Chapter 18 Oxidation-Reduction Reactions and Electrochemistry

Oxidation-Reduction Reactions • Reactions in which one or

more electrons are transferred are called oxidation-reduction reactions

▫ Also called redox reactions

▫ Almost always the case between metals and non-metals

▫ Also occurs sometimes between nonmetals

LEO goes GER

• Oxidation is defined as a loss of electrons

▫ Lose Electrons Oxidation (LEO)

• Reduction is defined as a gain of electrons

▫ Gain Electrons Reduction (GER)

Example

• Sodium reacting with chloride to produce sodium chloride

▫ 2Na(s) + Cl2(g) 2NaCl(s)

• Sodium loses an electron to become positive

▫ Oxidation or reduction?

• Chlorine gains an electron to become negative

▫ Oxidation or reduction?

Oxidation States • Also called oxidation numbers

▫ Assigning charges to various atoms in a compound helps us keep track of electrons in redox reactions

Examples of Oxidation States • Sodium metal

▫ Oxidation state of 0 (it’s an uncombined element) • Sodium ion & Chlorine ion in NaCl

▫ Na - Oxidation state of +1 (same as charge) ▫ Cl – Oxidation state of -1 (same as charge) ▫ Overall charge is 0 (as per rule #6)

• Sulfur & Oxygen in SO2 ▫ O: -2 (as per rule #3 for oxygen) ▫ S: +4 (to balance out 2 O’s)

• Nitrogen & Hydrogen in NH3 ▫ H: +1 (as per rule #4 for hydrogen) ▫ N: -3 (rule #5 – N has a greater electronegativity)

• Nitrogen & Hydrogen in NH4 ▫ H: +1 (rule #4) ▫ N: -3 ▫ Overall charge is +1 (rule #7)

Oxidation-Reduction Reactions

Between Nonmetals

• The combustion of methane ▫ CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

• CH4 ▫ H: +1 (rule 4) ▫ C: -4 (to balance out H as per rule 6)

• CO2 ▫ O: -2 (rule 3) ▫ C: +4 (to balance out O as per rule 6)

• So carbon lost 8 electrons in this process of all nonmetals

• Oxygen also gained 8 electrons if you were to look at it in detail

Oxidation & Reduction

• Oxidation can now be defined as an increase in oxidation state (loss of electrons)

• Reduction can now be defined as a decrease in oxidation state (gain of electrons)

• 2Na(s) + Cl2(g) 2NaCl(s)

• Sodium lost electrons so it was oxidized

• Chlorine gained electrons so it was reduced

Oxidizing and Reducing Agents • An oxidizing agent accepts electrons

▫ It contains the element that was reduced

• A reducing agent donates electrons

▫ It contains the element that was oxidized

• 2Na(s) + Cl2(g) 2NaCl(s)

▫ Na is the reducing agent (it donated electrons and was, itself, oxidized)

▫ Cl is the oxidizing agent (it accepted the electrons and was, itself, reduced)

Nonmetals

• CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

• Carbon’s oxidation state went from -4 to +4 ▫ So it’s oxidized (lost electrons)

• CH4 is the reactant that contains the carbon that was oxidized so it’s the reducing agent ▫ donated electrons that move around

• Oxygen’s oxidation state went from 0 to -4 ▫ So it’s reduced (gained electrons)

• The reactant that contained the oxygen is O2 ▫ Its the oxidizing agent (accepts the electrons)

Review 1. What is oxidation? 2. What is reduction? 3. In a reaction between a metal and nonmetal:

1. Which substance tends to be oxidized? Why? 2. Which substance tends to be reduced? Why?

4. Assign oxidation states to: 1. O2

2. HSO4-

5. 4 Fe + 3 O2 2Fe2O3

1. What is reduced? 2. What is oxidized? 3. What is the reducing agent? 4. What is the oxidizing agent?

Half-Reactions • Balancing redox reaction equations that occur in aqueous

solutions can be tedious • Half reactions are equations that have electrons as reactants

or products • Ce4+ (aq) + Sn2+ (aq) Ce3+ (aq) + Sn4+ (aq) • One half reaction represents a reduction process

