topic 12: acid - base equilibrium Flashcards
what is a bronsted lowry acid
a substance that can donate a proton
what is a bronsted lowry base
a substance that can accept a proton
what is each acid linked to
linked to a conjugate base on the other side of the equation
what is the definition of pH
-log[H+]
what do strong acids do
completely dissociate
what is the equation for finding [H+] from pH
10^-pH
what is the Kw expression
= [H+(aq) ][OH-(aq) ]
=10^-pkw
at 25 degrees celsius , what is the value of Kw
1x10^-14 mol2dm-6
what are the expression for pKw
=-logKw
what are weak acids
only slightly dissociate when dissolved in water, giving an equilibrium mixture
what is the weak acids dissociation expression (Ka)
[H+(aq)][A-(aq)]/[HA (aq)]
the larger the Ka
the stronger the acid
what are the 2 assumption made to simplify the Ka expression
1) [H+(aq)]eqm = [A(aq)] eqm because 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.
so [HA (aq) ] eqm = [HA(aq) ]
initial
what is the Ka expression
[H+(aq)]^2 /[HA (aq)]initial
using one of the assumptions, what is the other expression of pH
pKa
what is a buffer solution
one where the pH does not change significantly if small amounts of acid or alkali are added to it
what is a acidic buffer solution made from
made from a weak acid and a salt of that weak acid ( made from reacting the weak acid with a strong base)
example : ethanoic acid and sodium ethanoate
what is a basic buffer solution
made from a weak
base and a salt of that weak base (made from reacting the weak base with a strong acid)
example :ammonia and ammonium chloride
what happens when small amounts of acid is added to the buffer
the above equilibrium
will shift to the left removing nearly all the H+ ions added, CH3CO2 -(aq) + H+ (aq) <—> CH3CO2H (aq)
as there is a large concentration of the salt ion in the buffer the ratio
[CH3CO2H]/ [CH3CO2
-] stays almost constant, so the pH stays fairly
constant
what happens when a small amount of alkali is added to the buffer
the OH- will react with H+ ions to form water.
H+ + OH - H2O
the equilibrium will then shift to the right to produce more H+
ions.
CH3CO2H(aq) <—> CH3CO2-(aq) + H+(aq)
some ethanoic acid molecules are changed to ethanoate ions but as there is a large
concentration of the salt ion in the buffer the ratio [CH3CO2H]/[CH3CO2-] stays almost constant, so the pH stays fairly constant
what would happen to pH when the acid is diluted by 10
pH increases by 1 unit
H2CO3 (aq) ⇌ H+(aq) + HCO3–(aq)
what is the buffer that controls the blood’s pH
H2CO3/HCO3– buffer
what is the value of pH in blood the H2CO3
/HCO3– buffer maintains
7.35 and 7.45
what does the H2CO3
/HCO3– buffer do when the alkali is added
adding alkali reacts with H+ with the equation
H+ + OH- —> H2O
so the above equilibrium would shift right forming new H+ and more HCO3–
what does the H2CO3
/HCO3– buffer do when the acid is added
adding acid shifts the above equilibrium left.
the reaction is
H+ + HCO3– —> H2CO3
what happens to the moles if a small amount of alkali is added to a buffer
- if a small amount of alkali is added to a buffer then the moles of the buffer acid would reduce by the number of moles of alkali added and the moles of salt would increase by the same amount so a new
calculation of pH can be done with the new values
CH3CO2H (aq) +OH- —> CH3CO2-(aq) + H2O(l)
what happens to the moles if a small amount of acid is added to a buffer
if a small amount of acid is added to a buffer then the moles of the buffer salt would reduce by the number of moles of acid added and the moles of buffer acid would increase by the same amount so a
new calculation of pH can be done with the new values
CH3CO2-(aq) + H+ —> CH3CO2H (aq)
PRACTICAL: how do you construct a pH curve
- transfer 25cm3 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 (2cm3) noting the volume added
- stir mixture to equalise the pH
- measure and record the pH to 1 d.p.
- repeat steps 3-5 but when approaching endpoint add in smaller volumes of alkali
- add until alkali in excess
describe the curve of strong acid - strong base
equivalence point is 7
long vertical part around 3 to 9
initial and final pH
volume at neutralisation
general shape (pH at neutralisation)
describe the weak acid - strong base graph
equivalence point >7
steep part of curve >7 (around 7 to 9)
pH starts near 3
WHEN A STRONG BASE IS ADDED
describe the strong acid - weak base graph
vertical part of curve<7 (around 4 to 7)
equivalence point<7
starts at 1
WHEN ALKALI IS ADDED
what is the end - point of a titration
the point when the colour of the indicator changes colour
what happens when the end - point is reached
when [HIn] = [In-]
when would you use phenolphthalein
with strong bases but not weak bases
colour change: colourless acid —> pink alkali
when would you use methyl orange
with strong acids but not weak acids
colour change: red acid —> yellow alkali (orange end point)
what is the standard enthalpy change of neutralisation
the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water
why do weak acids have a less exothermic enthalpy change of neutralisation
because energy is absorbed to ionise the
acid and break the bond to the hydrogen in the un-dissociated acid
when will an indicator work
if the pH range of the indicator lies on the vertical part of the titration curve
- in this case the indicator will change colour rapidly and the colour change will correspond to the neutralisation point
describe and explain the behaviour of the solution formed in the region
circled on the sketch graph
any mention of buffer / buffering
(both) propanoic acid and potassium
propanoate present
CH3CH2COOH + OH− → H2O + CH3CH2COO−
On addition of OH− (in small quantities)
H+ ions react with (the added) OH−and
CH3CH2COOH ⇌ CH3CH2COO− + H+
shifts to the right
explain why the pH at the equivalence point of this titration is greater than 7
Propanoate (ions) react with water
Forming hydroxide ions
[OH−] > [H+]
CH3CH2COO− + H2O → OH− + CH3CH2COOH
the student made the following statement:
‘The pH of pure water is always 7.0’
is the student correct? use the following information to justify your answer.
- H2O(l) U H+(aq) + OH–(aq)
- Kw = 1.0 x 10–14 mol2dm–6 at 298 K
- ǻH is positive for the forward reaction in the equilibrium
No,
as T increases eqm moves to RHS / Kw
increases / ‘favours RHS’ / ∆Stotal increases
So [H+] ions increases / more H+ ions
[H+] > 1 x 10−7
Hence pH < 7 / pH decreases
state the three assumptions you have made in your calculations for weak acids
[H2S]equilibrium = [H2S]initial
[H3O+] = [HS−]
Ka2 very much smaller than Ka1
Measurements at 298 K / standard temperature
