What is an oxidizer?
Oxidation: losing electron(s), Oxidizer / oxidizing agent: a chemical that can oxidize another reagent. Reduction: gaining electron(s) - think of it as reduction of charge! Reducing agent: A chemical that can reduce another reagent.

And oxidation can never occur without a reduction, meaning if you use an oxidizing agent to oxidize a substance, when the oxidizing agent is in turn being reduced itself. It gains the electrons of the oxidized substance.

To describe how well a substance functions as an oxidizer, we use reduction potentials, measured in volts! A substance with a large, positive reduction potential is easy to reduce, which in turn means it is a good oxidizing agent. Similarly, if a substance has a large, negative reduction potential, it means that it is hard to reduce this substance. This is because it's good at reducing other substances itself - it's a reducing agent.

To figure out if a redox reaction is spontaneous (if it can potentially proceed on its own), we use standard reduction potentials, $E_0$. This is for concentrations of 1 M and a temperature of 25 degrees celcius. You look these up in your textbook or online.

An example: Is the following reaction spontaneous?
$$\ce{Cu^{2+}(aq) + 2Ag(s) -&gt; Cu(s) + 2Ag+(aq)}$$

To find out we first find the standard reduction potentials for each participating species:

$\ce{Cu^{2+}(aq) + 2e- -&gt; Cu(s)}, E_0=+0.34V$

$\ce{Ag^{+}(aq) + e- -&gt; Ag(s)}, E_0=+0.80V$

This describes how much they want to be reduced (GAIN electrons). The silver ion wants it the most! However, notice that in our reaction we're asking for it to LOSE electrons (be oxidized). THEREFORE: No, the reaction is not spontaneous under the aforementioned standard conditions. However, the REVERSE reaction is spontaneous. To calculate the potential for our reaction, we simply say:

$$E_0 = (+0.34V) - (+0.80V) = -0.46V$$

The reduction potential for the substance that we want to gain electrons (reduced), minus the potential for the substance we want to lose electrons (be oxidized). The fact that the result is negative, is what tells us the reaction will not proceed - and that it will in fact proceed in the OPPOSITE direction.

Take notice that these reduction potentials are not given the unit Volt for no reason! They are true electric potentials. You can think of electric potentials as "electron pressures". The pole with the highest electron pressure will be (-), because electrons are negatively charged, and the other will be (+) relative to it.

Remember that the reduction potential (the willingness to take electrons) for the silver ion is higher than that of the copper ion. Silver wants to take up electrons more than copper, meaning that electrons will travel from the copper to the silver. you could just as well say that the copper wants to get rid of its electrons more than the silver - copper has the higher electron pressure.

What about electronegativity?
Electronegativity does indeed correlate somewhat with reduction potentials, at least for simple reactions involving pure elements. And you can see that Ag does indeed have a higher electronegativity than Cu, which makes sense. Electronegativity doesn't take into account the oxidation states of different species etc., and quickly it becomes a better idea to just leave electronegativity out of the picture when dealing with redox reactions.