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Edexcel Chemistry Unit 4 > 4.7 Acid/Base Equilibria > Flashcards

4.7 Acid/Base Equilibria Flashcards

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1
Q

What are Brønsted-Lowry acids?

A

Proton donors that realise H+ ions when they’re mixed with water.

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2
Q

What ares Brønsted-Lowry bases?

A

Proton acceptors that take hydrogen ions from water molecules.

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3
Q

What happens to H+ ions when they are released into water.

A

They form hydroxonium ions, H3O+.

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4
Q

Give the equation for the dissociation of the general acid HA.

A

HA(aq) + H2O(l) → H3O+(aq) + A-(aq)

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5
Q

Give the equation of the base B accepting a proton.

A

B(aq) + H2O(l) → BH+(aq) + OH-(aq)

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6
Q

Give an example of a strong acid and a strong base and explain what makes them “strong”.

A

Hydrochloric acid and sodium hydroxide. They ionise almost completely in water.

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7
Q

What is a weak acid? Give an example.

A

Weak acids ionise only very slightly in water, setting up an equilibrium. Citric acid.

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8
Q

Give the equation for the dissociation of ethanoic acid.

A

CH3COOH(aq) ⇋ CH3COO-(aq) + H+(aq)

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9
Q

What is a weak base? Give an example.

A

Weak bases only slightly ionise in water. Ammonia is a weak base.

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10
Q

Give the equation for the dissociation of ammonia.

A

NH3(aq) + H2O(l) ⇋ NH4+(aq) + OH-(aq)

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11
Q

When can acids lose their protons?

A

When there is a base to accept them.

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12
Q

Give the equation for when the acid, HA, transfers a proton to base, B.

A

HA(aq) + H2O(l) ⇋ H3O+(aq) + A-(aq)

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13
Q

Name their two conjugate pairs in this equilibrium:

HCl(aq) + H2O(l) ⇋ H3O+(aq) + Cl-(aq)

A

HCl(aq) and Cl- (acid/base) and H2O(l) and H3O+(aq) (base/acid).

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14
Q

Give the equation for the dissociation of water.

A

2H2O(l) ⇋ H3O+(aq) + OH-(aq)

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15
Q

Write an expression for the equilibrium constant of the dissociation of water.

A

Kc = ([H+][OH-]) / [H2O]

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16
Q

What do you get if you multiply the constant Kc by [H2O], which is also constant.

A

[H+][OH-], the ionic product of water, Kw.

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17
Q

Give the expression for Kw.

A

Kw = [H+][OH-]

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18
Q

Kw always has the same value for…

A

An aqueous solution at a given temperature.

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19
Q

What is the value of Kw at the standard temperature?

A

At 298K, Kw = 1.0 x 10^-14 mol^2 dm^-6

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20
Q

What is pKw equal to?

A

-logKw

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21
Q

What is the advantage of expressing Kw as pKw?

A

They’re a reasonable size to work with.

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22
Q

What is pKw at 25ºC?

A

14

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23
Q

What is the definition of a neutral solution?

A

A solution in which [H+] = [OH-]

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24
Q

What is a solution where [OH-] is greater than [H+]?

A

Alkaline.

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25
Q

Give the equation for pH.

A

pH = -log[H+]

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26
Q

Describe the pH scale.

A

It goes from 0 (strongly acidic) to 14 (strongly alkaline) and pH 7 is neutral.

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27
Q

What is [H+] for strong monoprotic acids?

A

[Acid]

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28
Q

What is the pH of 0.1 mol dm^-3

A

pH 1.0

29
Q

How can you find [H+] of a strong base if you know its concentration?

A

[H+] = Kw / [OH-]

30
Q

Give the original expression for Ka.

A

Ka = ([H+][A-]) / [HA]

31
Q

Give the simplified expression for Ka.

A

Ka = [H+]^2 / [HA]

32
Q

Give two assumptions made when using the expression for Ka.

A

Only a small amount of HA dissociates so [HA] start = [HA] equilibrium. All of the H+ ions come from the acid so [H+] = [A-].

33
Q

What are the units of Ka?

A

mol dm^-3

34
Q

How would you use Ka to find the pH of a weak acid?

A

[H+] = √(Ka [HA]) then use -log[H+]

35
Q

How can you calculate Ka from pKa?

A

Ka = 10^(-pKa)

36
Q

What happens to the pH of a strong acid when you dilute it by a factor of 10?

A

It increases by 1.

37
Q

What happens to the pH of a weak acid when you dilute it by a factor of 10?

A

It increases by 0.5.

38
Q

Describe the shape of a strong acid, strong alkali titration curve.

