3.1.10 Acids and Bases Flashcards
1
Q
Define Brønsted–Lowry Acids
A
Proton donors
2
Q
When Brønsted–Lowry acids are mixed with water, they ______ ____
A
Release H+
3
Q
State the equation for when Brønsted–Lowry acids are mixed with water
A
4
Q
Define Brønsted–Lowry Bases
A
Proton acceptors
5
Q
When Brønsted–Lowry bases are mixed with water, they grab ___ from ____
A
grab H+ from H2O
6
Q
State the equation for when Brønsted–Lowry bases are mixed with water
A
7
Q
Define strong acids
A
- Dissociate (or ionise) completely in water
- Nearly all H+ ions released
8
Q
Where does the equilibrium of strong acids and bases reacting with water lie?
A
To the right
9
Q
Define Weak Acids
A
- Dissociate only very slightly in water
- Small no. of H+ ions formed
10
Q
Where does the equilibrium of weak acids and bases reacting with water lie?
A
To the left
11
Q
Acids only get rid of their protons if there’s a ____ to ____ them
A
Acids only get rid of their protons if there’s a base to accept them
12
Q
Write an equation for when HA (acid) reacts with B (base)
A
13
Q
What happens to the position of equilibrium if you add more HA or B?
A
Shifts to right
14
Q
What happens to the position of equilibrium if you add more BH+ or A-?
A
Shifts to left
15
Q
When acid is added to water, water acts as ____ and _____ the ____
A
When acid is added to water, water acts as base and accepts the proton
16
Q
Equilibrium’s far to the ___ for weak acids
A
left
17
Q
Equilibrium’s far to the ___ for strong acids
A
right
18
Q
What does water dissociate into?
A
Dissociates into hydroxonium ions and hydroxide ions
19
Q
Which side does the equilibrium lie on?
A
To the left
20
Q
State the ionic product of water (Kw) & the units
A
21
Q
Describe how Kw is derived from the equilibrium constant
A
- Can work out normal equilibrium constant from equation: H2O ⇌ H+ + OH-
- Kc = [H+] [OH-] / [H2O]
- So much more water compared to H+ & OH- that water is considered to have constant value
- ∴ if you multiply expression fo Kc (which is a constant) by [H₂O] (another constant), you get a constant
- New constant = Kw
22
Q
State what Kw is at 298K
A
1.00 x 10-14 mol² dm-6
23
Q
State what the Kw expression is at pure water
A
Kw = [H+]2
24
Q
Explain why Kw = [H+]2 in pure water
A
In pure water, there’s always one H+ ion for each OH- ion
25
What is the pH scale?
Measure of hydrogen ion concentration
26
The smaller the pH, the greater...
conc. of H+ ions
27
State the equation you can use to work out pH
(used for strong acids directly)
28
State the equation for calculating hydrogen ion concentration
29
What is meant by strong monoprotic acids?
e.g. HCL and HNO3 (nitric acid)
* 1 mole of acid produces 1 mole of hydrogen ions when it dissociates
* [H+] = [Acid]
30
What is meant by strong diprotic acids?
e.g. sulfuric acid
* Produces 2 mole of H+ for 1 mole of acid when it dissociates
* [H+] = 2[Acid]
31
What should you use to calculate pH of a strong base?
32
33
Why do you have to use Ka (acid dissociation constant) to work out the [H+] for weak acids?
* Weak acids (e.g. any that end in 'oic') dissociate only slightly in aq solution
* ∴ [H+] isn't equal to acid concentration
* ∴ have to use equilibrium constant Ka
34
Why can you assume?
As only a tiny amount of HA dissociates
35
State the expression for Ka
(used when calculating pH for weak acids)
36
State Ka expression you use for weak acids in aq solution with nothing else added. Include the units.
37
38
State the equation for calculating pKa
pKa = -logKa
39
State the equation for calculating Ka
Ka = 10-pKa
40
41
When a ___ acids reacts with a ___ base, for every mole of OH- added, ___ mole of HA is used and ___ mole of A- is formed
When a _weak_ acids reacts with a _strong_ base, for every mole of OH- added, _1_ mole of HA is used and _1_ mole of A- is formed
42
6 moles of HA with 1.3 moles of Ba(OH)2. State the moles before and after the reaction.
43
**Weak Acid + Strong Base**
Calculate the pH of the solution formed when 30 cm3 of 0.200 mol dm-3 ethanoic acid (pKa = 4.76) is added to 100 cm3 of 0.100 mol dm-3 NaOH.
44
**Weak Acid + Strong Base**
Calculate the pH of the solution formed when 50 cm3 of 0.500 mol dm-3 ethanoic acid (pKa = 4.76) is added to 75 cm3 of 0.200 mol dm-3 NaOH
45
When ___ of the HA molecules have reacted with OH-, ____ = _____ ∴ ___ = _____ or ____ = \_\_\_\_
When _half_ of the HA molecules have reacted with OH-, [HA] = [A-] ∴ Ka = [H+] or pKa = pH
46
Calculate the pH of the solution formed when 100 cm3 of 0.2 mol dm-3 ethanoic acid (pKa = 4.76) is added to 40 cm3 of 0.250 mol dm-3 KOH.
