OXIDATION AND

REDUCTION REACTIONS

EO #10 Students will identify parts of a redox reaction, and use the ½ method to balance redox equations in an acid solution.

OXIDATION AND REDUCTION INFO…..

An oxidation-reduction (redox) reaction is a type of chemical reaction that involves a transfer of electrons between two species. An oxidation-reduction reaction is any chemical reaction in which the oxidation number of a molecule, atom, or ion changes by gaining or losing an electron. Redox reactions are common and vital to some of the basic functions of life, including  photosynthesis, respiration, combustion, and corrosion or rusting.

DEFINITIONS

Oxidation: loss of electrons Reduction: gain of electrons Oxidation States [oxidation numbers]: lets us

keep track of electrons in oxidation-redox reactions by assigning charges to the various atoms in a compound.

Half-Reactions: equations that have electrons as reactants or products. One half-reaction represents a reduction process[electrons are shown on the reactant side – e’s are gained by a reactant] and the other half-reaction represents an oxidation process[electrons are shown on the product side – e’s are lost by a reactant].

WHAT THIS MEANS IS…

The number of electrons lost [on the reactant side] must be equal to the number of electrons gained [on the product side]

EXAMPLE EQUATION

Ce4+[aq] + Sn2+[aq] ---> Ce3+[aq] + Sn4+[aq]

REDUCED REACTION EQUATION

e- + Ce4+[aq] Ce3+

[aq]

OXIDIZED REACTION EQUATION

Sn2+[aq] Sn4+[aq] + 2e-

RULES FOR ASSIGNING OXIDATION STATES

The oxidation state (OS) of an element  corresponds to the number of electrons, e-, that an atom loses, gains, or appears to use when joining with other atoms in compounds. In determining the OS of an atom, there are seven guidelines to follow:

1. The OS of an individual atom is 0.2. The total OS of all atoms in: a neutral species

is 0 and in an ion is equal to the ion charge.3. Group 1 metals have an OS of +1 and Group

2 an OS of +24. The OS of fluorine is -1 in compounds

RULES FOR ASSIGNING OXIDATION STATES

5. Hydrogen generally has an OS of +1 in compounds

6. Oxygen generally has an OS of -2 in compounds

7. In binary metal compounds, Group 17 elements have an OS of -1, Group 16 of -2, and Group 15 of -3.

Note: The sum of the OSs is equal to zero for neutral compounds and equal to the charge for polyatomic ion species.

An atom is oxidized if its oxidation number increases, the reducing agent, and an atom is reduced if its oxidation number decreases, the oxidizing agent. The atom that is oxidized is the reducing agent, and the atom that is reduced is the oxidizing agent. (Note: the oxidizing and reducing agents can be the same element or compound).

ASSIGNING OS’S

Determine the Oxidation States of each element in the following reactions:

Fe(s)+O2(g)→Fe2O3(g)

Fe2+

Ag(s)+H2S→Ag2S(g)+H2(g)

SOLUTIONS

Fe and O2 are free elements; therefore, they each have an OS of 0 according to Rule #1. The product has a total OS equal to 0, and following Rule #6, O has an OS of -2, which means Fe has an OS of +3.

The OS of Fe corresponds to its charge; therefore, the OS is +2.

Ag has an OS of 0, H has an OS of +1 according to Rule #5, S has an OS of -2 according to Rule #7, and hence Ag in Ag2S has an OS of +1.

IDENTIFYING REDUCED AND OXIDIZED ELEMENTS

Determine which element is oxidized and which element is reduced in the following reactions (be sure to include the OS of each): 

Zn + 2H+ → Zn2+ + H2

2Al + 3Cu2+→2Al3+ +3Cu

CO32- + 2H+→ CO2 + H2O

SOLUTIONS

Zn is oxidized (Oxidation number: 0 → +2); H+ is reduced (Oxidation number: +1 → 0)

Al is oxidized (Oxidation number: 0 → +3); Cu2+ is reduced (+2 → 0)

This is not a redox reaction because each element has the same oxidation number in both reactants and products: O= -2, H= +1, C= +4.

(For further discussion, see the article on oxidation numbers).

HALF-REACTION METHOD FOR BALANCING EQUATIONS FOR OXIDATION-REDUCTION REACTIONS IN ACIDIC SOLUTIONS

1. Identify and write the equations for the oxidation and reduction half-reactions.

2. For each half-reaction: A. Balance all of the elements except hydrogen and

oxygen. Balance oxygen using H2O. Balance hydrogen using H+. Balance the charge using electrons.3. If necessary, multiply one or both balanced half-

reactions by an integer to equalize the number of electrons transferred in the 2 half-reactions.

4. Add the half-reactions, and cancel identical species that appear on both sides.

5. Check to be sure the elements and charges balance.