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A level Chemistry > Unit 3.9 - Acid-base equilibria > Flashcards

Unit 3.9 - Acid-base equilibria Flashcards

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1
Q

What is the Lowry-Bronsted theory of acids limited to?

A

Aqueous solutions

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2
Q

Acid

A

An ion or molecule which can donate a proton (H+ donor)

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3
Q

Base

A

An ion or molecule which can accept a proton (H+ acceptor)

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4
Q

What do acids do during a reaction?

A

Donate a proton to become a base

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5
Q

What do bases do during a reaction?

A

Accept a proton to become an acid

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6
Q

Equation to represent an acid-base reaction

A

A —><— B+ + H+

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7
Q

In the reaction between water and hydrochloric acid, what acts as the acid and what acts as the base and why?

A

Water has 2 lone pairs of electrons that can easily accept the proton from HCl to form a coordinate bond in H3O+ = behaves as a base
HCl behaves as an acid since it’s losing its proton to the H2O

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8
Q

Neutralisation reaction

A

One substance donates a proton to another substance - often water

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9
Q

Equation to represent a neutralisation reaction

A

Acid 1 + Base 2 —><— Base 1 + Acid 2

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10
Q

What type of pairs form during neutralisation reactions?

A

Conjugate acid-base pairs

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11
Q

Conjugate acid-base pairs

A

An acid and a base which differ only by he presence or absence of a proton

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12
Q

An acid and a base which differ only by he presence or absence of a proton

A

Conjugate acid-base pairs

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13
Q

Alkali

A

A base that’s soluble in water (most contain OH-)

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14
Q

A base that’s soluble in water

A

Alkali

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15
Q

Show the equation for the neutralisation reaction between hydrochloric acid and sodium hydroxide + explain

A

HCl (aq) + NaOH (aq) —> NaCl (aq) + H2O (l)
H+ + Cl- + OH- —> Na+ + Cl- + H2O
(Cancelling similar elements)
H+ (aq) + OH- (aq) —> H2O (l)

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16
Q

What do we assume with strong acids and bases?

A

That they completely dissociated in water into their ions

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17
Q

Problem with using Ka to indicate the acidic strength + what is used instead

A

Numbers are difficult to handle
pH (easier to follow)

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18
Q

Ka

A

Acid dissociation constant

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19
Q

Acid dissociation constant

A

Ka

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20
Q

What is pH a measure of?

A

The [H+(aq)] concentration

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21
Q

pH equation

A

pH = -log10[H+(aq)]

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22
Q

Hydrogen ion concentration equation

A

[H+] = 10^-pH

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23
Q

How do we use [H+] = 10^-pH ?

A

Shift log

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24
Q

pH scale?

A

Simple and widely applicable method for measuring the acidity/alkalinity of an aqueous solution

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25
Q

Can we get negative pH’s?

A

Yes

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26
Q

What would a negative pH indicate?

A

Highly strong acid, high concentration of H+ ions

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27
Q

Ways of measuring pH

A

pH meter
Universal indicator

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28
Q

Pros and cons of using a pH meter to measure pH

A

Pros = +-0.01, accurate
Cons = needs to be calibrated against a solution of known pH

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29
Q

Pros and cons of using a universal indicator to measure pH

A

Pros = quick, convenient, paper or solution, cheap
Cons = less accurate

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30
Q

Good use for universal indicators to measure pH

A

Soil pH

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31
Q

What temperature is pH normally quoted for?

A

25 degrees Celsius

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32
Q

What is pH dependent on?

A

Concentrate
Temperature
(Of the substance)

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33
Q

What type of acids can their pH be calculated directly and why?

A

Strong (if the concentrations of the solutions are known)
They are assumed to be totally dissociated into their ions

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34
Q

The pH of what type of acids can’t be found directly and why?

A

Weak acids
Only a small fraction of the molecules are dissociated into ions

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35
Q

Under which circumstances can the pH of strong acids be calculated directly?

