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Chemistry » Chemistry 111: Electrochemical Reactions » Galvanic And Electrolytic Cells

# Galvanic Cells

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## Galvanic cells

A galvanic cell (which is also sometimes referred to as a voltaic or wet cell) consists of two half-cells, which convert chemical potential energy into electrical potential energy.

### Definition: Galvanic cell

A galvanic cell is an electrochemical cell which converts chemical potential energy to electrical potential energy through a spontaneous chemical reaction.

In a galvanic cell there are two half-cells. Each half-cell contains an electrode in an electrolyte. The separation is necessary to prevent direct chemical contact of the oxidation and reduction reactions, creating a potential difference. The electrons released in the oxidation reaction travel through an external circuit (and do work) before being used by the reduction reaction.

A sketch of a galvanic cell.

#### Tip:

Remember that cells are not two-dimensional, although when asked to sketch a cell you should draw it as shown in the figure above.

In a galvanic cell (for example the cell shown above):

#### Tip:

The electrons released in the oxidation of $$\text{X}$$ remain on the anode, $$\text{X}^{+}$$ moves into solution.

The $$\text{Y}^{+}$$ ions are being reduced at the cathode and forming solid $$\text{Y}$$.

• The metal at the anode is $$\color{blue}{\textbf{X}}$$. $$\color{blue}{\textbf{Ox}}$$idation is loss of electrons at the $$\color{blue}{\textbf{An}}$$ode.

• The $$\color{blue}{\textbf{anode half-reaction}}$$ is $$\color{blue}{{\textbf{X(s)}} \to {\textbf{X}}^{+}{\textbf{(aq) + e}}^{-}}$$

• This half-reaction occurs in the half-cell containing the X(s) anode and the $$\text{X}^{+}(\text{aq})$$ electrolyte solution.

• The electrons released in the $$\color{blue}{\textbf{oxidation}}$$ of the metal remain on the $$\color{blue}{\textbf{anode}}$$, while the metal cations formed move into solution.

• The metal at the cathode is Y. $$\color{red}{\textbf{Red}}$$uction is gain of electrons at the $$\color{red}{\textbf{Cat}}$$hode.

• The $$\color{red}{\textbf{cathode half-reaction}}$$ is $$\color{red}{{\textbf{Y}}^{+}{\textbf{(aq) + e}}^{-} \to {\textbf{Y(s)}}}$$

• This half-reaction occurs in the half-cell containing the Y(s) cathode and the $$\text{Y}^{+}(\text{aq})$$ electrolyte solution.

• At the $$\color{red}{\textbf{cathode}}$$ metal ions in the solution are being $$\color{red}{\textbf{reduced}}$$ (accepting electrons) and deposited on the electrode.

• There are more electrons at the anode than at the cathode.

• Electrons will flow from areas of high concentration to areas of low concentration, therefore the electrons move $$\color{blue}{\textbf{from the anode}}$$, through the external circuit, $$\color{red}{\textbf{to the}}$$ $$\color{red}{\textbf{cathode}}$$

• Conventional current is measured as a flow of positive charge and so is in the opposite direction (from the cathode to the anode)

• The overall reaction is: $$\text{X}(\text{s}) + \text{Y}^{+}(\text{aq})$$ $$\to$$ $$\text{X}^{+}(\text{aq}) + \text{Y}(\text{s})$$.

• To represent this reaction using standard cell notation we write the following:

$$\color{blue}{\textbf{X(s)}}$$$$|$$$$\color{blue}{\textbf{X}^{+}{\textbf{(aq)}}}$$$$||$$$$\color{red}{\textbf{Y}^{+}{\textbf{(aq)}}}$$$$|$$$$\color{red}{\textbf{Y(s)}}$$

By convention:

• The $$\color{blue}{\textbf{anode}}$$ is always written on the $$\color{blue}{\textbf{left}}$$.

• The $$\color{red}{\textbf{cathode}}$$ is always written on the $$\color{red}{\textbf{right}}$$.

• The anode and cathode half-cells are divided by $$||$$ representing the salt bridge.

• The different phases within each half-cell (solid (s) and aqueous (aq) here) are separated by $$|$$.

