Chemistry

Electrochemical Cells

Laboratory #15

Henry Ko

AP Chemistry

Dulaney High School

March 12th, 2009

Abstract:

In this experiment, a standard table of reduction potentials of a series of metal ions is constructed using

copper, iron, lead, magnesium, silver, and zinc. These half cells are are connected by a salt bridge and all

potentials are measured with respect to the zinc electrode. Also, the measured voltage of a nonstandard

copper cell is calculated through the Nernst equation. The solubility product constant of AgCl is also

determined through the Nernst equation. The Ksp value for AgCl was determined to be 7.33 × 10−11 ,

yielding a percent error of 59.3%. The voltage for the cell reaction was experimentally determined to be 0.81

V.

Theory:

An electrochemical cell is produced when a redox reaction occurs. The resulting electron transfer

between the reaction runs through an external wire. Because the oxidation and reduction reactions are

physically separated from each other, these are called half-cell reactions. A half cell is prepared from

contact with the metal with its solution of ions. Each element’s unique electron configuration allows each to

develop a different electrical potential.

The standard reduction potential is the voltage that a half-cell, under standard conditions (1 M, 1

atm, and 25◦ C), develops when combined with the standard hydrogen electrode, that is arbitrarily assigned

◦

to a potential of zero volts. A positive Ecell value indicates that the redox reaction in that particular cell is

spontaneous.

Calculations of nonstandard potentials can be made using the Nernst Equation:

E

= E◦ −

RT

ln(Q)

nF

(1)

where E is the measured cell potential, E ◦ is the standard cell potential, R is the gas constant (8.314 J/mol

· K), T , is the temperature in K, n is the number of moles of electrons transferred as shown by the redox

reaction, and F is the Faraday constant (9.65 × 104 C/mol).

At STP, the Nernst equation...