 3 COOH Diprotic acids can donate two protons Diprotic acids can donate two protons Sulfuric acid H 2 SO 4 and carbonic acid H 2 CO 3 are diprotic acids Sulfuric acid H 2 SO 4 and carbonic acid H 2 CO 3 are diprotic acids Triprotic acids can donate three protons. Triprotic acids can donate three protons. Phosphoric acid H 3 PO 4 is a triprotic acid Phosphoric acid H 3 PO 4 is a triprotic acid

18. Polyprotic acids Polyprotic acids do not donate all their protons at once, but do so in steps when reacting with a base. Polyprotic acids do not donate all their protons at once, but do so in steps when reacting with a base. It also depends on the strength of the acid. It also depends on the strength of the acid. Sulfuric acid (H 2 SO 4 ) is diprotic. Sulfuric acid (H 2 SO 4 ) is diprotic. A diprotic acid ionises in two stages. A diprotic acid ionises in two stages.

19. Stage 1 H 2 SO 4 (aq) + H 2 O(l) HSO 4 - (aq) + H 3 O + (aq) Sulfuric acid is a strong acid in water so this stage occurs to completion. That is all the sulfuric acid molecules have ionised into hydrogen sulfate and hydronium ions.

20. Stage 2 HSO 4 - (aq) + H 2 O(aq) SO 4 2- (aq) + H 3 O + (aq) Hydrogen sulfate is only a weak acid so only a proportion ionise. A solution of sulfuric acid therefore contains hydrogen ions, hydrogen sulfate ions and sulfate ions.

21. Strength versus Concentration Strong and weak refer to acids. Strong and weak refer to acids. They are not the same as concentrated and dilute. They are not the same as concentrated and dilute. Concentrated and dilute describe the amount of acid or base dissolved in a given volume of solution. Concentrated and dilute describe the amount of acid or base dissolved in a given volume of solution.

22. Acidic, Basic and Neutral Solutions The acidity of a solution is a measure of the concentration of hydrogen ions present. The acidity of a solution is a measure of the concentration of hydrogen ions present. The higher the concentration of hydrogen ions, the more acidic the solution. The higher the concentration of hydrogen ions, the more acidic the solution. Quite often we use the H 3 O + instead of the hydrogen ion. Quite often we use the H 3 O + instead of the hydrogen ion.

23. Water Water is both an acid and a base Water is both an acid and a base Pure water undergoes self ionisation to a very small extent. Pure water undergoes self ionisation to a very small extent. H 2 O(l) + H 2 O(l) H 3 O + (aq) + OH - (aq) Water behaves as a very weak acid and a very weak base, producing one hydrogen ion (H 3 O + ) for every hydroxide ion (OH - ). Water behaves as a very weak acid and a very weak base, producing one hydrogen ion (H 3 O + ) for every hydroxide ion (OH - ). Acid Base

24. Acidic Solutions Pure water is neutral because the concentration of H 3 O + ions is equal to the concentration of OH - ions present. Pure water is neutral because the concentration of H 3 O + ions is equal to the concentration of OH - ions present. If an acid is added to water, more H 3 O + ions are produced. The concentration of H 3 O + ions becomes greater than that of OH - ions. If an acid is added to water, more H 3 O + ions are produced. The concentration of H 3 O + ions becomes greater than that of OH - ions. This results in an acidic solution. This results in an acidic solution.

25. Basic Solutions A basic solution is the opposite, if a base is added to water more OH - ions are produced and the concentration of OH - ions becomes greater than that of H 3 O + ions. A basic solution is the opposite, if a base is added to water more OH - ions are produced and the concentration of OH - ions becomes greater than that of H 3 O + ions.

26. Acid, basic and neutral solutions Therefore: Therefore: Acidic solutions contain a greater concentration of H 3 O + ions than OH - ions. Acidic solutions contain a greater concentration of H 3 O + ions than OH - ions. A neutral solution contains equal concentrations of H 3 O + and OH -. A neutral solution contains equal concentrations of H 3 O + and OH -. Basic solutions contain a lower concentration of H 3 O + ions than OH - ions. Basic solutions contain a lower concentration of H 3 O + ions than OH - ions.

27. Measuring Acidity Experimental measurements show that all aqueous solutions contain both H 3 O + ions and OH - ions and that the product of their molar concentrations is always 10 -14 at 25°C. Experimental measurements show that all aqueous solutions contain both H 3 O + ions and OH - ions and that the product of their molar concentrations is always 10 -14 at 25°C. This relationship is called the ionic product and can be represented by: This relationship is called the ionic product and can be represented by: [H 3 O + ] x [OH - ] = 10 -14 M 2 at 25°C The square brackets mean the concentration of the ions.

28. [H 3 O + ] x [OH - ] = 10 -14 M 2 at 25°C Pure water is neutral, so [H 3 O + ]=[OH - ] Pure water is neutral, so [H 3 O + ]=[OH - ] Since 10 -7 x 10 -7 = 10 -14 M 2 Since 10 -7 x 10 -7 = 10 -14 M 2 [H 3 O + ] = 10 -7 M and [OH - ] = 10 -7 M at 25°C [H 3 O + ] = 10 -7 M and [OH - ] = 10 -7 M at 25°C What happens to the [OH - ] as we increase [H 3 O + ]? What happens to the [OH - ] as we increase [H 3 O + ]?