• The electrodes in each half-cell are connected through a wire in the external circuit. There is also a salt bridge between the individual half-cells.

A galvanic cell uses the reactions that take place at at the two electrodes to produce electrical energy, i.e. the reaction occurs without the need to add energy.

#### Fact:

A spontaneous reaction is one that will occur without the need for external energy. Refer to the subsection on spontaneity for more information.

The zinc-copper reaction you performed in another lesson can be modified to make a galvanic cell. Bars of zinc and copper are used as electrodes, with zinc(II) sulfate and copper(II) sulfate solutions as the electrolytes.

In the galvanic cell experiment, make sure that the sodium chloride paste is highly concentrated and fills the U-tube for the best results.

## Optional Experiment: A galvanic cell

### Aim

To investigate the reactions that take place in a galvanic cell.

### Apparatus

• Zinc plate, copper plate, zinc(II) sulfate ($$\text{ZnSO}_{4}$$) solution ($$\text{1}$$ $$\text{mol.dm^{-3}}$$), copper(II) sulfate ($$\text{CuSO}_{4}$$) solution ($$\text{1}$$ $$\text{mol.dm^{-3}}$$), $$\text{NaCl}$$ paste

• Measuring balance, two $$\text{250}$$ $$\text{ml}$$ beakers, U-tube, cotton wool, zero-centered ammeter, connecting wire.

• Cleaning ethanol, ether (if available)

### Method

1. Weigh the copper and zinc plates and record their mass.

2. Pour $$\text{200}$$ $$\text{ml}$$ of the zinc sulfate solution into a beaker and place the zinc plate in the beaker.

3. Pour $$\text{200}$$ $$\text{ml}$$ of the copper(II) sulfate solution into the second beaker and place the copper plate in the beaker.

4. Fill the U-tube with the $$\text{NaCl}$$ paste and seal the ends of the tubes with the cotton wool (making a salt-bridge). The cotton will help stop the paste from dissolving in the electrolyte.

5. Connect the zinc and copper plates to the zero-centered ammeter and observe the ammeter.

6. Place the U-tube so that one end is in the copper(II) sulfate solution and the other end is in the zinc sulfate solution. Observe the ammeter.

7. Take the ammeter away and connect the copper and zinc plates to each other directly using copper wire. Leave to stand for about one day.

8. After a day, remove the two plates and rinse them: first with distilled water, then with alcohol, and finally with ether (if available). Dry the plates using a hair dryer.

9. Weigh the zinc and copper plates and record their mass.

### Note

A voltmeter can also be used in place of the zero-centered ammeter. A zero-centered voltmeter will measure the potential difference across the cell (not the flow of electrons), while an ammeter will measure the current.

### Discussion

• Did the ammeter record a reading before the salt-bridge was placed in the solutions?

• Did the ammeter record a reading after the salt-bridge was placed in the solutions? If yes, in what direction does the current flow?

• Fill in the table below:

 Plate Initial mass Final mass Zinc Copper
• How did the mass of the zinc and copper plates change?

• Based on what you know of oxidation and reduction, why did those mass changes take place?

• Which electrode is the anode and which is the cathode?

### Results

During the experiment, you should have noticed the following:

• When the salt bridge was absent, there was no reading on the ammeter.

• When the salt bridge was connected, a reading was recorded on the ammeter.

• The direction of electron flow is from the zinc plate towards the copper plate, meaning that conventional current flow is from the copper plate towards the zinc plate.

• After the plates had been connected directly to each other and left for a day, there was a change in their mass. The mass of the zinc plate decreased, while the mass of the copper plate increased.

• $$\color{blue}{\textbf{O}}$$xidation $$\color{blue}{\textbf{i}}$$s $$\color{blue}{\textbf{l}}$$oss of electrons, $$\color{red}{\textbf{R}}$$eduction $$\color{red}{\textbf{i}}$$s $$\color{red}{\textbf{G}}$$ain of electrons.

The zinc electrode lost mass. This implies that solid Zn metal atoms become ions and move into the electrolyte solution: $$\text{Zn}(\text{s})$$ $$\to$$ $$\text{Zn}^{2+}(\text{aq}) + 2\text{e}^{-}$$. Oxidation occurs at the zinc electrode.

