Electrochemistry Notes

Electrochemistry

Objectives

• Review the concepts of oxidation and reduction.
• Define the term "cell potential."
• Illustrate the Standard Hydrogen Electrode (SHE) and standard electrode (reduction) potentials.

What is Electrochemistry?

• Electrochemistry is the study of chemical reactions and their associated electrical changes.
• It also examines chemical changes that an electrical current can bring about.

Transfer of Electrons

• Electrochemistry involves the transfer of electrons.
• Oxidation: Compound A loses electrons.
  • A is a reducing agent and becomes oxidized.
• Reduction: Compound B gains electrons.
  • B is an oxidizing agent and becomes reduced.

Half Cell

• A half-cell consists of a metal electrode, $M$, partially immersed in an aqueous solution of its ions, $M^{n+}$.
• The anions required to maintain electrical neutrality in the solution are not shown.
• This is limited to metals that do not react with water.
• Oxidation: $M <br />\nrightarrow M^{r+}$
• Reduction: $M^{n+} <br />\nrightarrow M$

Cell Potential

• If a strip of metal (electrode) is placed in a solution of its ions, the metal loses electrons to form its cation: $M(s) <br />\nrightarrow M^{n+}(aq) + ne^{-}$
• The cations in solution may accept some of the lost electrons and get reduced to the metal atoms: $M^{n+}(aq) + ne^{-} <br />\nrightarrow M(s)$
• At equilibrium, the charge difference that develops between the metal strip and the solution is called a potential difference: $M^{n+}(aq) + ne^{-} <br />\nleftrightarrow M(s)$

Cell Potential Definition

• Cell potential (measured in Volts, V), also known as electromotive force (emf), is the tendency of species to lose or gain electrons.
• $E°$ represents the standard electrode potential, which is the potential of a species compared with the potential of a Standard Hydrogen Electrode.
• $2H^+ + 2e^- <br />\nleftrightarrow H<em>2$ $E°</em>{cell} = 0.00V$

Standard Hydrogen Electrode (SHE)

• The Standard Hydrogen Electrode (S.H.E.) consists of Hydrogen gas at 298K and 1 atm bubbling over a platinum electrode immersed in a solution of $H^+$ ions with concentration 1 $moldm^{-3}$.
• 2H^+(aq) + 2e^-
  rightleftharpoons H_2 (g, 1 atm) $E° = 0.00 V$

Relative Cell Potential

• More reactive metals lose electrons with greater ease, so the equilibrium position of the redox reaction lies more to the left: M^{n+}(aq) + ne^-
  rightleftharpoons M(s)
• More reactive metals possess more negative reduction potentials, and less reactive metals possess more positive reduction potential differences.

Standard Cell Potential of Zn Half Cell

• Zn (s)
  rightleftharpoons Zn^{2+}(aq) + 2e^- $E° = +0.76 V$
• Zn^{2+}(aq) + 2e^-
  rightleftharpoons Zn (s) $E° = -0.76 V$

Cell Potential and Oxidizing/Reducing Agents

• Species with a positive $E°$ value have a tendency to be reduced and are strong oxidizing agents (usually non-metals).
• Species with a negative $E°$ value have a tendency to be oxidized and are strong reducing agents (usually metals).

