5.1.3 Acids, Bases and Buffers Flashcards
1
Q
what is a Bronsted-Lowry acid
A
a proton (H+) donor
2
Q
what is a Bronsted-Lowry base
A
a proton (H+) acceptor
3
Q
what are conjugate base pairs
A
consist of 2 species that can be interconverted by transfer of a proton
4
Q
in the dissociation of HCl, what forms conjugate base and what forms conjugate acid
A
HCl → H+ + Cl-
1) in forward direction, HCl releases a proton and forms conjugate base Cl-
2) in backwards reaction, Cl- accepts a proton and forms HCl
5
Q
what reaction is always taking place in a neutralisation
A
H+ + OH- → H2O
6
Q
what would be the acid base equilibrium of the dissociation of HCl
A
HCl + OH- ⇌ H2O + Cl-
7
Q
what are the conjugate acid pairs in the acid base equilibrium of dissociation of HCl
A
HCl + OH- ⇌ H2O + Cl-
FORWARD DIRECTION:
- HCl donates H+, so acid 1
BACKWARD DIRECTION
- Cl- accepts H+, so base 1
FORWARD DIRECTION+
- OH- accepts H+, so base 2
BACKWARD DIRECTION:
- H2O donates H+, so acid 2
8
Q
what is a hydronium ion
A
H3O+(aq)
- ACTIVE INGREDIENT IN ANY AQUEOUS ACID
9
Q
where is the hydronium ion made
A
- in aqueous solutions of dissociation reactions, where a proton needs to be transferred from acid to base, you MUST have water present
- so a proton is transferred to water, forming H3O+
10
Q
rewrite the HCl acid-base equilibrium, but this time using the H3O+ ion
A
HCl + H2O ⇌ H3O+ + Cl-
11
Q
how is the equation for neutralisation simplified from using the H3O+ ion
A
H3O+ + OH- → 2H2O
H+ + OH- → H2O
12
Q
what does the ____basic bit of an acid tell you
A
the total number of H+ ions in the acid that can be replaced per molecule in an acid-base reaction
- e.g. monobasic, dibasic, tribasic
13
Q
why is the ____basic bit of an acid useful
A
- used to work out neutralisation equations
- each of the H+ must be replaced, so helps you find out how much base you need
- also able to find out volumes too
14
Q
what is the role of H+ ions in acids
A
- the active species of acids that is reacting
- can be emphasised through writing ionic equations, and you’ll find for any acid with the same basic, gives you the same ionic equation
15
Q
what is the ionic equation for 2HCl + Mg
A
2H+ + Mg → Mg2+ + H2
16
Q
what are spectator ions
A
ions that do not change in a reaction, always cancelled out
17
Q
what is a redox reaction that acids go through
A
acid + base → salt + hydrogen
18
Q
what is the ionic equation basic layout for acid + carbonate
A
2H+ + CO32- → salt + H2O + CO2
19
Q
what is the ionic equation basic layout for acid + base
A
2H+ + MgO → Mg2+ + H2O
20
Q
what is the ionic equation basic layout for acid + alkali
A
H+ + OH- → H2O
- as alkali is soluble in aqueous solutions, so its ions will cancel out too
21
Q
what is pH
A
a measure of hydrogen ion H+ concentration
22
Q
how is H+ ion concentration converted to pH
A
10⁻¹⁴ goes to pH 14
- so a low [H+(aq)] is a high pH
- and a high [H+(aq)] is a low pH
23
Q
what pHs indicate what
A
pH>7 : alkali
pH<7 : acid
pH=7 : neutral
24
Q
how can pH be measured
A
using a pH meter or indicators, either paper or universal
25
what type of scale is pH
logarithmic
- one pH number change is a x10 jump in concentration of H+ ions
- e.g. pH 1 has 10x more H+ ions than pH 2, and so on
26
how can you convert between pH and [H+]
pH = -log[H+]
[H+] = 10⁻ᵖᴴ
27
what do strong, monobasic acids do in aqueous solutions, and what does this mean
COMPLETELY dissociate
- HA → H+ + A-
- one mole of HA turns into one mole of H+
- so [HA] = [H+]
28
what is the equation for finding [H+] in strong acids
[H+]=[HA]
- concentration of acid is equal to concentration of H+ ions
29
how can you figure out pH changes based on dilution
- through using [H+]=[HA] formula for both
- we know pH 1 has x10 [H+] than pH2, but this is not a set scale
30
what is the difference between strong and weak acids
strong acids: completely dissociate in aqueous solutions (HCl)
weak acids: partially dissociate in aqueous solutions (CH3COOH) ⇌
31
what constants can be used for acid-base equilibria
- special equilibrium constants
- but all just versions of Kc
- where [products]/[reactants]
32
what is the dissociation of weak acids formula
HA ⇌ H+ + A-
33
what is the acid dissociation constant called
Ka
- used to find out [H+] of weak acids
34
what is the Ka
[H+][A-]/[HA]
35
