Oxidation and Reduction

loss of one or more electrons

For example:

Mg(s)+1/2O2(g) goes to MgO(s)

The magnesium ions was oxidized because it went from the atom for to the ion form of plus 2.


gain of one of more electrons

for example

Mg(s) + 1/2O2 goes to MgO(s)

The oxygen was reduced because it went from the oxygen gas atom to the oxide ion of -2

oxidation numbers
1. for simple compounds between two elements, each oxidation number is the same as the ion for the element in the compound

2. for covalent compounds, assume the compound is ionic

3. the sum of all the oxidation numbers is zero or equal to the charge of the ion

4. elements in their pure form have an oxidation of zero

5. oxygen= -2
hydrogen= +1

the same element can be both oxidized and reduced in the same reaction

For example: Cl2(g)+H2O(l) goes to HCl(aq)+ HClO (aq)

oxidizing agents
substances that are reduced and can readily gain electrons
reducing agents
these are oxidized and readily give up electrons
reactivity series
if an element is above another element in the reactivity series, then that element is more reactive

it can also determine how spontaneous a reaction is

voltaic cells
spontaneous reaction

simplest form is with a metal in a solution of its ions

salt bridge is used to prevent a charge build up

1.01×10^5 Pa pressure, 298K, 1M aqueous solution

flow of electrons is always from the negative to the positive


oxidation occurs at the negative electrode


reduction occurs at the positive electrode

electrochemical series
the standard hydrogen cell is used as the reference half- cell
electrolytic cell
makes a non-spontaneous reaction occur with an external source of power(electrolyte)

signified by molten of aqueous solution

electrolysis in an aqueous solution
in any aqueous solution, there will also be hydrogen ions and hydroxide ions present as well as the ions of the electrolyte

in some cases the electrodes that do not normally react, will react in this case

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