Acids, Bases and Buffers

Question 1

define a Brønsted–Lowry acid

Answer 1

a species that can donate a proton

Question 2

define a Brønsted–Lowry base

Answer 2

a species that can accept a proton

Question 3

describe and use the term conjugate acid– base pairs

Answer 3

A conjugate pair is a pair of acid and base that differ by H+. e.g. CH3COOH and CH3COO-
H2SO4 and HSO4-
H3O+ and H2O

Question 4

What is the difference between strong and weak acids?

Answer 4

Strong: fully dissociates in solution HA –> H+ + A-What
Weak: partially dissociates in solution to form equilibrium HA H+ + A-

Question 5

What is the significance of the acid dissociation constant, Ka?

Answer 5

Shows the extent of acid dissociation, and therefore the relative strength of weak acids: the higher the value of Ka the stronger the acid

Question 6

How can Ka values be more easily compared?

Answer 6

Use the pKa scale, -log10(Ka)

The lower the pKa the stronger the acid

Question 7

Define pH

Answer 7

pH=–log[H+];

[H+]=10^–pH

Question 8

How to find the pH of a strong acid?

Answer 8

Fully dissociated, so [H+] for a monobasic acid is equal to the concentration of that acid.
[HA] = [H+]
pH = -log[H+]

Question 9

find pH of a weak monobasic acid?

Answer 9

use Ka
[H+] = square root of Ka x [HA]
Ka = [H+]^2/[HA] (or [H+]x[A-])
pH = -log[H+]

Question 10

What assumptions are made in the calculation of the pH of a weak acid using Ka?

Answer 10

assumes negligible acid dissociation so that [HA] = concentration of acid
Assumes that water contributes very little H+ so that [A-]=[H+]

Question 11

What is Kw? when is it used?

Answer 11

the ionic product of water, 1x10^-14
Used to convert [OH-] into [H+] and thus find the pH of a strong base.
Kw = [H+][OH-]
[H+] = 1x10^-14/[OH-]

Question 12

What is a buffer solution?

Answer 12

a system that minimises pH changes on addition of small amounts of acid or base

Question 13

how can a buffer solution be set up?

Answer 13

by reacting excess weak acid with a salt of the weak acid. An equilibrium is set up
e.g.
CH3COOH CH3COO- + H+
Position of equilibrium shifts to cancel out pH changes when H+/OH- is added

Question 14

How is the pH of a buffer solution calculated?

Answer 14

[H+] = Ka x [SALT]/[ACID]
OR
pH = pKa + log([SALT]/[ACID])

Question 15

explain the role of carbonic acid– hydrogencarbonate in the body

Answer 15

acts as a buffer in the control of blood pH;

Question 16

When asked to sketch a pH titration curve, what are the marking points that must be covered?

Answer 16

• initial points at 1 for a strong acid, 3 for a weak acid and final points of the curve at 13 for a strong base or 11 for a weak base.
• line is vertical at the equivalence point/ at the end point volume
• vertical point of the grap hcovers the values: 3-11(SA/SB), 7-11(WA/SB) or 3-7(SA/WB)
• smooth S shaped curve

Question 17

which indicators can be used for which titrations?

Answer 17

strong acid/strong base: methyl orange or phenol phthalein
weak acid/strong base: phenol phthalein
strong acid/weak base: methyl orange
weak acid/weak base: neither - no rapid pH change

Question 18

wht criteria must an indicator have to be used in a titration?

Answer 18

In order to change colour at the equivalence point the pH range of the indicator must coincide with the sharp rise in pH at the equivalence point

Question 19

what is enthalpy of neutralisation?

Answer 19

the enthalpy change when 1 mol of water is formed during the neutralisation of an acid by an alkali
H+ + OH- –> H2O

It is always exothermic

Question 20

how is enthalpy change of neutralisation calculated?

Answer 20

q(J) = mc∆T
∆Hneut = q/mol water formed

Question 21

What is a thermometric titration?

Answer 21

temperature change during a titration is used to deduce the end-point.
The temp. will increase during neutralisation and decrease once the equivalence point has been reached.
Point at which temp. mixture is highest = end-point.