A favorite demonstration of mine is sticking a couple of zinc and copper electrodes in potatoes and magically running a clock. When I allow the students to place the electrodes into various pieces of produce and they create a charge they can read on a voltmeter, they are even more amazed.
Surely the science behind such magical occurrences should be horribly complex, but it is not. In fact it is really quite simple and occurs around us every day:  rusting iron, combustion, cooking, turning on a flashlight. Even breathing.These reactions involve the transfer of electrons from one element (or species) to another.  Giving up or donating an electron(s) is called oxidation and gaining an electron(s) is reduction. Oxidation and reduction always occur together and the reactions are called redox reactions.
Organic or biological chemists often refer to oxidation as the gaining of oxygen or loss of hydrogen while reduction is either the loss of oxygen or the gaining of hydrogen. While the terms definitely apply, not all oxidation reactions involve the element oxygen but it is how the term originated.
Every atom has the potential to participate in redox reactions, the transfer of electrons, but some are more willing or able than others. Every atom’s electrons have three possibilities:

(1)   do nothing,

(2)    share them between atoms, or

(3)   participate in transferring them, either as the donor or the receiver.

The first case deals with very stable atoms, the second is covalent bonds and the third is the world of oxidation and reduction.
When discussing the loss or gain of electrons, an atom’s electronegativity, the ability to draw electrons to itself, must be taken into account.  A bookkeeping system, called oxidation numbers or states, provides a way to keep track of the transfer of electrons.
Every atom has an oxidation number, even if it never changes.  Generally, elements in their natural state have an oxidation number of zero.  Let’s take a look at oxygen and hydrogen during the production of water.

 2 H₂ (g) + O₂ (g) → 2 H₂O (g)

Both oxygen and hydrogen occur as diatomic gases in nature and therefore have oxidation numbers of zero.  Once combined however, the neutral water molecule contains both.
The question is, who took the electrons and who donated them? A look at the periodic table says that oxygen is more electronegative so it is more likely to take the electrons and in fact, prefers to take two, hence the reason it bonds with two hydrogens.
In order to “see” the flow of electrons, half equations are written, one for the oxidized species and one for the reduced species.  These are only for convenience though, electrons do not pop out of elements and float around.
H₂ (g)  → 2 H⁺ (g) + 2 e^–            2(H₂ (g)  → 2 H⁺ (g) + 2 e^–) = 2 H₂ (g)  → 4 H⁺ (g) + 4 e^–
O₂ (g) + 4 e^– → 2 O^-2
Because oxygen needs four electrons, double the amount of hydrogen is used. The formation of water can be visualized as 2 H⁺ + O^-2 → H₂O and done twice because of the amount of the starting hydrogen and oxygen gases.
The oxidation number of each hydrogen, is therefore +1 by being oxidized and each oxygen is -2 by being reduced. A summary of common oxidation numbers can be seen in an interactive multimedia clip in section 8.1 in the Physical Chemistry book (see screengrab above).
Combustion, respiration, rusting and bleaching are all common examples of this type of reaction. The textbook is primarily concerned with redox reactions that can provide electrical energy, such as in batteries, that I will discuss in the future.