Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

A Level Chemistry > Topic 12: Acid Base Equilibria > Flashcards

Topic 12: Acid Base Equilibria Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

What are Bronsted-Lowry acids?

A
  • acids are proton donors
  • they release H+ ions when they’re mixed with water
  • H+ ions are always combined with H2O to form hydroxonium ions, H3O+
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

What is the reaction to form hydroxonium ions?

A
  • HA(aq) + H2O(l) → H3O+(aq) + A-(aq)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

What are Bronsted-Lowry bases?

A
  • bases are proton acceptors
  • when they’re in solution, they take hydrogen ions from water molecules
  • B(aq) + H2O(l) → BH+(aq) + OH-(aq)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

What are strong acids and strong bases?

A
  • strong acids dissociate almost completely in water (nearly all the H+ ions will be released
  • e.g. HCl is a strong acid
  • strong bases dissociate almost completely too
  • e.g. NaOH
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What are weak acids?

A
  • weak acids dissociate only very slightly in water
  • only small numbers of H+ ions are formed
  • an equilibrium is set up, which lies well over to the left
  • e.g. ethanoic acid
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

What are weak bases?

A
  • weak bases only slightly protonate in water
  • equilibrium lies on the left
  • e.g. ammonia
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

What are conjugate pairs?

A
  • conjugate pairs are species that are linked by the transfer of a proton
  • the species that has lost a proton is the conjugate base
  • the species that has gained a proton conjugate acid
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What the definition of the standard enthalpy change of neutralisation?

A
  • the standard enthalpy change of neutralisation is the enthalpy change when the acid and alkali neutralise each other under standard conditions to form one mole of water
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

How do you calculate pH?

A
  • pH = -log10 [H+]
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

How do you work out [H+]?

A
  • 10-pH
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

What is the H+ concentration of strong monoprotic acids?

A
  • the H+ concentration is the same as the acid concentration
  • they dissociate fully
  • monoprotic: each mole of acid produces one mole of hydrogen ions
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

How many protons are released per molecule of diprotic acid?

A
  • two protons
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

How do you find the pH of a weak acid?

A
  • use Ka (the acid dissociation constant)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What are the assumptions you have to make to find Ka?

A
  • [HA(aq)](start) ≈ [HA(aq)](equilibrium)
  • [H+(aq)] ≈ [A-(aq)]
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

What is the ionic product of water?

A
  • KW = [H+] [OH-]
  • units are always mol2dm-6
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

How do you use KW to find the pH of a strong base?

A
17
Q

What is KW at 25°C (298K)?

A
  • Kw = 1.0 x 10-14 mol2dm-6
18
Q

What is pKw at standard conditions?

A
  • always 14.00
19
Q

How do you calculate pKa and Ka from pKa?

A
  • pKa = -log10 Ka
  • Ka = 10-pKa
20
Q

How do you calibrate a pH meter?

A
  • place the bulb of the pH meter into deionised water and allow reading to settle
  • adjust reading so that it read 7.0
  • do the same with a standard solution of pH4 and pH 10
  • rinse the probe with deionised water between each reading
21
Q

How does the pH of a strong acid change when diluted?

A
  • diluting the acid by a factor of 10 increases the pH by 1
  • you can see by using pH = -log10[acid]
22
Q

How does the pH of a weak acid change when diluted?

A
  • diluting a weak acid by a factor of 10 increases the pH by 0.5
23
Q

Describe these titration curves

A
  1. strong acid/strong base
  2. strong acid/weak base
  3. weak acid/strong base
  4. weak acid/weak base
24
Q

How do you explain the shapes of titration curves?

A
  • the initial pH depends on the strength of the acid
  • a strong acid titration will start at a much lower pH than a weak acid
  • to start with, the addition of small amounts of the base have little impact on the pH of the solution
  • final pH depends on the strength of the base
  • the stronger the base, the higher the final pH
25
Q

What is the equivalence point?

A
  • the centre of the equivalence line
  • [H+] ≈ [OH-]
  • the acid is just neutralised
  • a tiny amount of base causes a sudden, big change in pH
26
Q

How do you find the pKa of a weak acid using titration curves?

A
  • find the pH at the half-equivalence point
  • half of the acid has been neutralised
  • at the half-equivalence point, [HA] = [A-]
  • therefore, Ka = [H+] and pKa = [pH] (see image)
27
Q

What is a buffer?

A
  • a buffer is a solution that minimises changes in pH when small amounts of acid or base are added
28
Q

What are the two ways of making acidic buffers?

A
  1. Mix a weak acid with the salt of its conjugate base:
    e. g. ethanoic acid and sodium ethanoate
  2. Mix an excess of weak acid with a strong base
    e. g. ethanoic acid and sodium hydroxide
29
Q

What is the equilibrium set up for acidic buffers?

A
30
Q

How do acidic buffers work?

A
  • the conjugate base accepts any extra H+ and conjugate acid release H+ if there is too much base
  • if you add a small amount of acid, the H+ concentration increases
  • most of the extra H+ ions combine with CH3COO- ions to form ethanoic acid
  • this shifts equilibrium to the left, reducing the H+ concentration close to original value
  • if a small amount of base is added, the OH- concentration increases
  • most of the extra OH- ions reaction with H+ ions to form water
  • this causes more CH3COOH to dissociate to form H+ ions, shifting equilibrium to the right
31
Q

What are alkaline buffers made of?

A
  • a weak base and one of its salts
  • e.g. a mixture of ammonia solution (a base) and ammonium chloride (a salt of ammonia)
32
Q

How do alkaline buffers work?

A
33
Q

How can buffer solutions been seen on a titration curve?

A
34
Q

Why are buffer solutions important in blood?

A
  • blood needs to be kept at pH 7.4
  • this is controlled by carbonic acid-hydrogencarbonate buffer system
  • carbonic acid dissociates into H+ ions and HCO3- ions

H2CO3(aq) ⇔ H+(aq) + HCO3-(aq)

  • if the concentration of H+ rises in the blood, then the HCO3- ions will react with the excess H+ ions
  • the equilibrium will shift to the left, reducing the H+ concentration to its original value
  • vice versa
  • H2CO3 concentration is controlled by respiration

H2CO3(aq) ⇔ H2O(l) + CO2(aq)

  • HCO3- levels are controlled by kidneys
35
Q

What assumptions need to be made to calculate the pH of a buffer solution?

A
  • the salt of the conjugate base is fully dissociated, so assume that the equilibrium concentration of A- is the same as the initial concentration of the salt
  • HA is only slightly dissociated, so assume that its equilibrium concentration is the same as its initial concentration
36
Q

How do you calculate the concentration of salt and acid or base to create a buffer with a specific pH?

A

A Level Chemistry (19 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions