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Chemistry A - Level > Topic 12 - Acid-base Equilibria > Flashcards

Topic 12 - Acid-base Equilibria Flashcards

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1
Q

What is the definition of a Bronsted-Lowry Acid?

A

Proton donor

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2
Q

What is the definition of a Bronsted-Lowry Base?

A

Proton acceptor

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3
Q

What happens in acid-base reactions?

A

Protons are transferred between substances

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4
Q

Describe what happens in the forwards reaction and backwards reaction when HCl dissociates in water

A

HCl + H2O <=> Cl- + H3O+

In the forwards reaction HCl acts as an acid and donates a proton to H2O. H2O acts as a base and accepts that proton.

In the backwards reaction H3O+ acts as an acid and donates a proton to Cl-. Cl- acts as a base and accepts that proton.

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5
Q

What are the conjugate pairs in the dissociation of HCl in water?

A

HCl and Cl-
H2O and H3O+

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6
Q

Conjugate pairs differ by…?

A

H+

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7
Q

How would you describe the conjugate pairs in the dissociation of HCl and water?

A

HCl is the conjugate acid of Cl-
Cl- is the conjugate base of HCl
H3O+ is the conjugate acid of H2O

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8
Q

What are the conjugate pairs in the dissociation of NH3 in water?

A

NH3 + H2O <=> NH4+ + OH-

NH3 and NH4+
H2O and OH-

NH4+ is the conjugate acid of NH3
OH- is the conjugate base of H2O

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9
Q

What is the definition of a strong acid?

A

One that fully dissociates into its ions

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10
Q

What is the definition of a weak acid?

A

One that partially dissociates into its ions

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11
Q

What type of conjugate pairs do strong acids, strong bases, weak acids and weak bases have?

A

Strong acids - weak conjugate base
Strong bases - weak conjugate acid
Weak acids - strong conjugate base
Weak bases - strong conjugate acid

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12
Q

What is the definition of pH?

A

-log(base 10)[H+]

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13
Q

What value do you use to calculate the pH of a solution?

A

Hydrogen ion concentration

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14
Q

How do you calculate the hydrogen ion concentration from pH?

A

10^-pH

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15
Q

A high concentration of H+ ions means…?

A

A low pH

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16
Q

A low concentration of H+ ions means…?

A

A high pH

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17
Q

A decrease in pH of 1 on the pH scale is equivalent to…?

A

The concentration of H+ ions increasing by 10 fold

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18
Q

What is the assumption made when strong acids dissociate?

A

They dissociate fully into their ions, therefore:

Concentration of acid = concentration of H+ ions when dissociated

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19
Q

What is the definition of the ionic product of water?

A

The product of the molar concentrations of H+ and OH- ions at a specified temperature

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20
Q

What is the value of the ionic product of water equal to at 298K?

A

1 x 10^-14

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21
Q

Write the equation for the ionic product of water

A

Kw = [H+][OH-]

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22
Q

How will temperature affect the ionic product of water?

A

At higher temperatures the forwards reaction, H2O dissociating into H+ and OH- is favoured as the forwards reaction is endothermic. Therefore, the POE shifts to the right and the value of Kw increases. [H+] and [OH-] increase and so the pH decreases.

Vice versa for when the temperatures are lower

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23
Q

What are the units for the ionic product of water?

A

mol^2dm^-6

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24
Q

If you have [OH-] how do you calculate the pH of the solution?

A
  • Use Kw to find [H+]
  • -Log(base 10)[H+]
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25
Q

What do you need to watch out for when doing acid-base calculations?

A

Diprotic/Dibasic compounds

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26
Q

What is the definition of a diprotic/dibasic compound?

A

One that contains 2 potential H+ ions to donate per molecule

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27
Q

If 0.1 moldm^-3 of sulfuric acid dissociates in water, what concentration of H+ ions do you have?

A

0.2 moldm^-3

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28
Q

Does pure water dissociate into its ions?

A

Yes, but only slightly

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29
Q

What is the assumption made when deriving Kw?

A

The expression for Kc for water is:

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

Assumption: Since [H2O] is very large and changes by only a very small amount we can consider it to be a constant

Kc[H2O] = [H+][OH-]

Kc[H2O] = Kw

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30
Q

For the dissociation of water, is the forwards reaction endothermic or exothermic?

A

Endothermic

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31
Q

At higher temperatures the pH of water is…?

A

Lower (more acidic)

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32
Q

At lower temperatures the pH of water is…?

A

Higher (more basic)

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33
Q

Is water always neutral?

A

Yes

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34
Q

When will a solution be neutral?

A

When the [H+] = [OH-]

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35
Q

If you add an acid or alkali to water, will that change the value of Kw?

