# 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:

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

## Standards¶

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)

E.g.

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

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

\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 $$>0$$