A

If the concentrations of the solutions are known

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36
Q

What do we need to remember to do with Diprotic acids?

A

Multiply the concentration with 2 for calculating pH

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37
Q

How do we work out the concentration of a specific substance from pH?

A

[H+] = 10^-pH
Symbol equation to represent the dissociation
Molar ratio

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38
Q

Strong cid

A

One that almost totally dissociates into its ions in solution

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39
Q

Example of a strong acid

A

Hydrochloric acid

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40
Q

HCl dissociating equation

A

HCl (aq) —> H+ (aq) + Cl-(aq)

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41
Q

What do acids dissociate into?

A

A proton and an anion

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42
Q

Weak acid

A

One that only partially dissociates into its ions in solution

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43
Q

Example of a weak acid

A

Ethanoic acid (all organic acids tend to be weak)

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44
Q

What type of acids all tend to be weak?

A

Ethanoic acids

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45
Q

Ethanoic acid dissociation equation

A

CH3COOH (aq) —><— CH3COO- (aq) + H+ (aq)

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46
Q

Acid dissociation constant

A

Ka

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47
Q

Ka

A

Acid dissection constant

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48
Q

What type of process is it when acid dissociates in solution?

A

An equilibrium process

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49
Q

What does each acid dissociation reaction have?

A

An equilibrium constant, Ka

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50
Q

How would you write out Ka for this reaction?
HA (aq) —> H+ (aq) + A- (aq)

A

Ka = [H+][A-]
————
[HA]

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51
Q

How do we write out the expression for Ka (the acid dissociation constant)?

A

Just as we were with KC with products over reactants

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52
Q

Ka unit

A

moldm^-3 (every time)

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53
Q

Why do stronger acids dissociate more?

A

The more dissociated an acid is, the more H+ ions and ions there will be, so the stronger the acid

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54
Q

Value of Ka for a weak acid

A

Low

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55
Q

Value of Ka for a strong acid

A

High

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56
Q

What does a high value of Ka indicate?

A

A strong acid

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57
Q

pKa expression

A

pKa = -log10Ka

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58
Q

Ka expression

A

Ka = 10^-pKa

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59
Q

What does a higher value of pKa indicate?

A

A weaker acid

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60
Q

What does a lower value of pKa indicate?

A

A strong acid

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61
Q

What type of acids have the highest pKa values?

A

Weak acids

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62
Q

Explain why HNO3 is a strong acid based on its Ka and pKa values

A

Very high Ka value
Low pKa value

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63
Q

Proof that HNO3 is a very strong acid

A

Even silver can react with it, and silver is usually inert

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64
Q

What can water be described as?

A

A weak electrolyte

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65
Q

Why is water described as a weak electrolyte?

A

It partially dislocates into its ions in solution

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66
Q

Equation for water dissociating into its ions in solution

A

H2O (l) —><— H+ (aq) + OH- (aq)

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67
Q

Enthalpy change for the dissociation of water in solution

A

Positive

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68
Q

Kw expression

A

Kw = [H+(aq)][OH-(aq)]

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69
Q

How is the Kw expression obtained?

A

Applying the equilibrium law and assuming that the concentration of water is effectively constant
Kc = [H+(aq)][OH-(aq)]
————————
[H2O (l)]
(Eliminating H2O)
Kw = [H+(aq)][OH-(aq)]

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70
Q

Kw

A

The ionic product of water

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71
Q

The ionic product of water symbol

A

Kw

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72
Q

Value of Kw at 25 Celsius

A

1x10^-12mol2dm^-6 (in db)

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73
Q

Relationship between the H+ (aq) and OH- (aq) ions in pure water

A

Equal concentrations

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74
Q

When are there equal concentrations of H+ and OH- ions?

A

When pure water dissociates

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75
Q

Concentration of H+ or OH- ions from dissociated pure water + explanation

A

1x10^-7 (for both - concentrations are equal)
Sqrt of Kw

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76
Q

What type of process it the self ionisation of water?

