Acid and Base equillibrium Flashcards

1
Q

Define, by the bronsted and lowry definition, what is meant by an acid.

A
  • a substance that donates H+ ions
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2
Q

Define, by the bronsted and lowry definition, what is meant by a base

A
  • a substance that accepts H+ ions
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3
Q

in which case will you decide whether it is an acid or alkali?

A
  • the one that has a higher Ka value
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4
Q

give the formula for pH

A

pH= -log[H+]

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5
Q

give the formula to find the concentration of H+ ions

A
  • 1x10^pH
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6
Q

how to calculate Kw

A

= [H+] [OH-]

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7
Q

what is the value of Kw at 25 degrees c for most solutions? give units

A

1x10^-14 mol^2dm^-6

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8
Q

what does pKw mean?

A

-log[Kw], so the value would be -log1x10^-14, which would

= 14.

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9
Q

How to find pH of pure water?

A
  • Kw^2 = H+ concentration
  • so use this to find the pH
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10
Q

why does water’s pH change at temperatures different to 25 degrees?

A
  • 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.

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11
Q

How do we work out the pH of a strong base?

A
  • rearrange Kw expression to find H+ ions (so Kw/OH-)
  • find pH using log expression
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12
Q

what is the expression for the weak acid dissociation expression?

A

Ka= [H+][A-]/ [HA]

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13
Q

what does a larger Ka value signify?

A
  • that the acid/base is stronger
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14
Q

what 2 assumptions are made when calculating pH of a weak acid?

A
  • 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.
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15
Q

taking into account these assumptions, what does the Ka expression simplify to?

A

Ka= [H+]^2/[HA] initial

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16
Q

What needs to happen in order for us to assume that pKa=pH?

A
  • when a weak acid has reacted with exactly half the neutralisation volume of alkali, so half the point of neutralisation.
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17
Q

how to work out the pH of a diluted acid?

A

H+ (diluted) = old vol/new vol x H+ (old)
- then use the log expression to work out pH

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18
Q

how to work out the pH of a diluted alkali?

A

-OH- (diluted) = old vol/new vol x OH- (old)
- [H+] = Kw/OH-
- work out pH using log expression.

19
Q

How do you calculate the H+ ion conc in a buffer solution?

A

[H+] = Ka x [HA]/[A-]

20
Q

How do you produce a buffer solution?

A
  • made form combining a weak acid and a salt of that weak acid (a conjugate base)
21
Q

Define a buffer solution

A
  • a solution in which pH does not change significantly if small amounts of acid or alkali are added to it.
22
Q

for a buffer where H+ ions are on the right, where will the equilibrium shift if small amount of acid is added?

A
  • 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.
23
Q

In a buffer solution where H+ ions are on the right side, if small amounts of alkali are added, where would the equilibrium shift?

A
  • 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.
24
Q

what do you need to create a buffer solution with a pH BELOW 7?

A
  • mix together a weak acid in excess and a strong base.
25
By what factor will the pH increase by if we dilute a strong acid by a factor of 10?
by a factor of 1
26
How do we calculate the pH of buffer solutions?
- [H+] = Ka x [HA]/[A-]
27
What type of solution would a weak acid and strong base make?
alkaline
28
What type of solution would a strong acid and a weak base make?
- acidic
29
What type of solution would a strong acid and a strong base make?
neutral
30
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,
31
What does the enthalpy change of neutralisation depend on?
- the acid or base being used - varies from acid to acid
32
Define what is meant by a monoprotic acid
means that 1 mole of the acid will produce 1 mole of H+ ions
33
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
34
What are two main assumptions that are made when calculating pH of a weak acid?
- [HA] initial = [HA] eq - [H+]=[A-]
35
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
36
what two pieces of info do we need to calculate pH of a buffer
- Ka value, conc of weak acid and salt
37
what assumption is false in a buffer solution
that H+ = A-
38
what 2 assumptions are made when calculating pH of a buffer
[HA] eq = [HA] initial [A-] = [salt]
39
give the equation to calculate concs in a buffer
pH = pKa + log([A-]/[HA]) Henderson-Hasselbach equation
40
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
41
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.
42
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.
43
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.
44