Oxidation Numbers and Redox Reactions

Section 7.2

Oxidation Numbers

• Oxidation number: the number of electrons that must be added to or removed from an atom in a combined state to convert the atom into the elemental form

• AKA oxidation state• Do not have an exact physical meaning

Usage

• Useful for naming compounds, writing formulas, and balancing chemical equations

• There are specific rules for assigning oxidation numbers to a species

• Shared electrons are assumed to belong to the more electronegative atom in each bond

Rules

1. Atoms in a pure element have an oxidation number of 0

2. F has an oxidation number of -13. O has an oxidation number of -2, except

in peroxides (-1) or bonded with F (+2)4. H is +1 when in a compound where the

other element is more electronegative than it is

Rules Continued

5. H is -1 when it is bonded with metals6. The more-electronegative element in a

binary molecular compound is assigned the number equal to the negative charge it would have as an anion, the other gets the positive charge it would get as a cation

Rules Continued

7. The algebraic sum of the oxidation numbers of all atoms in a neutral compound is 0

8. The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion

9. A monoatomic ion has an oxidation number equal to its charge

Practice

• Assign oxidation numbers to each atom:• UF6

• We know F has an oxidation number of -1, so the oxidation number for F6 is -6

• Since the compound is neutral, the U must be +6

• See the board for how to write this

Practice

• Assign oxidation numbers to each atom:• H2SO4

• H is bonded with more electronegative atoms, so it is +1 for each (total +2)

• O is -2 for each for a total of -8• S must be +6 to balance the neutral

formula

Practice

• Assign oxidation numbers to each atom:• ClO3

-

• O is -2 for a total of -6• The charge on the ion is -1, so Cl must be

+5• Practice problems on page 234 in the

book at school.

Redox Reaction• Reduction and oxidation occurs• This is a reaction in which electrons are

transferred from one atom to another• Oxidation: loss of electrons • Also the increase in oxidation number• Reduction: gain of electrons• Also the decrease in oxidation number• LEO goes GER or OILRIG

Example of Redox Reaction

• Zn(s) + I2(aq) → Zn2+(aq) + 2 I-(aq)

• Half reactions are used to show redox• Zn(s) → Zn2+

(aq) + 2e- (oxidized)

• I2 (aq) + 2e- → 2 I-(aq) (reduced)

Oxidation Number in Redox Reactions

• 2KBr(aq) + Cl2(aq) 2KCl(aq) + Br2(aq)

• Assign oxidation numbers to all elements in the balanced equation

• K in KBr = +1, K in KCl = +1 (no change)• Br in KBr = -1, Br in Br2 = 0 (increase)

• Cl in Cl2 = 0, Cl in KCl = -1 (decrease)

• Bromine is oxidized, chlorine is reduced

Reducing and Oxidizing Agents

• The species that is reduced is called the oxidizing agent

• The species that is oxidized is called the reducing agent.

• 2Zn + O2 2ZnO

• Zn in 2Zn = 0, Zn in ZnO = +2, Zn is oxidized and is the reducing agent

Continued• O in O2 = 0, O in ZnO = -2, O is reduced

and is the oxidizing agent• Practice: identified which species is

reduced, oxidized, the oxidizing agent, and the reducing agent

• N2(g) + 3H2(g) 2NH3(g)

• N is reduced, oxidizing agent, H is oxidized, reducing agent

Disproportionation

• Disproportionation: a process by which a substance acts as both the reducing agent and the oxidizing agent

• 2 Cu+(aq) → Cu2+

(aq) + Cu(s)

• +1 +2 0