A

NO, it will only change the [H+] and [OH-] in solution

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36
Q

What is the definition of a strong base?

A

One that fully ionises in solution

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37
Q

What is the definition of a weak base?

A

One that only slightly ionises in solution

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38
Q

Definition of strength of an acid, its Ka?

A

Measure of the degree of ionisation/dissociation that occurs in solution

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39
Q

A high Ka value means…?

A

Strong acid

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40
Q

A low Ka value means…?

A

Weak acid

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41
Q

Which three strong acids should you know?

A
  • HNO3
  • H2SO4
  • HCl
42
Q

Which weak acid should you know?

A

Carboxylic acids

43
Q

Why is pKa used sometimes when plotting a graph instead of Ka?

A

When there is a large range of Ka values

44
Q

pKa = …?

A

-log(base 10)Ka

45
Q

A high pKa value means…?

A

weak acid

46
Q

A low pKa value means…?

A

strong acid

47
Q

pKw = …?

A

-log(base 10)Kw

48
Q

Write the dissociation equation and Ka equation for the weak acid HA

A

HA <=> H+ + A-

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

49
Q

What are the assumptions that we make for when doing calculations with the Ka of weak acids?

A

For Ka = [H+][A-]/[HA]

  1. Since acid is only slightly ionised/dissociated we assume that the equilibrium concentration of acid [HA] is equal to the original acid concentration
  2. We assume equilibrium [H+] = [A-] so we can rewrite it as [H+]^2. This is because we ignore any H+ present int he solution due to the dissociation of water molecules, which is small
50
Q

Write the Ka equation for the weak acid HA after we have followed the assumptions

A

Ka = [H+]^2/ acid concentration

51
Q

How would the pH of a strong acid change if you diluted it 10 times?

A

pH would increase by 1 unit

52
Q

How would the pH of a strong acid change if you diluted it 100 times?

A

pH would increase by 2 units

53
Q

How would the pH of a strong acid change if you diluted it 1000 times?

A

pH would increase by 3 units

54
Q

How would the pH of a weak acid change if you diluted it 1000 times?

A

pH would increase by 1.5 units

55
Q

How would the pH of a weak acid change if you diluted it 100 times?

A

pH would increase by less than 1 unit

56
Q

How would the pH of a weak acid change if you diluted it 10 times?

A

pH would increase by less than 0.5 units

57
Q

Why is the change in pH different for weak acids versus strong acids?

A

Weak acids do not fully dissociate in solution (you have an equilibria) when water is added the equilibrium shifts to the right to oppose the change. Whilst there is a higher fraction of the acid dissociated this is countered by the dilution of the dissociated products.

58
Q

Can pH be negative?

A

Yes

59
Q

What are equimolar solutions and how can they be used?

A

Equimolar solutions are solutions that contain the same number of moles. They can be used to determine whether a substance is an acid, base or salt and whether it is strong or weak.

60
Q

How do you determine the pH of solutions?

A
  • Use a pH meter to measure the pH of solutions
  • Calibrate it first by placing the probe into deionised water, allowing it to settle, and then adjusting the reading so it reads 7.0. Do the same with a standard solution of pH 4 and pH 10 ensuring that you rinse the probe with deionised water between readings
61
Q

How do you prove that your weak acid pH decreases by 0.5 as you dilute the acid by factors of 10, 100 and 1000?

A

You rearrange your Ka equation:

[H+] = root(Ka[acid])

When diluting a weak acid the acid concentration decreases by 10 fold so you sub the values into the equation

62
Q

How do you produce a titration curve?

A

Measure the pH change of a solution of acid or alkali as alkali or acid is added to it

63
Q

Describe the shape of a titration curve for a strong acid-strong base neutralisation reaction

A

Plot of pH against volume of strong base added (could be volume of strong acid added)

Initial pH = pH of acid
Final pH = pH of base

pH gradually changes at first
pH rapidly changes between pH 3.5 and 10.5 -> the curve is almost vertical in this region
One drop of alkali causes a pH change of 7 units

Equivalence point = pH 7

64
Q

Describe the shape of a titration curve for a strong acid-weak base neutralisation reaction

A

Plot of pH against volume of weak base added (could be volume of weak acid added)

Initial pH = pH of strong acid
Final pH = pH of weak base -> Determine using Kb (probably would not have to do this), normally between 9 or 10

pH gradually changes at first
pH rapidly changes between pH 3.5 and 7.0 -> the curve is almost vertical in this region
One drop of alkali causes a pH change of 3.5 units

Equivalence point < 7

This is because your weak base has a strong conjugate acid

65
Q

Describe the shape of a titration curve for a weak acid-strong base neutralisation reaction

