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Chemistry A level (Edexcel) > TOPIC 12 - Acid-base equilibria > Flashcards

TOPIC 12 - Acid-base equilibria Flashcards

1
Q

What is a Bronsted-Lowry acid?

A

A proton donor

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2
Q

Do actual H+ ions really exist in solution?

A

No. They tend to form hydroxonium ions. (H3O+)

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3
Q

What is a Bronsted-Lowry base.

A

Proton acceptor

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4
Q

What is the general equation for the reaction between an acid and water?

A
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5
Q

What is the general equation for the reaction between an alkali and water?

A
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6
Q

What is the difference between strong bases/acids and weak bases/acids?

A

Weak base/acid dissociate partly while strong bases/acids dissociate completely.

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7
Q

What is the conjugate acid of a base?

A

The species of a base with an added proton.

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8
Q

What is the conjugate base of an acid?

A

The species of that acid without the proton.

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9
Q

What is the conjugate base of HA?

A

A-

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10
Q

What is the conjugate acid of B?

A

BH+

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11
Q

What is the conjugate acid of H2O?

A

H3O+

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12
Q

What is the conjugate base of H2O?

A

OH-

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13
Q

What do acids and bases give when they react together?

A

Salt + Water.

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14
Q

What is the equation for the reaction between NH3 with H2O?

A
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15
Q

What is the equation for the reaction between ammonia and sulfuric acid?

A
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16
Q

What isthe standard enthalpy change of neutralisation?

A
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17
Q

Is the standard enthalpy of neutralisation normally endothermic or exothermic?

A

exothermic

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18
Q

Why does enthalpy of neutralisation vary with different weak acids/bases?

A
  • Energy released by OH- and H+ is the same.
  • However weak acids and bases are constantly dissociating during neutralisation due to equilibrium) and energy needed to dissociate different weak bases/acids varies.
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19
Q

Why does the enthalpy of neutralisation of strong acids/bases

A
  • Strong acids and bases are already fully dissociated in solution so no energy is needed.
  • The OH- + H+ enthalpy is the same.
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20
Q

What is the equation for pH?

A
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21
Q

What is a polyprotic / polibasic acid?

A

Acids that have more than one proton

(diprotic, triprotic etc.)

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22
Q

What assumption is made when calculating the pH of strong acids/bases?

A

They dissociate fully.

(i.e. [HA] = [H+] )

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23
Q

What is Kw?

A

Ionic product of water. (1.0x10^-14 under standard conditions)

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24
Q
A
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25
Q

What is the general equation for Ka?

A
26
Q

What is the assumption made when calculating Ka for weak acids with respect to HA?

A
27
Q

What is the assumption made when calculating Ka for weak acids with respect to H+?

A
28
Q
A
29
Q

Does Kw change if temperature changes?

A

Yes.

30
Q

Why are logs used for pH / pKa / pKw ?

A

So that values can be displayed on smaller scales, which makes it easier to use.

31
Q

What is pKa?

A
32
Q
A
33
Q

How does pH of a strong acid solution vary if you consecutively dilute it by a factor of 10?

A

pH increases by 1 each time.

34
Q

How does the pH of a weak acid solution vary if you consecutively dilute it by a factor of 10?

A

pH increases by 0.5 each time.

35
Q

What is the titration curve for strong acid and strong base?

A
36
Q

What is the titration curve for strong acid and weak base?

A
37
Q

What is the titration curve for weak acid and strong base?

A
38
Q

What is the titration curve for weak acid and weak base?

A
39
Q

What is the half neutralisation point? What is important about it? Why?

A

The point halfway between zero and the end point.

At this point, pKa = pH.

This is because half the H+ have reacted and so [HA] = [A-].

So they cancel each other out of the Ka equation to leave Ka = [H+]. So pKa = pH.

40
Q

What is the working range of methyl orange?

A

pH 3 - 4.5

41
Q

What is the working range of phenolpthalein?

A

pH 8.2 - 10

42
Q

What is the colour of methyl orange in acids and bases?

A

Yellow in alkaline.

Red in acid.

43
Q

What is the colour of phenolpthalein in acids and bases?

A

Colourless in acid

Pink with alkaline

44
Q

What should be taken into account in weak acid / weak base titrations?

A

They have no sharph pH change and so no indicator is suitable. A pH meter should be used.

45
Q

What is a buffer?

A
46
Q

How do acidic buffers resist a change in pH when an acid (H+) is added?

A
  1. Acidic buffers are a solution of a weak acid and its salt, so there are lots of HA and A-.
  2. H+ react with A- to give HA.
  3. Since there are large quantities of A- and HA, the changes have little effect on the concentration.
  4. Ka = [H+][A-] / [HA] so if the ratio of [A-] to [HA] is more or less the same, and Ka is a constant under the same temperature, then [H+] has to remain more or less the same.
  5. Therefore pH remains the same.
47
Q

How do acidic buffers resist a change in pH when a base (OH-) is added?

A
  1. OH- reacts with the H+ in solution from the weak acid.
  2. Following Le Chatelier’s principle, equilibrium shifts to the right (HA > H+ + A-) to counteract the change and replaced the lost H+.
  3. Therefore pH remains the same.
48
Q

How could an acidic buffer be made from excess weak acid and a strong base?

A

Strong base reacts with weak acid to give the salt for the acid.

Since the acid is in excess we still have large quantites of acid left.

Therefore an acidic buffer solution is cretaed (large quantites of weak acid and its salt).

49
Q

Around what pH do basic buffers maintain a solution?

A

Above 7.

50
Q

What does a basic buffer consist on?

A

Weak base and its salt.

51
Q

How do basic buffers resist a change in pH when a base is added?

A
  1. Increase in OH- shift the equilibrium to produce more base (BOH > B+ + OH- )
  2. OH- ions react with the conjugate acid (i.e. NH4 -) from the salt to form the base.
  3. Thus the OH- concentration is more or less the same and so pH doesn’t change.
52
Q

How do basic buffers resist a change in pH when acid (H+) is added?

A
  1. H+ react with the OH- ions in solution (from the weak base).
  2. Equilibrium shifts to the right to replace the OH- that have reacted. (BOH > B+ + OH-)
  3. Therefore pH remains constant.
53
Q

What happens at A?

A

pH raises quickly because plenty of OH- ions from the base react with the H+ ions.

54
Q

What happens at B?

A

A buffer solution is established. Some of the acid has reacted, forming the salt, and some of the acid remains.

Therefore the pH barely changes.

55
Q

What happens at C?

A

Equivalence point. All of the acid is used up so the buffer is no longer present and so pH raises quickly.

56
Q

How does breathing out CO2 reduce the pH in the blood?

A

H2O(aq) + CO2(aq) <> H2CO3(aq) <> H+(aq) + HCO3-(aq)

When CO2 is exhaled, equilibrium shifts to the left to replace it, so more H2CO3 is used, so equilibrium shifts to the left, so more H+ ions are used up, decreasing the pH.

57
Q
A
58
Q

What assumptions are made when calculating the pH of a buffer solution?

A
59
Q

What is the Heddeson-Heselbalch equation for concentrations in a buffer solution? What assumptions are made?

A
60
Q
A
61
Q
A

Chemistry A level (Edexcel) (19 decks)

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