# Acid and Base equillibrium Flashcards

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1
Q

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

A
• a substance that donates H+ ions
2
Q

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

A
• a substance that accepts H+ ions
3
Q

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

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

give the formula for pH

A

pH= -log[H+]

5
Q

give the formula to find the concentration of H+ ions

A
• 1x10^pH
6
Q

how to calculate Kw

A

= [H+] [OH-]

7
Q

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

A

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

8
Q

what does pKw mean?

A

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

= 14.

9
Q

How to find pH of pure water?

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

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
12
Q

what is the expression for the weak acid dissociation expression?

A

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

13
Q

what does a larger Ka value signify?

A
• that the acid/base is stronger
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.
15
Q

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

A

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

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.
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

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
Q

By what factor will the pH increase by if we dilute a strong acid by a factor of 10?

A

by a factor of 1

26
Q

How do we calculate the pH of buffer solutions?

A
• [H+] = Ka x [HA]/[A-]
27
Q

What type of solution would a weak acid and strong base make?

A

alkaline

28
Q

What type of solution would a strong acid and a weak base make?

A
• acidic
29
Q

What type of solution would a strong acid and a strong base make?

A

neutral

30
Q

why doesn’t the pH of a an acidic buffer change when a few drops of NaOH are added to it?

A
• Because the acid dissociates to replace used up H+ ions
• and the NaOH reacts to form water,
31
Q

What does the enthalpy change of neutralisation depend on?

A
• the acid or base being used - varies from acid to acid
32
Q

Define what is meant by a monoprotic acid

A

means that 1 mole of the acid will produce 1 mole of H+ ions

33
Q

What do we use the Kw expressions for, and How?

A

to calculate pH of a strong base - rearrange to find H+ ion conc, then use log expression to find pH

34
Q

What are two main assumptions that are made when calculating pH of a weak acid?

A
• [HA] initial = [HA] eq
• [H+]=[A-]
35
Q

Explain how pH is maintained in the blood. Include any relavant equations

A

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
Q

what two pieces of info do we need to calculate pH of a buffer

A
• Ka value, conc of weak acid and salt
37
Q

what assumption is false in a buffer solution

A

that H+ = A-

38
Q

what 2 assumptions are made when calculating pH of a buffer

A

[HA] eq = [HA] initial
[A-] = [salt]

39
Q

give the equation to calculate concs in a buffer

A

pH = pKa + log([A-]/[HA])

Henderson-Hasselbach equation

40
Q

Give the following features of a titration curve:
- axes labels
- interpretation of curve at each point in time

A
• 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
Q

A mixture of Ammonia and ammonium ions acts as a buffer.
using ionic equations, explain how.

A
• 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
Q

describe how buffer solutions maintain the pH relatively constant when an acid is added.

A
• 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
Q

describe how buffer solutions maintain the pH relatively constant when an alkali is added.

A
• 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
Q
A