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a level chemistry > 3.9 acid-base equilibria > Flashcards

3.9 acid-base equilibria Flashcards

1
Q

an acid is a proton (donor/acceptor)?

A

donor
(H+ donor)

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2
Q

a base is a proton (donor/acceptor)?

A

acceptor
(H+ acceptor)

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3
Q

it is possible that a species that behaves like an acid in one reaction will behave like a base in another

A
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4
Q

in acid-base reactions we get conjugate acid-base pairs

A

e.g NH3 + H2O —> NH4+ + OH-
base acid CA CB

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5
Q
  • nitric acid, HNO3, is a strong Bronsted-Lowry acid
  • nitrous acid, HNO2, is a weak Bronsted-Lowry acid
  • A student suggests that an acid-base equilibrium is set up when nitric acid is mixed with nitrous acid
  • complete the equation for the equilibrium that would be set up and label the conjugate acid-base pairs for:

HNO3 + HNO2 ⇌ ____ + _____

A

HNO3 + HNO2 ⇌ NO3 - + H2NO2 +

HNO3 = acid
HNO2 = base
NO3 - = conjugate base
H2NO2 + = conjugate acid

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6
Q

what are strong acids?

A
  • acids that fully dissociate
  • HA —> H+ + A-
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7
Q

what are weak acids?

A
  • acids that only partially dissociate (in a reversible reaction)
  • HA ⇌ H+ + A-
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8
Q

what does the degree of dissociation of a weak acid depend on?

A

its acid dissociation constant, Ka

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9
Q

express Ka for HA ⇌ H+ + A-

A

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

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10
Q

for strong acids, HA —> H+ + A- and the concentration of HA = H+

A
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11
Q

for weak acids, HA ⇌ H+ + A-, with conc of HA 0.1M, how do you find out the concentration of H+?

A

Ka = [H+][A-] / [HA]
- [H+] = [A-]
- [HA] - approximately 0.1M because some has split already

so: (for weak acids)
Ka = [H+]^2 / [HA]

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12
Q

what are some examples of weak acids?

A
  • ethanoic acid
  • any carboxylic acid
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13
Q

need to remember the pH calculations for AS

A
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14
Q

what are the 2 assumptions we make to say that for weak acids, Ka = [H+]^2 / [HA]?

A
  • assume [H+] = [A-] because they have dissociation according to a 1:1 ratio
  • assume the initial concentration of the undissociated acid remains constant because the amount of dissociation of a weak acid is so small
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15
Q
  • ethanoic acid CH3COOH
  • ethanoic acid is a weak acid with an acid dissociation constant, Ka of 1.75×10^-5 moldm^-3 at 25°C
  • the student uses a pH meter to measure the pH of a solution of CH3COOH at 25°C
  • the measured pH is 2.440
  • calculate the concentration of ethanoic acid in the solution [3]
A

Ka = 1.75x10^-5
Ka = [H+] / [HA]
[H+] = 10 ^ -2.44 = 3.63 x10^-3
1.75x10^-5 = (3.63)^2 / [HA]
[HA] = 0.753

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16
Q

the dissociation of water equation:

A

H2O ⇌ H+ + OH-
- water is always slightly dissociated into hydrogen ions and hydroxide ions

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17
Q

the concentration of water [H2O] is constant, such that it can be incorporated into a new constant Kw, where at 298K:
Kw=

A

Kw = [H+][OH-]
= 1.0 x10^-14 mol^2dm^-6

  • this is the ionic product of water
  • this means that if the conc of H+ ions increases (e.g by adding acid to the solution), the conc of OH- ions must decrease proportionally
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18
Q

any solution in which the concentrations of H+ and OH- are equal are said to be what?

A

neutral

[H+] = [OH-]

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19
Q

the dissociation of water is an (endothermic/exothermic) reaction?

A

endothermic

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20
Q

1cm^3 = 1g H2O

Mr= 18.02
mass = 1g
n = 0.055
v = 0.001dm^3
so conc = 55.5

[H2O] = 55.5 moldm^-3
- constant in any aq solution

A
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21
Q

what is the value of the Kw of water?

A

1 x10^-14 mol^2dm^-6
- tiny so equilibrium very to RHS
- ∴ H2O hardly dissociated

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22
Q

neutrality rests on the ____ of the water

A

temperature

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23
Q

Kw = [H+][OH-]
- if non-neutral i.e basic, Kw = ?
- if neutral bc [H+]=[OH-], Kw = ?

A
  • Kw = [H+][OH-]
  • Kw = [H+]^2 (pure)
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24
Q

what is a strong base?

A

one in which fully dissociated into its ions in water

(strong bases in water are rare as water is a very weak acid and will not readily give up protons to other species)