A

S shaped curve starting around pH 1 and finishing around pH 13.

39
Q

Describe the shape of a strong acid, weak alkali titration curve.

A

S shaped curve starting around pH 1 and finishing around pH 8.

40
Q

Describe the shape of a weak acid, strong alkali titration curve.

A

S shaped curve starting around pH 4 and finishing around pH 13, with a small kink at the beginning.

41
Q

Describe the shape of a weak acid, weak alkali titration curve.

A

S shaped curve starting around pH 4 and finishing around pH 9, with no vertical region.

42
Q

Give the colours of the indicator methyl orange and state the pH range where the colour changes.

A

Red at low pH, changes around 3.1 - 4.4 to yellow at a high pH.

43
Q

Give the colours of the indicator Phenolphthalein and state the pH range where the colour changes.

A

Colourless at low pH, changes around 8.3 - 10 to pink at a high pH.

44
Q

Which indicator would you use for strong acid, strong alkali titration?

A

Methyl orange or phenolphthalein.

45
Q

Which indicator would you use for strong acid, weak alkali titration?

A

Methyl orange.

46
Q

Which indicator would you use for weak acid, strong alkali titration?

A

Phenolphthalein.

47
Q

Which indicator would you use for weak acid, weak alkali titration?

A

Neither methyl orange nor phenolphthalein. A pH meter would be more suitable.

48
Q

What is the half-equivalence point?

A

Where half of a weak acid has been neutralised by a strong alkali.

49
Q

What happens to [A-] at the half-equiclance point?

A

[A-] = [H+]

50
Q

What is Ka equal to at the half-equivalence point?

A

Ka = [H+]

51
Q

How can you find pKa at the half-equivalence point?

A

pKa = pH

52
Q

What is a buffer solution?

A

A solution that resists changes in pH when small amounts of acid of alkali are added.

53
Q

What are acidic buffers made from?

A

A weak acid and one of its salts.

54
Q

Give an example of an acidic buffer.

A

Ethanoic acid and sodium ethanoate.

55
Q

What happens to the acid salt of a buffer when it dissolves?

A

It fully dissociates into its ions.

56
Q

What is in an ethanoic acid buffer solution?

A

A large amount of ethanoate ions and a large amount of undissociated ethanoic acid molecules.

57
Q

What happens to an ethanoic acid buffer solution when you add small amounts of acid?

A

Most of the extra H+ ions combine with the ethanoate ions to form undissociated acid, shifting the equilibrium and reducing the h+ concentration to near its original value.

58
Q

What happens to an ethanoic acid buffer solution when you add small amounts of alkali?

A

The OH- ions react with H+ to form water, removing h+ ions form the solution. This causes more acid to dissociate, replacing most of the H+ ions lost.

59
Q

What are alkaline buffers made form?

A

A weak base and one of its salts.

60
Q

Give an example of an alkaline buffer.

A

Ammonia solution and ammonium chloride.

61
Q

What happens when a small amount of acid is added to an ammonia buffer.

A

The acid adds H+ ions, but most of them react with the NH3, shifting the equilibrium and reducing the H+ concentration to near its original value.

62
Q

What happens when a small amount of alkali is added to an ammonia buffer?

A

The OH- concentration increases and the OH- ions react with H+ ions to form water. The NH4+ molecules dissociate to generate the lost H+ ions, resisting the change in pH.

63
Q

What causes the distinctive shape of a weak acid, strong base titration curve?

A

Buffer action.

64
Q

In terms of buffer action, what happens when sodium hydroxide is added to ethanoic acid?

A

The pH initially starts to change because NaOH is a strong alkali but then sodium ethanoate is formed and the ethanoic acid gradually dissociates to replenish the H+ ions, causing the curve to level off.

65
Q

Give two examples of buffer action in biological environments.

A

Constant pH in cells. Blood needs to be kept at pH 7.4.

66
Q

Give the equation for the buffer equilibrium occurring in cells to maintain a constant pH.

A

H2PO4- ⇋ H+ + HPO4`2-

Dihydrogen phosphate and hydrogen phosphate.

67
Q

Give the equations for the buffer equilibriums occurring in the blood to maintain pH 7.4.

A

H2CO3(aq) ⇋ H+(aq) + HCO3-(aq)

H2CO3(aq) ⇋ H2O(l) + CO2(g)

68
Q

Give three examples of buffers used in foods.

A

Sodium citrate, phosphate ions and benzoate ions.

69
Q

Why are buffers used in foods?

A

Changes in pH can be caused by bacteria and fungi and this causes the food to deteriorate.

Edexcel Chemistry Unit 4 (6 decks)

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