Physical Unit 1.12: Acids & Bases (A2) Flashcards
what is the Bronsted-Lowry definition for an acid, base & acid-base reaction/equilibria?
acid: proton donor
base: proton acceptor
acid-base reaction/equilibria: involves the transfer of protons
what is the pH scale?
logarithmic scale
a measure of hydrogen ion concentration
formula for pH (strong acids)
pH = -log[H+]
always to 2dp
[H+] =
10^-pH
define monobasic & dibasic acid & give e.g.s
monobasic acid: acids where each molecule dissociates to form 1H+ ion e.g. HNO3, HCl
= monoprotic
dibasic = diprotic: each molecule dissociates to form 2 H+ ions e.g. H2SO4
define strong acid
(in solution) all molecules dissociate to form H+ ions
[H+] in dilution of a strong acid
[H+] = (volume of acid/total volume) x conc. acid
it’s the proportion of the volume x conc.
what is Kw, the ionic product of water?
water dissociates slightly:
H2O <–> H+ + OH- endothermic
conc. water is 55.6 moldm-3
Kw is derived from Kc of this dissociation
what is the formula for Kw? & derivation
Kc = [H+][OH-] / [H2O]
Kw = [H+][OH-]
define neutral
[H+] = [OH-]
how does temperature affect pH of water & neutrality of water?
pure water is neutral
increasing temp. shifts equilibrium in endothermic direction, in forwards direction
= which increases [H+] & [OH-]
= pH changes
but pure water stays neutral as [H+] = [OH-]
as temp. increases, Kw increases
what is Kw @ room temp.?
1x10^-14
@ stp in pure water, Kw =
pH =
Kw = [H+]^2 = 1x10^-14
[H+] = root 1x10^-14 = 10^-7
pH = 7.00
define monoprotic & diprotic base
monoprotic base: 1 mole of base accepts 1 mole of H+ ions
/
each molecule accepts 1 H+ ion
diprotic base: 1 mole of base accepts 2 moles of H+ ions
each molecule accepts 2 H+ ions
how do you calculate pH of a strong base?
use Kw = 1x10^-14 @stp
use [H+] = Kw/[OH-]
to calculate conc. of OH-, multiply conc. of substance by # moles of OH- in one molecule
to calculate conc. of base, divide conc. of substance by # moles of OH- in one molecule
how do you calculate the pH of a solution formed from the reaction b/w a strong acid & strong base?
- calculate moles H+
- calculate moles OH-
- calculate leftover moles at end, either of H+ or OH-
- calculate [H+] or [OH-] of leftover at end by mol/total vol.
- calculate pH
define weak acid
only a small fraction of the molecules dissociate to form H+ ions
e.g. carboxylic acids
weak acid equilibrium
HA <–> H+ + A-
HA = weak acid
A- = salt of weak acid
describe the salt of a weak acid
slightly basic bc proton accepted by A- to form HA
what is the formula for Ka, the acid dissociation constant for a weak acid?
what are the assumptions made?
Ka = [H+][A-] / [HA]
assumptions:
1. [HA] = original [HA], the conc. doesn’t change much due to equilibrium
2. [H+] only comes from the dissociation of HA & not from water/Kw
what is the formula for pKa?
pKa = -log(Ka)
Ka = 10^-pKa
important A- info
in a solution of pure weak acid dissolved in water, A- comes from the dissociation of weak acid (HA)
so [A-]=[H+]
if any ionic salt is dissolved in water, [A-] = [salt] & A- from the dissociation of weak acid is negligible
how do you calculate the pH of solution formed from the reaction b/w weak acid & strong base?
HA + OH- –> A- + H2O
1. calculate moles HA
2. calculate moles OH-
3. calculate excess moles (at end) or HA or OH-
if excess HA:
4. calculate moles HA left & moles A- formed
5. calculate [HA] leftover & [A-] formed
6. use Ka to find [H+]
7. find pH
if excess OH-:
4. calculate [OH-]
5. use Kw to find [H+]
6. find pH
if excess base, it is irrelevant whether acid was weak or strong bc it has all reacted
what is half neutralisation of weak acid?
when half of HA molecules have reacted with OH-
what is true at half neutralisation of weak acid?
[HA]=[A-]
so Ka = [H+]
pKa = pH
define indicator
weak acids where HA & A- are different colours
(the pH at which the indicator changes colour varies b/w indicators)
for indicators:
at low pH
at high pH
at low pH, HA is the main species present
at high pH, A- is the main species present
what does universal indicator contain?
it is a mixture of different indicators so shows several colours at each pH
table of methyl orange & phenolphthalein, colour of HA, colour of A-, pH range of colour change
methyl orange:
red
yellow
3.2-4.4
phenolphthalein:
colourless
pink
8.2-10.0
for an indicator to change colour where moles acid = moles base,
it must change colour within the range of rapid pH change at the end point of the titration
define equivalence point
point when moles acid = moles alkali - but pH not always 7
rapid pH change around equivalence point
define end point
when indicator changes colour - should coincide with equivalence point if correct indicator is used
what do pH curves look like for: strong acid-strong base
strong acid-weak base
weak acid-strong base
weak acid-weak base
see booklet
weak acid-weak base curve has no vertical section
define buffer solution
solution that resists changes in pH when small amounts of acid or alkali are added
so an approximately constant pH is maintained
define acidic buffer
& define basic buffer
acidic buffer solutions contain a weak acid & the salt of that weak acid
pH < 7
[acid] & [salt] are much higher than [H+]
basic buffer solutions contain a weak base & the salt of that weak base
pH > 7
[base] & [salt] are much higher than [OH-]
how can an acidic buffer also be made?
basic?
excess HA & strong alkali –> HA & A- mixture
excess weak base & strong acid –> weak base + its salt
[H+] =
Ka[HA] / [A-]
[H+] α HA/A-
what is the effect of adding small amounts of acid or alkali on the ratio of HA:A-
the ratio stays roughly constant so pH hardly changes
describe what happens when a little H+ is added to acidic buffer
write equilibrium equation
HA <–> H+ + A- related to Q
the added H+ is removed by reaction with A- to form HA
this shifts the position of equilibrium to the left
describe what happens when a little H+ is added to basic buffer
write equilibrium equation
related to Q e.g. NH3 + H2O <–> NH4+ + OH-
the added H+ is removed by reaction with OH-
so some NH3 reacts to replace OH
so the position of equilibrium shifts right
describe what happens when a little OH- is added to acidic buffer
write equilibrium equation related to Q
HA <–> H+ + A-
the added OH- is removed by reaction with H+
so some HA breaks down/dissociates to replace H+
so the position of equilibrium shifts right
describe what happens when a little OH- is added to basic buffer
write equilibrium equation related to Q
e.g. e.g. NH3 + H2O <–> NH4+ + OH-
the added OH- is removed by reaction with NH4+ to form NH3
so the position of equilibrium shifts left
how do you calculate the pH of acidic buffers?
- write out:
Ka = [H+][A-] / [HA]
[H+] = Ka[HA] / [A-] - write out equation
- calculate moles HA left & moles A- formed
- calculate [HA] leftover & [A-] formed
- use Ka to find [H+]
- find pH
explain why the expression for Kw does not include the concentration of water
the concentration of water is a constant
water only slightly dissociates so its concentration is much greater than [H+] or [OH-]
how does the value of Kw change as temperature increases?
Kw increases
the forward reaction is endothermic
so as temperature increases, the position of equilibrium shift right to decrease temperature
why is [H+] squared in Ka & Kw expressions?
for Kw, in pure water [H+] = [OH-]
for Ka, [H+] = [A-] because the ratio is 1:1
