1
Q
What does the Bronsted-Lowry theory define acids as
A
- proton donors
2
Q
What does the Bronsted-Lowry theory define bases as
A
- proton acceptors
3
Q
What are polyprotic acids
A
- acids that are able to donate more that one H+ ion in reaction
4
Q
What are polybasic bases
A
- bases that accept more that one proton
5
Q
What are monoprotic acids and what is an example
A
- when 1 mole of an acid produces 1 mole of H+ ions
E.g: HNO3
6
Q
What are diportic acids and what is an example
A
- when 2 mole of an acid produces 2 mole of H+ ions
E.g: H2SO4
7
Q
What are triportic acids and what is an example
A
- when 3 mole of an acid produces 3 mole of H+ ions
E.g: H3PO4
8
Q
What are the products from a acid base reaction
A
- salt and water
9
Q
How does ammonia produce OH- ions
A
- by reacting with water first accepting a proton and forming ammonium ions and OH- ions
10
Q
Metals + acids=
A
Salts + hydrogen
11
Q
Metal oxides + acids=
A
- salt + water
12
Q
Metal hydroxides + acids=
A
- salt + water
13
Q
Metal carbonates + acids=
A
- salt + water + carbon dioxide
14
Q
Define the term conjugate pairs
A
- a pair of species that are related to each other by the difference of a proton
15
Q
Complete the sentence:
Any species that has gained a proton is the ….
A
- conjugate acid
16
Q
Complete the sentence:
A species that has lost a proton is the ….
A
- conjugate base
17
Q
Identify the conjugate pairs in this reaction:
HA + B <=> BH+ + A-
A
- ACID 1: HA
- BASE 1: A-
- ACID 2: BH+
- BASE 2: B
18
Q
What are strong bases
A
- bases that dissociate almost completely
19
Q
What are weak bases
A
- bases that partially dissociate
20
Q
What are strong acids
A
- acids that dissociate completely
21
Q
What are weak acids
A
- acids that partially dissociate forming equilibrium reactions
22
Q
What are examples of weak acids
A
- ethanoic acid & other carboxylic acids
23
Q
What are examples of strong acids
A
- HCl
-H2SO4 - HNO3
24
Q
What are examples of strong bases
A
- KOH
- NaOH
25
What are examples of a weak base
- NH3
26
What reaction is favoured in weak acid/base reaction
- the backwards reaction
27
What reaction is favoured in strong acid/base reaction
- the forward reaction
28
When acids and bases react what do they exchange
- protons
29
Does water behave as an acid or a base
- base
30
What does water dissociate into
2H2O <=> H3O+ + OH-
31
How does water dissociate into its ions
- very weakly
- very little OH- and H+ ions compared to water molecules
32
What is Kw
- the ionic product of water
33
What are the units for Kw
- mol^2 dm^-6
34
What is the Kw expression
Kw = [H+] [OH-]
35
What is the assumption we make with water dissociation, that we use in Kw
- that water has a constant value
36
Complete the sentence:
Kw is the same in a solution at a given….
- temperature
37
What is the value of Kw
- 1.00 x 10^-14 mol^2 dm^-6
38
When does the value of Kw change
- when here is a change in temperature
39
What is the concentration of H+ and OH- ions in pure water
- [H+] = [OH-]
- equal amount
40
What Kw expression can we use for pure water and why
Kw = [H+]^2
- because we can assume that there is the same amount of OH- ions and H+ ions
41
What does the pH scale measures
- the concentration of H+ ions in solution
42
What equation does we use to find the pH
= -log10[H+]
43
What is the equation to work out the [H+] when give a pH
[H+]= 10^-pH
44
When calculating the pH of strong acids what do we assume
- that they fully dissociate
45
What equation do you sue to calculate the pH of a strong base
- ionic water equation
- pH equation
46
What do we assume when working out the pH of strong bases
- that bases dissociate fully
47
What 2 assumptions do we make when dealing with weka acid
1) only a small amount of the acid dissociates so we assume that [HA] = [HA]
2) all the H+ ions come from the acid-> [H+]=[A-]
48
What equation do we use when working out the concentration of weak acids
Ka = [H+]^2
—————
[HA]
49
What are the units for Ka
- moldm^-3
50
What is pKa
- another way of measuring the strength of an acid (similar to pH)
51
What is the equation for pKa
-log10Ka
52
What is the rearranged formula of pKa for Ka
Ka= 10^-pka
53
How do we measure pH experimentally
- use pH meters that are calibrated correctly
54
What do titrations allow us to do
- to work out concentration of an acid or base
55
Briefly describe how you would do a titration reaction
1) you have an acid or alkali in a burette with a known concetration
2) you also have an acid or alkali with an unknown concentration but known volume in the conical flask
3) you add a few drops of indicator (usually phenolphthalein)
4) you turn the knob of the burette horizontally to open it and the chemical will drop into the conical flask until the indicator changes colour (end point)
56
