3.9 Acid-Base Equilibria

Question 1

What is the Lowry-Brønsted theory?

Answer 1

The Lowry-Brønsted 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-Brønsted acid and give an example

Answer 2

A Lowry-Brønsted acid is a proton donor.
Example: Ammonium ions (NH +)

Question 3

Define a Lowry-Brønsted base and give an example

Answer 3

A Lowry-Brønsted 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

Hydrochloric acid (HCl), sulfuric acid (H2SO4), nitric acid (HNO3)
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

Ethanoic acid (CH3COOH), hydrogen sulfide (H2S), any 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

Sodium hydroxide (NaOH), potassium hydroxide (KOH), calcium hydroxide (Ca(OH)2)
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

Ammonia (NH3), methylamine (CH3NH2)
pH range for weak bases: just above 7 up to 11

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

Answer 14

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 a 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

Question 19

Given that Ka for propanoic acid at 298 K is 1.3 x 10-5 mol dm-3, what is the pH of 0.02 mol dm-3 of propanoic acid at 298 K?

Answer 19

Question 20

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

Answer 20

‘Concentrated’ implies there are many moles per dm3.
‘Strong’ relates to the dissociation of the substance and implies the acid/base almost completely dissociates in water.

Question 21

Give the formulas used to convert between Ka and pKa

Answer 21

pKa = -log10[Ka]
Ka = 10(-pKa)

Question 22

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

Answer 22

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

Question 23

Give the equation for the ionic product of water

Answer 23

Question 24

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

Answer 24

Question 25

What is the pH of pure water at room temperature?

Answer 25

7

Question 26

How does the pH at which water is neutral ([OH] = [H+]) change as temperature increases?

Answer 26

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 27

Given that Kw at 298 K is 1.0 x 10-14 mol2 dm-6, what is the pH of 0.10 mol dm-3 NaOH at 298 K?

Answer 27

Question 28

What is a pH curve?

Answer 28

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 29

What is the equivalence point on a pH curve?

Answer 29

Question 30

Answer 30

Question 31

Answer 31

Question 32

Answer 32

Question 33

What is a buffer solution?

Answer 33

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 a 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 34

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

Answer 34

Question 35

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 35

Question 36

Explain the significance of buffer solutions in nature

Answer 36

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 37

Why are buffers used in industrial processes?

Answer 37

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

Question 38

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

Answer 38

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 39

Define salt hydrolysis

Answer 39

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

Question 40

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

Answer 40

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

Question 41

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

Answer 41

Question 42

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

Answer 42

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.