Buffers and neutralisation

Question 1

What is a buffer solution?

Answer 1

A system that minimises pH changes when small amounts of an acid or base are added.
It contains two components:
The weak acid, removes added alkali.
The conjugate base, removes added acid.

Question 2

How do you prepare weak acid buffer solutions from a weak acid and its salt?

Answer 2

E.g. Add CH3COOH to one of its salts, CH3COONa (sodium ethanoate).
When ethanoic acid is added to water, it partially dissociates and the amount of ethanoate ions is very small.
When salts are added to water, the salt completely dissolves, and so provides a source of the conjugate base.

Question 3

How do you prepare weak acid buffer solutions by partial neutralisation of the weak acid?

Answer 3

Add an aqueous solution of the alkali, e.g. NaOH to an excess of weak acid.
The weak acid is partially neutralised by the alkali, forming the conjugate base.
Some of the weak acid is left over unreacted.
The resulting solution contains a mixture of the salt and weak acid.

Question 4

What are the two reservoirs to remove added acid and alkali?

Answer 4

The equilibrium position lies well to the left.
When the conjugate base ions are added to the acid, the position shifts even further to the left, reducing the already small concentration of H+ ions.
The solution left mainly is the acid and conjugate base.

Question 5

How does the buffer react when acid is added?

Answer 5

[H+] increases.
H+ ions react with the conjugate base, A-.
The equilibrium shifts to the left, removing most of the H+ ions.
HA <–> H+ + A-
<———————

Question 6

How does the buffer react when alkali is added?

Answer 6

[OH-] increases.
The small concentration of H+ ions reacts with the OH- ions.
H+ + OH- –> H2O
HA dissociates, shifting the equilibrium position to the right to restore most of the H+ ions.
HA <–> H+ + A-
———————->

Question 7

How do you choose the components of the buffer solution?

Answer 7

A buffer is most effective when there are equal concentrations of weak acid and its conjugate base.
When [HA] = [A-]:
The pH of the buffer is the same as the pKa value of HA.
The operating pH is typically over about 2 pH units, centred at the pH of the pKa value.

Question 8

How do you calculate the pH of a buffer solution?

Answer 8

[H+] = Ka x [HA] / [A-]
Ka / [H+] = [HA] / [A-] can be used to work out ratios.
-log [H+] to find pH

Question 9

How do you work out pH of a buffer solution if there is partial neutralisation in a question?

Answer 9

Find out the moles of both the conjugate base (the base at the start) and the acid.
Find out how many moles are unreacted to find out the acid moles in the buffer.
Divide these by the volume of the buffer to find the concentrations in the buffer.
Use the equation [H+] = Ka x [HA] / [A-]
pH = -log[H+]

Question 10

Why does blood pH need controlling?

Answer 10

Different parts of the body need specific pH values for effective functioning.
Enzymes are very sensitive and each has an optimum pH.
Blood plasma needs to be maintained at a pH of 7.35-7.45.

Question 11

What happens if pH falls outside the range?

Answer 11

If it falls below 7.35, a condition called acidosis develops, which causes fatigue, shortness of breath, and can cause shock or death.
If it rises above 7.45, alkalosis can cause muscle spasms, light headedness and nausea.

Question 12

What is the carbonic acid-hydrogencarbonate buffer system for addition of acids?

Answer 12

[H+] increases.
H+ ions react with the conjugate base HCO3-.
The equilibrium position shifts to the left, removing most of the H+ ions.
H2CO3 <–> H+ + HCO3-
<———————————

Question 13

What is the carbonic acid-hydrogencarbonate buffer system for addition of alkali?

Answer 13

[OH-] increases.
The small concentration of H+ ions reacts with the OH- ions.
H+ + OH- –> H2O
H2CO3 dissociates, shifiting the equilibrium position to the right to restore most of the H+ ions.
H2CO3 <–> H+ + HCO3-
———————————->

Question 14

How do you calculate the concentration ratio of HCO3- / H2CO3?

Answer 14

Ka = [H+] [HCO3-] / [H2CO3]
rearrange [H2CO3] / [HCO3-] = Ka / [H+]
Convert pH into [H+] and pKa into Ka.
Calculate the ratio using the above equation.

Question 15

What is the equivalence point?

Answer 15

The equivalence point of the titration is the volume of one solution that reacts exactly with the volume of the other solution.
It matches the stoichiometry of the reaction.
It is the centre of the vertical section of the pH titration curve.

Question 16

What are the features of a titration curve?

Answer 16

Excess of acid: pH increases slowly as basic solution is added.
Vertical section: pH increases rapidly on addition of a very small volume of base.
Equivalence point: the centre of the vertical section.
Excess of base: pH increases slowly as basic solution is added.
(opposite for adding acid to base)

Question 17

What are the different types of titration curves?

Answer 17

Adding strong base to weak acid - small bump at pH 3, ends at pH 13.
Adding weak acid to weak base - wiggly curve with little vertical section.

Question 18

What are strong/weak acids/bases?

Answer 18

Strong acid - HCl, H2SO4
Strong base - NaOH
Weak acid - CH3COOH
Weak base - NH3

Question 19

What is the end point of a titration?

Answer 19

The indicator contains equal concentrations of weak acid and conjugate base.
The colour will be inbetween the two extreme colours.

Question 20

What are acid-base indicators?

Answer 20

A weak acid, that has a distinctively different colour from its conjugate base.
E.g. methyl orange, the acid is red, the conjugate base is yellow.
Phenolphthalein, the acid is colourless, conjugate base is pink.

Question 21

What happens to indicators when base is added?

Answer 21

OH- ions react with H+ in the indicator.
H+ + OH- –> H2O
The weak acid, HA, dissociates, shifting the equilibrium position to the right.
The colour changes, first to the end point, then to the final colour.
e.g. methyl orange goes from red to orange to yellow.

Question 22

What happens to indicators when acid is added?

Answer 22

H+ ions react with the conjugate base A-.
The equilibrium position shifts to the left.
The colour changes, first to the end point then to the final acidic colour.

Question 23

What happens in an end point?

Answer 23

[HA] = [A-] and Ka = [H+] (HA and A- cancel out).
Ka = [H+] and pKa = pH

Question 24

What is the relationship between Ka and pKa?

Answer 24

The stronger the acid, the larger the Ka value and the smaller the pKa value.
The weaker the acid, the smaller the Ka value and the larger the pKa value.

Question 25

How are indicators chosen?

Answer 25

Choose 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, but not always possible.
Indicators change colour over about 2 pH units.

Question 26

Which indicators are suitable for which titrations?

Answer 26

Methyl orange operates about pH 3-4, phenolphthalein about 8.5-10.
HCl and NaOH, both suitable.
CH3COOH and NaOH, only phenolphthalein suitable.
HCl and NH3, only methyl orange suitable.
CH3COOH and NH3, no indicator suitable, as there is no vertical section.