what are the units of Ka
moldm-3
36
when will Ka change
- ONLY with temperature, as will all Kc
- set at standard 25 degrees celsius
37
what does a larger KA value show
- the equilibrium is further to the right, so greater [H+]
- greater dissociation
- STRONGER the acid
38
why is Ka often difficult to deal with
- often get numbers with negative indices
- difficult to compare, so convert to the negative log instead of pKa
- is more manageable than Ka
- and easier to compare relative strengths
39
what is the conversion of Ka to pH
pKa = logKa
Ka = 10⁻ᵖᴷᵃ
40
what is the relationship between strength of acid, Ka and pKa
- strong acid
- large Ka
- small pKa
41
at equilibrium, how is [HA] and [H+] related for strong and weak acids
strong acid : [H+]=[HA]
weak acid : [H+]≠[HA] , but rather HA ⇌ H+ + A-
42
what does the concentration of H+ formed at equilibrium depend on
- the concentration of the acid
- the Ka acid dissociation constant
43
rewrite Ka using equilibrium values
[H+]equi[A-]equi/[HA]equi
- where [HA] equi can be rewritten as [HA]start-[H+]equi
- where when KA dissociates, H+ and A- are mad in equal quantities
44
what are the two approximations you make when identifying the value of Ka of an acid
1) H+ and A- are equal, so [H+]=[A-]
(although are the same, some H+ is formed via the dissociation of water, but will be extremely small and can be neglected
2) assume [HA]>>[H+], as dissociation of weak acid is small, so can neglect any decrease in concentration of HA from dissociation, so [HA]start=[HA]equi
45
what is the Ka formula produced once approximations have been accounted for
[H+]²/[HA]
46
how can you work backwards from Ka to figure out [H+]
[H+] = √Kax[HA]
- from this, can calculate pH via -log[H+]
47
how can you determine Ka experimentally
- prepare a standard solution of a weak acid of known concentration
- measure the pH using a meter
- work out [H+] using 10⁻ᵖᴴ
48
what are the limits of the approximations we make when using Ka formula
1) assume dissociation of water is negligible, so [H+]=[A-]
- at 25 degrees, dissociation of water = 10⁻⁷, so if acid has pH above 6, it becomes significant
- so approximation breaks down for: VERY WEAK OR DILUTE ACIDS
2) assume [HA]>>[H+], so only works for weak acids with a small Ka value, as otherwise [HA] would be drastically different from [HA] at equilibrium ([HA]-[H+])
- so approximation breaks down for stronger weak acids with Ka>10⁻² or very dilute solutions
49
what happens during the ionisation of water
- ionises slightly, and acts as both an acid and a base
- H2O + H2O ⇌ H3O+ + OH-
- simplifies to H2O ⇌ H+ + OH-
50
what would Ka be if we treat water as a weak acid
[H+][OH-]/[H2O]
51
whats special about the dissociation of water
- it is VERY SMALL, so [H2O] can be treated as a constant
52
what does Ka of water rearrange to
Ka x [H2O] = [H+][OH-]
Ka x [H2O] is Kw
53
what is Kw
the ionic product of water
54
what is the equation of Kw
Kw = [H+][OH-]
55
when does Kw vary
varies with temperature, but at 25 degrees celsius, is equal to 1x10⁻¹⁴
56
how can the pH of pure water be identified at 25 degrees celsius
- on dissociation, is neutral, so produces same number of [OH-] as [H+]
- so [OH-]=[H+]
- Kw=[H+]²
- = 1 x 10⁻¹⁴
- rearrange for H+ and put into pH formula and you get answer of 7
57
what does Kw control
equilibrium constant that controls the concentrations of H+ and OH- in aqueous solutions
- if acid: [H+]>[OH-]
- if alkali: [H+]<[OH-]
- if neutral: [H+]=[OH-]
- Kw has a value of 1 x 10⁻¹⁴, which is able to control these relative concentrations
58
for whole numbered pHs, how can Kw be used quickly
the indices just need to add up to -14, so if you know one, you know the other
59
what is an alkali
a soluble base that releases OH- ions in aqueous solutions
60
what is a strong base
COMPLETELY DISSOCIATES
- AOH → OH- + H+
- AOH and OH are both one mole
- so [AOH]=[OH-]
- where A is any metal
61
how can you find pH of a strong base
1) from concentration of acid, you know concentration of OH- as they are equal
2) can sub this into Kw along with Kw's value
3) find out concentration of H+ from this
4) sub into pH=-log[H+]
62
what is a weak acid
equilibrium is positioned well to the left, only partially dissociate and release OH- ions in aqueous solutions
63
what is a buffer solution
