Topic 12: Acid-Base Equilibria

Revision cards for the 2015 spec Edexcel A Level Chemistry, topic 12.

(For some reason some reversible arrows have come out as --> rather than <--> , sorry :/ I can't change it!)

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  • Created by: LouiseG
  • Created on: 16-10-16 11:28
Define a Bronstead-Lowry acid
A proton (i.e. hydrogen ion) donor (e.g. HCl --> H+ + Cl- )
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Define a Bronstead-Lowry base
A proton acceptor (e.g. NH3 + H+ ---> NH4+ )
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What is the conjugate base of HCl?
Cl- . The conjugate base of an acid is the species formed from an acid when it looses a/some protons. In the "reverse" reaction, it would act as the "base" and gain a proton to become the original species again.
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What is the conjugate acid of HSO4- ?
H2SO4 (the species after it has gained a proton)
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What is a diprotic acid?
An acid that can donate two H+ ; e.g. H2SO4
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What is another name for a diprotic acid?
A dibasic acid
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CO3 2- is a _____ base
Diacidic (it can accept two protons from an acid)
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Why is H2O amphoteric?
Amphoteric species can act as both acids and bases. Water can accept a proton to become H3O+ or loose one to become OH-
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Define a strong acid
An acid which dissociates completely in solution: HCl --> H+ + Cl-
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Define a weak acid
An acid which only partly dissociates, represented by a reversible arrow
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What do square brackets [ ] denote?
Concentration: [H+] = concentration of hydrogen ions
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What is the equation to work out pH?
pH = -log[H+]
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A strong acid is assumed to be __(a)___ therefore the concentration of the acid will be __(b)__to [H+]
(a) Fully dissociated (b) equal to (for a monobasic acid - its a bit more complecated with dibasic, as the second dissociation doesn't go to completion)
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Therefore, what is the pH of 0.02moldm-3 HCl?
[H+] = [HCl] = 0.02 ; pH = -log0.02 = 1.70
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How many d.p. is pH given to?
Two
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Can pH be negative?
Yes! A strong acid could have a pH less than 0 in calculations. (However, this isn't seen in real practice due to the difference between H+ ion activity and concentration)
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How is [H+] calcuated from pH?
10^-pH
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What is the expression for the partial dissociation of a weak acid?
HA < --> H+ + A-
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What is the simplified expression of Ka?
[H+]^2 / [HA]
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What assumption has been made in this expression?
That [H+] = [A-] at eq. ; therefore the expression Ka = [H+][A-]/[HA] is simplified
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Therefore, how can you calcuate the pH of a weak acid?
Find [H+] from [H+] = sqrt(Ka x [HA]) (Rearrange the Ka equation) and then put into -log[H+]
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What other constant can you calcuate from Ka? How?
pKa, which is like the pH of Ka (it gives a smaller range of values, and a smaller number = stronger acid). Thus it is pKa = -logKa
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What is Kw?
The ionic product of water, Kw = [OH-][H+] (water's dissociation constant)
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What is it's value at 298K?
1.0x10^-14 mol2dm-6
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Therefore what is the pH of water at this temperature?
If [OH-]=[H+] at eq, then Kw=[H+]^2 ; therefore [H+]=sqrt(Kw) = 1.0x10^-7. Put into -log[H+] gives a pH of 7.00
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What happens to the pH of water at its temperature rises?
It decreases. This is as the reaction H2O --> H+ + OH- is endothermic in the forward direction, therefore more [H+] ions are released at high temps (hence lower pH) -Note it is still CHEMICALLY NEUTRAL, as [H+] = [OH-] still.
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How is the pH of a strong base calculated?
As it is strong, the base's conc is directly proportional to its [OH-]. Therefore, [H+] can be found from [H+] = Kw/[OH-] ; and then put into -log[H+]
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How many times stronger (for a strong acid) is pH 1.00 than 2.00?
Ten times (it's a logarithmic scale)
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The pH value of a weak acid increases by __(a)__ for every 10-fold increase in concentration
0.5. This is as the addition of water affects the equilibrium position and produces more [H+] (it's more acidic than you would expect).
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How do you find a pH of a mixture (STRONG acid-base)?
1. Find the moles of acid and base 2.Considering the ratio they react in, take the smaller mole value from the larger. 3.Find the new conc of this excess acid/alkali 4. Calculate pH as you would for an acid or base
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What is important about the half equivalence point on a titration curve?
It's the point where [A-] = [HA]; therefore Ka=[H+] and pH=pKa (can find Ka from the pH, 10^-pH = Ka)
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What is the equivalence point?
The point where the acid and base have reacted in exactly the ratio as given in the stoichiometric equation. It is NOT necessarily neutral or pH 7.00
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What happens around the equivalence point on a titration curve?
The pH changes rapidly (steep section of the graph) (except for weak acid - weak base titrations)
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In a weak acid - strong base titration curve, roughly where should equivalence be?
pH 8-9 (ABOVE pH 7)
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Which type of titration is Phenolphthalein suitable for?
Weak acid-Strong base, as its pH range is 8.2 - 10.0. Also maybe strong acid-strong base.
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Why is methyl orange not suitable as an indicator for the titration between NaOH and CH3COOH?
Because its colour change (pH range) does not fall near the equivalence point (in the steep section of the graph).
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Why does the pH in a strong base - weak acid titration only change gradually at first?
A buffer solution has formed (from the weak acid and its conjugate base).
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What is a buffer solution?
A solution that resists changes in pH (pH only changes by a small amount) when small amounts of acid or base are added.
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Give an acid buffer.
Any weak acid and its conjugate base, (in the form of its K or Na salt): e.g. ethanoic acid and sodium ethanoate.
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Give the most common basic buffer.
Ammonia and ammonium chloride.
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How does an acidic buffer of sodium ethanoate and ethanoic acid oppose the addition of an acid (H+ ions?)
CH3COO- + H+ ---> CH3COOH , thus "using up" the H+ and shifting the position of equilibrium
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Write two equations for the reaction of a base, OH- , with this acid buffer.
CH3COOH + OH- --> CH3COO- + H2O and OH- + H+ ---> H2O
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What is the equation for calcuating the [H+] of an acidic buffer?
[H+] = Ka x ( [Acid]/[Salt] )
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What is the important buffer in the blood?
H2CO3, Carbonic acid, with HCO3- , the carbonate ion: H2CO3 --> HCO3- + H+
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Give another equilibrium this acid is found in, with CO2
H2CO3 --> H2O + CO2
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Why do some weak acids have a less exothermic enthalpy of neutralisation than strong acids?
Because energy is required to disassociate the acid (endothermic) which reduces the overall energy released.
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Why does HF release more energy on neutralisation than many strong acids?
Because the exothermic hydration process (when the ions of H+ and A- become (aq) ) releases more energy than is needed for dissociation
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Other cards in this set

Card 2

Front

Define a Bronstead-Lowry base

Back

A proton acceptor (e.g. NH3 + H+ ---> NH4+ )

Card 3

Front

What is the conjugate base of HCl?

Back

Preview of the front of card 3

Card 4

Front

What is the conjugate acid of HSO4- ?

Back

Preview of the front of card 4

Card 5

Front

What is a diprotic acid?

Back

Preview of the front of card 5
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