Table of Standard Reduction Potentials

• The table lists various reduction half-reactions and their corresponding standard reduction potentials ($E°$ in Volts).
• Stronger oxidizing agents are at the top of the table (more positive $E°$ values), while stronger reducing agents are at the bottom (more negative $E°$ values).
• Examples of half-reactions and their $E°$ values:
  • $F_2(g) + 2e^- <br />\nrightarrow 2F^-(aq)$, $E° = 2.87 V$
  • $H<em>2O</em>2(aq) + 2H^+(aq) + 2e^- <br />\nrightarrow 2H_2O(l)$, $E° = 1.78 V$
  • $MnO<em>4^-(aq) + 8H^+(aq) + 5e^- \nrightarrow Mn^{2+}(aq) + 4H</em>2O(l)$, $E° = 1.51 V$
  • $Cl_2(g) + 2e^- <br />\nrightarrow 2Cl^-(aq)$, $E° = 1.36 V$
  • $Cr<em>2O</em>7^{2-}(aq) + 14H^+(aq) + 6e^- <br />\nrightarrow 2Cr^{3+}(aq) + 7H_2O(l)$, $E° = 1.36 V$
  • $O<em>2(g) + 4H^+(aq) + 4e^- \nrightarrow 2H</em>2O(l)$, $E° = 1.23 V$
  • $Br_2(aq) + 2e^- <br />\nrightarrow 2Br^-(aq)$, $E° = 1.09 V$
  • $Ag^+(aq) + e^- <br />\nrightarrow Ag(s)$, $E° = 0.80 V$
  • $Fe^{3+}(aq) + e^- <br />\nrightarrow Fe^{2+}(aq)$, $E° = 0.77 V$
  • $O<em>2(g) + 2H^+(aq) + 2e^- \nrightarrow H</em>2O_2(aq)$, $E° = 0.70 V$
  • $I_2(s) + 2e^- <br />\nrightarrow 2I^-(aq)$, $E° = 0.54 V$
  • $O<em>2(g) + 2H</em>2O(l) + 4e^- <br />\nrightarrow 4OH^-(aq)$, $E° = 0.40 V$
  • $Cu^{2+}(aq) + 2e^- <br />\nrightarrow Cu(s)$, $E° = 0.34 V$
  • $Sn^{2+}(aq) + 2e^- <br />\nrightarrow Sn^{2+}(aq)$, $E° = 0.15 V$
  • $2H^+(aq) + 2e^- <br />\nrightarrow H_2(g)$, $E° = 0 V$
  • $Pb^{2+}(aq) + 2e^- <br />\nrightarrow Pb(s)$, $E° = -0.13 V$
  • $Ni^{2+}(aq) + 2e^- <br />\nrightarrow Ni(s)$, $E° = -0.26 V$
  • $Cd^{2+}(aq) + 2e^- <br />\nrightarrow Cd(s)$, $E° = -0.40 V$
  • $Fe^{2+}(aq) + 2e^- <br />\nrightarrow Fe(s)$, $E° = -0.45 V$
  • $Zn^{2+}(aq) + 2e^- <br />\nrightarrow Zn(s)$, $E° = -0.76 V$
  • $2H<em>2O(l) + 2e^- \nrightarrow H</em>2(g) + 2OH^-(aq)$, $E° = -0.83 V$
  • $Al^{3+}(aq) + 3e^- <br />\nrightarrow Al(s)$, $E° = -1.66 V$
  • $Mg^{2+}(aq) + 2e^- <br />\nrightarrow Mg(s)$, $E° = -2.37 V$
  • $Na^+(aq) + e^- <br />\nrightarrow Na(s)$, $E° = -2.71 V$
  • $Li^+(aq) + e^- <br />\nrightarrow Li(s)$, $E° = -3.04 V$

Cell Potential Calculation

• The total cell potential can be expressed as: $E°<em>{cell} = E°</em>{ox} + E°_{red}$
• Example: $Zn(s) + Cu^{2+}(aq) \nleftrightarrow Zn^{2+}(aq) + Cu(s)$
  • E°_{Zn^{2+}(aq) + 2e^-
    rightleftharpoons Zn (s)} = -0.76V
  • E°_{Cu^{2+}(aq) + 2e^-
    rightleftharpoons Cu (s)} = +0.34V
  • Thus: $E°_{cell} = (0.76 + 0.34)V = 1.10V$

Practice Problem

• What is the standard cell potential for the reaction:
  2 Ag^+(aq) + Cu (s)
  rightleftharpoons 2 Ag (s) + Cu^{2+}(aq)