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chemistry > topic 12 > Flashcards

topic 12 Flashcards

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1
Q

Bronsted-Lowry acid

A

a substance that can donate a proton

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2
Q

Bronsted-Lowry base

A

a substance that accepts a proton

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3
Q

Ka formula

A

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

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4
Q

calculating pH

A

pH = -log [H+]

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5
Q

define the term pH

A

a figure expressing the acidity or alkalinity of a solution on a logarithmic scale on which 7 is neutral, lower values are more acid and higher values more alkaline

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6
Q

what happens to strong acids in water

A

they completely dissociate

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7
Q

finding [H+] from pH

A

[H+] = 1 x 10–pH

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8
Q

what happens to weak acids in water

A

only slightly dissociate - separate into ions

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9
Q

Ka of a weak acid

A

simplifies to : Ka = [H+ (aq)]^2
[HA (aq)]

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10
Q

what assumptions are made when calculating pH of a weak acid

A
  1. [H+] = [A-] as they have dissociated according to a 1:1 ratio.
  2. as the amount of dissociation is small we assume that the initial concentration of the undissociated acid has remained constant.i
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11
Q

ionic product of water definition

A

in all aqueous solutions and pure water the following equilibrium occurs:

H2O ⇌ H+ + OH-

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12
Q

Kw equation

A

Kw = [H+][OH-]

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13
Q

deducing Kw

A

Kc = [H+] [OH-]
[H2O]

Kc x [H2O] = [H+] [OH-]

Kw = [H+][OH-]

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14
Q

calculating the pH of strong base

A

use Kw to work out [H+]
use pH equation

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15
Q

pKw

Kw from pKw

A
  • pKw = -log Kw
  • Kw = 10^-pKw
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16
Q

Kw of all aqueous solutions

A

1 x 10^-14

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17
Q

finding pH of pure water

A

pure water and neutral solutions are neutral because the [H+] = [OH-]

so Kw = [H+]^2 which means [H+] is the square root of Kw

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18
Q

what happens to Kw as temp increases

A

increases

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19
Q

weak acid dissociation expression

A

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

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20
Q

what does Ka measure

A

Ka measures acid strength

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21
Q

pKa

A

pKa = -log Ka

so

Ka = 10 ^-pKa

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22
Q

large Ka strong or weak acid?

A

strong

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23
Q

small Ka strong or weak acid?

A

weak

24
Q

working out pH of weak acid at half equivalence

A

we can assume that [HA] = [A-]
so [H+] = Ka and pH = pKa

25
Q

pH of diluted strong acid

A

[H+] = [H+]old x old volume / new volume

pH = -log [H+]

26
Q

pH of diluted base

A

[OH-] = [OH-]old x old volume
new volume
[H+] = Kw
[OH-]
pH = -log [H+]

27
Q

comparing the pH of a strong acid and a weak acid after dilution 10 times

A

diluting a strong acid 10 times will increase its pH by one unit

28
Q

comparing the pH of a strong acid and a weak acid after dilution 100 times

A

diluting a strong acid 100 times will increase its pH by two units

29
Q

what would happen to weak acids when they are diluted

A

they would not change in the same way. they increase by less than 1 unit

30
Q

what happens to equilibrium when diluting a weak acid

A

pushes equilibrium to the right so the degree of dissociation increases and more H+ ions are produced meaning pH increases less than expected

31
Q

buffer solution definition

A

one where the pH does not change significantly if SMALL amounts of acid or alkali are added to it

32
Q

what is an acidic buffer made from

A

a weak acid and a salt of that weak acid - weak acid reacting with a strong base

33
Q

e.g. of an acidic buffer

A

ethanoic acid and sodium ethanoate

CH3CO2H and CH3CO2-Na+

34
Q

what is a basic buffer made from

A

from a weak base and a salt of that weak base - weak base reacts with strong acid

35
Q

e.g. of a basic buffer

A

ammonia and ammonium chloride

NH3 and NH4+Cl-

36
Q

ethanoic acid buffer

A

CH3CO2H ⇌ CH3CO2- + H+
acid……………….conjugate
……………………… base

37
Q

what happens to buffer solution if small amounts of acid are added

A

equilibrium will shift to the left removing nearly all the H+ ions added
moles of buffer acid would increase by the same amount so a new calculation of pH can be done with new values

38
Q

what happens to buffer solution if small amounts of alkali are added

A

the OH- ions react with H+ ions to form water
moles of salt would increase by the same amount so a new calculation of pH can be done with new values

39
Q

calculating the pH of buffer solutions

A

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

40
Q

controlling pH in blood

A

carbonic acid-hydrogencarbonate eqb acts as a buffer in the control of blood pH
H2CO3 is present in blood plasma - maintaining a pH between 7.35 and 7.45

41
Q

constructing a titration/pH curve

A
  1. transfer 25 cm³ of acid to a conical flask with a volumetric pipette
    1. measure initial pH of the acid with a pH meter
    2. add alkali in small amounts (2 cm³) noting the volume added
    3. stir mixture to equalise it
    4. measure and record the pH to 1 d.p.
    5. repeat steps 3-5 but when approaching end point add in smaller volumes of alkali
    6. add until alkali is in excess
42
Q

calibrating pH meter

A

measure known pH of a buffer solution

43
Q

why is calibrating pH meter important

A

pH meters can lose accuracy on storage

44
Q

calibrating a pH probe

A

putting a probe in a set buffer (often at pH 4) and pressuring a calibration button/setting that for that pH

45
Q

strong acid - strong acid curve

e.g.

A

long vertical line is between pH 3 to 9
pH at equivalence point = 7
initial pH ~1

HCl & NaOH

46
Q

weak acid - strong base curve

e.g.

A

the equivalence point is more than 7
initial pH ~ 3
final pH ~ 13

CH3CO2H & NaOH

47
Q

strong acid - weak base curve

e.g.

A

vertical line less than 7 (4 to 7)
the equivalence point is less than 7
initial pH ~ 1
final pH ~ 9

HCl & NH3

48
Q

weak acid - weak base

e.g.

A

no vertical part

CH3CO2H & NH3

49
Q

pH range of phenolphthalein &
colour change

A

~ 7.5 to 9
use for strong bases
colourless to pink

50
Q

pH range for methyl orange & colour change

A

~ 2 to 4
use with strong acids
red acid to yellow alkali

51
Q

half neutralisation volume - for weak acids - equation

spec; 22.ii

A

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

at 1/2 neutralisation volume the [HA] = [A-]

so, Ka = [H+] and pKa = pH

52
Q

enthalpy change of neutralisation

A

the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce one mole of water

53
Q

are enthalpy changes of neutralisation endothermic and exothermic

A

always exothermic

54
Q

enthalpy changes of neutralisation for reactions with strong acids and bases

A

values are similar (about -56 to -58 kJ mole^-1) because the same reaction is occurring

55
Q

enthalpy changes of neutralisation for reactions with weak acids and bases

A

less exothermic enthalpy change because energy is absorbed to ionise the acid and break the bond to the hydrogen in the un-dissociated acid

chemistry (28 decks)

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