pH + buffers Flashcards
bronsted-lowry acid definition
proton donor
bronsted-lowry base definition
proton acceptor
conjugate acid-base pair definition + 3 examples
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 +
monobasic definition + 2 examples
acids which can only donate 1 proton/ H+
e.g. HCl and CH3COOH
dibasic definition + 2 examples
acids which can donate 2 protons
e.g. H2SO4 and H2CO3
tribasic definition + 2 examples
acids which can donate 3 protons
e.g. H3PO4 and H3BO4
how can monobasic/dibasic acids change features of a practical e.g. a titration
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-
give the 2 pH equations
pH = -log[H+(aq)]
[H+] = 10^-pH
what is the effects of logs on the pH scale
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
strong acid definition
an acid which fully dissociates into ions in aqueous solution
weak acid definition
an acid which only partially dissociates into ions in aqueous solution
give the equation for the dissociation of a weak acid
HA(aq) <> H+(aq) + A-(aq)
ka definition + equation
ka = acid dissociation constant
ka = [H+][A-]/[HA]
what does a high ka value mean
high ka means equilibrium position is further to the right so greater acid dissociation, meaning greater acid strength
give the 2 pka equations
pka = -log(ka)
ka = 10^-pka
what are 3 assumptions made when calculating ka or using the acid dissociation equation
- [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
give equation and value for the ionic product of water kw
kw = [H+][OH-]
kw = 1 x 10^-14 mol2dm-6 at 25C
what is the pH of pure water at 25C
7
how does temperature affect kw
ionisation of water is endothermic so increased temp = increased kw
give a the reaction equation for the ionisation of water and the ka equations relative to kw
H2O(l) <> H+(aq) + OH-(aq)
ka = [H+][OH-]/[H2O]
ka x [H2O] = [H+][OH-] = kw
what is the assumption made for strong acids
as strong acids ionise fully, [HA] = [H+]
why can the pH of 1x10^-8 moldm-3 HCl(aq) not be 8
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
buffer definition
a solution that resists/minimises pH changes when small amounts of acids/bases are added
what are buffers made of + 1 example
weak acids + conjugate base (salt)
e.g. ethanoic acid + -ethanoate
give an equation for the action of a buffer and describe how it minimises pH changes
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
explain why the pH of a weak acid rises when its salt is added to it
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
what are 2 ways a buffer can be made
- weak acid + base
- weak acid + its salt
describe a method to make a buffer from a weak acid + a base
- 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
describe a method to make a buffer from a weak acid + its salt
- 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
outline how to calculate the pH of buffer solutions
- 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
give an example of a buffer in the body
H2CO3 <> HCO3- + H+
found in the blood
explain how the carbonic acid-hydrogencarbonate buffer system controls blood pH
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
at around what pH would the end point be reached in a strong acid-strong base titration
around pH 3-11
at around what pH would the end point be reached in a strong acid-weak base titration
around pH 3-7
at around what pH would the end point be reached in a weak acid-strong base titration
around pH 7-11
at around what pH would the end point be reached in a weak acid-weak base titration
around pH 7
describe the process of using a pH meter
- 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
on a titration curve, what indicates the end point
the vertical section of the graph
what is the point in the centre of the vertical section called
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
how is the curve different if it shows a titration where acids are added to bases
the curve will go down from left to right, still the same shape
how can you use a titration curve to tell if in indicator is suitable
if the pH range of the indicator lies within the vertical region of the curve
how do indicators work
indicators are basically coloured buffers, and work in the same way
can be HA <> A- + H+
or AOH <> A+ + OH-