The copper electrode gained mass. This implies that the Cu metal ions in the electrolyte solution become metal atoms and deposit on the electrode: $$\text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-}$$ $$\to$$ $$\text{Cu}(\text{s})$$. Reduction occurs at the copper electrode.

• $$\color{blue}{\textbf{Ox}}$$idation is loss of electrons at the $$\color{blue}{\textbf{an}}$$ode. Oxidation occurs at the zinc electrode, therefore the zinc plate is the anode.

$$\color{red}{\textbf{Red}}$$uction is gain of electrons at the $$\color{red}{\textbf{cat}}$$hode. Reduction occurs at the copper electrode, therefore the copper plate is the cathode.

• When $$\text{Zn}(\text{s})$$ $$\to$$ $$\text{Zn}^{2+}(\text{aq}) + 2\text{e}^{-}$$ the electrons are deposited on the anode, which becomes negatively charged.

When $$\text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-}$$ $$\to$$ $$\text{Cu}(\text{s})$$ the electrons are taken from the cathode, which becomes positively charged.

### Conclusions

When a zinc(II) sulfate solution containing a zinc plate is connected by a salt bridge to a copper(II) sulfate solution containing a copper plate, reactions occur in both solutions. The decrease in mass of the zinc plate suggests that the zinc metal electrode has been oxidised to form $$\text{Zn}^{2+}$$ ions in solution. The increase in mass of the copper plate suggests that reduction of $$\text{Cu}^{2+}$$ ions has occurred here to produce more copper metal.

The important thing to notice in this experiment is that:

• the chemical reactions that take place at the two electrodes cause an electric current to flow through the external circuit

• the overall reaction must be a spontaneous redox reaction

• chemical energy is converted to electrical energy

• the zinc-copper cell is one example of a galvanic cell

#### Fact:

It was the Italian physician and anatomist Luigi Galvani who marked the birth of electrochemistry by making a link between chemical reactions and electricity. In 1780, Galvani discovered that when two different metals (copper and zinc for example) were connected to each other and then both touched to different parts of a nerve of a frog leg at the same time, they made the leg contract. He called this ‘animal electricity’.

In the zinc-copper cell, the copper and zinc plates are the electrodes. The salt bridge plays a very important role in a galvanic cell:

• An electrolyte solution consists of metal cations and spectator anions.

$$\text{NaCl}(\text{aq})$$ is $$\text{Na}^{+}(\text{aq})$$ and $$\text{Cl}^{-}(\text{aq})$$ in the paste.

• There is a build up of positive charge in the $$\color{blue}{\textbf{anode half-cell compartment}}$$ as solid metal is $$\color{blue}{\textbf{oxidised}}$$ and the positive ions move into solution. So there are more postive metal ions in the electrolyte than negative ions.

$$\color{blue}{\textbf{Zn(s)} \to \textbf{Zn}^{2+}\textbf{(aq)}}$$, while the number of $$\text{SO}_{4}^{2-}$$ ions remains the same.

• To balance the charge, negative ions from the salt bridge move into the anode half-cell compartment.

$$\text{Cl}^{-}$$ ions move from the salt bridge to the anode half-cell compartment.

• There is a decrease in positive charge in the $$\color{red}{\textbf{cathode half-cell compartment}}$$ as metal ions are $$\color{red}{\textbf{reduced}}$$ and form solid metal. So there are more negative ions in the electrolyte than positive metal ions.

$$\color{red}{\textbf{Cu}^{2+}\text{(aq)} \to \textbf{Cu(s)}}$$, while the number of $$\text{SO}_{4}^{2-}$$ ions remains the same.

• To balance the charge, positive ions from the salt bridge move into the cathode half-cell compartment.

$$\text{Na}^{+}$$ ions move from the salt bridge to the cathode half-cell compartment.

The salt bridge acts as a transfer medium that allows ions to flow through without allowing the different solutions to mix and react directly. It allows a balancing of the charges in the electrolyte solutions, and allows the reactions in the cell to continue.

Without the salt bridge, the flow of electrons in the outer circuit stops completely. This is because the salt bridge is needed to complete the circuit.