47
Draw the pH curve for a strong acid/strong base. Include where the graph starts and where it levels off.
48
Draw the pH curve for a strong acid/weak base. Include where the graph starts and where it levels off.
49
Draw the pH curve for a weak acid/strong base. Include where the graph starts and where it levels off.
50
Draw the pH curve for a weak acid/weak base. Include where the graph starts and where it levels off.
51
What is the name given to the section that is vertical on the pH curve?
Equivalence point or end point
52
At equivalence point or end point, a tiny amount of base/acid causes a ...
sudden, big change in pH
53
What happens to the pH curves when you titrate a base with an acid instead?
Shapes of the curves stay the same but they flip over
54
How do you deicide which indicator to use?
Need to pick one that changes colour over a narrow pH range that lies entirely on the vertical part of the pH curve
55
Name 2 indicators
* Methyl orange
* Phenolphthalein
56
State the colour of methyl orange at low pH (in acid)
Red
57
State the colour of methyl orange at high pH (in base)
Yellow
58
State the approx. pH of colour change for methyl orange
3.1-4.4
59
State the colour of phenolphthalein at low pH (in acid)
Colourless
60
State the colour of phenolphthalein at high pH (in base)
Pink
61
State the approx. pH of colour change for phenolphthalein
8.3-10
62
State an indicator you can use for a strong acid/strong base titration
Methyl orange or phenolphthalein
(Rapid pH change over range for both indicators)
63
State an indicator you can use for a strong acid/weak base titration
Methyl orange
64
State an indicator you can use for a weak acid/strong base titration
Phenolphthalein
65
State an indicator you can use for a weak acid/weak base titration
Can't use indicator
66
Why can't you use a indicator for weak acid/weak base titrations?
Don't get sharp change in weak acid/weak base titration
67
What should you use instead of an indicator for weak acid/weak base titrations?
pH meter
68
Name 3 things you can do to make your titration results as accurate as possible
1. Measure neutralisation volume as precisely as possible
* (i.e. to nearest 0.05 cm3)
2. Repeat titration at least 3 times and take mean titre value
* Makes result reliable
3. Don't use anomalous results
* Results should be within 0.1 cm3 of each other
69
If you use a pH meter, describe how you can work out how much acid or base is needed for neutralisation.
1. Draw pH curve
2. Find equivalence point (mid-point of the line of rapid pH change) and draw a vertical line downwards until it meets the x-axis
3. Value at x-axis = volume of acid or base needed
70
71
When a diprotic acid reacts with a base, the reaction occurs in __ \_\_\_\_
2 stages
72
Why is it that when you react a diprotic acid with a base, the reaction occurs in 2 stages?
∵ 2 protons are removed from acid separately
73
Diprotic acid (e.g. ethanedioic acid) + strong base results in a pH curve with __ equivalence points
2
74
75
Write the equation for when hydrogen ions react with hydroxide ions
H+ + OH- → H2O
76
Write the equation for when hydrogen ions react with carbonate ions
2 H+ + CO32- → H2O + CO2
77
Write the equation for when hydrogen ions react with bicarbonate ion
H+ + HCO3- → H2O + CO2
78
Write the equation for when hydrogen ions react with ammonia
H+ + NH3 → NH4+
79
What is a buffer?
Solution that resists changes in pH when small amounts of acid or base are added, or when it's diluted
80
How are acidic buffers made?
Made by mixing weak acid with one of its salts
e.g. ethanoic acid and sodium ethanoate
81
Explain what happens when add a small amount of acid to this acidic buffer solution
* H+ conc. ↑
* Most of extra H+ ions combine with CH3COO- ions to form CH3COOH
* Shifts equilibrium to left = reducing H+ conc. close to its original value
* ∴ pH doesn't change
82
Explain what happens when add a small amount of base to this acidic buffer solution
* OH- conc. ↑
* Most of extra OH- ions react with H+ ions to form water = removing H+ ions from solution
* Causes more CH3COOH to dissociate to form H+ ions = shifting equilibrium to the right
* H+ conc. increases until it's close to original value ∴ pH doesn't change
83
How are basic buffers made?
Made by mixing weak base with one of its salt
e.g. solution of ammonia and ammonium chloride acts as a basic buffer
84
In a solution of ammonia and ammonium chloride acts a ____ \_\_\_, the salt ...