A

Endothermic

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77
Q

What happens to the value of Kw as the temperature increases?

A

Increases

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78
Q

When is Kw a constant value?

A

At a particular temperature, even though the values of H+ and OH- concentrations may not be equal

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79
Q

Does altering the temperature alter Kw? How?

A

Higher temperature = higher Kw

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80
Q

What happens to strong bases in water?

A

Completely dissociated

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81
Q

Example of a strong base

A

NaOH

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82
Q

Equation for NaOH (a strong base) completely dissociating in water

A

NaOH (aq) —> Na+ (aq) + OH- (aq)

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83
Q

Why does water act as both an acid and a weak base?

A

It’s amphoteric

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84
Q

When does the concentration of OH- not equal the concentration of H+?

A

For a strong base dissolving in water

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85
Q

Why does the concentration of OH= ions not equal the concentration of H+ ions when a strong base dissociates in water?

A

Due to the ions from the base

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86
Q

The concentration of which ion is equal to the concentration of the base when a strong base dissociates and why?

A

The concentration of OH- ions
As the base is fully dissociated

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87
Q

How can we work out pH for a strong base dissociating in water since the concentration of H+ is not equal to the concentration of OH- ions?

A

Concentration of OH- is equal to the concentration of the base
If we re given this and Kw, we can work out [H+] and thus pH

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88
Q

What might we have to do when attempting to work out the pH of a strong base?

A

Work out how it dissociates

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89
Q

What’s important to note about conjugate acid-base pairs?

A

Has to be in the order “acid-base”
The others would be base-acid conjugates

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90
Q

How can we work out [H+] and [OH-] with pure water?

A

Square root of Kw

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91
Q

What do we do if we asked to work out the pH of a final solution, and we only have some values for the initial volumes and concentrations?

A

C1V1= C2V2

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92
Q

Why is the ionisation of water and endothermic process?

A

As the value of Kw increases as the temperature increases

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93
Q

How do you write an expression for Ka the acid dissociation constant?

A

[H+][anion-]
——————
[acid]

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94
Q

What do you do to get the concentration of a strong base with (OH)2?

A

Divide the concentration by two (not multiply)

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95
Q

What is Ka specific for?

A

A specific temperature

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96
Q

What is the assumption made for weak acids?

A

[H+] = [A-]

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97
Q

What is the assumption [H+] = [A-] made for?

A

Weak acids

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98
Q

[A-]

A

Concentration of the anion of the acid

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99
Q

What can the equilibrium law only be applied to?

A

Weak electrolytes

100
Q

Example of a weak electrolyte

A

Ethanoic acids

101
Q

Why can the equilibrium law only be applied to weak electrolytes?

A

If the electrolytes were stronger, they would fully dissociate, so they would be no equilibrium symbol

102
Q

How do we calculate the pH of a weak acid? Explain the process

A

Ka = [H+][A-]
————
[acid]

Since [H+] = [A-], we can think of it as being [H+]^2
So, [H+] = square root of Ka x [acid]
Then use the normal pH expression

103
Q

How can we consider [H+] in weak acids and why?

A

[H+] since [H+] = [A-] on the top of the Ka expression

104
Q

Steps for calculating the pH of a weak acid with Ka

A
  1. Balanced symbol equation
  2. Equation for Ka
  3. Insert values
105
Q

Why can we assume that the concentration of a weak acid is the same at equilibrium as when undissociated?

A

Because the degree of dissociation is so small

106
Q

What can we assume about the concentration of a weak acid at equilibrium compared to when undissociated and why?

A

The same
The degree of dissociation is very small

107
Q

How would we calculate Ka from pH?

A

[H+] = 10^-pH
Ka = [H+]^2
———
[acid]

108
Q

Buffer solution

A

A solution whose pH does not change to any appreciable extent on addition of small amounts of acid or alkali

109
Q

What does a typical buffer solution contain?

