Chemistry - 3.1.12: Acids and Bases

Bronsted-Lowry acid

Proton donor

Give a safety precaution when using a burette.

Fill it below eye level

What must you do before recording an initial burette reading? (2)

1. Remove the funnel

2. Ensure no air bubbles

Why is it necessary to calibrate the pH metre?

Over time / after storage, it does not give accurate readings

Describe a method to produce a pH curve (6)

1. Measure pH of acid (with a pH metre)

2. Add alkali in in known small amounts eg. 2-3cm3

3. Stir mixture

4. Measure pH

5. Repeat until alkali is in excess

6. Add in smaller increments near end point

Bronsted-Lowry base

Proton acceptor

Oxonium ion

H3O+

Alternative names of H3O+

Hydronium ion, hydroxonium ion

What do acid-base equilibria involve?

Protons are transferred from the acid to the base

Equation for the ionisation of water

H2O (l) ⇌ H+ (aq) + OH- (aq)
OR: H2O (l) + H2O (l) ⇌ H3O+ (aq) + OH- (aq)

Kw meaning

The ionic product of water

What is the symbol for the ionic product of water?

Kw

Kw expression

Kw = [H+][OH-]

Where does the Kw expression come from?

Kc = [H+][OH-]/[H2O] - water only slightly dissociates, so there are many more H2O molecules than the ions and [H2O] doesn't change so can be removed

When does the value of Kw change?

When temperature changes

Formula to calculate pH

pH = -log[H+]

What is the pH scale?

A logarithmic scale used as a measure of hydrogen ion concentration

Why is the pH scale used to measure [H+]? (2)

[H+] in aqueous solution covers a wide range of very small concentrations.
Using only log[H+] would result in negative numbers.

How do you calculate the pH of a strong base? (3)

1. Find [OH-] (adjust based on ratio based on empirical formula of the base)
2. Use [OH-] and Kw to find [H+]
3. Use [H+] to find pH

How do you calculate the concentration of a strong base from pH? (3)

1. Use pH to find [H+]
2. Use [H+] and Kw to find [OH-]
3. Use [OH-] to find concentration (adjust based on empirical formula of the base)

What does it mean if pH < 7?

The solution is acidic - it contains more H+ than OH-

What does it mean if pH > 7?

The solution is alkaline - it contains more OH- than H+

How do you calculate [H+] from pH?

10^(-pH)

Strong acid

An acid that fully dissociates into ions in solution

Weak acid

An acid that is only slightly dissociated into ions in solution

Ka meaning

Acid dissociation constant

Formula for Ka?

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

Where does the formula for Ka come from?

It is the same as the equilibrium constant for the dissociation of weak acids

What does it mean if Ka increases?

As Ka increases, more acid is dissociated so the acid is stronger

How do you calculate the pH of a weak acid using Ka and [HA]? (2)

1. Use Ka= [H+]^2 / [HA] to find [H+]
2. Then, use pH = -log[H+]

Units of Ka?

moldm^-3 - find by cancelling units from the Ka expression

Give an alternate formula for Ka and explain why you can use it?

Ka = [H+]^2 / [HA] - at equilibrium, [H+] = [A-]

How do you calculate the concentration of a weak acid using pH and Ka? (2)

1. Calculate [H+] from pH using [H+] = 10^(-pH)
2. Use Ka = [H+]^2/[HA] to find [HA]

How do you calculate the pH of a strong acid from its concentration? (2)

1. From equation, [HA] = [H+]
2. Then, pH = -log[H+]

Formula for pKa

pKa = -log[Ka]

To how many decimal places is pH or pKa given to?

2 decimal places

From Ka, how can you tell which acid is stronger?

The one with the higher Ka

From pKa, how can you tell which acid is stronger and why?

The one with the lower pKa - means they have a higher Ka

Buffer

A solution that resists change of pH when small amounts of acid or base are added or on dilution

2 types of buffers and what they're made from?

Acidic buffers - weak acid + soluble salt of conjugate base
Basic buffers - weak base + soluble salt of conjugate acid

What do buffer solutions do?

They maintain an approximately constant pH despite dilution or addition of small amounts of acid or base

How does a buffer resist changes in pH upon addition of a small amount of acid? (4)

1. Write down reversible reaction for buffer
2. Adding acid increases [H+]
3. So the anion reacts with the H+ and the equilibrium shifts to the left to lower [H+]
4. So overall, [H+] and the ratio of [HA]:[A-] remains roughly constant

Why would a weak acid with a chlorine atom attached be stronger than a weak acid without that?