_basic_ _buffer_, the salt _fully dissociates in solution_
85
In a solution of ammonia and ammonium chloride, some of NH3 will...
react with water molecules
86
Explain what happens when you add a small amount of acid to this basic buffer solution
* H+ conc. ↑
* Some of H+ ions react with OH- ions to make H2O
* ∴ equilibrium position moves to right to replace OH- ions that have been used up
* Some of H+ ions react with NH3 = NH4+
* These reactions remove most of extra H+ ions ∴ pH won't change much
87
Explain what happens when you add a small amount of base to this basic buffer solution
* OH- conc. ↑
* Most of extra OH- ions will react with NH4+ ions to form NH3 and H2O
* Equilibrium will shift to left, removing OH- ions from solution
* Stops pH from changing much
88
**Making a buffer by adding a salt solution**
Calculate the pH of a buffer made from 45cm3 of 0.1 mol dm-3 ethanoic acid and 50cm3 of 0.15 mol dm-3 sodium ethanoate (Ka = 1.7 x 10-5)
89
**Making buffer by mixing weak acids and strong bases**
55cm3 of 0.5 mol dm-3 CH3CO2H is reacted with 25cm3 of 0.35 mol dm-3 NaOH. Calculate the pH of the resulting buffer solution. Ka = 1.7 x 10-5 mol dm-3
CH3CO2H + NaOH → CH3CO2Na + H2O
90
**Calculating change in pH of buffer on addition of acid/alkali**
2cm3 of 0.10 mol dm-3 NaOH is added to 100cm3 of a buffer solution containing 0.15 mol dm-3 ethanoic acid and 0.10 mol dm-3 sodium ethanoate (Ka ethanoic acid = 1.74 x 10-5 mol dm-3). Calculate the change in pH of the buffer solution.
91
Draw pH curve when...
## Footnote
Flask: 25 cm3 0.10 mol dm-3 HNO3
Burette: 50 cm3 0.20 mol dm-3 NaOH
92
**Dilution of a Strong Acid**
Calculate the pH of the solution formed when 250 cm3 of 0.300 mol dm-3 H2SO4 is made up to 1000 cm3 solution with water
* [H+] in original H2SO4 solution
* 2 x 0.300 = 0.600
* [H+] in diluted solution
* 0.600 x 250/1000 = 0.150
* pH = -log(0.150)
= 0.82
93
**Reaction between strong acid & strong base**
Calculate the pH of the solution formed when 50 cm3 of 0.100 mol dm-3 H2SO4 is added to 25 cm3 of 0.150 mol dm-3 NaOH
94
**Reaction between strong acid & strong base**
Calculate the pH of the solution formed when 25 cm3 of 0.250 mol dm-3 H2SO4 is added to 100 cm3 of 0.2 mol dm-3 NaOH
95
5 cm3 of 0.10 mol dm-3 hydrochloric acid is added to 1 dm3 of a buffer solution containing 2.35 x 10-2 mol of methanoic acid and 1.84x10-2 mol of sodium methanoate (Ka methanoic acid = 1.78 x 10-4 mol dm-3). Calculate the pH of the buffer solution after this addition. (4)
* Mol H+ added: 5 x 10-3 x 0.1 = 5 × 10–4
* Mol HCOOH: 2.35 x 10-2 + 5 x 10-4 = 2.40 × 10–2
* Mol HCOO–: 1.84 x 10-2 – 5 x 10-4 = 1.79 × 10–2
* [H+] = Ka × [HA] / [A-]
* 1.78 × 10-4 × 2.40 × 10-2 / (1.79 × 10-2 ) = 2.39 × 10-4
* pH = 3.62
96
Describe how to investigate of how the pH of a solution of ethanoic acid changes as sodium hydroxide solution is added
1. Calibrate the pH meter
1. Place pH probe in standard buffer solutions (e.g. pH 4, 7, 9) & record pH readings
2. Plot graph of pH reading (y-axis) against the pH of the buffer solution
2. Rinse pipette with ethanoic acid & fill it with this
* Transfer 20cm3 of ethanoic acid to beaker
3. Rinse burette with NaOH & then fill it with this
4. Rinse pH probe with distilled water and clamp it so bulb is immersed in ethanoic acid
* Use to rod to stir solution & record pH
5. Add NaOH to ethanoic acid in small regular intervals
* e.g. 2cm3 and then change to smaller intervals around 22cm3 e.g. 0.2cm3
* Repeat until alkali in excess
* Add in smaller increments near endpoint
6. Stir solution and take a reading after each addition
97
Explain briefly why a pH meter should be calibrated before use (1)
Over time/after storage meter doesn't give accurate readings
98
State why water at 50°C is neutral (1)
[H+] = [OH−]
99
Describe breifly how you would ensure that a reading from a pH meter is accurate (2)
* Calibrate meter with solution(s) of known pH/buffer(s)
* Adjust meter/plot calibration curve
100
Two solutions, one with a pH of 4.00 and the other with a pH of 9.00, were left open to the air. The pH of the pH 9.00 solution changed more than that of the other solution. Suggest what substance might be present in the air to cause the pH to change. Explain how and why the pH of the pH 9.00 solution changes. (3)
* CO2
* pH decreases
* acidic (gas) reacts with alkali / OH−
* or CO2 + 2OH− → CO32− + H2O
* CO2 + OH− → HCO3−
101
Use information from the curve in the figure above to explain why the end point of this reaction would be difficult to judge accurately using an indicator. (2)
* The change in pH is gradual at the end point
* An indicator would change colour over a range of volumes of sodium hydroxide / indicator would not change colour rapidly with a few drops of NaOH