A

Plot of pH against volume of strong base added (could be volume of weak acid added)

Initial pH = pH of acid -> you can determine this using Ka and concentration
Final pH = pH of strong base

pH changes more rapidly than with SA-SB
pH rapidly changes between pH 7 and 10.5 -> the curve is almost vertical in this region
One drop of alkali causes a pH change of 3.5 units

Equivalence point > 7

This is because your weak acid has a strong conjugate base

66
Q

Describe the shape of a titration curve for a weak acid-weak base neutralisation reaction

A

Plot of pH against volume of weak base added (could be volume of weak acid added)

pH initially changes in the same way as with S.B - W.A
pH does not change rapidly at any point so there is no vertical section of the graph

Equivalence point = depends on the relative strength of the acid and base but will be around 7

67
Q

Definition of equivalence point

A

Point at which moles of H+ = moles of OH-

68
Q

What does the volume of acid/alkali needed to be added to reach the equivalence point depend on?

A

Concentrations of your acid + alkali and the volume of whichever one is getting the solution added to it

69
Q

What are the differences between the strong acid-strong base titration curve and the strong base-weak acid titration curve?

A

W.A - S.B
- Higher equivalence point
- Higher starting pH
- Graph is steeper initially as you add alkali
- The steep part is shorter comparative to S.A - S.B

70
Q

When is the half equivalence point?

A

When you have added half the volume of acid/alkali compared to the equivalence point

71
Q

For any weak acid, what is true at the equivalence point? Why is this true?

A

pH = pKa

Because:

HA <=> [H+][A-]/[HA]

At the half equivalence point only half of the acid has been neutralised by the strong base, half the molecules of HA have dissociated and been neutralised so [A-] = [HA]

Therefore:

Ka = [H+] at half equivalence point

72
Q

Why do suitable indicators need to be chosen for neutralisation reactions?

A

As they can only indicate a pH within a certain range

Range of indicator needs to lie completely within the range of the vertical section of the titration curve

73
Q

How do you determine the pH at the equivalence point?

A
  • Determine moles of your acid and base
  • Minus the smaller moles from the larger (which one is in excess)
  • Turn moles back to concentration
  • If [H+] then -log(base 10)[H+]
  • If [OH-] then Kw to find [H+] then -log(base 10)[H+]
74
Q

In what titration curve and which part can buffer behaviour be observed?

A

weak acid - strong base

  • In the initial section of the curve
75
Q

What is the definition of a buffer solution?

A

A solution that resists changes in pH when small amounts of acid or alkali are added to it

76
Q

What are the two ways that buffer solution can be formed?

A
  • Partial neutralisation
  • Weak acid/base and conjugate base/acid salt
77
Q

How is an acidic buffer solution made?

A

Mixture of weak acid and salt of its conjugate base

e.g CH3COOH and CH3COONa

78
Q

How is a basic buffer solution made?

A

Mixture of weak base and salt of its conjugate acid

e.g NH3 and NH4Cl

79
Q

Describe how a basic buffer solution works?

A

1st dissociation:
Weak acid or weak base that dissociates partially into its ions

2nd dissociation:
The conjugate base/acid salt that dissociates fully into its ions

  • The buffer solution will contain a high concentration of the conjugate base/acid as the salt from the 2nd dissociation dissociates fully
  • This means that the POE will lie very far to the left in the first dissociation
    -Therefore, you have a high concentration of your undissociated weak acid/base and a relatively low concentration of H+ ions

For any buffer solution there is:
- A RESERVOIR of UNDISSOCIATED WEAK ACID/BASE MOLECULES
- A RESERVOIR OF CONJUGATE BASE/ACID IONS FROM THE DISSOLVED SALT

80
Q

For a basic buffer, describe how the buffer resists changes in pH when a small amount of acid (H+) or alkali (OH-) is added?

A

H+:
Your weak base is partially dissociated in water
The H+ ions react with the OH- ions, low concentration of these, forming water. As you have a reservoir of NH3 and H2O this change is reversed by equilibrium 1 shifting to the right which replaces the OH- ions that reacted, keeping the pH almost constant

OH-:
Your weak base is partially dissociated in water
The OH - ions react with the large reservoir of NH4+ ions reforming your original products, so equilibrium shifts to the left. Your OH- concentration hardly changes so your pH almost remains constant

81
Q

For an acid buffer, describe how the buffer resists changes in pH when a small amount of acid (H+) or alkali (OH-) is added?

