Buffers and Neutralisation (Chapter 21)

Question 1

What is a buffer solution?

Answer 1

A system that minimises pH changes on addition of small amount of an acid or base

Question 2

What are the two components of a buffer solution and what are their roles?

Answer 2

1) weak acid (HA) - to remove added alkali

2) its conjugate base (A-) - to remove added acid

Question 3

What are the two ways to prepare a buffer solution?

Answer 3

1) from a weak acid and its salt

2) by partial neutralisation of the weak acid

Question 4

Describe how you would prepare a buffer solution from a weak acid and its salt

Answer 4

Mix a solution of a weak acid (e.g. CH3COOH) with a solution of one of its salts (e.g. CH3COONa)

Question 5

Describe how mixing a solution of a weak acid (e.g. CH3COOH) with a solution of one of its salts (e.g. CH3COONa) makes a buffer solution

Answer 5

1) in water, CH3COOH partially dissociates and the amount of ethanoate ions (CH3COO-) is very small
2) ∴ CH3COOH is the source of the weak acid component of the buffer solution: CH3COOH <=> CH3COO- + H+
3) in water, CH3COONa completely dissolves (as it is an ionic compound)
4) ∴ CH3COONa is the source of the conjugate base component of the buffer solution component of the bugger solution: CH3COONa + aq => CH3COO- + Na+

Question 6

Describe the preparation of a buffer solution by partial neutralisation of the weak acid

Answer 6

1) add an aqueous solution of an alkali e.g. NaOH to an excess of the weak acid
2) the weak acid is partially neutralised by the alkali, forming the conjugate base
3) some of the weak acid is left over unreacted ∴ the resulting solution contains a mixture of the salt and any unreacted weak acid

Question 7

How does the CH3COOH/CH3COO- buffer solution become almost a 100% solution of CH3COOH and CH3COO-?

Answer 7

1) in the CH3COOH equilibrium, the equilibrium position lies well towards the acid
2) ∴ when CH3COO- ions are added to CH3COOH, the equilibrium position shifted even further to the left, reducing the already small [H+] and leaving a solution containing mainly CH3COOH and CH3COO-

Question 8

Describe the CH3COOH/CH3COO- buffer solution

Answer 8

• CH3COOH and CH3COO- act as 2 reservoirs that are able to act independently to remove acid and alkali
• This is achieved by shifting the buffer’s equilibrium system either to the right or left

Question 9

How can a buffer solution lose its buffering ability?

Answer 9

1) when acids and alkalis are added to a buffer, the two components in the buffer solution react and will eventually be used up
2) as soon as one component has all reacted, the solution loses its buffering ability towards acids or alkalis - as the
3) as the buffer works, pH does change but only by a small amount - it does not stay constant

Question 10

What controls the pH in an acid buffer solution?

Answer 10

• The conjugate acid-base pair HA/A- in an acid buffer solution controls the pH
• This control can be explained by shifts in equilibrium position

Question 11

Explain what happens when acid (H+) is added to a buffer solution

Answer 11

1) [H+] increases
2) H+ ions react with the conjugate base A-
3) ∴ equilibrium shifts to the left, removing most of the H+ ions - the conjugate base removes the added acid

Question 12

Explain what happens when alkali (OH-) is added to a buffer solution

Answer 12

1) [OH-] increases
2) the small [H+] reacts with the OH-, decreasing the [H+]
3) HA dissociates, shifting the equilibrium position to the right to restore most of the H+ ions

Question 13

How does the weak acid used in a buffer solution affect the buffer solution?

Answer 13

Different weak acids result in buffer solutions that operate over different pH ranges

Question 14

When is a buffer most effective a remove either added acid or alkali?

Answer 14

When there are equal concentrations of the weak acid and its conjugate base in the buffer solution

Question 15

What happens to the pH of a buffer solution when [HA]=[A-]?

Answer 15

• Ka = [H+] ∴ the pH of the buffer solution = the pKa value of HA (pH=pKa)
• The operating pH is typically over about 2 pH units, centred at the pH of the pKa value
• The ratio of concentrations of the weak acid and its conjugate base can then be adjusted to fine-tune the pH of the buffer solution

Question 16

If the pKa of CH3COOH is 4.76, what is the typical operating pH range of its buffer solution?

Answer 16

3.76 to 5.76

Question 17

How do you calculate the [H+] for a buffer solution?

Answer 17

• [H+] = Ka x [HA]/[A-] ([salt])

Question 18

Why, for a buffer solution, does [H+] does not equal [A-]?

Answer 18

Bc A- has been added as one of the components of the buffer

Question 19

What does the pH of a buffer solution depend on?

Answer 19

1) pKa of the weak acid

2) the ratio of [HA]:[A-]

Question 20

Describe the pH of blood plasma

Answer 20

• Blood plasma needs to be maintained at a pH of 7.35 to 7.45
• the pH is controlled by a mixture of buffers with the carbonic acid/hydrogencarbonate (H2CO3/HCO3-) buffer system being the most important
• pH of normal healthy blood = 7.40
• pH < 7.35 = acidosis pH>7.45 = alkalosis

Question 21

Explain what happens when acid (H+) is added to the H2CO3/HCO3- buffer solution

Answer 21

1) [H+] increases
2) H+ ions react with the conjugate base HCO3-
3) ∴ the position of equilibrium shifts to the left, removing most of the H+ ions

Question 22

Explain what happens when alkali (OH-) is added to the H2CO3/HCO3- buffer solution

Answer 22

1) [OH-] increases
2) the small [H+] reacts with the OH-, reducing the [H+]
3) ∴ H2CO3 dissociates, shifting the equilibrium position to the right to restore most of the H+ ions

Question 23

Why is there a risk of H2CO3 building up in the body and how does the body prevent this?