25
substances which dissolve in water to produce an excess of OH- ions are said to be ____
alkaline
26
what is the pH of the strong base 0.1M NaOH?
- NaOH —> Na+ + OH- 0.1M 0.1M 0.1M Kw= [H+][OH-] 1x10^-14 = 0.1[H+] [H+] = 1x10^-13 pH = 13
27
pKa is sometimes used instead of Ka when discussing relative acidities of acids
- this is because acids can have very small values of Ka and using pKa is simply a more convenient way of comparing the degree of acid dissociation or strenght
28
pKa equation:
pKa = -log10 Ka (Ka = 10^ -pKa)
29
- the more positive Ka, the (stronger/weaker) the acid - the more negative the pKa, the (stronger/weaker) the acid
- stronger - stronger (lower pKa = stronger acid higher Ka = stronger acid)
30
what is a buffer solution?
- one that minimises pH changes on addition of small amounts of acid or alkali
31
buffers are used to keep the pH constant in lots of everyday solutions including:
- shampoos - biological washing powders - blood
32
what are the two types of buffer systems?
- acidic buffers (*wjec focuses on this one) - basic buffers
33
what are acidic buffers made from?
- weak acid and its salt e.g: - ethanoic acid and sodium ethanoate - propanoic acid and lithium propanoate
34
weak acid and its salt: CH3COOH ⇌ CH3COO- + H+ high conc low low (bc small dissociation = any weak acid) if add CH3COONa (salt of weak acid) (salts fully dissolve in solutions) high high high (= any buffer)
35
what happens if acid is added to an acidic buffer? (ethanoic acid + sodium ethanoate)
- H+ added - CH3COO- reacts with H+ - added H+ all reacts - so no extra H+ - since [H+] stays fairly constant, pH stays fairly constant
36
what happens if a base is added to an acidic buffer? (ethanoic acid + sodium ethanoate)
- OH- added - reacts with H+ (producing H2O) - decreases H+ ions - so equilibrium shift to the right - such that more CH3COOH is dissociated - so increases [H+] to similar level again, keeping pH fairly constant since weak acids are only slightly dissociated there is plenty of CH3COOH in the solution to replace the H+ ions that were removed by the OH- ions
37
what are basic buffers made from?
- a weak base and its salt e.g - ammonia and ammonium chloride (has pH > 7)
38
what happens if acid is added to a basic buffer? (ammonia and ammonium chloride)
- H+ added - H+ reacts with the OH- to form H2O - equilibrium shifts to the right in order to replace the OH- ions that have been removed - [OH-] remains fairly constant, so [H+] stays fairly constant - so pH remains constant
39
what happens if a base is added to a basic buffer? (ammonia and ammonium chloride)
- OH- added - equilibrium shifts to left to reduce the concentration of OH- again - [OH-] remains fairly constant - so pH remains constant
40
what are some limitations to buffer systems?
- can only resist pH changes when small amounts of H+ or OH- are added because there is a limit to the amounts of weak acid and its salt (or weak base and its salt) available to react - once all the acid molecules in an acidic buffer have dissociated, any further addition of alkali will cause pH to increase rapidly. at this point the buffer is saturated (in same way, once all ions from salt have combined with H+ to form acid molecules again, any further addition of H+ will cause pH to decrease rapidly)
41
what is the hydrogen ion concentration of a solution of hydrochloric acid which has a pH of 2.0?
[H+] = 10^-pH = 10^-2 = 0.01 moldm^-3
42
give examples of strong acids and state the pH range which indicates a strong acid
- hydrochloric acid HCl - sulfuric acid H2SO4 - nitric acid HNO3 - pH range of strong acids: 0-3
43
give examples of strong bases:
- sodium hydroxide (NaOH) - potassium hydroxide (KOH) - calcium hydroxide (Ca(OH)2) pH 12-14
44
give an example of a weak base:
- ammonia (NH3) pH just above 7-11
45
what is the acid dissociation constant, Ka?
is a measure of how strong an acid is in a solution
46
what are the units for Ka?
moldm^-3
47
derive the ionic product of water using the equation for the ionisation of water:
H2O ⇌ H+ + OH- - so Kc = ([H+][OH-]) / [H2O] - since [H2O] is very large compared to [H+] and [OH-], [H2O]Kc can be considered to be constant - then [H2O]Kc = Kw - and so Kw = [H+][OH-]
48
what is the pH of pure water at room temperature?
7
49
how does the pH at which water is neutral ([OH-] = [H+]) change as temperature increases?
- the forwards reaction is endothermic - and is therefore favoured when the temperature of the water is increased - so the pH at which water is neutral decreases when the temperature is increased
50
what is the equivalence point on a pH curve?
- point at which the pH curve is vertical - the point at which the solution has been neutralised
51
what does a pH curve of a strong base added to a strong acid look like?
- pH is initially at around 1 as the strong acid is in excess - thr pH ends up being very high, indicating a strong base is in excess
52
what does a pH curve of a strong base added to a weak acid look like?
- pH starts about 5 whrre theres an excess of weak acid - it finishes with a high pH where theres an excess of strong base
53
what does a pH curve of a weak acid added to a weak base look like?
- pH starts about 8-9 where theres an excess of weak base - it finishes with a pH around 5 when theres an excess of weak acid
54
why are buffers used in industrial processes?
- to maintain the optimum reaction conditions for large scale manufacturing
55
what needs to be considered when selecting an indicator for a titration?
- the indicator chosen for titration must change colour at the neutralisation point - the indicator should change colour over a narrow pH range which coincides with the vertical section of the pH curve
56
indicator: phenolphthalein colour at low pH: pH of colour change: colour at high pH:
colourless 8.3-10 pink
57
indicator: methyl orange colour at low pH: pH of colour change: colour at high pH:
red 3.1-4.4 yellow
58
why should a pH meter be used instead of an indicator in weak acid/weak base titrations?
- there is no sharp pH change so very difficult to successfully use an indicator to get an accurate point of neutralisation - a pH meter should be used instead since it will be more accurate and precise

a level chemistry (21 decks)

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