A

A weak acid and one of its salts (an acid and its conjugate base)

110
Q

If our buffers are made up of a weak acid and one of its salts, what is largely dissociated into its ions and what’s not?

A

Acid largely undissociated into its ions (weak)
Salt dissociated into its ions (ionic)

111
Q

Example of a weak acid and one of its salts (its conjugate base) that make up a buffer solution

A

Ethanoic acid
Sodium ethanoate

112
Q

How do we know if a buffer is acidic?

A

Negative pH

113
Q

2 processes at work within a buffer solution

A

Partial dissociation of the acid that makes the buffer into its ions
Complete dissociation of the salt of the acid

114
Q

In our example of ethanoic cid and sodium ethanoate making up a buffer solution, what is partially dissociating into its ions and what is completely dissociating into its ions?

A

Partial dissociation of ethanoic acid
Complete dissociation of sodium ethanoate

115
Q

Partial dissociation of ethanoic acid into its ions equation

A

CH3COOH —><— CH3COO- + H+

116
Q

What type of reactions are partial dissociations and what does this mean?

A

Equilibrium reactions - can manipulate the position

117
Q

Equation for the complete dissociation of sodium ethanoate into ions

A

CH3COONa —> CH3COO- + H+

118
Q

What is the same about the processes of the partial dissociation of the acid and the complete dissociation of the salt of the acid at work in a buffer solution?

A

Same anion given out in both processes

119
Q

What is different about the anion given out in the processes of the partial dissociation of the acid and the complete dissociation of the salt in a buffer solution?

A

Much more of the anion is given out during the complete dissociation of the salt

120
Q

Which equilibrium law can we apply with buffers too?

A

The same equation for the acid dissociation constant

121
Q

Acid dissociation constant equation

A

Ka = [A-][H+]
————
[acid]

122
Q

Which concentrations are assumed to be the same in a buffer made up of ethanoic acid and sodium ethanoate solution?

A

The concentration of the ethanoic acid and the undissociated ethanoic acid
The concentration of the ethanoate ions and the sodium ethanoate

123
Q

The concentrations of what are large and the concentrations of what are small in a buffer solution made up of ethanoic acid and sodium ethanoate?

A

Large concentrations —> ethanoic acid and ethanoate ions
Small concentrations —> hydrogen ions

124
Q

How is the [CH3COO-]/[CH3COOH] ratio kept steady in a buffer solution of ethanoic acid and sodium ethanoate?

A

The concentrations of the ethanoic acid and ethanoate ions are large and that of the hydrogen ions is small

125
Q

Expression for [H+] in a buffer solution and explain exactly which values would be inputted

A

[H+] = Ka x [acid]/[anion]

[acid] = undissolved
[anion] = assumed to be equal to the concentration of the salt

126
Q

2 assumptions in buffer solutions

A

[H+] is not equal to [A-] (much higher [A-])
[A-] = [salt]

127
Q

Explain what happens when acid is added to a buffer solution

A

The system wishes to decrease the concentration of hydrogen ions using Le Chatelier’s Principle
the position of equilibrium moves to the left
Thus the pH remains constant

128
Q

Explain how the concentration of hydrogen ions in a buffer solution is decreased when the position of equilibrium moves

A

The negative ions combine with the hydrogen ions and the majority of the added hydrogen ions are removed

129
Q

Why is it possible for a buffer solution to remain at a constant pH when acid is added?

A

Because the negative ion concentration is large

130
Q

Which equation do we write out to explain how buffer solutions work? Why?

A

The partial dissociation of the acid
This is the equilibrium process

131
Q

Explain how buffer solutions work when alkali is added to the system

A

The hydroxide ions react with the hydrogen ions to form water
The system wishes to increase the hydrogen ion concentration
The position of equilibrium moves to the right

132
Q

Why does the concentration of hydrogen ions in a buffer solution decrease when an alkali is added?