Answer 23

• The body produces far more acidic materials than alkaline, which the conjugate base HCO3- converts to H2CO3
• The body prevents H2CO3 from building up by converting it to CO2 gas which is then exhaled by the lungs

Question 24

What is the Henderson-Hasselbach equation?

Answer 24

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

Question 25

What do acid-base titrations use?

Answer 25

Indicators - to monitor neutralisation reactions accurately

Question 26

What is a pH meter used for?

Answer 26

To monitor pH changes that take place during a titration

Question 27

Explain the shape of the strong-acid strong-base titration curve (volume of base on x-axis)

Answer 27

1) when the base is first added, the acid is in great excess and the pH increases very slightly
2) as the vertical section is approached, the pH starts to increase more quickly as the acid is used up more quickly
3) eventually, the pH increases rapidly during addition of a very small volume of base, producing the vertical section
4) after the vertical section, the pH will rise very slightly as the base is now in great excess
(Z shape for volume of acid on x-axis)

Question 28

What is the equivalence point of a titration?

Answer 28

• The volume of one solution that exactly reacts with the volume of the other solution
• The solutions have then exactly reacted with one another and the amounts use match the stoichiometry of the reaction
• It is the centre of the vertical section of the pH titration curve

Question 29

What is an acid-base indicator?

Answer 29

A weak acid that has a distinctively different colour from its conjugate base e.g. for methyl orange, HA is red and A- is yellow

Question 30

What happens at the end-point of a titration?

Answer 30

The indicator contains equal concentrations of HA and A- ∴ the colour will be in between the two extreme colours e.g. for methyl orange, the end point is orange

Question 31

How does an indicator work?

Answer 31

An indicator is a weak acid ∴ the equilibrium position is shifted towards the weak acid in acidic conditions and towards the conjugate base in basic conditions, changing the colour as it does so

Question 32

Describe what happens to methyl orange during a titration in which a strong base is added to a strong acid

Answer 32

1) initially, methyl orange is red as the presence of H+ ions forces the equilibrium position well to the left
2) when the OH- ions are added, they react with H+ in the indicator
3) ∴ HA dissociates, shifting the equilibrium position to the right
4) the colour changes, first to orange at the end point and finally to yellow as the equilibrium position is shifted to the right

Question 33

Describe what happens to methyl orange during a titration in which a strong acid is added to a strong base

Answer 33

1) initially, methyl orange is yellow as the presence of OH- ions forces the equilibrium position well to the right
2) when the H+ ions are added, they react with the conjugate base A- in the indicator
3) ∴ the position of equilibrium shifts to the left, removing most of the H+ ions
4) the colour changes, first to orange at the end point and finally to red when the equilibrium position has shifted to the left

Question 34

What is true about different indicators?

Answer 34

They have different Ka values and change colour over different pH ranges

Question 35

What happens at the end point?

Answer 35

• [HA] = [A-] ∴ Ka = [H+] ∴ pKa = pH

- ∴ pH of end point = pKa of HA

Question 36

What does the sensitivity of an indicator depend on?

Answer 36

The indicator itself and eyesight - most indicators change colour over a range of roughly 2 pH units

Question 37

How do you choose an indicator for a titration?

Answer 37

• You must use an indicator that has a colour change which coincides with the vertical section of the pH titration curve
• Ideally, the end point and equivalence point would coincide
• However, this may not be possible and the end point may give a volume that is slightly different from the equivalence point - but any different will be very small in the order of 1 or 2 drops (0.05-0.1cm3)

Question 38

Why is no indicator suitable for a weak acid-weak base titration?

Answer 38

Bc there is no vertical section on the titration curve and even at its steepest, the pH requires several cm3 to pass through a typical pH indicator range of 2 pH units

Question 39

What are three common indicators?

Answer 39

Phenolphthalein, methyl orange and bromophenol blue

Question 40

What does a pH meter consist of?

Answer 40

An electrode, that is dipped into a solution and connected to a meter the displays the pH reading

Question 41

Why is pH meter more accurate than indicator paper?

Answer 41

A pH meter typically records pH values to 2 d.p whereas indicator paper is usually matched from colour charts t give more accurate measurements of pH during a titration

Question 42

Describe how you would use a pH meter to plot a titration curve in an acid-base titration

Answer 42

1) place the electrode of the pH meter into a flask of acid
2) after each addition of base, swirl and record the pH and volume of base added
3) reduce the volume of bases added each time as the pH changes less rapidly until an excess of base has been added and the pH has been basic with little change for several additions and then plot pH vs volume of base added

Question 43

What is an alternative automatic method to use a pH meter to plot a titration curve?

Answer 43

1) attach the pH meter to a datalogger and use a magnetic stirrer in the flask
2) the pH titration curve can then be plotted automatically using the datalogger or appropriate software on a computer