Chapter 21 - Buffers and Neutralisation Flashcards Preview

A level Chemistry (OCR) > Chapter 21 - Buffers and Neutralisation > Flashcards

Flashcards in Chapter 21 - Buffers and Neutralisation Deck (20)
Loading flashcards...
1
Q

What is a buffer solution?

A

A buffer solution is a solution that resists and minimises changes in pH when small amounts of acid or base are added.

2
Q

What components do buffer solutions contain?

A

a. A Buffer solution contains two components to remove added acid or alkali – a weak acid (component 1) and its conjugate base (component 2).
i. The weak acid, HA, removes added alkali.

ii. The conjugate base, A-, removes added acid.

3
Q

How does a buffer work?

A

a. When alkalis and acids are added to a buffer, the two components in the buffer solution will react and will eventually be used up. As soon as one component has all reacted, the solution loses its buffering ability towards added acid or alkalis.
b. As the buffer works, the pH does change but only by a small amount – you should not assume that the pH stays completely constant.

4
Q

What are the two ways you can prepare a weak acid buffer solution?

A

a. Preparation from a weak acid and its salt.
i. Mix a solution of weak acid with a solution of one of its salts.
ii. When weak acid is added to water, it partially dissociates, and the amount of ethanoate ions in the solution is very small. Therefore, the weak acid is the source of the weak acid component in the buffer solution.
iii. Salts of weak acids are ionic compounds and when added to water, the salt completely dissolves. Dissociation into ions is complete and so the salt is the source of the conjugate base component of the buffer solution.
b. Preparation by partial neutralisation of the weak acid
i. Add an aqueous solution of an alkali, such as NaOH, to an excess of weak acid.
ii. The weak acid is partially neutralised by the alkali, forming the conjugate base. Some of the weak acid is left over unreacted as it is in excess. Therefore, the resulting solution contains a mixture of the salt of the weak acid (component 2) and any unreacted weak acid (component 1).

5
Q

How does conjugate base remove any added acid?

A

a. On addition of an acid, H+(aq):
1. [H+] increases.

  1. H+ ions react with the conjugate base, A-.
  2. The equilibrium position shifts to the left, removing most of the H+ ions.
6
Q

How does weak acid remove any added alkali?

A

a. On addition of an alkali, OH-(aq):
1. [OH-] increases.

  1. The small concentration of H+ ions reacts with the OH- ions:
    a. H+ + OH- = H20
  2. HA dissociates, shifting the equilibrium position to the right to restore most of the H+ ions.
7
Q

When is a buffer the most effective at removing either added acid or alkali?

A

a. When there are equal concentrations of the weak acid and its conjugate base (When [HA] = [A-]):
i. The pH of the buffer solution is the same as the pKa value of HA.

ii. The operating pH is typically over about two pH units, centred at the pH of the pKa value.
b. So, if the concentrations of HA and A- in a buffer are equal then the pH of the buffer is equal to the pKa value of the weak acid.

8
Q

At what pH does blood plasma need to be maintained at and what is the most important buffer system used to control the pH.

A

a. Blood plasma needs to be maintained at a pH between 7.35 and 7.45.
b. The pH is controlled by a mixture of buffers, with the carbonic acid-hydrogen carbonate (H2CO3/HCO3-) buffer system being the most important. Normal healthy blood should have a pH of 7.40.

9
Q

What happens if the pH slips outside this range?

A

a. If the pH falls below 7.35, people can develop a condition called acidosis, which can cause fatigue, shortness of breath, and in extreme cases, shock or death.
b. If the pH rises above 7.45, the condition is called alkalosis, which can cause muscle spasms, light-headedness, and nausea.
c. The body produces far more acidic materials than alkaline, which the conjugate base HCO3- converts to H2CO3. The body prevents H2CO3 building up by converting it to carbon dioxide gas, which is then exhaled by the lungs.

10
Q

What is the Henderson-Hasselbalch equation?

A

pH = pKa + log(A-/HA)

11
Q

How would you monitor the pH as an aqueous base is added to an acid solution?

