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Galvanic Cells - Standard Conditions

In galvanic cells, a redox reduction is used to create an electrical current.

\[ \ce{Zn_{(s)} + Cu_{(aq)}^{2+} <=> Cu_{(S)} + Zn_{(aq)}^{2+}} \]

Here, Zn is being reduced in an oxidation reaction and \(Cu^{2+}\) is being oxidised in a reduction reaction.

  • Electrons are moving from the zinc to the copper.

We can make these electrons flow through a wire to produce a current


This process can be written as:

\[ \ce{Zn_{(s)} | Zn^{2+} (1 M) || Cu^{2+} (1 M) | Cu_{(s)}} \]



The measured potential energy of the reaction (E°_cell) is the potential energy of the two half reaction combined. This potential is based on comparison to the standard of:

\[ \ce{2e− + 2H+ (1M) - >H2 (1 atm)} \]

Calculating Potential

The equation for the potential calculation is:

\[ E^\circ_{cell}=E^\circ_{reduction}−E^\circ_{oxidation} \]

These \(E^\circ\) values can be found in refernce tables and refer to the reduction potential of the half reaction at standard conitions (25°C, 1 atm, 1 M)

  • The formula above accounts for this by inverting (−) the oxidation reaction.
  • Lower \(E^\circ\) values are more likely to reduce (be involved in oxidation)


\(\hskip{1cm} \ce{Cu^{2+}} + 2e− \ce{-> Cu_{(s)}} \hskip{1cm} E^\circ = +0.34\:V \hskip{1cm} \text{Reduction}\)

\(\hskip{1cm}\ce{Zn^{2+}} + \ce{2e− <-Zn_{(s)}} \hskip{1cm} E^\circ = −0.76 \: V \hskip{1cm} \text{Oxidation}\)

\[ \begin{align} E^\circ_{cell}&=E^\circ_{reduction}−E^\circ_{oxidation}\\ E^\circ_{cell}&=+0.34−(−0.76)\\ E^\circ_{cell}&=+1.10V \end{align} \]

Spontaneous Reactions

For a spontaneous reactions to occur you need three things:

  1. An oxidant (to undergo reduction and give away electrons)
  2. A reductant (to undergo oxidation and take electrons)
  3. An \(E^\circ_{cell}\) value \(>0\)