A

H+:
Your weak acid is partially dissociated in water
The H+ ions react with the reservoir of CH3COO- ions causing equilibrium 1 to shift to the left
Your H+ ions concentration hardly changes and therefore, the pH remains almost constant

OH-:
Your weak acid is partially dissociated in water
The OH- ions react with the low concentration of H+ ions present. Equilibrium 1 then shifts to the right to replace the the H+ ions which reacted, can do this because of the reservoir of CH3COOH, therefore the pH remains almost constant

82
Q

Definition of an acidic buffer?

A

A solution that resists changes in pH to keep the solution below pH 7

83
Q

Definition of a basic buffer?

A

A solution that resists changes in pH to keep the solution above pH 7

84
Q

What two ways can a buffer solution be made by partial neutralisation?

A

Weak base + strong acid
Weak acid + strong base

85
Q

How does partial neutralisation work?

A

Partial neutralisation is when your weak acid/base is partially neutralised by a strong base/acid.

The weak acid/base is only partially neutralised because it is in EXCESS

Therefore, all the strong base/acid reacts to form salt + water but there is STILL SOME WEAK ACID/BASE LEFT IN SOLUTION

Therefore, you have your weak acid/base and their conjugate base/acid salt. The salt fully dissociates forming a reservoir of the conjugate base/acid. This causes the POE of the weak acid/base to shift very far to the left so there is a reservoir of undissociated weak acid/base

86
Q

How does buffer action occur in titration curves?

A

Partial neutralisation:

When you have a weak acid/base and strong base/acid initially/finally in the curve you have excess of your weak acid/base and so you have a buffer solution

87
Q

Do concentrations and volumes of substances matter when making a buffer from weak acid/base and conjugate base/acid salt?

A

No

88
Q

Do concentrations and volumes of substances matter when making a buffer from partial neutralisation?

A

Yes

89
Q

Describe how to determine pH of a buffer solution formed from partial neutralisation

A

Partial Neutralisation:
1. Calculate moles of acid/alkali
2. Calculate moles of salt
3. As salt is limiting reagent, moles of salt = conjugate base/acid salt
4. Calculate concentration of acid/base
5. Calculate concentration of conjugate base/acid salt
6. Ka = [H+][A-] -> salt concentration/ [HA] -> acid/base concentration
7. Rearrange for H+
8. pH

90
Q

Describe how to determine pH of a buffer solution formed from the normal method

A
  1. . Ka = [H+][A-] -> salt concentration/ [HA] -> acid/base concentration
  2. Rearrange for H+
  3. pH
91
Q

What are the assumptions we make during buffer calculations?

A
  1. [HA] - the original acid concentration since very little is dissociated
  2. [A-] - the salt concentration since the salt is fully ionised and a very small amount of A- comes from the acid
92
Q

State the Henderson Hasselbalch equation

A

pH = pKa + log(base 10) [A-]/[HA]

93
Q

Definition of enthalpy change of neutralisation?

A

The enthalpy change when one mole of water is formed when solutions of acid and alkali are reacted together

94
Q

Are enthalpies of neutralisation exothermic/negative or endothermic/positive?

A

Exothermic/negative - more bonds being made than broken

95
Q

For a strong acid - strong base neutralisation what is always the value for the enthalpy change of neutralisation?

A

-57 kjmol^-1

96
Q

Will weak acid/base neutralisations be more or less exothermic than strong acid/base neutralisations? Why?

A

Less exothermic - This is because weak acids/bases do not fully dissociate into their ions this means that energy is used ionising the compounds first

97
Q

The more weak the acid/base the … the value of the enthalpy of neutralisation?

A

less negative/exothermic - more energy required to ionise the acid/base

98
Q

What pH does the blood need to be kept around?

A

pH 7.4

99
Q

What buffer system is present in the blood

A

Carbonic acid - hydrogencarbonate buffer system

100
Q

How does the carbonic acid - hydrogencarbonate buffer system in the blood behave as a buffer?

A

H2CO3 ions dissociate following the equation:

H2CO3 <=> H+ + HCO3-

There are reservoirs of H2CO3 and HCO3-

[H+] rises:
When the concentration of H+ ions rises in the blood the reservoir of HCO3- ions reacts with the H+ ions and the equilibrium will shift to the left reducing the H+ ion concentration to almost its original value. This stops the pH of the blood form dropping

[H+] falls:
When [H+] falls the reservoir of H2CO3 molecules oppose the change by dissociating, this means the equilibrium will shift to the right increasing the [H+] to its original value

101
Q

How are the levels of H2CO3 controlled in the blood?

A

By respiration: as this involves breathing out CO2

H2CO3 <=> H2O + CO2

102
Q

How are the levels of HCO3- controlled in the blood?

A

By the kidneys: excess is excreted out in urine

Chemistry A - Level (26 decks)

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