A

a. Using a pipette, add a measured volume of acid to a conical flask.
b. Place the electrode of the pH meter in the flask.
c. Add the aqueous base to the burette and add to the acid in the conical flask, 1cm^3 at a time.
d. After each addition, swirl the contents. Record the pH and the total volume of the aqueous base added.
e. Repeat steps c and d until the pH starts to change more rapidly. Then add the aqueous base 1cm^3 at a time again until an excess has been added and the pH has been basic, with little change, for several additions.

12
Q

What is an alternative method you could use to monitor the pH as an aqueous base is added to an acid solution?

A

You could attach the pH meter to a datalogger and use a magnetic stirrer in the flask. The aqueous base would then be added from the burette to the flask slowly, and the pH titration curve could be plotted automatically using the datalogger or an appropriate software on a computer.

13
Q

Describe pH titration curves.

A

a. When the base is first added, the acid is in great excess and the pH increases very slightly. As the vertical section is approached, the pH starts to increase more quickly as the acid is used up more quickly.
b. Eventually, the pH increases rapidly during addition of a very small volume of base, producing the vertical section. Only drops of solution will be needed for the whole vertical section.
c. After the vertical section, the pH will rise very slightly as the base is now in great excess.
d. The equivalent point is the centre of the vertical section of the pH titration curve.

14
Q

What is the equivalence point?

A

a. The equivalence point of the titration is the volume of one solution that exactly reacts with the volume of the other solution.
b. The solutions have then exactly reacted with one another and the amounts used to match the stoichiometry of the reaction.

15
Q

What is an acid-base indicator?

A

a. An acid-base indicator is a weak acid, HA, that has a distinctively different colour from its conjugate base, A-, for example, for the common indicator methyl orange: the weak acid is red, and the conjugate base is yellow.
b. At the end point of a titration, the indicator contains equal concentrations of HA and A- and the colour will be in between the two extreme colours. For methyl orange, the colour at its end point is orange.

16
Q

Explain indicator colour changes.

A

a. An indicator is a weak acid. The equilibrium position is shifted towards the weak acid in acidic conditions or towards the conjugate base in basic conditions, changing the colour as it does so.
b. In a titration in which a strong base is added to a strong acid, methyl orange is initially red as the presence of H+ ions forces the equilibrium position well to the left.
i. On addition of a basic solution containing OH- ions:

  1. OH- ions react with the H+ in the indicator:
    a. H+ + OH- = H20.
  2. The weak acid, HA, dissociates, shifting the equilibrium position to the right.
  3. The colour changes, first to orange at the end point and finally yellow as the equilibrium position is shifted to the right.
    ii. If methyl orange is added initially to a basic solution and acid is added.
  4. H+ ions react with the conjugate base, A-.
  5. The equilibrium position shifts to the left.
  6. The colour changes, first to orange at the end point and finally to red when the equilibrium has shifted to the left.
17
Q

How would you choose the right indicator in a titration?

A

In a titration, 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.

18
Q

How would you draw rough sketches of acid-base titration curves?

A

a. Weak acid starts at a curved line from pH of 3 and weak base curves off at around pH 9.
b. Strong acid starts at a sharp line from pH of 1 and strong base curves off at around pH 13.

19
Q

How would you find the Ka using a titration curve where a strong acid is being titrated by a weak base?

A

Half way between the equivalence point and the start of the titration, a buffer is created where half as much acid has been titrated by base. At this point, base will be equal to acid and a buffer has been set up. Due to the henderson-hasselbalch equation pH=pKa as log(1) = 0. At the buffer pH on the titration curve, that will be equal to pKa and thus you will be able to work out the Ka of the acid.

20
Q

How would you find the Ka using the equivalence point?

A

Work out the reaction at the equivalence point and see if the final solution will be acidic, basic or neutral. Using this, create an equilibrium equation of the ion affecting the final solution when dissolved in water. Calculate the [H+] at the equivalence point using the pH. Then use this [H+] or calculate [OH-] using the ionic product of water and then finally calculate the Ka.