A

The hydroxide ions react with the hydrogen ions to form water

133
Q

Explain what happens in the example of the ethanoic acid and sodium ethanoate buffer solution when acid is added

A

The ethanoate ions combine with the hydrogen ions and the majority of the added hydrogen ions are removed

134
Q

Explain what happens in the example of the ethanoic acid and sodium ethanoate buffer solution when alkali is added

A

The ethanoic acid molecules dissociate to restore the concentration of the hydrogen ions

135
Q

Can buffers reach a breaking point? How and why?

A

Yes
If we keep adding more acid
The ions would eventually run out

136
Q

What are buffer solutions important for in the real world?

A

Biological systems
Industrial processes

137
Q

Example of a biological system that buffer solutions are important in?

A

Maintaining the pH of blood in the range 7.35 to 7.5 (weak alkali)

138
Q

how is the pH of blood maintained?

A

By hydrogencarbonate and phosphate buffers

139
Q

What would happen if the pH of blood were to fall outside the optimum range? Explain

A

Serious health implications
E.g - acidosis, where enzymes denature. Can occur during an athsma attack due to the inability to get rid of CO2

140
Q

What do many processed foods contain and why?

A

Buffers
So that they can safely be eaten without changing the pH of the blood

141
Q

Example of an industrial process that buffers are important in + explain

A

Fermentation
Keeps enzyme pH at optimum

142
Q

What are basic buffers made up of?

A

A weak base and a salt

143
Q

Example of a basic buffer solution that maintains an alkaline pH

A

Ammonium chloride and ammonia solution

144
Q

What happens in a basic buffer solution of ammonium chloride and ammonia solution?

A

The ammonium chloride dissociated completely, releasing all the ammonium ions

145
Q

Equation for ammonium chloride dissociating completely to release all the ammonium ions

A

NH4Cl (aq) —> NH4+ (aq) + Cl- (aq)

146
Q

Key equilibrium in the basic buffer made up of ammonium chloride and ammonia solution

A

NH4+ —><— NH3 + H+

147
Q

Explain what happen when we add an acid to a basic buffer

A

Causes the position of equilibrium to shift to the left
This removes additional H+ ions provided by the acid

148
Q

Explain what happens when we add a base to a basic buffer

A

Adding a base removes the H+ ions which causes the position of equilibrium to shift to the right to make more H+ ions

149
Q

How do we explain the new pH of a solution when acids/bases are added to ammonia?

A

Use the key equilibrium of ammonia
NH4 —><— NH3 + H+

150
Q

What do we use to prove that the pH of water is 7?

A

The fact that [H+] = [OH-] in pure water
sqrt kw = [H+] and use this in pH equation

151
Q

What does the pH of a salt solution depend on?

A

The reactants used to make the salt (i.e - whether the acid or alkali are strong or weak)

152
Q

What is it due to if the pH of a salt solution is not 7?

A

Hydrolysis

153
Q

What does hydrolysis cause in a salt solution?

A

The pH is no longer 7

154
Q

Nature of a solution made from a strong acid and a strong base

A

Neutral, pH 7

155
Q

Nature of a solution made from weak acid and strong base

A

Alkaline, pH above 7

156
Q

Nature of a solution made from strong acid and weak base

A

Acidic, pH less than 7

157
Q

Example of a salt consisting of a strong acid and strong base

A

NaCl

158
Q

Example of a salt consisting of a weak acid and a strong base

A

CH3COONa

159
Q

Example of a salt consisting of a strong acid and a weak base

A

NH4Cl

160
Q

In which situation is there no hydrolysis in a salt?

A

In a salt consisting of a strong acid and a strong base

161
Q

What happens as a result of there being no hydrolysis when a strong acid and a strong base react?

A

The number of hydrogen and hydroxide ions remains constant and the pH is 7

162
Q

In what kind of salts does hydrolysis occur?

