Fundamentals of Electronics III: Batteries

Introduction

As was said in the first lesson, Electronics is firstly a study of chemistry, which is a study of applied physics. But as in all applied science, some things must be assumed and simplified in order to get anywhere fast. For now, we will be ignoring the physical underpinnings of chemistry that do not relate directly to electrical charges.

That being said, batteries deal with some fairly involved chemistry. We are just going to touch some basic concepts,and also assume that if you have some interest you will find other resources for an in depth analysis or more complex examples. For example, the Chemistry section is a good start.

As a consequence of keeping the visual aids unplagarized, I sacrificed the quality. As you can see I made them in paint again, I appoligize if they are difficult to read.

Setup of a cell

The cell that we will be looking at today is an example of a galvanic cell. That means that it has two components, One is called a cathode, the other an anode. They represent the positive and negative terminals of the battery, respectively.

cathode: The + connection of a battery, out of which (conventional) current flows.
anode: The - connection of a battery, which (conventional) current flows into.

The battery we continue to describe below works through a chemical process called a Redox reaction. We can force this reaction to occur based on what elements we use and how we set them up to react.

Both will have metalic ions in sulfate solution, and a some solid material of that same metal. Ions are molecules that have charge, either by a surplus or deficiency of electrons to balance them. The metal ions will be spontaneously undergoing what is known as oxidation and reduction reaction reactions. Reduction reactions are chemical reactions that involve an atom gaining electrons, in this case to balance the charge of ions to neutral and therefore make them insoluble, so they become part of the solid metal. Oxidation reactions are the opposite, and these reactions involve a loss of electrons to form an ion, which then leaves the metal to the solution. When isolated, these reactions occur in equilibrium; they occur toward a balance of electrical neutrality, then occur at identical rates. In other words, a few molecules will leave the solid to solution to balance the negative solution, then no change will occur.

Two electrons attach to a Zn ion and it leaves the solution

Two electrons leave the Zn atom and it becomes an ion in solution

ion: A molecule that has attained an overall charge by loss or gain of electrons
Reduction: A reaction in which a molecule gains electrons, thus reducing its positive charge
Oxidation: A reaction in which a molecule loses electrons, thus increasing its positive charge
equilibrium: Any reaction in which two opposing reactions are occuring at equal rates, where on average no change is observed.

Activation of a cell

What we want to do is setup a situation in which the reduction and oxidation do not occur at identical rates, and therefore the system is not at equilibrium. By connecting the two nodes of our battery, we can use differences in electronegitivity of the two chosen metals, and bias the situation so one will undergo reduction while the other undergoes oxidation. This is known as a Redox reaction.

Electronegitivity: A tendency of an atom to attract electrons
Redox: A reaction in which a combonation of an oxidation and a reduction results in a transfer of electrons between two molecules.

The standard, classic textbook example uses Copper (elment Cu) and Zinc (Zn) as the two metals, simply because of the ease of expressing the redox reaction. Both form ions of 2+ charge when in solution, so a single redox reaction could be expressed directly as the transfer of two electrons.

If the solid material in each node will be connected by a conductor, such as a copper wire, The equilibrium would be offset and a different type of reaction would occur. If the element chosen in one cell were more electronegative than the other, reduction would be favored in one cell, and oxidation in the other. The corresponding loss and gain half reactions would form one whole reaction, and it is classified as a redox reaction. This would result in electron flow down the copper wire to favor one direction, which is exactly what we want. However at this point, the side losing the electrons would quickly become electrically positive, and oppose the new flow down the wire until equilibrium was once again achieved.

With the anode and cathode connected, the cathode quickly becomes electrically negitive due to the acceptance of electrons. This opposes the redox reaction and equilibrium is achieved.

We don't want this, because we want to be able to harness the power of the electrons moving until the redox reaction is completely done in that direction. We need to balance it out electrically, and we do that with a salt bridge. A salt bridge is a porous barrier that allows the movement of ions from one cell to the other to balance out the excess charge from the redox reaction, so it can continue.

Salt Bridge: A porus barrier, filled with salt solution, that allows the flow of electrical charge to offset the loss of electrons occuring in the anode. The salt is typically pottasium (element Ka). It allows the reaction to continue.

The reaction will continue untill all the ions in solution in the cathode have moved onto the solid, or all the solid metal in the anode has moved into solution. Then the battery is used up, until those ingredients are replaced.

Use of a cell

We could place a load (such as a light bulb) in between the connection of the two nodes and use the movement of electrons to our advantage. I'm sure you would like more detail, but we will get into that later.

Moving on….

And we're done with the battery example. This is an example as you would learn in introductory chemistry at a high school level. Most modern batteries do not use this setup, but do use the underlying concepts of redox reactions. Other examples of batteries are button batteries (used in watches), lead acid batteries (used in your car), or rechargable AA batteries.

Next: AC generator (alternator)

Previous lesson: Ohm's Law