Module 5: pH

Question 1

Give the 3 reactions of acids.

Answer 1

Acid + Base —> Salt + Water
Acid + Metal —> Salt + Hydrogen
Acid + Carbonate —> Salt + Water + CO2

Question 2

Define strong acids

Answer 2

Fully dissociate in water.
Examples:
HCl
HNO3 (nitric acid)
H2SO4 (sulfuric acid)

Question 3

What do acids do in water?

Answer 3

Acids split apart in water (dissociate in water) and an equilibrium is set up.

Question 4

What are acids called that donate 1 proton?

Answer 4

Monobasic acids.
E.g:
HCl —> H+ + Cl-

Question 5

What are dibasic acids?

Answer 5

Acids that donate 2 protons.
Done in 2 stages. E.g:
H2SO4 —> H+ + HSO4-
HSO4- —> H+ + SO4^2-

Question 6

What are acids called that donate 3 protons?

Answer 6

Tribasic acids.
Done in 3 stages. E.g:
H3PO4  -->  H+  +  H2PO4-
H2PO4-  -->  H+  +  HPO4^2-
HPO4^2-   -->  H+  +  PO4^3-

Question 7

Define weak acids.

Answer 7

Partially dissociated.
Tend to be organic acids (ethanoic acid).
The equilibrium position lies massively on the left.

Question 8

Define conjugate acid-base pairs.

Answer 8

An acid-base pair is a set of two related species that transform into each other by the gain or loss of a proton (H+).

Question 9

Is an acid a proton donor or acceptor?

What about a base?

Answer 9

An acid is a proton donor.

A base is a proton acceptor.

Question 10

Can you get a conjugate acid-base pair when 2 acids are mixed?

Answer 10

YES.

When 2 acids are mixed, the stronger acid will donate a proton to the weaker acid.

Question 11

What is the equilibrium constant when acids from an equilibrium in water?

Answer 11

Ka

Question 12

Do we include water in the equilibrium constant for acid dissociations?

Answer 12

Water is in excess, so is not included in the equilibrium constant.

Question 13

What does a large value of Ka represent?

Answer 13

A large value of Ka means a high degree of dissociation.

The bigger the value of Ka, the stronger the acid.

Question 14

What does a small value of Ka represent?

Answer 14

A small value of Ka means a small level of dissociation.

The smaller the value of Ka, the weaker the acid.

Question 15

Give the equations for calculating:

1) pKa
2) Ka

Answer 15

pKa = -log10Ka (The bigger the value of pKa, the weaker the acid).

Ka = 10^-pKa

Question 16

Give the equation for calculating:

1. pH
2. [H+]

Answer 16

pH = -log[H+]

[H+] = 10^-pH

Question 17

How do you calculate the pH of strong acids?

Answer 17

As strong acids are fully dissociated into H+, this means the [reactants] = [H+]

Example:
HCl —–> H+ + Cl-

Ka= [H+][Cl-] / [HCl]

[HCl] = [H+]

pH=-log10([HCl]

Question 18

How do you calculate the pH of weak acids?

Answer 18

A weak acid is partially ionised. We need to have the value of Ka to calculate the pH.

Example:
CH3CO2H —–> CH3CO2- + H+

Ka = [CH3CO2-][H+] / [CH3CO2H]

Ka = [H+]^2 / [CH3CO2H]

[H+] = square root (Ka x [CH3CO2H])

pH = -log[H+]

Question 19

How do you calculate the pH of a strong base?

Answer 19

To do this we use an equilibrium called “the ionic product of water”

Example:
H2O —–> H+ + OH- [H2O] is regarded as a constant,
so is removed from the Kw
equation.
Kw = [H+][OH-] / [H2O]
Kw = [H+][OH-]

Data sheet: Kw = 1x10^-14 mol^2dm^6 at 298K

1x10^-14 = [H+][OH-]

[H+] = 1x10^-14 / [OH-]

pH = -log[H+]

Question 20

Describe the neutralisation- titration curves for a strong acid and a strong base.

Answer 20

1.  Vertical section covers a large change in pH:
start at pH3 
end at pH 12
2.  Equivalence point at pH 7
3. Starts at approx pH 1

Question 21

Describe the neutralisation - titration curve for a strong acid and a weak base.

Answer 21

1.   Vertical section covers a smaller change in pH:
start at pH 2
end at pH 8
2.  Equivalence point at pH 7
3.  Starts at approx pH 1
Ends at approx pH 9

Question 22

Describe the neutralisation - titration curve for a weak acid and a strong base.

Answer 22

1.  Vertical section covers a smaller change in pH:
start at pH 6
end at pH 12
2.  Equivalence point at pH 7
3.  Starts at approx pH 3

Question 23

Describe the neutralisation - titration curve for a weak acid and a weak base.

Answer 23

1. There is no vertical section on the graph. This means:
   no indicator would work
   impossible to carry out an accurate titration.
2. Starts at approx pH 3
   Ends at approx pH 9

Question 24

What does the equivalence point represent?

Answer 24

The same no. of mols of acid is added to the same no. of mols of alkali.

Question 25

What should you consider when choosing the correct indicator?

Answer 25

Indicator must have a colour change that lies in the pH of the vertical part of the graph.

Question 26

Name 4 common strong acids.

Answer 26

HCl = Hydrochloric acid
H2SO4 = Sulfuric acid
HNO3 = Nitric acid
HBr = Hydrobromic acid

Question 27

Why would no indicator work for a weak acid/ weak base titration?

Answer 27

There is no vertical part on the graph.

Question 28

What is a buffer?

Answer 28

A solution that minimises changes in pH, upon the addition of H+ and OH-

Question 29

How is a buffer made?

Answer 29

The acid must be in excess.

Question 30

How does a buffer work?

Answer 30

1. Equilibrium of acid dissociating (HA —> A- + H+)

Add acid (H+)
Equilibrium shifts to LHS.
H+ reacts with A- to make HA. H+ + A- –> HA
Reservoir of A-, so solution can absorb a significant amount of H+, minimising change in pH.

Add alkali (OH-)
Equilibrium shits to RHS.
OH- reacts with H+ to make water. H+ + OH- –> H2O
Reservoir of undissociated acid (HA), so solution can absorb significant amount of OH-, minimising changes in pH.

Question 31

What is the buffer system in blood?

Answer 31

H2CO3 —> HCO3- + H+
(Equilibrium)
Carbonic acid —> Hydrogen carbonate

Question 32

How do you calculate the ratio of Hydrogen carbonate:Carbonic acid in blood.

Answer 32

Ka [HCO3-]
—– = ———-
[H+] [HCO3]

1. Calculate healthy [H+] and Ka.
2. Use figures to calculate the ratio.

Question 33

State a compound that can be added to ethanoic acid to make a buffer solution.

Answer 33

CH3COONa OR Na