A

Salts made from…
Weak acid + strong base
Or
Strong acid + weak base

163
Q

Explain with the use of equations the hydrolysis of sodium ethanoate with water (a salt made from a weak acid and a strong base)

A

When sodium ethanoate is dissolved in water it dissociates into its ions
CH3COONa —> CH3COO- + Na+
The anion of the salt then gets hydrolysed as it reacts with water molecules as follows
CH3COO- + H2O —><— CH3COOH + OH-

164
Q

Why is the final solution alkaline when hydrolysis of a salt made from a weak acid and a strong base occurs?

A

From the reaction of the anion with water….
CH3COO- + H2O —><— CH3COOH + OH-
The ethanoate ion is acting as a base and the water as an acid. The solution is alkaline as OH- has been released, which is a strong base. It outweighs the effect of the weak acid.

165
Q

Explain the hydrolysis of ammonium chloride with water (a salt made from a strong acid and a weak base)

A

When ammonium chloride dissolves in water, it dissociates fully into its ions
NH4Cl —> NH4+ + Cl-
The ammonium ions are hydrolysed by the after as follows:
NH4+ + H2O —><— NH3 + H3O+

166
Q

Why is the final solution acidic when hydrolysis of a salt made from a strong acid and a weak base occurs?

A

From the reaction of the anion with water molecules:
NH4+ + H2O —><— NH3 + H3O+
The ammonium ion is acting as the acid and the water as a base. The solution is acidic due to the formation of hydrogen ions during the hydrolysis. H3O+ being a sting acid outweighs the weak base ammonia

167
Q

What is an indicator?

A

A substance which is used to measure the end points in titrations

168
Q

What are indicators usually?

A

Weak acids which change their colour depending on the pH of the solution in which they are present

169
Q

Equilibrium equation of methyl orange

A

Hme (aq) —><— H+ (aq) + Me- (aq)
Red. Yellow.
Acid solution. Alkaline solution.

170
Q

What does the equilibrium of the indicator methyl orange tell us?

A

Undissociated acid = red
Dissociated acid = yellow

171
Q

When will a methyl orange indicator turn colour from yellow to red?

A

When the pH is 3.7

172
Q

Methyl orange indicator colour change when the pH is 3.7

A

From yellow to red

173
Q

What colour will methyl orange appear at the exact end point?

A

Orange

174
Q

When will the methyl orange indicator appear orange?

A

At the exact end point

175
Q

Why does methyl orange appear orange at its exact end point?

A

It’s a mixture of red and yellow

176
Q

What is the pH range in which an indicator changes its colour over?

A

2 units

177
Q

pH range and colour range of the methyl orange indicator

A

2.7 —> 4.7
Red —> yellow

178
Q

pH range and colour range of Congo red indicator

A

3.0 —> 5.0
Blue —> red

179
Q

pH range and colour range of litmus indiciator

A

5.0 —> 8.0
Red —> blue

180
Q

pH range and colour range of phenolphthalein indicator

A

8.4 —> 10.4
Colourless —> red

181
Q

What happens during a titration as the alkali is added to the acid?

A

Change in pH

182
Q

What must an indicator do at the end point of a titration is it is to be effective?

A

It must change colour quickly

183
Q

List the reactions that can be followed successfully using an indicator to determine the end point

A
  1. Strong acid v.s strong alkali
  2. Strong acid v.s weak alkali
  3. Weak acid v.s strong alkali
184
Q

Weak alkali example

A

NH3

185
Q

Weak acid example

A

CH3COOH

186
Q

Strong alkali example

A

NaOH

187
Q

Strong acid example

A

HCl

188
Q

In what type of reaction is it not possible to determine its end point?

A

Between a weak acid and a weak alkali

189
Q

Why is it not possible o determine the end point of a reaction between a weak acid and a weak alkali?

A

The change in pH at the end point is small and gradual

190
Q

What experiment must we do in order to be able to plot a graph of pH v.s volume?