▫ Electrons are shown on the reaction side e- + Ce4+ (aq) Ce3+ (aq)

• One half reaction represents an oxidation process ▫ Electrons are shown on the product side

Sn2+ (aq) Sn4+ (aq) + 2e-

• To balance, the number of electrons lost must equal the number of electrons gained ▫ To do this, we must “double” the reduction reaction

2e- + 2Ce4+ (aq) 2 Ce3+ (aq) ▫ Electrons then cancel each other out and we get:

2Ce4+ (aq) + Sn2+ (aq) 2Ce3+ (aq) + Sn4+ (aq)

The Half-Reaction Method for Balancing

Equations for Redox reactions Occurring in

Acidic Solutions

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

Step 2 For each half-reaction: • Balance all of the elements except hydrogen and oxygen • Balance oxygen using H2O • Balance hydrogen using H+

• Balance the charge using electrons

Step 3 If necessary, multiply one or both balanced half-reactions by an integer to equalize the number of electrons transferred in the two half-reactions.

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

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

Electrochemistry

• A battery uses the energy from an oxidation reduction reaction to produce an electric current

• Electrochemistry is the study of the interchange of chemical and electrical energy.

• Electrochemistry involves 2 types of processes:

1. The production of an electric current from a chemical (redox) reaction

2. The use of an electric current to produce a chemical change

5Fe+(aq) + MnO4-

(aq) + 8H+(aq) 5Fe3+

(aq) + Mn2+(aq) + 4H2O(l)

• To produce a current, you need to separate the oxidizing and reducing agents & connect them with a wire

5Fe+(aq) + MnO4-

(aq) + 8H+(aq) 5Fe3+

(aq) + Mn2+(aq) + 4H2O(l)

• This setup doesn’t work – one side would become very negative and requires a lot of energy

5Fe+(aq) + MnO4-

(aq) + 8H+(aq) 5Fe3+

(aq) + Mn2+(aq) + 4H2O(l)

• We can solve this problem simply by connecting the solutions to keep the overall charge zero

Connecting oxidizing & reducing agents

• Both ways allow only ions to flow back and

forth

Electrochemical Battery

• Also called a galvanic cell is a device powered by an oxidation-reduction reaction where the oxidizing agent is separated from the reducing agent so that the electrons must travel through a wire from the reducing agent to the oxidizing agent

Electrochemical Battery

• The electrode where oxidation (losing electrons) occurs is called the anode

• The electrode where reduction occurs (gaining of electrons) is called the cathode

Batteries • Lead storage

battery

▫ Used in automobiles

▫ The reducing agent is lead metal

▫ The oxidizing agent lead(IV)oxide

Batteries • The tendency for electrons

to flow from an electron to the cathode depends on how well the reducing agent releases electrons and how well the oxidizing agent accepts those electrons

• The “pressure” on electrons to flow from one electrode to another in a battery is called the potential

Dry Cell Batteries

• Batteries that do not contain a liquid electrolyte

• In watches, mp3 players, etc.

Corrosion • Most metals are found in nature in compounds with

nonmetals such as oxygen and sulfur ▫ Ex: iron exists as iron ore

• Corrosion can be viewed as the process of returning metals to their natural state – the ores from which they were originally obtained ▫ Most oxidize – reaction with oxygen does this

• So we connect or coat metals with a metal that furnishes electrons more easily than iron (such as magnesium) ▫ This is called cathodic protection

Electrolysis • Electrolysis

involves forcing a current through a cell to produce a chemical change that would otherwise not occur

Review

1. What are half reactions? 2. What is electrochemistry? 3. What are the 2 major types of electrochemical

processes? 4. To create an electrochemical cell, what do you

have to separate? 5. Why can’t you have just one wire connecting the

oxidizing and reducing agents? 6. What happens at the anode? The cathode? 7. What is a dry cell battery? 8. What is corrosion? 9. What is electrolysis?