Topic 19 - Acids and Bases
From KstructIB
[edit] 19.1 Lewis theory
[edit] 19.1.1
A lewis acid is defined as a species which accepts an electron pair to form a dative/co-ordinate covalent bond. A lewis base is a species which donates an electron pair to form such a bond. This is a special type of covalent bond because the bond is formed by two electrons from one species and none from the other. This often occurs in the formation of complex ions. In BL acid/base reactions a H+ accepts a pair of electrons since it has not electrons of it's own to form a bond. Note, however, that Lewis theory is more general than BL.
[edit] 19.2 Calculations involving acids and bases
[edit] 19.2.1
H2O(l) <=> H+(aq) + OH-(aq). Kw=[H+][OH-]
The value of Kw is 1 x 10-14 at 25c, but varies with temperature.
[edit] 19.2.2
pH = -log[H+] (pH is the negative log of the concentration of H+ ions)
pOH = -log[OH-]
pKw = -log([H+][OH-]) (is to 14 at 25c)
[edit] 19.2.3
Use the above equations to calculate other values.
Note, [H+] or [OH-] are directly related to the concentration of the acid/base. Doubling the concentration of the acid will double [H+] and halve [OH-], and the vice versa for bases.
[edit] 19.2.4
In general HA(aq) <=> H+(aq) + A-(aq) or B(aq) + H2O(l) <=> BH+ + OH-(aq).
Therefore Ka = [H+][A-]/[HA] and Kb = [BH+][OH-] / [B]
[edit] 19.2.5
Ka is a constant which describes the ionisation of an acid (i.e. how strong it is) and Kb does the same for bases.
pKa is the log for of Ka, defined as pKa = -log(Ka) and pKb = -log(Kb). As with the pH scale, a 1 fold change in pKa or pKb will signify a ten fold change in Ka or Kb.
[edit] 19.2.6
Ka x Kb = Kw (both sides equal 1 x 10-14 at 25c)
pKa + pKb = pKw (both sides equal 14 at 25c)
[edit] 19.2.7
Strong acids have weak conjugate bases and strong bases have weak conjugate acids. A strong acid has a large Ka value (or a small pKa value) and a strong base, likewise, has a large Kb value and a small pKb value.
[edit] 19.2.8
All the above equations need to be applied as appropriate given the required input data. The quadratic formula will not be required. I'd suggest using it however, since it's often easier than messing around with assumptions and so on (especially when your calculator can solve it in a second).
[edit] 19.3 Buffer solutions
[edit] 19.3.1
A buffer solution is composed of a weak acid/base and it's conjugate base/acid. We'll assume that we have weak acid for what follows, so reverse it for base.
A solution of weak acid is made, and this forms a equilibrium with the water: HA + H2O <=> A- + H3O+.
To this solution, some of the acid's conjugate base (A-) is added, resulting in an increase in the concentration of A-. Some of this reacts with the H3O+.
The result of this is, when equilibrium is re-established, there is a considerable amount of both HA and A- present in the solution, in an dynamic equilibrium. If some other acid is added, this will react with the A-, but this causes the equilibrium to shift to the right, almost completely counteracting any pH change. The addition of a base, which reacts with the HA, cause the equilibrium to be shifted to the left, again resulting in very little pH change. This continues until one of the two components, either HA or A- are completely used up, at which time the pH then changes normally.
[edit] 19.3.2
The pH of a buffer solution can be found via the following expression Ka = [H+][A-] / [HA] ). This expression can first be rearranged to [H+] = Ka x [HA]/[A-]. Given the concentration of both the Acid and its conjugate base, and the Ka value of the acid, the concentration of H+ can be calculated and this can be converted into a value for pH.
[edit] 19.4 Salt solutions
[edit] 19.4.1
For cations : 2A+ + H2O -> A2O + 2H+, A2+ + H2O -> AO + 2H+ etc.
For anions : A- + H2O -> AH + OH-, A2- + 2H2O -> AH2 + 2OH- etc.
[edit] 19.4.2
Cations (positive ions) will be acidic, Anions (negative ions) will be basic in water. There are, just to make it fun, the following exceptions.
Neutral cations : Group 1, Be and Mg
Neutral anions : Cl-, Br-, I-, SO42-, NO3-, ClO4-.
[edit] 19.5 Acid-base titrations
[edit] 19.5.1
Titration curves are much easier to show with diagrams. If you have it, see page 57-58 of Advanced chemistry by Michael Lewis.
Strong acid, strong base : The curve starts off very low and is initially very flat until, at equivalence it is almost vertical. After equivalence, the curve becomes very flat again. The curve starts very low and finishes very high on the graph due to low/high pH values of strong acid and base respectively.
Weak acid, strong base : The curve begins comparatively high on the graph, and rises sharply initially. After a period it reaches a region where the solution acts as a buffer, still rising continually, but not as steep. the curve then then turns up sharply at equivalence and then tapers off to the strong base's pH value.
Strong acid, weak base : Identical to the strong acid strong base curve only the eventual point is lower since the weak base will have a lower pH.
Weak acid, weak base : The graph starts sharply up, but then tapers off, reaching only a somewhat steep section in center, before flattening off to the weak base pH. There is no steep section and so it is not possible to find a suitable indicator.
One other thing : At the point halfway to equivalence pH=pKa or pOH=pKb and pH+pOH = 14. This allows you to find pKa or pKb from the curve. This fact can be derived, but it's easier to just remember it.
[edit] 19.6 Indicators
[edit] 19.6.1
Indicators work by setting up a weak acid/base equilibrium where the acid and its conjugate base have different colors. HIn(aq) <=> H+(aq) + In-(aq). Where HIn is one colour and In is the other.
This equilibrium can be adjusted by the concentration of H+ through the addition of acids or bases, which results in a colour change.
[edit] 19.6.2
The pH range of the indicator falls around it's pKa value, and so to be useful, the pKa must fall within the inflection of the titration curve.
[edit] 19.6.3
As I think we just said, the value of pKa for the indicator must fall around the equivalence point of the titration to work effectively.