a system that minimises the pH change when small amounts of an acid or base are added
64
what are the 2 components of a buffer
1) weak acid (removes the alkali)
2) and its conjugate base (removes the acid)
65
what happens to the components of a buffer when you add and acid or alkali
- the 2 components in the buffer solution react
- will eventually become used up
- as soon as one component has been fully used up
- the solution will lose its buffering ability towards the added acid/alkali
66
what is true about the pH overtime of a buffer
- as the buffer works
- the pH changes by a small amount
- doesn't stay completely constant throughout
67
when you are making a buffer from a weak acid, what 2 components do you need to form
- the weak acid
- and its conjugate base
68
what are 2 methods of forming weak acid buffer solutions
- from a weak acids and its salt
- from the partial neutralisation of a weak acid
69
how do you form a buffer from a weak acid and its salt
- mix solutions of the weak acid
- and a solution of its salt together
- e.g. ethanoic acid solution and sodium ethanoate solution
70
when forming a buffer from the solution of a weak acid and its salt, where do you get the acid component of the buffer from
1) when you add the acid to water, it partially dissociates
- e.g. CH3COOH ⇋ CH3COO- + H+
- the CH3COOH is the component (there is only a very small amount of the CH3COO- ion, as a weak acid so only partially dissociates)
71
when forming a buffer from the solution of a weak acid and its salt, where do you get the conjugate base component of the buffer from
2) when you add the salt to the water, it dissolves and completely dissociates into ions, providing the source of the conjugate ion
- e.g. CH3COONa + aq → CH3COO- + Na+
- the CH3COO- is the conjugate base component (and the salt has completely dissociated into this)
72
what do you need to get a buffer through the neutralisation of a weak acid
- excess of a weak acid
- aqueous solution of an alkali, like NaOH
73
how do you get a buffer through the partial neutralisation of a weak acid
- the weak acid is partially neutralised by the alkali
- forms the conjugate base
- but weak acid is in excess, so some is left over unreacted
- so left with the salt of a weak acid (containing the conjugate base) and the unreacted weak acid
74
how is a buffer set up, so supply high concentrations of the 2 components
equilibrium: CH3COOH ⇋ CH3COO- + H+
- initially, the equilibrium lies well to the left, with the CH3COOH is in excess as we have a weak acid which only partially dissociates
- when we add extra CH3COO- ions (conjugate base ions), the concentration of them increases
- the equilibrium position will shift even further to the left
- so now we have CH3COOH and CH3COO- in excess
- but a VERY LOW concentration of H+ ions
- means the solution is mainly just CH3COOH and CH3COO-
75
how do the 2 components of a buffer act
- the weak acid and its conjugate base
- act as 2 reservoirs
- act independently to eachother to remove added acid/alkali
- via shifting the buffer's equilibrium system to the left and the right
76
what do buffers control about a solution
- the pH of it
- via work of the conjugate acid-base pair
77
when you add an acid, how does the conjugate base help restore pH
- if you add acid
- [H+] increases
- the H+ ions will react with the conjugate base
- the equilibrium position will shift to the left
- removing most of the H+ ions
- WHEN YOU ADD ACID, EQUILIBRIUM SHIFTS TO THE LEFT, AWAY FROM THE H+
78
when you add an acid, how does the weak acid act to restore pH
- when you add an alkali OH-
- the small concentration of H+ ions present react with the OH - ions (H+ + OH- → H2O)
- the weak acid will dissociate, shifting the equilibrium to the right
- to restore most of the H+ ions which have been lost
- WHEN YOU ADD ALKALI, EQUILIBRIUM SHIFTS TO THE RIGHT, TO FORM MORE H+
79
when are buffer solutions the most effective
- most effective at removing added acid and alkali
- when there are equal concentrations of weak acid and conjugate base
80
what is true about a buffer when [HA] = [A-]
- Ka = [H+] (from the Ka equation)
- so pKa of the weak acid = pH of the buffer
- and the operating pH of the buffer will be one below and one above the pH/pKa value of it
