ELC_Galvanic Cells Lesson

Galvanic Cells

Have you ever felt a "shock" to your teeth when a piece of tin foil or a metal spoon touches a sliver or mercury filling in your mouth? That shock was actually a small electrical current. The energy that produced the electric current was created by a spontaneous redox reaction between the two different metals. It was like you created a small battery in your mouth!

Galvanic Cell diagram with the following labeled: anode
salt bridge
cathode
oxidation half-cell
reduction half-cell The device that produces electricity in this manner is called a galvanic cell (sometimes called a voltaic cell). The key to creating the electric current, or flow of electrons, is to separate the oxidation and reduction half-reactions, connecting them by a wire. This way the current flows through the wire which can be attached to a device like a radio or a clock that will use that electricity. This is how a battery works.

A galvanic cell is made of two containers, each called a half-cell. The oxidation reaction occurs in one half-cell and the reduction reaction occurs in the other half-cell. Electrons released during oxidation flow through a wire or other external circuit and are then accepted by the reduction half-cell. The separation of the half-reactions of the spontaneous reaction is what causes the electric current to flow.

Red Cat image Each half-cell holds a solution and an electrode. An electrode is a solid electric conductor through which the current flows. The two electrodes are called the anode and the cathode. Oxidation always occurs at the anode, and reduction always occurs at the cathode. An easy way to remember this is with the mnemonic device red cat. This will help you remember that red uction occurs at the cat hode, and, by default, oxidation occurs at the anode. We know that oxidation is the loss of electrons, meaning that electrons are products. So, the anode, where oxidation occurs, is negative. The cathode is positive. Each electrode is immersed in a solution containing a dissolved salt of the corresponding metal.

Finally, the half-cells are connected by a salt bridge. The salt bridge connects the two half-cells, completing the circuit. It contains a salt, like KCl or NaNO₃, that won't interfere with the reaction taking place. It is necessary to maintain the overall neutral charge of each half-cell. As electrons are transferred from the oxidation half-cell to the reduction half-cell, a negative charge builds in the reduction half-cell and a positive charge in the oxidation half-cell. The salt bridge prevents this build of charge by supplying ions to keep each solution neutral.

Galvanic Cell Notation

Instead of having to draw the galvanic cell each time, a shorthand method of describing them was devised by chemists.

The oxidation process is listed first, starting with the anode solid then the ion in solution for the oxidation half reaction. A vertical line, |, is used to separate the anode from the ions.
Next, a double vertical line, ||, is drawn to represent the salt bridge.
The reduction is listed last. This time the ion is listed first, followed by |, and finally the cathode solid.

Oxidation
Reduction
anode solid anode ion in solution || cathode ion in solution | cathode solid

The cell in the slideshow above would be described as:

$LaTeX: Zn_(s)\vert Zn^{2+}_{(aq)} \vert\vert Cu^{2+}_{(aq)} \vert Cu_{(s)}$ $Zn_(s)\vert Zn^{2+}_{(aq)} \vert\vert Cu^{2+}_{(aq)} \vert Cu_{(s)}$

Galvanic Cell Example diagram In some cases, the substance oxidized or the product of reduction is not a solid. In those cases, platinum serves as an inert electrode. It is just a physical site for the oxidation or reduction to take place. In the following example of a galvanic cell, the cathode is platinum. It is not a part of the actual redox reaction. Here, the anode is zinc in a solution containing zinc ions. The platinum cathode is in a solution containing two different iron ions, Fe²⁺ and Fe³⁺. The oxidation reaction is Zn → Zn²⁺ + 2e^-. We know that reduction is a gain of electrons. So, we must write the reduction half-reaction to show electrons as a reactant. Fe³⁺ + 1e^- → Fe²⁺

The number of electrons is balanced, and the overall reaction is:

Zn_(s) + 2Fe³⁺ _(aq) → Zn²⁺ _(aq) + 2Fe²⁺ _(aq)

In the cell notation, the substance written on each end must be the solid electrodes. So, both iron ions will be written to the right of the salt bridge symbol.

$LaTeX: Zn_{(s)} \vert Zn^{2+}_{(aq)} \vert \vert Fe^{2+}_{(aq)}, Fe^{3+}_{(aq)} \vert Pt$ $Zn_{(s)} \vert Zn^{2+}_{(aq)} \vert \vert Fe^{2+}_{(aq)}, Fe^{3+}_{(aq)} \vert Pt$

Remember to work on the module practice problems as you complete each section of content.

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