Redox Reactions Chapter 18 + O 2 . Oxidation-Reduction (Redox) Reactions “redox” reactions: rxns...

Redox Reactions

Chapter 18

+ O2

Oxidation-Reduction (Redox) Reactions

“redox” reactions: rxns in which electrons are transferred from one species to another

oxidation & reduction always occur simultaneously

we use OXIDATION NUMBERS to keep track of electron transfers

Rules for Assigning Oxidation Numbers:

1) the ox. state of any free (uncombined) element is zero. Ex: Na, S, O2, H2, Cl2, O3

Rules for Assigning Oxidation Numbers:

2) The ox. state of an element in a simple ion is the charge of the ion.

Mg2+ oxidation of Mg is +2

Rules for Assigning Oxidation Numbers:

3) the ox. # for hydrogen is +1

(unless combined with a metal, then it has an ox. # of –1)

Ex: NaOH (H bonded to O) v. NaH (H bonded to Na)

Rules for Assigning Oxidation Numbers:

4) the ox. # of fluorine is always –1.

Rules for Assigning Oxidation Numbers:

5) the ox. # of oxygen is usually –2.

Why USUALLY? Not -2 when it’s in a peroxide, such as hydrogen peroxide:

H2O2

Rules for Assigning Oxidation Numbers:

6) in any neutral compound, the sum of the oxidation #’s = zero.

Rules for Assigning Oxidation Numbers:

7) in a polyatomic ion, the sum of the oxidation #’s = the overall charge of the ion.

Rules for Assigning Oxidation Numbers:

**use these rules to assign oxidation #’s; assign known #’s first, then fill in the #’s for the remaining elements:

Examples: Assign oxidation #’s to each element:

a) NaNO3

Examples: Assign oxidation #’s to each element:

b) SO32-

Examples: Assign oxidation #’s to each element:

c) HCO3-

Examples: Assign oxidation #’s to each element:

d) H3PO4

Examples: Assign oxidation #’s to each element:

e) Cr2O72-

Examples: Assign oxidation #’s to each element:

f) K2Sn(OH)6

Definitions

Oxidation: the process of losing electrons (ox # increases)

Reduction: the process of gaining electrons (ox # decreases)

Oxidizing agents: species that cause oxidation (they are reduced)

Reducing agents: species that cause reduction (they are oxidized)

To help you remember…

OIL RIG Oxidation Is Loss Reduction Is Gain

Are all rxns REDOX rxns?

a reaction is “redox” if a change in oxidation # happens; if no change in oxidation # occurs, the reaction is nonredox.

Examples

MgCO3 MgO + CO2

Examples

Zn + CuSO4 ZnSO4 + Cu

Examples

NaCl + AgNO3 AgCl + NaNO3

Examples

CO2 + H2O C6H12O6 + O2

Balancing Redox Equations

Balancing Redox Equations

In balancing redox equations, the # of electrons lost in oxidation (the increase in ox. #) must equal the # of electrons gained in reduction (the decrease in ox. #)

There are 2 methods for balancing redox equations:1. Change in Oxidation-Number Method2. The Half-Reaction Method

1. Change in Oxidation-Number Method:

based on equal total increases and decreases in oxidation #’s

Steps:

1. Write equation and assign oxidation #’s.

2. Determine which element is oxidized and which is reduced, and determine the change in oxidation # for each.

3. Connect the atoms that change ox. #’s using a bracket; write the change in ox. # at the midpoint of each bracket.

4. Choose coefficients that make the total increase in ox. # = the total decrease in ox. #.

5. Balance the remaining elements by inspection.

Example

S + HNO3 SO2 + NO + H2O

If needed, reactions that take place in acidic or basic solutions can be balanced as follows:

Acidic: Basic:

add H2O to the side needing oxygen

balance as if in acidic sol’n

then add H+ to balance the hydrogen

add enough OH- to both sides to cancel out each H+ (making H2O) & then cancel out water as appropriate

Example: Balance the following equation, assuming it takes place in acidic solution.

ClO4- + I- Cl- + I2

2. The Half-Reaction Method:

separate and balance the oxidation and reduction half-reactions.

Steps:1. Write equation and assign oxidation #’s.2. Determine which element is oxidized and which is reduced,

and determine the change in oxidation # for each.3. Construct unbalanced oxidation and reduction half reactions.4. Balance the elements and the charges (by adding electrons

as reactants or products) in each half-reaction.5. Balance the electron transfer by multiplying the balanced half-

reaction by appropriate integers.6. Add the resulting half-reaction and eliminate any common

terms to obtain the balanced equation.

Example: Balance the following using the half-reaction method:

HNO3 + H2S NO + S + H2O