- can then adjust the ratio of the [HA] / [A-] to fine tune the pH
81
what is the equilibrium and the Ka equation of a buffer solution
- HA ⇌ H+ + A-
- so Ka = [H+][A-]/[HA]
82
what must you know in order to calculate the pH of a buffer
- the [H+] value from the Ka equation
- BUT you can't approximate [H+]=[A-] anymore
- as you have added the [A-] (conjugate base) in as one of the components
- so you need to know all three other components of the equation to work out
- including the Ka value of the weak acid, and the ratio of [HA]/[A-] (ratio of acid to conjugate base)
83
how do you calculate the pH of a buffer
- need to find the concentrations of HA and A-, and use the value of Ka given
- if you do not have concentrations already given, you need to work them out (work out the moles present, and then their concentration in the TOTAL volume of solution)
84
how do you calculate the pH of a buffer made from partial neutralisation of a weak acid
- calculate the moles of A- (equal to the amount of alkali (NaOH) used)
- calculate the amount of moles of HA that you began with
- calculate the moles of HA present in the buffer, which will the the moles to begin with - the moles of conjugate base made (as some of the weak acid is used up to make the base, and the remaining amount (as in excess) is the weak acid) [n(HA)-n(NaOH)]
85
why do we need to control the pH of the human body
- needs to be precisely controlled
- and different parts need different pHs
- e.g. to provide the optimum pH for enzymes
- pH of normal healthy blood is 7.4 (ranges from 7.35 to 7.45)
86
how is pH controlled in the human body
- controlled by a mixture of buffer solutions
- e.g. in the blood plasma
- most important buffer is the carbonic acid/hydrogencarbonate buffer
- H2CO3/HCO3-
87
what happens if your pH changes even by a small amount outside the range
- below 7.35 = acidosis (fatigue, shortness of breath, shock, death)
- above 7.45 = alkalosis (muscle spasms, lightheadedness, nausea)
- although small differences in pH, make a very big difference as pH is such a sensitive scale
88
what is the equilibrium for the carbonic acid/hydrogencarbonate buffer system
H2CO3 ⇌ H+ + HCO3-
89
what happens when you add an acid to the carbonic acid/hydrogencarbonate buffer system
- add acid
- [H+] increases
- H+ react with the conjugate base HCO3-
- equilibrium shifts to the left, removing most of the H+ ions
90
what happens when you add alkali to the carbonic acid/hydrogencarbonate buffer system
- add alkali
- [OH-] increases
- the small concentration of H+ present react with the OH- (make water)
- H2CO3 dissociates, shifting the equilibrium to the right
- restoring most of the H+ ions
91
how can you find the concentration ratio of HCO3-/H2CO3 in the blood
- find the ratio you need by manipulating the Ka equation
- find Ka and H+ values (e.g. if they give you pH and pKa, just work backwards)
- input into the equation you need, and get the ratio
92
as the body produces more acidic products than alkaline, what does this mean for the components of the carbonic acid/hydrogencarbonate buffer system
- more HCO3- is converted into H2CO3
- so it can build up
- but to prevent this
- H2CO3 is converted into CO2
- and exhaled by the lungs
93
how can we monitor the neutralisation reaction taking place in an acid-base titration
- using indicators
- using pH meters to monitor the pH changes taking place throughout the titration
94
what is a pH meter, and why may it be used over indicator paper
- an electrode dipped into a solution and connected to a meter displaying the pH reading
- more accurate than indicator paper
- gives value in 2dp, as opposed to colour comparison to chart of whole numbers
95
PRACTICAL: how can you use a pH meter to monitor the pH throughout a acid-base titrations
- add a measured volume of acid to a conical flask
- place the pH meter electrode inside
- add a base to the burette, and add it into the acid solution 1cm3 at a time
- after each addition, swirl and record the pH and the total volume of base you have added
- repeat until the pH change is more rapid, and then add the base dropwise
- add in 1cm3 readings again until you have added an excess of base (the pH has stayed basic and relatively the same for a while)
- plot a graph of pH against total base(aq) added
96