Hoe do you read how much chemical was added from the burette
- by reading form the bottom of the meniscus
57
How many decimal places do you always record everything in when doing a titration
- 2 decimal places
58
Which results do you look at when working out your average titration
- the two concordant results (results within 0.10cm^3 of each other)
59
When using phenolphtalein when the acid is in the conical flask what colour does it change to
- colourless
60
When using phenolphtalein when the alkali is in the conical flask what colour does it change to
- pink
61
When using methyl orange when the acid is in the conical flask what colour does it change to
- yellow
62
When using methyl orange when the alkali is in the conical flask what colour does it change to
- Red
63
What are titration curves and what do they show
- the pH against the volume of bases added from titration
- different combinations of a weak & strong acids and bases
64
What do these titration graphs show
1) strong acid/ strong base ( starts at pH 1 ends at pH 13)
2) strong acid/ weak base (starts at pH 1 ends at pH 9)
3) weak acid/ strong base (starts at pH 5 ends at pH 13)
4) weak acid/ weak base (starts at pH 5 ends at pH 9)
65
What is an equivalent point on a titration curve
- the point where the acid has been fully neutralised
- at this point we assume that the [H+] = [OH-]
66
What does the sharp vertical rise show in a titration curve
- a rapid change in pH
67
What is the colour change for methyl orange
- Colour change at pH3- 4.5
68
What is the colour change for phenolpthalein
- colour chnage at pH 8.2- 10
69
What reactions do we use methyl orange for
1) strong acids/strong base
2) strong acid/weak base
70
What reactions do we use phenolpthalein for
- weak acid/strong base
71
What do we use when trying to detect a colour change with weak acids/weak bases
- pH meter
72
Define the term buffers
- A system that MINIMISES pH changes when small amounts of acid or a base are added
73
What are two types of buffers
1) Acidic buffers
2) Basic buffers
74
What do acid idc buffers do
- they RESIST the change in pH in order to keep the solution BELOW pH7
75
How do you make acidic buffers
- from weak acids and its salt
76
What is an example of an acidic buffer
- ethanoic acid (weak acid)
- and its salt (sodium ethanoate)
77
In any buffer solution how many equilibrium equations are there
- 2
78
What do the 2 equilibrium equations exist in buffer solutions
- co-exist in the same beaker
79
What are the 2 equilibrium equations for ethanoic acid and its salt sodium ethanoate
1) CH3COOH <=> CH3COO- + H+
(HIGH) (LOW) (LOW)
*because weak acids partially dissociate so the equilibrium favours the backwards reaction and lies to the left
2) CH3COO- Na+ <=> CH3COO- + Na
(LOW) (HIGH) (HIGH)
*salts dissociate fully so equilibrium favours the forwards reaction and lies well over to the right
80
What happens when we add an acid(H+) to an ethanoic buffer
- The 2 buffer equilibrium equations link
- H+ ions react with the CH3COO- ions in solution (opposite ions attract)
- high concentration of these from the salt (as the salt dissociates)
- more CH3COOH is produced which means equilibrium shifts to the LEFT
81
What happens when we add a base (oH-) to an ethanoic buffer
- OH- ions react with the H+ ions in solution
- there is a low concentration of H+ ions however they can be REPRODUCED from a high concentration of CH3COOH to counteract the change (Le Chatelier’s principle)
- equilibrium shifts to the RIGHT to replace the reacted H+ ions
82
Why can’t we use Ka= [H+] in buffer calculations
———-
[HA]
- because [H+] doesn’t equal [A-] because there is a salt and an acid from two different sources
83
What is an assumption we make when calculating buffers
1) that salts dissociate fully and weak acids dissociate poorly
[salt]= [A-]
[HA] at the start= [HA] at equilibrium
84
Answer this question:
pH= 3.64
85
Answer this question:
pH change: 4.73 to 4.68
86
Where do we use buffers
-in household products
- naturally occurring in living things (e.g blood)
87
How and Why does ours blood act as a buffer
- it is vital to make sure blood pH is maintained as close to pH 7.4 as possible
- our body systems rely on this so a buffer is present in our body to help. Carbon dioxide plays a big role
88
What are the 2 equilibrium equations for blood
- carbonic acid (H2CO3)
- hydrogen carbonate (HCO3 -)
H2CO3 (aq) <=> H+(aq) + HCO3 -(aq)
H2CO3(aq) <=> H2O(l) + CO2 (aq)
89
How id carbonic acid controlled
- via respiration in our cells
- when we breathe out C02 the level of carbonic acid reduces as equilibrium shifts RIGHT to attempt to replace them
90
How is hydrogen carbonate controlled
- via the kidneys
- excess is removed by them
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