3.9 Acid-base equilibria

Question 1

What is the Lowry-Bronsted theory?

Answer 1

The Lowry-Bronsted theory states that acid-base equilibria involves the transfer of protons between substances and substances can be classified as acids or bases depending on their interaction with protons.

Question 2

Define a Lowry-Bronsted acid and give an example

Answer 2

A Lowry-Bronsted acid is a proton donor.
Example: Ammonium ions (NH4^+)

Question 3

Define a Lowry-Bronsted base and give an example

Answer 3

A Lowry-Bronsted base is a proton acceptor.
Example: Hydroxide ions (OH-)

Question 4

Describe the difference between a strong acid and a weak acid

Answer 4

A strong acid dissociates almost completely in water which means nearly all the H+ ions are released.

A weak acid only partially dissociates in water so only a small number of H+ ions are released.

Question 5

Describe the difference between a strong base and a weak base

Answer 5

A strong base dissociates almost completely in water so nearly all the OH- ions are released.

A weak base only partially dissociates in water so only a small number of OH- ions are released.

Question 6

Give an expression for pH in terms of [H+]

Answer 6

pH= -log10[H+]
Equivalently:
[H+]=10^-pH

Question 7

What is the relationship between pH and hydrogen ion concentration, [H+]?

Answer 7

The pH scale is a measure of hydrogen ion concentration. The lower the pH, the higher the concentration of hydrogen ions.

Question 8

What is the hydrogen ion concentration of a solution of hydrochloric acid which has a pH of 2.0?

Answer 8

[H+] = 10^-pH
= 10^-2
=0.01 mol dm^-3

Question 9

Give examples of strong acids and state the pH range which indicates a strong acid

Answer 9

Examples:
Hydrochloric acid, Sulfuric acid, Nitric acid

pH range of strong acids: 0-3

Question 10

Give examples of weak acids and state the pH range which indicates a weak acid

Answer 10

Examples:
Ethanoic acid, hydrogen sulfide, and organic carboxylic acid

pH range of weak acids: 4 to just below 7

Question 11

Give examples of strong bases and state the pH range which indicates a strong base

Answer 11

Examples:
Sodium hydroxide, potassium hydroxide, calcium hydroxide

pH range of strong bases: 12-14

Question 12

Give examples of weak bases and state the pH range which indicates a weak base

Answer 12

Examples:
Ammonia, methylamine

pH range for weak bases: just above 7 up to 12

Question 13

What is the acid dissociation constant, Ka?

Answer 13

The acid dissociation constant, Ka, is a measure of how strong an acid is in a solution.

Question 14

Give the formula used to calculate Ka for a reaction of the form
HA(aq) <—> H+(aq) + A-(aq)

Answer 14

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

Question 15

What are the units for Ka?

Answer 15

mol dm^-3

Question 16

How does the strength of an acid relate to the value of Ka?

Answer 16

Ka is the equilibrium constant for the dissociation of an acid: HA <—> H+ + A-

The stronger the acid, the further to the right the equilibrium lies so there is a higher concentration of products. This causes Ka to increase.

Question 17

Why is Ka used to find the pH of weak acid?

Answer 17

Weak acids only partially dissociate in water so the concentration of H+ ions is not the same as the acid concentration (as with strong acids). This means the pH cannot be found using [H+] and so Ka is used instead.

Question 18

Give the formula used to find Ka of a weak acid

Answer 18

For a weak acid you can assume that all the H+ ions in solution come from the acid so that [H+]=[A-] and you can assume that [HA] equilibrium = [HA] start

So for a weak acid:
Ka = [H+]^2 / [HA]

Question 19

What is the difference between describing an acid/ base as ‘concentrated’ compared to ‘strong’?

Answer 19

‘Concentrated’ implies there are many moles per dm^3.
‘Strong’ relates to the dissociation of the substance and implies the acid/base almost completely dissociates in water.

Question 20

Give the formulas used to convert between Ka and pKa

Answer 20

pKa = -log10(Ka)
Ka = 10^-pKa

Question 21

What is the pKa of an acid which has a Ka value of 1.60 x 10^-2 mol dm^-3?

