•The pH scale is a convenient way to represent solution acidity. The pH is a log scale based on 10, where

pH = -log[H+]

Thus for a solution where[H+] = 1.0 x 10-7 MpH = -log(1.0 x 10-7)pH = -(-7.00) = 7.00

•Significant figures for logarithms.•The rule is that the number of decimal places in the log is equal to the number of significant figures in the original number.

[H+] = 1.0 x 10-9 M

pH = 9.00

2 significant figures

2 decimal places

•Since pH is a log scale based on 10, the pH changes by 1 for every power of 10 change in [H+].•For example, a solution of pH 3 has a H+ concentration 10 times that of a solution of pH 4 and 100 times that of a solution of pH 5.

•Since pH is defined as –log[H+], the pH decreases as [H+] increases.

•Consider the log form of Kw:Kw = [H+][OH-]log Kw = log[H+] + log[OH-]-log Kw = -log[H+] – log[OH-]

•Thus pKw = pH + pOH

•Since pKw = 14.00pH + pOH = 14.00

•When dealing with acid-base equilibria, the focus is on the solution components and their chemistry.•For example, what species are present in a 1.0 M solution of HCl? •Since HCl is a strong acid, we assume it is completely dissociated; thus, a 1.0 M solution of HCl exists as H+ and Cl- ions rather than HCl molecules.•Next we must consider which components are significant and which can be ignored.•In 1.0 M HCl, the major species are H+, Cl-, and H2O. •Since the solution is very acidic, OH- is present only in tiny amounts and can be ignored.

•To illustrate, let’s calculate the pH of 1.0 M HCl.•List the major species involved: H+, Cl-, and H2O.•Since we want to calculate the pH, focus on those major species that furnish H+.•H+ from dissociation of HCl •H+ from autoionization of H2O.

H2O(l) ⇌ H+(aq) + OH-(aq)•From Le Chatelier’s Principle the H+ from the dissociated HCl will drive the position of this equilibrium to the left.•Thus the amount of H+ contributed by water is negligible and can be neglected.•Thus, the [H+] in the solution is 1.0 M.•The pH is then

pH = -log[H+] = -log(1.0) = 0

•Remember a weak acid is one for which the equilibrium lies far to the left.•Most of the acid originally placed in the solution is still present as HA at equilibrium. •That is, a weak acid dissociates only to a very small extent in aqueous solution.•Calculating the pH of a weak acid involves looking at the equilibrium occurring in aqueous solution.•We will also look at solving for the pH of weak acid mixtures.

•Note: Percent dissociation is also known as percent ionization.•It is often useful to specify the amount of weak acid that has dissociated to achieve equilibrium.•Percent dissociation is defined as:

•For a given weak acid, the percent dissociation increases as the acid becomes more dilute.

%100x)L/mol(ionconcentratinitial)L/mol(ddissociateamount

onDissociatiPercent

•For solutions of any weak acid HA, [H+] decreases as [HA]0 decreases, but the percent dissociation increases as [HA]0 decreases.

•As an acid is diluted ([HA]0 decreases) the system must shift right to reach the new equilibrium position.•In other words as [HA]0 decreases Q is less than Ka and the system shifts right with more HA dissociating.