1.2.1 Standard Electrode Potential

1. Metal atoms have greater tendency to lose electron, hence they are good reducing agents.
2. Each electrode has a different electrode potential.
3. Electrode potential is the potential difference produced when an equilibrium is established between metal \(M\) and the aqueous solution containing metal \(M ^{ nt }\) ions in a half-cell.
4. The standard electrode potential is also known as standard reduction potential. It is measured using the standard hydrogen electrode as the reference electrode.
5. Diagram below shows a chemical cell that is used to obtain the value of the standard electrode potential, \(E ^{\circ}\) of zinc.

   

   Standard conditions for the cells are:
   1. Concentration of ions in an aqueous solution is \(1.0 mol dm ^{-3}\)
   2. Temperature at \(25^{\circ} C\) or \(298 K\)
   3. Gas pressure at \(1 atm\) or \(101 kPa\)
   4. Platinum is used as an inert electrode

Oxidising Agents and Reducing Agents Based on the Value of Standard Electrode Potential

1. Table below shows the various electrode potentials are arranged in increasing order of their reduction potential. This series is called standard electrode potential series.

\[ \begin{array}{|l|r|} \hline {\text { Half-cell equation }} & E ^{\circ}( V ) \\ \hline K ^{+}( aq )+ e ^{-} \rightleftharpoons K ( s ) & -2.92 \\ \hline Ca ^{2+}( aq )+2 e ^{-} \rightleftharpoons Ca ( s ) & -2.87 \\ \hline Na ^{+}( aq )+ e ^{-} \rightleftharpoons Na ( s ) & -2.71 \\ \hline Mg ^{2+}( aq )+2 e ^{-} \rightleftharpoons Mg ( s ) & -2.38 \\ \hline Al ^{3+}( aq )+3 e ^{-} \rightleftharpoons Al ( s ) & -1.66 \\ \hline Zn ^{2+}( aq )+2 e ^{-} \rightleftharpoons Zn ( s ) & -0.76 \\ \hline Fe ^{2+}( aq )+2 e ^{-} \rightleftharpoons Fe ( s ) & -0.44 \\ \hline Sn ^{2+}( aq )+2 e ^{-} \rightleftharpoons Sn ( s ) & -0.14 \\ \hline Pb ^{2+}( aq )+2 e ^{-} \rightleftharpoons Pb ( s ) & -0.13 \\ \hline 2 H ^{+}( aq )+2 e ^{-} \rightleftharpoons H ( g ) & 0.00 \\ \hline Cu ^{2+}( aq )+2 e ^{-} \rightleftharpoons Cu ( s ) & +0.34 \\ \hline O _2( g )+2 H _2 O ^2( l )+4 e ^{-} & +0.40 \\ \hline O _2( s )+2 e ^{-} \rightleftharpoons 2 l ^{-}( aq ) & +0.54 \\ \hline Ag ^{+}( aq )+ e ^{-} \rightleftharpoons Ag ^{-}( s ) & +0.80 \\ \hline Br _2( l )+2 e ^{-} \rightleftharpoons 2 Br ^{-}( aq ) & +1.07 \\ \hline Cl _2( g )+2 e ^{-} \rightleftharpoons 2 Cl ^{-}( aq ) & +1.36 \\ \hline S _2 O _8{ }^{2-}( aq )+2 e ^{-} \rightleftharpoons & +2.01 \\ 2 SO _4{ }^{2-}( aq ) & + \\ \hline \end{array} \]

2. \(E ^{\circ}\) value is a measure of the tendency of a substance to accept or donate electrons.
3. \(E ^{\circ}\) value is used to predict the:
   1. atom or ion that will undergo oxidation or reduction.
   2. strength of oxidising agent or reducing agent.
   3. products of electrolysis.
4. The more positive the value of standard electrode potential, \(E ^{\circ}\), the easier for the atom or ion to undergo reduction. Example: The value of standard electrode potential, \(E ^{\circ}\) for chlorine, \(Cl _2\) is positive, hence chlorine undergoes reduction easier. Therefore, chlorine is a good oxidising agent.
5. The more negative the value of standard electrode potential, \(E ^{\circ}\), the easier for the atom or ion to undergo oxidation. Example: The value of standard electrode potential, \(E^{\circ}\) for zinc, \(Zn\) is negative, hence zinc undergoes oxidation easier. Therefore, zinc is a good reducing agent.