# Acids & Bases

An Acid Base introductuon video featuring Arrhenius, Bronsted Lowry, and Lewis Acids and Bases.

A Brønsted-Lowry acid is a proton donor

A Brønsted-Lowry base is a proton acceptor

pH is a measure of the strength of acid

pH = -log10[H+]

[H+] is concentration of H+ in units of mol dm-3.

For a strong acid, total ionisation means that it is easy to tell [H+] ....it is simply the same as the acid concentration you are given.

The only exception to that is if the strong acid happens to be dibasic (like sulfuric acid) and gives two H+ for every acid which dissociates.

For a weak acid, very few of the acid molecules dissociate and so an equilibrium is set up.

For the weak acid HA, there is an equilibrium between        HA,  H+ and  A-

The acid dissociation constant, Ka is defined as

Ka = [H+][A-] / [HA]

This statement is always true. Under certain conditions we can make a modification to this Ka expression. If we have a weak acid simply dissolving in water (ie no alkali present), [H+] = [A-].

Since the weak acid is only very slightly ionised in water, [HA] at equilibrium is almost exactly the same as the initial [HA]. We can say that it is the same.

So......for this special condition, Ka expression becomes Ka = [H+]2 / [HA].

It turns out that this can be converted to

pH = ½ pKa – ½ log10[HA]

(if you find this hard to remember, try singing it to Baa-Baa Black Sheep!)

A separate set of special conditions can apply if [A-] = [HA].

This occurs when the acid is “half-neutralised”.

In this particular case,

pH = pKa

Dealing with OH-:

Water itself is slightly dissociated into H+ and  OH-

A special dissociation constant ( Kw) is defined as

Kw = [H+][OH-]

At 25 oC, Kw = 1 x 10-14 . This allows us to find the pH of an alkaline solution. If we know the [OH-], we can easily find [H+].

Titration curve calculations:

There are six stages to consider as we add alkali to a weak acid:

1.   Pure dilute acid solution: Ka = [H+]2 / [HA] or pH = ½ pKa – ½ log10[HA]

2.   Some alkali but not enough to neutralise the acid: Ka = [H+][A-] / [HA]

3.   At half equivalence: pH = pKa

4.   Some alkali but not enough to neutralise the acid: Ka = [H+][A-] / [HA]

5.   At the equivalence point: Use standard titration calculations to calculate the molarity of acid or molarity of alkali as required. The “weakness of the acid” does not matter other than in the choice of indicator.

6.   Sufficient alkali to go beyond equivalence point: Assume that the acid is all used up, calculate the number of moles of excess alkali, calculate the total volume of the solution, calculate the [OH-] and then use Kw to find the [H+]

Equivalence Points and Choice of Indicator:

pKin is the pH at which the indicator has equal concentrations of its two forms. You should choose an indicator which has pKin within the steep part of the titration curve so that there is a sharp colour change when one drop of alkali (or acid) is added.

Strong acid / strong alkali It will have pH = 7 at equivalence (many indicators will do because there is a long steep section to the titration curve)

Weak acid / strong base pH > 7 at equivalence (use phenolphthalein indicator)

Strong acid / weak base pH < 7 at equivalence (use methyl orange indicator)

Weak acid / weak base Can’t predict the pH (no suitable indicator)

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