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Revision:Acids and Bases - 19

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TSR Wiki > Study Help > Subjects and Revision > Revision Notes > Chemistry > Acids and Bases - 19

19.1 Lewis theory

19.1.1

A lewis acid is defined as a species which accepts an electron pair to form a dative/coordinate 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 (or in BL acid/base reactions because the H⁺ ion has no electrons, and so must accept a pair to bond with anything. Lewis theory is more general than BL).

19.2 Calculations involving acids and bases

19.2.1

$\mathsf{H_2O_{(l)} \leftrightarrows H^+_{(aq)} + OH^-_{(aq)} :} K_w=[H^+][OH^-]$

The value of K_w is 1 x 10-14 at 25^oC, but varies with temperature.

19.2.2

$pH = -\log [H^+]$ (pH is the negative log of the concentration of H⁺ ions),

$pOH = -\log [OH^-]$ ,

$pK_w = -\log ([H^+][OH^-])$ ( is equal to 14 at 25^oC)

19.2.3

Use the above equations to calculate other values.

Nb: [H⁺] or [OH^-] are directly related to the concentration of the acid/base....therefore doubling the concentration of the acid will double [H⁺] and halve [OH^-] (and the reverse is true for bases).

19.2.4

In general:

$\mathsf{ HA_{(aq)} \leftrightarrows H^+_{(aq)} + A^-_{(aq)}}$ -- or -->

$B_{(aq)} + H_2O_{(l)} \leftrightarrows BH^+ + OH^-_{(aq)}}$ ,

therefore:

$\displaystyle K_a = \frac{[H^+][A^-]}{[HA]}$ and $\displaystyle K_b = \frac{[BH^+][OH^-]}{[B]}$ .

19.2.5

K_a is a constant which describes the ionisation of an acid (ie how strong it is) and K_b does the same for bases. pK_a is the log for of K_a, defined as:

$pK_a = -\log (K_a)$ and $pK_b = -\log (K_b)$ .

Like previously with the pH scale, a 1 fold change in pK_a will signify a ten fold change in K_a and the same for K_b.

19.2.6

$K_a \times K_b = K_w$ (ie they equal 1 x 10^-14 at 25^oC) and

pK_a + pK_b = pK_w (ie 14 at 25^oC).

19.2.7

Strong acids have weak conjugate bases. Strong bases have weak conjugate acids...A strong acid has a large K_a value (or a small pK_a value). A strong base, likewise, has a large K_b value and a small pK_b value.

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 (but it is often easier than messing around with assumptions etc, especially when your calculator can solve it in a second :)

19.3 Buffer solutions

19.3.1

A buffer solution is composed of a weak acid/base and it's conjugate base/acid. (assuming weak acid for what follows, reverse for base). A solution of weak acid is made, and this forms a equilibrium with the water as follows:

$\mathsf{HA + H_2O \leftrightarrows A^- + H_3O^+}$ .

To this solution, some of the acid's conjugate added (A^-), resulting in an increase in the concentration of A^-, some of which reacts with the H3O⁺. The result of this is, when equilibrium is reestablished, 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.

19.3.2

The pH of a buffer solution can be found from the expression for K_a ( $K_a = \frac{[H^+][A^-]}{[HA]}$ ). This can be rearranged to:

$[H^+] = K_a \times \frac{[HA]}{[A^-]}$ .

Given the concentration of both the Acid and its conjugate base, and the K_a value of the acid, the [H⁺] can be calculated and this can then be converted into a value for pH.

19.4 Salt solutions

19.4.1

For cations:

$\mathsf{2A^+ + H_2O \longrightarrow A_2O + 2H^+}$ ,

$\mathsf{A^{2+} + H_2O \longrightarrow AO + 2H^+}$ etc...

For anions:

$\mathsf{A^- + H_2O \longrightarrow AH + OH^-}$ ,

$A^{2-} + 2H_2O \longrightarrow AH_2 + 2OH^-}$ etc...

19.4.2

Cations (positive ions) will be acidic, Anions (negative) will be basic in water...with the following exceptions

Neutral cations: Group 1, Be and Mg

Neutral Anions: Cl^-, Br^-, I^-, SO₄^2-, NO₃^-, ClO₄^-.

19.5 Acid-base titrations

19.5.1

Titration curves... (see page 57-58 of Advanced chemistry by Michael Lewis)

Strong acid, strong base...The curve starting off very low is initially very flat until at equivalence it is almost vertical then very flat again (starting and finishing very low and high respectively one the graph due to low and high pH 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 = pK_a or pOH = pK_b and pH + pOH = 14. So you can find pK_a or pK_b from the curve. It can be derived, but it's easier to remember it)