topic 12 Flashcards
Bronsted-Lowry acid
a substance that can donate a proton
Bronsted-Lowry base
a substance that accepts a proton
Ka formula
Ka = [H+][A-]
[HA]
calculating pH
pH = -log [H+]
define the term pH
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
what happens to strong acids in water
they completely dissociate
finding [H+] from pH
[H+] = 1 x 10–pH
what happens to weak acids in water
only slightly dissociate - separate into ions
Ka of a weak acid
simplifies to : Ka = [H+ (aq)]^2
[HA (aq)]
what assumptions are made when calculating pH of a weak acid
- [H+] = [A-] as they have dissociated according to a 1:1 ratio.
- as the amount of dissociation is small we assume that the initial concentration of the undissociated acid has remained constant.i
ionic product of water definition
in all aqueous solutions and pure water the following equilibrium occurs:
H2O ⇌ H+ + OH-
Kw equation
Kw = [H+][OH-]
deducing Kw
Kc = [H+] [OH-]
[H2O]
Kc x [H2O] = [H+] [OH-]
Kw = [H+][OH-]
calculating the pH of strong base
use Kw to work out [H+]
use pH equation
pKw
Kw from pKw
- pKw = -log Kw
- Kw = 10^-pKw
Kw of all aqueous solutions
1 x 10^-14
finding pH of pure water
pure water and neutral solutions are neutral because the [H+] = [OH-]
so Kw = [H+]^2 which means [H+] is the square root of Kw
what happens to Kw as temp increases
increases
weak acid dissociation expression
Ka = [H+] [A-]
[HA]
what does Ka measure
Ka measures acid strength
pKa
pKa = -log Ka
so
Ka = 10 ^-pKa
large Ka strong or weak acid?
strong
small Ka strong or weak acid?
weak
working out pH of weak acid at half equivalence
we can assume that [HA] = [A-]
so [H+] = Ka and pH = pKa
pH of diluted strong acid
[H+] = [H+]old x old volume / new volume
pH = -log [H+]
pH of diluted base
[OH-] = [OH-]old x old volume
new volume
[H+] = Kw
[OH-]
pH = -log [H+]
comparing the pH of a strong acid and a weak acid after dilution 10 times
diluting a strong acid 10 times will increase its pH by one unit
comparing the pH of a strong acid and a weak acid after dilution 100 times
diluting a strong acid 100 times will increase its pH by two units
what would happen to weak acids when they are diluted
they would not change in the same way. they increase by less than 1 unit
what happens to equilibrium when diluting a weak acid
pushes equilibrium to the right so the degree of dissociation increases and more H+ ions are produced meaning pH increases less than expected
buffer solution definition
one where the pH does not change significantly if SMALL amounts of acid or alkali are added to it
what is an acidic buffer made from
a weak acid and a salt of that weak acid - weak acid reacting with a strong base
e.g. of an acidic buffer
ethanoic acid and sodium ethanoate
CH3CO2H and CH3CO2-Na+
what is a basic buffer made from
from a weak base and a salt of that weak base - weak base reacts with strong acid
e.g. of a basic buffer
ammonia and ammonium chloride
NH3 and NH4+Cl-
ethanoic acid buffer
CH3CO2H ⇌ CH3CO2- + H+
acid……………….conjugate
……………………… base
what happens to buffer solution if small amounts of acid are added
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
what happens to buffer solution if small amounts of alkali are added
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
calculating the pH of buffer solutions
Ka = [H+] [A-]
[HA]
[H+] = Ka x [HA]
[A-]
controlling pH in blood
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
constructing a titration/pH curve
- transfer 25 cm³ of acid to a conical flask with a volumetric pipette
- measure initial pH of the acid with a pH meter
- add alkali in small amounts (2 cm³) noting the volume added
- stir mixture to equalise it
- measure and record the pH to 1 d.p.
- repeat steps 3-5 but when approaching end point add in smaller volumes of alkali
- add until alkali is in excess
calibrating pH meter
measure known pH of a buffer solution
why is calibrating pH meter important
pH meters can lose accuracy on storage
calibrating a pH probe
putting a probe in a set buffer (often at pH 4) and pressuring a calibration button/setting that for that pH
strong acid - strong acid curve
e.g.
long vertical line is between pH 3 to 9
pH at equivalence point = 7
initial pH ~1
HCl & NaOH
weak acid - strong base curve
e.g.
the equivalence point is more than 7
initial pH ~ 3
final pH ~ 13
CH3CO2H & NaOH
strong acid - weak base curve
e.g.
vertical line less than 7 (4 to 7)
the equivalence point is less than 7
initial pH ~ 1
final pH ~ 9
HCl & NH3
weak acid - weak base
e.g.
no vertical part
CH3CO2H & NH3
pH range of phenolphthalein &
colour change
~ 7.5 to 9
use for strong bases
colourless to pink
pH range for methyl orange & colour change
~ 2 to 4
use with strong acids
red acid to yellow alkali
half neutralisation volume - for weak acids - equation
spec; 22.ii
Ka = [H+][A-]
[HA]
at 1/2 neutralisation volume the [HA] = [A-]
so, Ka = [H+] and pKa = pH
enthalpy change of neutralisation
the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce one mole of water
are enthalpy changes of neutralisation endothermic and exothermic
always exothermic
enthalpy changes of neutralisation for reactions with strong acids and bases
values are similar (about -56 to -58 kJ mole^-1) because the same reaction is occurring
enthalpy changes of neutralisation for reactions with weak acids and bases
less exothermic enthalpy change because energy is absorbed to ionise the acid and break the bond to the hydrogen in the un-dissociated acid
