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Galvanic cell

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A galvanic cell is a device in which chemical energy is converted into electric energy through the transfer of electrons. This is accomplished through a redox reaction.

The reduction half-reaction of the redox reaction occurs at the cathode (RED CAT)

The oxidation half reaction occurs at the anode (AN OX).

To maintain the flow of electrons something is needed to transfer positive charge. This can be accomplished in two ways:

(1) A salt bridge. This allows the transfer of positive charge through the movement of positive ions. In the example below

Copper is the cathode in the cathode half-cell.

Here is where the reduction of Cu2+ ion to Cu metal occurs.

Zn metal is the anode in the anode half-cell.

Here is where the oxidation of Zn to Zn2+ ion occurs.

Other ions are present for charge neutralization, ionic conduction, and completion of the circuit.

This is the basis for most batteries.

You can usually see the + marked on the battery’s cathode, while the other end is the anode.

electrochemical half-cell or Galvanic

As we said above, to maintain the flow of electrons something is needed to transfer positive charge.

Another way to do this is to use a porous disk:

porous disc Galvanic cell electrochemistry

Electrodes are not stable

Electrodes slowly corrode.

Here is an example from a Zinc and Copper Galvanic cell

Here, electrons flow in the wire (above the solution) from Zn to Cu.

Zinc Copper electrochemical cell Galvanic

because Zn is a more active metal than Cu , it tends to lose e-

So the Zn electrode is oxidized: a Zinc ion and 2 free e- are made per original Zn atom

This Zn ion breaks apart from the electrode and floats off into the solution

 

Zinc electrodes corrode electrochemistry Galvanic

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Pure metals corrode because they aren’t stable

Why do the electrodes corrode? Well the real question is “Why don’t all metals corrode”?

Look around you – what metals don’t corrode (rust)? Only gold, platinum and a few others. Every other metal does.

Look for pure metals… good luck – you won’t find any.  They’re all already chemically bound to other substances.  Instead of finding copper, we find copper ore. Same for iron, or anything else.

How do we get pure metals, then? We need to expend a lot of energy to separate the metal that we want from the other atoms.

Here’s the physics explanation of why this is so. It has been excerpted and adapted from Corrosion of metals  (author unknown.)

Pure metals contain more bound energy, representing a higher energy state than that found in the nature as sulphides or oxides.

Corrosion and energy needed to get pure metal

All material in the universe strives to return to its lowest energy state.

Same for metals. They tend to revert to their lowest energy state which they had as sulphides or oxides. They revert to a low energy level by corrosion.

For batteries, we see electrochemical corrosion. Takes place in an aqueous environment.

All metals in dry air are covered by a very thin layer of oxide, about 100Å (10-2µm) thick. This layer is built up by chemical corrosion with the oxygen in the air. At very high temperatures, the reaction with the oxygen in the air can continue without restraint and the metal will rapidly be transformed into an oxide.

metals oxide corrosion.gif

At room temperature the reaction stops when the layer is thin. These thin layers of oxide can protect the metal against continued attack, e.g. in a water solution. In actual fact, it is these layers of oxide and/or products of corrosion formed on the surface of the metal that protect the metal from continued attack to a far greater extent that the corrosion resistance of the metal itself.

These layers of oxide may be more or less durable in water, for instance. We know that plain carbon steel corrodes faster in water than stainless steel. The difference depends on the composition and the penetrability of their respectively oxide layers. The following description of the corrosion phenomenon will only deal with electrochemical corrosion, i.e. wet corrosion.

Corrosion cells

How do metals corrode in liquids? Let us illustrate this, using a corrosion phenomenon called bimetal corrosion or galvanic corrosion. The bimetal corrosion cell can e.g. consist of a steel plate and a copper plate in electrical contact with one another and immersed in an aqueous solution (electrolyte).

The electrolyte contains dissolved oxygen from the air and dissolved salt. If a lamp is connected between the steel plate and the copper plate, it will light up. This indicates that current is flowing between the metal plates. The copper will be the positive electrode and the steel will be the negative electrode.

Galvanic cell start

The driving force of the current is the difference in electrical potential between the copper and the steel. The circuit must be closed and current will consequently flow in the liquid (electrolyte) from the steel plate to the copper plate. The flow of current takes place by the positively charged iron atoms (iron ions) leaving the steel plate and the steel plate corrodes.

The corroding metal surface is called the anode. Oxygen and water are consumed at the surface of the copper plate and hydroxyl ions (OH-), which are negatively charged, are formed. The negative hydroxyl ions “neutralize” the positively charged iron atoms. The iron and hydroxyl ions form ferrous hydroxide (rust).

Galvanic cell corrosion corroding

In the corrosion cell described above, the copper metal is called the cathode. Both metal plates are referred to as electrodes and the definition of the anode and the cathode are given below.

Anode: Electrode from which positive current flows into an electrolyte.
Cathode: Electrode through which positive electric current leaves an electrolyte.

When positive iron atoms go into solution from the steel plate, electrons remain in the metal and are transported in the opposite direction, towards the positive current.

Galvanic cell corrosion rust

Videos

https://www.youtube.com/watch?v=C26pH8kC_Wk

Learning Standards

HS-PS1-10(MA). Use an oxidation-reduction reaction model to predict products of reactions given the reactants, and to communicate the reaction models using a representation that shows electron transfer (redox). Use oxidation numbers to account for how electrons are redistributed in redox processes used in devices that generate electricity or systems that prevent corrosion.*

 

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