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chemistry module 5 > pH + buffers > Flashcards

pH + buffers Flashcards

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1
Q

bronsted-lowry acid definition

A

proton donor

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2
Q

bronsted-lowry base definition

A

proton acceptor

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3
Q

conjugate acid-base pair definition + 3 examples

A

2 species which differ by 1 H+ ion
each species can either act as an acid and donate the proton to become a base, or act as a base and accept the proton to become an acid
e.g. Cl - and HCl
NH3 and NH4 +
H2O and H3O +

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4
Q

monobasic definition + 2 examples

A

acids which can only donate 1 proton/ H+
e.g. HCl and CH3COOH

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5
Q

dibasic definition + 2 examples

A

acids which can donate 2 protons
e.g. H2SO4 and H2CO3

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6
Q

tribasic definition + 2 examples

A

acids which can donate 3 protons
e.g. H3PO4 and H3BO4

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7
Q

how can monobasic/dibasic acids change features of a practical e.g. a titration

A

changes volumes needed for neutralisation
e.g. titrating acids with 0.1moledm-3 NaOH
25cm3 of 0.1moldm-3 HCl would need 25cm3 NaOH to neutralise as 1:1 ratio of H+:OH-
25cm3 of 0.1moldm-3 H2SO4 would need 50cm3 to neutralise as 2:1 ratio of H+:OH-

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8
Q

give the 2 pH equations

A

pH = -log[H+(aq)]
[H+] = 10^-pH

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9
Q

what is the effects of logs on the pH scale

A

as logs are used, the pH scale is 10 fold
as it is negative logs, high [H+] = low pH and low [H+] = high pH

so pH 6 > 7 = [H+] x 10^-1

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10
Q

strong acid definition

A

an acid which fully dissociates into ions in aqueous solution

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11
Q

weak acid definition

A

an acid which only partially dissociates into ions in aqueous solution

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12
Q

give the equation for the dissociation of a weak acid

A

HA(aq) <> H+(aq) + A-(aq)

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13
Q

ka definition + equation

A

ka = acid dissociation constant
ka = [H+][A-]/[HA]

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14
Q

what does a high ka value mean

A

high ka means equilibrium position is further to the right so greater acid dissociation, meaning greater acid strength

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15
Q

give the 2 pka equations

A

pka = -log(ka)
ka = 10^-pka

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16
Q

what are 3 assumptions made when calculating ka or using the acid dissociation equation

A
  • [H+] = [A-]
    therefore ka = [H+]^2/[HA]
  • initial [HA] = [HA] at equilibrium, we assume that since weak acids don’t dissociate much the dissociation is small enough that concentrations can be taken as equal
  • the contribution of H+ from water is negligible
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17
Q

give equation and value for the ionic product of water kw

A

kw = [H+][OH-]
kw = 1 x 10^-14 mol2dm-6 at 25C

18
Q

what is the pH of pure water at 25C

A

7

19
Q

how does temperature affect kw

A

ionisation of water is endothermic so increased temp = increased kw

20
Q

give a the reaction equation for the ionisation of water and the ka equations relative to kw

A

H2O(l) <> H+(aq) + OH-(aq)
ka = [H+][OH-]/[H2O]
ka x [H2O] = [H+][OH-] = kw

21
Q

what is the assumption made for strong acids

A

as strong acids ionise fully, [HA] = [H+]

22
Q

why can the pH of 1x10^-8 moldm-3 HCl(aq) not be 8

A

pure water has a pH of 7, the 1x10^-8moldm-3 of H+ adds more H+ to water, so it cannot have a pH greater than 7

23
Q

buffer definition

A

a solution that resists/minimises pH changes when small amounts of acids/bases are added

24
Q

what are buffers made of + 1 example

A

weak acids + conjugate base (salt)
e.g. ethanoic acid + -ethanoate

25
Q

give an equation for the action of a buffer and describe how it minimises pH changes

A

HA <> H+ + A-

when pH falls due to the addition of acid, H+ from the acid reacts with the conjugate base A- to form HA, moving position of equilibrium to the left, as [H+] decreases, pH rises therefore counteracting the effect of the acid

when pH rises due to the addition of base, OH- reacts with H+ to make water, causing the equilibrium to move to the right as HA ionises to replace lost H+, as [H+] increases, pH falls therefore counteracting the effect of the base

26
Q

explain why the pH of a weak acid rises when its salt is added to it

A

the acid exists in an equilibrium: HA <> A- + H+
when salt (e.g. XA) is added, it ionises completely, producing more A- and pushing the equilibrium to the left, as [H+] decreases this causes pH to increase

27
Q

what are 2 ways a buffer can be made

A
  • weak acid + base
  • weak acid + its salt
28
Q

describe a method to make a buffer from a weak acid + a base

A
  • in a beaker add excess weak acid and a small amount of a base e.g. NaOH
  • acid + base > salt + water
  • however, as the acid is in excess, this will ensure that a partial neutralisation occurs as there is not enough base for a full neutralisation
  • therefore the solution will contain an acid and its salt/conjugate base
29
Q

describe a method to make a buffer from a weak acid + its salt

A
  • equilibrium of weak acid: HA <> A- + H+
  • this dissociation of a weak acid is very small so mostly HA is left
  • equilibrium of salt: XA <> A- + X+
  • salts dissolve in water completely therefore ionise completely, so mostly A- is left
  • this gives a mix of weak acid + conjugate base
30
Q

outline how to calculate the pH of buffer solutions

A
  • ka = [H+][A-]/[HA] can be used
  • the assumption that [H+]=[A-] is not true for buffers as A- is one of the components
  • therefore [H+] = ka x [HA]/[A-] should be used
  • when [HA]=[A-] then ka=[H+] therefore pka=pH
31
Q

give an example of a buffer in the body

A

H2CO3 <> HCO3- + H+
found in the blood

32
Q

explain how the carbonic acid-hydrogencarbonate buffer system controls blood pH

A

if acid is added, H+ ions from acid react with HCO3- to form H2CO3, moving position of equilibrium to the left causing [H+] to decrease, counteracting effect of acid
if base is added, OH- ions react with H+ ions to form water, causing more H2CO3 to dissociate so position of equilibrium moves to the right, causing [H+] to increase, counteracting effect of the base

33
Q

at around what pH would the end point be reached in a strong acid-strong base titration

A

around pH 3-11

34
Q

at around what pH would the end point be reached in a strong acid-weak base titration

A

around pH 3-7

35
Q

at around what pH would the end point be reached in a weak acid-strong base titration

A

around pH 7-11

36
Q

at around what pH would the end point be reached in a weak acid-weak base titration

A

around pH 7

37
Q

describe the process of using a pH meter

A
  • probe of pH meter must be calibrated using buffer solutions of known pH
  • temperature setting of pH probe must match temp of unknown solution
  • probe must be placed vertically in solution and the pH shouldn’t be read until steady
  • probe should be washed with deionised/distilled water between each measurement
38
Q

on a titration curve, what indicates the end point

A

the vertical section of the graph

39
Q

what is the point in the centre of the vertical section called

A

the equivalence point
- this would line up with a neutral pH and with the rough volume of base that would be needed to neutralise acid

40
Q

how is the curve different if it shows a titration where acids are added to bases

A

the curve will go down from left to right, still the same shape

41
Q

how can you use a titration curve to tell if in indicator is suitable

A

if the pH range of the indicator lies within the vertical region of the curve

42
Q

how do indicators work

A

indicators are basically coloured buffers, and work in the same way
can be HA <> A- + H+
or AOH <> A+ + OH-

chemistry module 5 (5 decks)

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