29. At 25°C A Solutions is: A Solutions is: Acidic if [H 3 O + ] > 10 -7 M and [OH - ] 10 -7 M and [OH - ] < 10 -7 M Neutral is [H 3 O + ] = 10 -7 M = [OH - ] Neutral is [H 3 O + ] = 10 -7 M = [OH - ] Basic if [H 3 O + ] 10 -7 M Basic if [H 3 O + ] 10 -7 M

30. Worked Example 0.1 mol of hydrogen chloride (HCl) gas was bubbled into sufficient water to produce 1L of solution. Calculate the solution concentration of: (a) H 3 O + ions (b) OH - ions

31. Worked Example In a 5.6x10 -6 M HNO 3 solution at 25°C, calculate the concentration of: In a 5.6x10 -6 M HNO 3 solution at 25°C, calculate the concentration of: (a) H 3 O + ions (b) OH - ions

32. The pH Scale The pH scale is a useful way of indicating the acidity of a solution. The pH scale is a useful way of indicating the acidity of a solution. pH is defined as: pH is defined as: pH = -log 10 [H 3 O + ] Where [H 3 O + ] is measured in mol L -1. Where [H 3 O + ] is measured in mol L -1. The pH of a solution decreases as the concentration of hydrogen ions increases The pH of a solution decreases as the concentration of hydrogen ions increases

33. pH Since pH is a logarithmic scale, increasing the concentration of H + by a factor of 10 results in a decrease of one pH unit. Since pH is a logarithmic scale, increasing the concentration of H + by a factor of 10 results in a decrease of one pH unit. For example [H + ] = 0.001 M at 25°C For example [H + ] = 0.001 M at 25°C Then the pH = -log [H + ] Then the pH = -log [H + ] = -log [0.001] = -log [0.001] = -log [10 -3 ] = -log [10 -3 ] = -(-3) = -(-3) = 3 = 3

34. pH If [H + ] = 0.01 M at 25°C If [H + ] = 0.01 M at 25°C What would the pH be? What would the pH be? If [H + ] = 10 -7 M at 25°C If [H + ] = 10 -7 M at 25°C What would the pH be? What would the pH be?

35. pH

36. Calculating the pH of aqueous solutions In the following examples [H + ] is used represent [H 3 O + ], since the terms can be used interchangeably. In the following examples [H + ] is used represent [H 3 O + ], since the terms can be used interchangeably. In order to calculate the pH of an aqueous solution, you must first calculate the concentration of the H + ions and apply the formula: In order to calculate the pH of an aqueous solution, you must first calculate the concentration of the H + ions and apply the formula: pH = = -log 10 [H 3 O + ]

37. If the OH - ion concentration is given then the equation [H 3 O + ] x [OH - ] = 10 -14 M 2 If the OH - ion concentration is given then the equation [H 3 O + ] x [OH - ] = 10 -14 M 2 Must be used first to determine the hydrogen ion concentration in the solution at 25°C Must be used first to determine the hydrogen ion concentration in the solution at 25°C Calculating the pH of aqueous solutions

38. Worked Example What is the pH of a solution in which What is the pH of a solution in which [H + ] = 0.0135M?

39. Worked examples What is the pH of a 0.0050 M solution of Ba(OH) 2 ? What is the pH of a 0.0050 M solution of Ba(OH) 2 ? What is the pH of a solution, at 25°C, that contains 1.0g NaOH in 100mL solution? What is the pH of a solution, at 25°C, that contains 1.0g NaOH in 100mL solution? 30.0mL of 0.100M HNO 3 is added to 50.0mL water. What is the pH of the diluted solution. 30.0mL of 0.100M HNO 3 is added to 50.0mL water. What is the pH of the diluted solution.

40. Calculating the concentration of H + in a solution of a given pH If a pH of a solution is known, it can be used to determine the concentration of hydronium ions. If a pH of a solution is known, it can be used to determine the concentration of hydronium ions. The pH relationship can be used in the form: The pH relationship can be used in the form: [H + ] = 10 -pH If pH = 5.00, [H + ] = 10 -5. = 0.0000100 M

41. Worked Example What is [H + ] in a solution of pH 3.47? What is [H + ] in a solution of pH 3.47? What is the concentration of OH - ions in a solution of pH 10.4? What is the concentration of OH - ions in a solution of pH 10.4?

42. Strong acid/Strong base The steepest slant of the curve demonstrates the equivalence point on the titration curve Small volume of strong acid added produces a large changes in pH. This is demonstrating a sharp end point.

43. pH Curves pH curves showing change of pH during a titration of a a strong base with a strong acid, and b a weak base with a strong acid. Phenolphthalein, which changes colour in the pH range 8.2–10, gives a sharp end point in a but a broad end point in b. Methyl orange, which changes colour between pH 3.1 and 4.5, would be a more suitable indicator for the second titration.

44. Which indicator would be best to identify the equivalence point? Chapter 4 Q6. The graphs in Figure 4.7 show the pH curves for titrations involving combinations of acids and bases of various strengths. You have a choice of phenolphthalein and methyl orange indicator. Phenolphthalein changes colour over a pH range 8.2 to 10.0. Methyl orange changes colour between pH 3.2 and 4.4. Decide which indicator’s would be suitable to identify the equivalence point for each reaction. Provide reasons for your selections.

45. Which indicator would be best to identify the equivalence point? Change in pH during a titrations of: a a strong acid with a strong base; b a strong acid with a weak base; c weak acid with a strong base; d weak acid with a weak base.

46. Answer Chapter 4 A6. aThe equivalence point occurs in the range pH 3 to pH 11. Both indicators will change colour over this pH range. Both indicators will provide a sharp end point, i.e. they will change colour at the equivalence point with the addition a small volume, 1 drop, of acid. bThe equivalence point occurs in the pH range 3 to 7. Methyl orange provides a sharper end point over this pH range. cThe equivalence point occurs in the pH range 7 to 11. Phenolphthalein provides the sharper end point. dBoth indicators will provide a broad end point and neither would be suitable.