3.9 - Acid - base equilibria Flashcards
1
Q
Acid?
A
proton donor
2
Q
Base?
A
proton acceptor
3
Q
Strong acid?
A
fully disassociates
4
Q
E.g?
A
HNO3 —– H+ + NO3-
H3PO4 ——- 3H+ + PO43-
5
Q
Weak acid?
A
partially disassociates
Ethanoic acid
CH3COO-
6
Q
pH def?
A
potential hydrogen
- log10 H+
7
Q
a change in 1 pH unit?
A
increases the H+ concentration by a factor of 10
8
Q
a chage in 2 pH units?
A
increases the H+ concentration by a factor of 100
9
Q
the negative sign?
A
the higher the pH, the lower the hydrogen ion concentration
10
Q
Ka?
A
disassociation constant
H2SO4 —- 2 H+ SO42-
11
Q
Ka equation?
A
(H+)2 (SO42-)
///////////////
H2SO4
12
Q
A weak acid?
A
low Ka value
13
Q
A strong acid?
A
high Ka value
14
Q
Nitric acid?
A
24 mol dm -3
15
Q
Sulfurous acid?
A
1.4 x 10 ^-3 mol dm -3
16
Q
Methanoic acid?
A
1.8 x 10 ^-4 mol dm -3
17
Q
Ethanoic acid?
A
1.7 x 10 ^ -5 mol dm-3
18
Q
Carbonic acid?
A
4.5 x 10 ^-7 mol dm-3
19
Q
Range of Ka = so large?
A
they re converted to pKa values to give a linear scale
20
Q
Pka of Nitric acid?
A
-log10(-24) = -1/38
21
Q
Pka of sulfurous acid?
A
-log10(1.3 x 10 ^-2) = 1.85
22
Q
Pka of methanoic acid?
A
- log 10 (1.8 x 10 ^-4) = 3.74
23
Q
Pka of ethanoic acid?
A
-log10 ( 1.7 x 10 ^-5) = 4.77
24
Q
Pka of carbonic acid?
A
-log10 = 4.5 x 10 ^- 7 = 6.35
25
Ionic product of water?
Kw
Water can be purified to remove all the impurities to obtain 100 % water
even after purification, water can still conduct electricity as they are still dissolved ions
These ions cannot be removed because they are produced by water itself.
26
Why does the reaction lie to the left?
water exists as molecules not ions
27
strong acid?
pH = -loga(H+)
28
weak acid?
H_ = square root of ka x acid
29
strong acids?
will only be monobasic
1 hydrogen
30
Buffer?
a solution that resists change in pH when small volumes of acid or alkali is added?
31
if large volumes of acid or alkali is added?
pH doesn't remain constant
only small volumes should be added
32
how does a buffer work?
maintains the constant pH by removing or adding H+ or OH- ions
33
making a buffer?
using ethanoic acid and salt of the same acid so sodium ethanoate
Disassociation of
CH3COONA -- CH3COO- + Na+
CH3COOH -- CH3COO- + H+
adding a small volume of acid to the buffer
HCl --- H+ + Cl-
34
The number of H+ ions increased?
the equilibria shifts to the left, so the numbers of H+ ions, decreases as they react to the anion CH3COO-
35
Adding an alkali to a buffer?
CH3COONa---- CH3COO- + Na+
CH3COOH --- CH3COO - +H +
NaOH --- Na+ + OH-
36
As the OH- increased?
it removes the H+ ions to form water, this causes the equilibrium for ethanoic acid to shift to the right
37
Uses of a buffer?
for storing enzymes as their pH needs to remain constant
Pharmaceutical industry to store biological molecules which could become denatured in the wrong pH
used in the brewing or baking industry to keep yeast at a pH of 7.4
38
Basic buffers?
maintain an alkaline pH, ammonium chloride = mixed with ammonia solution, allowing them both to disassociate.
NH4Cl (aq) ----- NH4(aq)
NH4+(aq) ---- NH3(g) + H+(g)
Adding OH- ions removes the H+ ions so the equilibrium shifts to the right to make more OH - ions
NH4Cl ( aq) ---- NH4+(aq) + Cl- (aq)
NH4 + (aq) ---- NH3(g) + H+ ( g)
Adding H + ions shifts the equilibrium to the left to try and decrease the number of H+ ions
39
work out the pH buffer that has equal concentrations of ethanoic acid + sodium ethanoate if the ka value of ethanoic acid?
CH3COOH ---- CH3COO- + H+
Ka = (CH3COOH)(H+)
/////////////////////////
(CH3COOH)
but if concentration is the same = 1
ka = ( H+)
1,7 x 10 ^-5 (H+)
pH = -log10(H+)
pH = -log10(1.7x10^-5)
pH = 4.77
40
Henderson - Hasselbach Equation?
pH buffer = pKa + log
( salt/acid)
41
pH of a neutral salt?
NaCl - made from a strong acid and a strong alkali
42
pH of an acidic salt?
Na4Cl
made from a strong acid + weak alkali
43
Basic salt?
example = CH3COONa
weak acid with strong alkali
44
pH graph for strong acid / strong base ?
If the concentration of NaOH = 0.1mol dm-3 and is added to an acid of the same concentration
( monobasic) HCl or HNO3
The graph starts at a pH of 1 because the acid = strong
There is a gradual increase in pH as the 20 cm3 is added.
There is a sudden increase in pH from 3-12 when the volume of the base = the volume of the acid ( this is the vertical region and must be drawn with a ruler)
when the last 20cm3 is added, there a gradual increase in pH which ends at 13.
45
The vertical region?
Known as the equivalence point when the concentrations of the 2 solutions are the same.
This will occur when the volumes are equal
If acid is weak, the graph changes as the H+ ions do not fully disassociate
46
pH graph for weak acid/strong base
pH gradually increases to 4 as the base = added
When the volume of NaOH = 1/2 the volume of the acid, the pH levels off
Known as the buffering region because the unreacted acid reacts with the salt formed to make a buffer
the pH reaches 5 when 25 cm3 of NaOH is added
the vertical region occurs between 5 and 11
As the remaining 25cm3 is added, the pH increases to 13.
47
pH graph for strong acid - weak base
the pH up to 7, is similar to the strong acid, strong base
Graph levels off between 35 and 45 cm3 due to the buffering effect
the pH increases gradually to 12.
48
pH graph for weak acid - weak base
There is only a small vertical region between 6 and 8
a pH probe should be used so the points can be plotted more accurately
the equivalence point is always a point of inflection
( where the gradient of the graph = 0 or infinity)
49
Indicator?
substances that change colour as pH changes and is made of weak acids
50
Indicators?
can appear different colours in acid or alkaline conditions
51
Examples of indicators?
Phenolphthalein( 8.3-10)
Bromothymol - (6.0-7.5)
Litmus ( 4.0-6.5)
Methyl orange ( 3.2-4.4)
all in vertical region