Answer 21

pKa = -log10(Ka)
= -log10(1.6 x 10^-2)
= 1.80 ( 3.s.f)

Question 22

Give the equation for the ionic product of water

Answer 22

Kw = [H+][OH-]
Or equivalently:
Kw = [H3O+][OH-]

Question 23

Derive the ionic product of water using the equation for the ionisation of water

Answer 23

In water, the following equilibrium is set up:
H2O <—> H+ + OH-

So, Kc = ([H+][OH-]) / [H2O]. Since [H2O] is very large compared to [H+] and [OH-], [H2O]Kc can be considered to be constant. Then [H2O]Kc = Kw and so Kw = [H+][OH-].

Question 24

What is the pH of pure water at room temperature?

Answer 24

7

Question 25

How does the pH at which water is neutral change as temperature increases?

Answer 25

In the equilibrium of water, the forwards reaction is endothermic and is therefore favoured when the temperature of the water is increased. So, the pH at which water is neutral (when the concentration of H+ and OH- is the same) decreases when the temperature is increased.

Question 26

What is a pH curve?

Answer 26

Graphs which plot pH against volume of acid or base added are called pH curves and can be used to help identify the point of neutralisation of a solution.

Question 27

What is the equivalence point on a pH curve?

Answer 27

The equivalence point is also called the end point and is the point at which the pH curve is vertical. This is the point at which the solution has been neutralised.

Question 28

What is a buffer solution?

Answer 28

A buffer solution is a solution which is able to resist changes in pH when small volumes of acid or base are added.

A buffer solution is commonly formed from a weak acid and its salt or weak base and its salt. This produces a mixture containing H+ ions and a large pool of OH- ions which helps to resist any change in pH.

Question 29

How is an acidic buffer solution made by mixing sodium ethanoate with ethanoic acid?

Answer 29

Ethanoic acid is a weak acid so will partially dissociate so lots of undissociated ethanoic acid molecules will remain in solution. The sodium ethanoate will fully dissociate, producing lots of ethanoate ions, CH3COO-. Therefore the following equilibrium is set up which will move to counteract changes in pH:
CH3COOH(aq) <—> H+(aq) + CH3COO-(aq)

Question 30

Consider the buffer solution made by mixing sodium ethanoate with ethanoic acid. How does the pH change when a small amount of acid is added?

Answer 30

The following equilibrium is set up within the buffer solution:
CH3COOH <—> H+ + CH3COO-

If a small amount of acid is added, the concentration of H+ increases. Most of the extra H+ ions combine with CH3COO- ions to form CH3COOH so the equilibrium shifts to the left. This reduces the concentration of H+ to near to its original value so the pH does not change.

Question 31

Explain the significance of buffer solutions in nature

Answer 31

Buffer solutions are common in nature in order to keep systems regulated. Enzymes in living organisms often require an optimum pH and this can be maintained by a buffer solution.

Question 32

Why are buffers used in industrial processes?

Answer 32

Industrial processes use buffer solutions to maintain the optimum reaction conditions for large scale manufacturing.

Question 33

Give an example of where buffer solutions are used in industrial processes

Answer 33

Buffers are used in fermentation. Buffer solutions are added before fermentation begins to prevent the solution becoming too acidic.
They are also used in fabric dyeing processes and to perform chemical analysis.

Question 34

Define salt hydrolysis

Answer 34

A reaction where one of the ions from a salt reacts with water to form an acidic or basic solution.

Question 35

What needs to be considered when selecting an indicator for titration?

Answer 35

The indicator chosen for a titration must change colour at exactly the end point of titration. The indicator should change colour over a narrow pH range, which coincides with the vertical section of the pH curve.

Question 36

Give an example of an indicator which can be used in a titration and describe its colour change

Answer 36

Phenolphthalein
Colour at low pH: colourless
Colour at high pH: Pink
pH of colour change: 8.3-10

Methyl Orange
Colour at low pH: Red
Colour at high pH: Yellow
pH of colour change: 3.1-4.4

Question 37

Why should a pH meter be used instead of an indicator in weak acid/ weak base titrations?

Answer 37

In weak acid/ weak base titrations there is no sharp pH change so it is very difficult to successfully use an indicator to get an accurate point of neutralisation. A pH meter should be used instead since it will be more accurate and precise.