As a result of inter-ionic effects the effective concentration is often actually less then the actual concentration and the effective concentration is termed the activity of the ion. The activity of all ions at very dilute concentrations (i.e. the activity of all the reactant ions and products) can be assumed to be unity. At higher concentrations, however, the effective concentration is defined of by its molar concentration (C_i) as, which is given by,

                                                                          (15)

Where (γ_i) is the activity coefficient and the molar concentration is related to the ionic strength (Ζ) as follows,

                                                                                      (16)

where (Z_i) is the charge on the ion (i). Thus, activities that differ from unity will result in a different emf for the cell.

The general Debye and Huckel equation give a value for (g_I) in the following form,

                                                                         (17)

_{where (a) and (b) are constants determined from experimental data.}

For dilute solutions, i.e., the ionic strength is < 0.01, then

                                                                        (18)

_{For solutions having ionic strengths one order of magnitude higher,}

                                                                        (19)

For solutions of greater ionic strengths the equation fails and activity coefficients must be determined experimentally. Using the equations discussed, consider the calculation of the pH, (i.e., -log₁₀ C⁺_H) of a solution containing 0.01 N HCl and 0.09N KCl.

If the pH of the solution is assumed to depend solely on the HCl present then,

                   pC_H   =  -log₁₀C_H  =   -log₁₀(0.01) =   2.00                   (20)

Alternatively using the equations previously developed a more accurate figure can be arrived at. Calculating the ionic strength from equation (16),

                                            _{= [(0.01 x 1)+(0.01 x 1) +(0.09 x 1) + (0.09 x 1)}

                                            _{= 0.10N From equation (19),}

                                                     γ _{H+  =  0.753} 

                               _{Thus, pαH  =  -log(αH+) = -log(γH+)(CH+)}                       

                                                                        _{=  log(0.753)x(0.1)}                  

                              Thus,                          paH = 2.12

The emf of an electrochemical cell can be calculated from the activity or concentration of the salts involved as follows.

 Consider the general half reaction, involving (n) electrons having charge (e)

                                                 Ox  +   ne  ↔ Red

Where the reaction involves the reduction of the ions Ox to the ions Red that involves (n) electrons during the process.. The potential of the half-cell is given by the Nernst equation, namely.

                                                                     (21)

or                                 (22)                                                                                                                                                                                                                                                                                        

where (E_o) is the standard electrode potential, (R ) is the gas constant, (T) the absolute temperature and (F) is faradays constant (96,457 coulombs per equivalent), and (a_Ox and a_Red) are the activities of the oxidized and reduced forms respectively.