A

If we have acid in a conical flask and base in a burette (or vice versa) and measure the pH for every 5cm^3 of base added, a graph of pH v.s volume of alkali added to a fixed acid volume can be plotted

191
Q

Equation for the neutralisation reaction that occurs when an acid and base react

A

H+ + OH- —> H2O

192
Q

Is the general shape of a titration curve when an acid and a base react the same?

A

Yes

193
Q

Describe the general shape of a titration curve when an alkali is added to an acid

A

Slow rise
Vertical rise
Slow rise

194
Q

Describe the general shape of a titration curve when an acid is added to an alkali

A

Slow drop
Vertical drop
Slow drop

195
Q

Describe the pH change around the equivalence point

A

Large change in pH for only a small amount of titrant being added

196
Q

Equivalence point

A

The point when the number of moles of acid added is equal to the number of moles of alkali

197
Q

End point of a titration

A

The point in a titration at which the indicator changes colour

198
Q

What type of end point do we need in a titration?

A

One that changes quickly at the equivalence point

199
Q

What do we need for a colour change to occur at the equivalence point?

A

Only 1 drop of titrant

200
Q

At which pH will the equivalence point be adjacent to in a strong acid-strong base titration curve?

A

7

201
Q

What do different equivalence points depend on?

A

There being different acids or bases

202
Q

What happens at the equivalence point in terms of pH?

A

Sudden change

203
Q

In the case of adding HCl to NaOH, what is the equivalence point representing?

A

It’s the vertical point where nHCl = nNaOH

204
Q

How do we work out the pH of the solution using a titration curve?

A

We take the mean number (the centre point) of the vertical line

205
Q

How do we work out which indicator to use for a reaction using a titration curve?

A

The pH of the indicator must fit perfectly within the vertical region of the titration plot

206
Q

If we used an indicator that didn’t fit perfectly within the vertical region of the titration plot, what would happen?

A

Outside of the region would give an undershoot (too far before) or overshoot (too far after the vertical region)

207
Q

How would we work out the initial pH to plot if we had to draw a titration curve for the opposite reaction?

A

From the concentrations given

208
Q

Compare how a titration curve for adding strong acid to strong alkali and adding strong alkali to strong acid would be different

A

Acid to alkali —>slow drop, vertical drop, slow drop
Alkali to acid —> slow rise, vertical rise, slow rise

209
Q

Equation for adding hydrochloric acid to ammonium hydroxide

A

HCl + NH4OH —> NH4Cl + H2O

210
Q

Describe the titration curve for a strong acid and a weak base

A

Sharper decrease in pH but then the gradient flattens off

211
Q

Why does the gradient flatten off on the titration curve of a strong acid and a weak base?

A

Due to buffering

212
Q

Explain how buffering is happening in the reaction between HCL and NH4OH

A

the salt (NH4Cl) and the base (NH4OH) are reacting together to reduce the effect of the change

213
Q

What is it that causes the difference in the shape of the curve with a strong acid-weak base titration curve?

A

Buffering

214
Q

Why is there a really slow change in pH to begin with when reacting a strong acid and a weak base compared to strong acids and bases?

A

There’s no buffering when reacting strong acids and strong bases

215
Q

Why is the equivalence point at a pH that is below 7 when reacting a strong acid a weak base?

A

Since the salt formed is acidic

216
Q

Describe the change in pH when reaching a strong acid and a weak base?

A

Rapid fall of pH to start with but the rate of fall slows down

217
Q

Describe the titration curve when adding a weak acid to a strong base

A

pH rises quite quickly and then flattens off due to buffering
Sharp rise and then the gradient decreases markedly

218
Q

When is the buffering when reacting a strong base and weak acid and why?

A

When the solution is acidic as it is weak

219
Q

Where is the equivalence point when reacting a strong base and a weak acid and why?

