Acid and Base equillibrium Flashcards
Define, by the bronsted and lowry definition, what is meant by an acid.
- a substance that donates H+ ions
Define, by the bronsted and lowry definition, what is meant by a base
- a substance that accepts H+ ions
in which case will you decide whether it is an acid or alkali?
- the one that has a higher Ka value
give the formula for pH
pH= -log[H+]
give the formula to find the concentration of H+ ions
- 1x10^pH
how to calculate Kw
= [H+] [OH-]
what is the value of Kw at 25 degrees c for most solutions? give units
1x10^-14 mol^2dm^-6
what does pKw mean?
-log[Kw], so the value would be -log1x10^-14, which would
= 14.
How to find pH of pure water?
- Kw^2 = H+ concentration
- so use this to find the pH
why does water’s pH change at temperatures different to 25 degrees?
- Le Chatelier’s
principle can predict the change. - dissociation of water is endothermic so
increasing the temperature would push the
equilibrium to the right
giving a bigger
concentration of H+ ions
resulting in a lower pH.
How do we work out the pH of a strong base?
- rearrange Kw expression to find H+ ions (so Kw/OH-)
- find pH using log expression
what is the expression for the weak acid dissociation expression?
Ka= [H+][A-]/ [HA]
what does a larger Ka value signify?
- that the acid/base is stronger
what 2 assumptions are made when calculating pH of a weak acid?
- initial conc of the undissociated acid stays constant, as it’s very small, so eq HA = Initial HA
- [H+] = [A-] , as the dissociation has taken place in a 1:1 ratio.
taking into account these assumptions, what does the Ka expression simplify to?
Ka= [H+]^2/[HA] initial
What needs to happen in order for us to assume that pKa=pH?
- when a weak acid has reacted with exactly half the neutralisation volume of alkali, so half the point of neutralisation.
how to work out the pH of a diluted acid?
H+ (diluted) = old vol/new vol x H+ (old)
- then use the log expression to work out pH
how to work out the pH of a diluted alkali?
-OH- (diluted) = old vol/new vol x OH- (old)
- [H+] = Kw/OH-
- work out pH using log expression.
How do you calculate the H+ ion conc in a buffer solution?
[H+] = Ka x [HA]/[A-]
How do you produce a buffer solution?
- made form combining a weak acid and a salt of that weak acid (a conjugate base)
Define a buffer solution
- a solution in which pH does not change significantly if small amounts of acid or alkali are added to it.
for a buffer where H+ ions are on the right, where will the equilibrium shift if small amount of acid is added?
- to the left
- as conc of H+ ions increases, the equillibrium will shift to the side with a higher number of OH- ions (or any other basic ions)
- so since the concentration of the salt ion in the buffer solution is large, and stays relatively constant, the pH does as well.
In a buffer solution where H+ ions are on the right side, if small amounts of alkali are added, where would the equilibrium shift?
- will shift to the right
- as equillibrium will shift to side with more H+ ions.
- large conc of salt in buffer, so ratio of the reactant stays constant, and so does the pH.
what do you need to create a buffer solution with a pH BELOW 7?
- mix together a weak acid in excess and a strong base.
By what factor will the pH increase by if we dilute a strong acid by a factor of 10?
by a factor of 1
How do we calculate the pH of buffer solutions?
- [H+] = Ka x [HA]/[A-]
What type of solution would a weak acid and strong base make?
alkaline
What type of solution would a strong acid and a weak base make?
- acidic
What type of solution would a strong acid and a strong base make?
neutral
why doesn’t the pH of a an acidic buffer change when a few drops of NaOH are added to it?
- Because the acid dissociates to replace used up H+ ions
- and the NaOH reacts to form water,
What does the enthalpy change of neutralisation depend on?
- the acid or base being used - varies from acid to acid
Define what is meant by a monoprotic acid
means that 1 mole of the acid will produce 1 mole of H+ ions
What do we use the Kw expressions for, and How?
to calculate pH of a strong base - rearrange to find H+ ion conc, then use log expression to find pH
What are two main assumptions that are made when calculating pH of a weak acid?
- [HA] initial = [HA] eq
- [H+]=[A-]
Explain how pH is maintained in the blood. Include any relavant equations
H2CO3–> H+ + HCO3-
H2CO3 —-> H20 + CO2
- CO2 in blood forms carbonic acid, increasing its conc on the left.
- so eq will shift to the right, forming more H+ ions.
- excess of hydrogen carbonate ions combine with H+ ions to control blood pH
what two pieces of info do we need to calculate pH of a buffer
- Ka value, conc of weak acid and salt
what assumption is false in a buffer solution
that H+ = A-
what 2 assumptions are made when calculating pH of a buffer
[HA] eq = [HA] initial
[A-] = [salt]
give the equation to calculate concs in a buffer
pH = pKa + log([A-]/[HA])
Henderson-Hasselbach equation
Give the following features of a titration curve:
- axes labels
- interpretation of curve at each point in time
- x axis: volume of base/acid added
- y axis: pH
- start: pH is relatively stable, which shows buffering capacity, resists changes in pH with small volume of acid/alkali added
- then dramatic increase/decrease in pH as the buffering capacity is diminished, buffer solution no longer exists.
- end - levels out again as full point of neutralisation is reached.
- middle of starting line - half point of neutralisation where pKa = pH
A mixture of Ammonia and ammonium ions acts as a buffer.
using ionic equations, explain how.
- large reservoir of ammonia and ammonium ions
- H+ reacts with ammonia to form ammonium ions (H+ + NH3 —> NH4)
- OH- from water reacts with H+ to form water (OH- + H+ —-> H2O)
- so pH stays relatively constant since ratio of ammonia and ammonium ions also stays relatively constant.
describe how buffer solutions maintain the pH relatively constant when an acid is added.
- Large reservoir of HA and A-
- so H+ added shifts eq to the right
- pH stays constant because ratio of HA to A- also remains constant.
describe how buffer solutions maintain the pH relatively constant when an alkali is added.
- Large resevoir of HA and A-
- OH- added will react with H+ to form water
- pH stays constant cus ratio of HA:A- stays constant as well.
