Topic 12: Acid Base Equilibria

Question 1

What are Bronsted-Lowry acids?

Answer 1

• acids are proton donors
• they release H⁺ ions when they’re mixed with water
• H⁺ ions are always combined with H₂O to form hydroxonium ions, H₃O⁺

Question 2

What is the reaction to form hydroxonium ions?

Answer 2

• HA_(aq) + H₂O_(l) → H₃O⁺_(aq) + A^-_(aq)

Question 3

What are Bronsted-Lowry bases?

Answer 3

• bases are proton acceptors
• when they’re in solution, they take hydrogen ions from water molecules
• B_(aq) + H₂O_(l) → BH⁺_(aq) + OH^-_(aq)

Question 4

What are strong acids and strong bases?

Answer 4

• strong acids dissociate almost completely in water (nearly all the H⁺ ions will be released
• e.g. HCl is a strong acid
• strong bases dissociate almost completely too
• e.g. NaOH

Question 5

What are weak acids?

Answer 5

• weak acids dissociate only very slightly in water
• only small numbers of H⁺ ions are formed
• an equilibrium is set up, which lies well over to the left
• e.g. ethanoic acid

Question 6

What are weak bases?

Answer 6

• weak bases only slightly protonate in water
• equilibrium lies on the left
• e.g. ammonia

Question 7

What are conjugate pairs?

Answer 7

• conjugate pairs are species that are linked by the transfer of a proton
• the species that has lost a proton is the conjugate base
• the species that has gained a proton conjugate acid

Question 8

What the definition of the standard enthalpy change of neutralisation?

Answer 8

• the standard enthalpy change of neutralisation is the enthalpy change when the acid and alkali neutralise each other under standard conditions to form one mole of water

Question 9

How do you calculate pH?

Answer 9

• pH = -log₁₀ [H⁺]

Question 10

How do you work out [H⁺]?

Answer 10

• 10^-pH

Question 11

What is the H⁺ concentration of strong monoprotic acids?

Answer 11

• the H+ concentration is the same as the acid concentration
• they dissociate fully
• monoprotic: each mole of acid produces one mole of hydrogen ions

Question 12

How many protons are released per molecule of diprotic acid?

Answer 12

• two protons

Question 13

How do you find the pH of a weak acid?

Answer 13

• use K_a (the acid dissociation constant)

Question 14

What are the assumptions you have to make to find K_a?

Answer 14

• [HA_(aq)]_(start) ≈ [HA_(aq)]_{(equilibrium)}
• [H⁺_(aq)] ≈ [A^-_(aq)]

Question 15

What is the ionic product of water?

Answer 15

• K_W = [H⁺] [OH^-]
• units are always mol²dm^-6

Question 16

How do you use K_W to find the pH of a strong base?

Answer 16

Question 17

What is K_W at 25°C (298K)?

Answer 17

• K_w = 1.0 x 10^-14 mol²dm^-6

Question 18

What is pK_w at standard conditions?

Answer 18

• always 14.00

Question 19

How do you calculate pK_a and K_a from pK_a?

Answer 19

• pK_a = -log₁₀ K_a
• Ka = 10^-pK_a

Question 20

How do you calibrate a pH meter?

Answer 20

• place the bulb of the pH meter into deionised water and allow reading to settle
• adjust reading so that it read 7.0
• do the same with a standard solution of pH4 and pH 10
• rinse the probe with deionised water between each reading

Question 21

How does the pH of a strong acid change when diluted?

Answer 21

• diluting the acid by a factor of 10 increases the pH by 1
• you can see by using pH = -log₁₀[acid]

Question 22

How does the pH of a weak acid change when diluted?

Answer 22

• diluting a weak acid by a factor of 10 increases the pH by 0.5

Question 23

Describe these titration curves

Answer 23

1. strong acid/strong base
2. strong acid/weak base
3. weak acid/strong base
4. weak acid/weak base

Question 24

How do you explain the shapes of titration curves?

Answer 24

• the initial pH depends on the strength of the acid
• a strong acid titration will start at a much lower pH than a weak acid
• to start with, the addition of small amounts of the base have little impact on the pH of the solution
• final pH depends on the strength of the base
• the stronger the base, the higher the final pH

Question 25

What is the equivalence point?

Answer 25

- the centre of the equivalence line - [H⁺] ≈ [OH^-] - the acid is just neutralised - a tiny amount of base causes a sudden, big change in pH

Question 26

How do you find the pK_a of a weak acid using titration curves?

Answer 26

- find the pH at the half-equivalence point - half of the acid has been neutralised - at the half-equivalence point, [HA] = [A^-] - therefore, K_a = [H⁺] and pK_a = [pH] (see image)

Question 27

What is a buffer?

Answer 27

- a buffer is a solution that minimises changes in pH when small amounts of acid or base are added

Question 28

What are the two ways of making acidic buffers?

Answer 28

1. Mix a weak acid with the salt of its conjugate base: e. g. ethanoic acid and sodium ethanoate 2. Mix an excess of weak acid with a strong base e. g. ethanoic acid and sodium hydroxide

Question 29

What is the equilibrium set up for acidic buffers?

Answer 29

Question 30

How do acidic buffers work?

Answer 30

- the conjugate base accepts any extra H⁺ and conjugate acid release H⁺ if there is too much base - if you add a small amount of acid, the H⁺ concentration increases - most of the extra H+ ions combine with CH₃COO^- ions to form ethanoic acid - this shifts equilibrium to the left, reducing the H+ concentration close to original value - if a small amount of base is added, the OH- concentration increases - most of the extra OH^- ions reaction with H⁺ ions to form water - this causes more CH₃COOH to dissociate to form H+ ions, shifting equilibrium to the right

Question 31

What are alkaline buffers made of?

Answer 31

- a weak base and one of its salts - e.g. a mixture of ammonia solution (a base) and ammonium chloride (a salt of ammonia)

Question 32

How do alkaline buffers work?

Answer 32

Question 33

How can buffer solutions been seen on a titration curve?

Answer 33

Question 34

Why are buffer solutions important in blood?

Answer 34

- blood needs to be kept at pH 7.4 - this is controlled by carbonic acid-hydrogencarbonate buffer system - carbonic acid dissociates into H⁺ ions and HCO₃^- ions H₂CO_3(aq) ⇔ H⁺_(aq) + HCO₃^-_(aq) - if the concentration of H+ rises in the blood, then the HCO₃^- ions will react with the excess H+ ions - the equilibrium will shift to the left, reducing the H+ concentration to its original value - vice versa - H₂CO₃ concentration is controlled by respiration H₂CO_3(aq) ⇔ H₂O_(l) + CO_2(aq) - HCO₃^- levels are controlled by kidneys

Question 35

What assumptions need to be made to calculate the pH of a buffer solution?

Answer 35

- the salt of the conjugate base is fully dissociated, so assume that the equilibrium concentration of A^- is the same as the initial concentration of the salt - HA is only slightly dissociated, so assume that its equilibrium concentration is the same as its initial concentration

Question 36

How do you calculate the concentration of salt and acid or base to create a buffer with a specific pH?

Answer 36