A

pH above 7 since this is where neutralisation has occurred and the salt formed is an alkaline salt

220
Q

Why sit here term “neutral point” not appropriate and the term neutralisation isn’t always helpful?

A

Because, for example, when adding a weak acid to a strong base, the salt formed will be alkaline and this is where neutralisation has occurred

221
Q

Describe the titration curve of adding a weak acid to a weak base

A

The titration is rarely performed as there is no clear end-point, it is more of an inflexion at the equivalence point
There is, however, indications of buffering at either side of the equivalence point

222
Q

pH of the equivalence point when adding a weak base and a weak acid

A

7

223
Q

What is often used when reacting a weak acid and a weak base?

A

A pH probe

224
Q

Example of a reaction where there is more than one equivalence point

A

The reaction between ethanedioic acid and sodium hydroxide

225
Q

Equation for the reaction between ethanedioic acid and sodium hydroxide

A

H2C2O2 + 2NaOH —> Na2C2O2 + 2H2O

226
Q

When we have a reaction where there’s more than one equivalence point, what do the different equivalence points represent?

A

When the first and second reactions have completely finished

227
Q

Why does a reaction with ethanedioic acid have 2 equivalence points?

A

Ethanedioic acid is dibasic (or Diprotic) - it has 2 acidic hydrogens

228
Q

How does ethanedioic acid give its different equivalence points?

A

It loses the acidic groups individually

229
Q

What do we have to do if we have 2 different end-points for a reaction?

A

Would have to switch indicators - a double titration

230
Q

Henderson-Hasselbalch equation

A

pH(buffer) = pKa + log[salt]/[acid]

231
Q

What is the Henderson-Hasselbalch equation used for?

A

It’s an alternative way to work out buffer pHs

232
Q

Which radio do we use with the Henderson-Hasselbalch equation?

A

The ratio of salt to acid concentrations

233
Q

What is true if the concentration of the salt is equal to the concentration of the acid (Henderson-Hasselbalch equation)?

A

Since [salt]/[acid] = 1 and log(1) = 0

pHbuffer = pKa

234
Q

When is pHbuffer = pKa true? Why?

A

Half way towards the equivalence point
Half of the acid will have turned into salt

235
Q

Half-equivalence point

A

Where half of the acid has been neutralised

236
Q

Which two things are true at the half-equivalence point?

A

Amount of salt formed = amount of acid remaining
pH = pKa

237
Q

How could we work out Ka using the half-equivalence point?

A

The half-equivalence point would be halfway towards the equivalence point, so half of the volume at the equivalence point
We then extrapolate at this volume to find pH
Here, pH = pKa
Ka = 10^-pKa

238
Q

How do we explain why buffering occurs in the buffer zone?

A

-write the equilibrium between the conjugate acid and base
-explain what happens to the position of equilibrium when acids or bases are added

239
Q

How is the pH of a buffer solution kept constant when acid is added?

A

The anions react with the H+ ions to remove them from the solution to keep the pH constant

240
Q

How do we get the reversible reaction for a basic buffer solution?

A

The salt dissociates completely, releasing ammonium ions which set up the reversible reaction

241
Q

Ka in a titration plot

A

Use the pH at the half value for neutralisation (not half way of the vertical region) in ka = 10^-pH

242
Q

Explain why we get 2 vertical regions on a titration curve of a dibasic acid

A

2 acidic protons = the Ka of each proton is different = 2 vertical regions

243
Q

Explain the location of the 2 vertical regions for a dibasic acid

A

The first vertical region occurs at half the volume of the second vertical region as each occurs after removing the same number of protons

244
Q

What happens when we have an indicator that doesn’t change colour within the vertical region of a titration curve?

A

The indicator will change colour gradually when the acid/base is added

245
Q

Important use for buffer solutions

A

Storing enzymes at constant pH

246
Q

What has to be true about an indicator for it to be suitable to use?

A

Its colour change needs to be complete in the vertical region

A level Chemistry (39 decks)

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