PRACTICAL: how can you use a pH meter to track a titration completely automatically
- use a data logger on a probe
- and a magnetic stirrer for the flask
97
explain the 3 sections of a pH titration curve, where a strong base has been added to a strong acid
1: pH increases slowly as the acid is still in great excess (as it approaches the vertical section, increases quicker as acid is used up more quickly)
2: the pH increase rapidly even when a very small volume of base is added, forming the vertical section, where the concentrations of base and acid are similar
3: the pH rises very slowly as you add the basic solution, as the base is now in great excess
- equivalence point is found in the centre of the vertical section of a pH titration curve
98
explain the pH titration curve where you are adding an acid to a base
- same as before
- but just flipped around
99
what is the equivalence point
gives the volume of one solution that reacts EXACTLY with the volume of another solution
- so the point where they have exactly reacted with eachother matching the stoichiometry of the reaction
100
what are acid-base indicators made of
- a weak acid HA that has a distinctively different colour from its conjugate base A-
101
what is an example of an acid-base indicator
- methyl orange: red when a weak acid and yellow when its conjugate base
102
what's the colours of an indicator at its end point
- colour in between the 2 extremes
- as contains equal concentrations of HA and A-
- e.g. orange for methyl orange
103
why do indicators show a colour change
- in acidic conditions, the equilibrium of the indicator shits towards the weak acid
- in alkali conditions, the equilibrium of the indicator shifts towards the conjugate base
- and changes colour as it does so
104
what is the equilibrium of indicators
- just weak acids
- so HA ⇌ H+ + A-
105
what happens to methyl orange indicator when you are adding a strong base to a strong acid
- methyl orange is initially red (as H+ ions, so equilibrium shifts to the left, towards the weak acid)
- when you add base, you are adding OH-
- this reacts with the H+ (form water)
- the weak acid then dissociates, shifting the equilibrium towards the right
- so the colour changes, as now more A- present
- orange at end point
- yellow once fully shifted to the right
106
what happens to the colours of methyl orange when you add a strong acid to a strong base
- methyl orange is initially yellow
- acid H+ is added
- the H+ added react with the A- conjugate base
- and the equilibrium shifts to the left
- the colour changes, for orange at end point and red once fully shifted towards the acid
107
what will be the pH of an indicator at its end point
- different indicators have different Ka values
- and change colour at different pH ranges
- at the end point, [HA]=[A-]
- so Ka=[H+]
- so pH at end point = Ka of the weak acid
108
what does the sensitivity of an indicator depend on
- the indicator used itself
- and the eyesight of those observing it
109
what indicator do you need to choose for a reaction
- must have an indicator that produces a colour change during the vertical section of the pH titration curve
- ideally at the equivalence point
- but not always possible, and the end point might actually occur when the volume is slightly out of the equivalence point
- but doesn't really matter, as you will only be out dropwise
110
what is important to see graphically when choosing indicators
- where the indicator does a colour change
- and whether this is in the vertical section of the pH titration curve
111
what do the different pH titration curves look like
- strong acid-strong base: straight lines at the bottom and the top, equivalence point in the middle
- weak acid-strong base: the bottom straight line doesn't go all the way to the bottom, so starts higher up
- strong acid-weak base: the top straight line doesn't go as high, so shifted down
- weak acid-weak base: neither lines go straight to the bottom or the top, and no real vertical section in between
112
what is true about choosing indicators for weak acid-weak base titrations
- no indicator is suitable
- as no real vertical section
- so no real equivalence point
- even at its steepest section, you still need a considerable amount of base (a few cm3) to even increase the pH by 2 units