1. Cl is very electronegative so withdraws electrons (negative inductive effect)

2. So it weakens the O-H bond

How does a buffer resist changes in pH upon addition of a small amount of base?

1. Write down reversible reaction for buffer
2. Adding base decreases [H+] as the OH- and H+ react
3. So the equilibrium shifts to the right to increase [H+]
4. So overall, [H+] and the ratio of [HA]:[A-] remains roughly constant

In buffers, why do you need to add the soluble salt of the conjugate acid/base?

The salt fully ionises, increasing the supply of H+/A- (respectively) so the pH remain relatively constant

What happens if a lot of acid or alkali is added to a buffer solution?

The buffer is saturated - the pH changes by a lot

What is important about the half-neutralisation point of a reaction?

pH = pKa

Give an example of a buffer

The blood

How do you calculate the pH of a buffer? (3)

1. Find moles of acid, base and salt at the end of the reaction
2. Calculate concentration of reactant in excess and of salt
3. Use Ka to calculate [H+]. then use that to calculate pH

Draw the titration curve of a strong base added to a strong acid (3)

pH starts around 1
Equivalence point from around pH 3 to pH 11, long vertical line
Final pH at around 12 or 13

Draw the titration curve of a weak base added to a strong acid (3)

pH starts around 1
Equivalence point from around pH 3 to pH 7 or 8, short vertical line
Final pH at around 9 or 10

Draw the titration curve of a strong base added to a weak acid (3)

pH starts around 3
Equivalence point from around pH 6 to pH 11, short vertical line
Final pH at around 12 or 13

Draw the titration curve of a weak base added to a weak acid (3)

pH starts around 3
Equivalence point from around pH 6 to pH 7 or 8, very short vertical line
Final pH at around 9 or 10

Equivalence point

The point in a titration at which sufficient base has been added to just neutralise the acid or vice-versa - the reaction is just complete

End-point

The point in a titration at which an indicator changes colour

What is special about the equivalence point?

Moles H+ = moles OH-

3 things to consider when selecting an indicator?

1. Colour change must be sharp rather than gradual at end point - only one drop of acid/alkali needed for complete colour change
2. End point of titration given by indicator must be the same as the equivalence point
3. Indicator should give a distinct colour change

Half-equivalence point

Point in a titration halfway between the start and the equivalence point

Alternative name for half-equivalence point

Half-neutralisation point

Why is [H₂O] not included in the Kw expression?

[H₂O] is almost constant

When calibrating a pH probe using different buffer solutions, why do you need to wash the pH probe with distilled water between each measurement? (2)

1. So the different solutions don’t contaminate each other

2. To wash off any residual solution, which could interfere with the reading

When measuring the pH during a titration, why do you add a smaller volume of NaOH between each pH measurement as the end point of the titration is reached? (2)

1. To avoid missing the end point

2. Large change in pH near end point

From a titration curve and the pH range of an indicator, how can you tell if the indicator would be suitable for the titration?

Its colour change/pH range is within the steep/vertical part of the titration curve

How do you calculate Ka from a titration curve? (2)

1. Look at half-equivalence point - read off pH value and volume added

2. pH = pKa so use Ka = 10^-pKa to calculate Ka

What mathematical expression can you use to show that a buffer solution has a constant pH despite dilution?

Ratio of [X^-]:[HX]

Describe a procedure to measure the pH during a titration (3)

1. Place fixed volume of acid/alkali in beaker

2. Add alkali/acid in small portions from burette

3. Stir and use pH metre to record pH after each addition of alkali/acid

Describe a practical procedure you could carry out to identify the most suitable indicator from a selection for a titration (2)

1. Repeat procedure of adding acid/alkali to sample and measuring pH with each indicator

2. Select indicator that changes colour rapidly during pH range of equivalence point

Why is it difficult to judge the end point of a neutralisation reaction between a weak acid and base accurately with an indicator?

The change in pH is gradual at the end point - the indicator would change colour over a range of volumes of alkali/acid

How do you make a buffer solution with a strong base and weak acid solution?

Add excess weak acid to the strong base (so mixture of acid anions and acid)

When acids dissociate in solution, why is the H atom bonded to O the one that dissociates rather than the one bonded to C?

O is significantly more electronegative than H, so the O-H bond is already polarised and the H is already in the process of becoming an ion