chapter 18

Chapter 18

Oxidation-Reduction Reactions

18.1 Electron Transfer Reactions

1. To learn about metal-nonmetal oxidation–reduction reactions

2. To learn to assign oxidation states

Oxidation-Reduction Reactions

• Oxidation-Reduction (redox) reaction – a reaction in which one ore more electrons are transferred– Oxidation – loss of electrons– Reduction – gain of electrons

Oxidation-Reduction Reactions

– Which element is oxidized? – Which element is reduced?

• In the following reactions, identify which element is oxidized and which element is reduced:2Mg(s) + O2(g) 2MgO(s)

2Al(s) + 3I2(s) 2AlI3(s)

Oxidation States

• Lets us keep track of electrons in oxidation-reduction reactions by assigning charges to the various atoms in a compound– Binary ionic compounds: oxidation state = the

charge of the ion– NaCl– MgO

• An atom in a pure element has an oxidation number of 0– Na– Cl2

• Oxidation states in covalent compounds – equal to the imaginary charges we determine by assuming that the most electronegative atom in a bond possesses both of the shared electrons– water

• The most electronegative elements are given oxidation states equal to the charge of their anion– F– O– N– Cl

• The sum of the oxidation states for an electrically neutral compound must be 0

• NO2

• Assign oxidation state to all atoms in the following:– CO2

– SF6

– NO3-

18.2 Balancing Oxidation-Reduction Reactions

1. To understand oxidation and reduction in terms of oxidation states

2. To learn to identify oxidizing and reducing agents

3. To learn to balance oxidation-reduction equations using half reactions

Oxidation-Reduction Reactions Between Nonmetals

• Oxidation- increase in oxidation state (loss of electrons)

• Reduction – decrease in oxidation state (gain of electrons)

• 2Na(s) + Cl2(g) NaCl

• Na oxidized – Na is also called the reducing agent (electron

donor).

• Cl2 reduced – Cl2 is also called the oxidizing agent (electron

acceptor).

• CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

• C oxidized – CH4 is the reducing agent.

• O2 reduced – O2 is the oxidizing agent.

• Identify the atoms that are oxidized and those that are reduced and specify the oxidizing and reducing agents.

2Al(s) + 3I2(s) 2AlI3(s)

PbO(s) + CO(g) Pb(s) + CO2(g)

Balancing redox reactions by the half reaction method

• Half reaction – equation that have electrons as reactants or products– One half reaction represents a reduction process

• Electrons are reactants

– One half reaction represents an oxidation process• Electrons are products

MnO4-(aq) + Fe2+(aq) Fe3+(aq) + Mn2+(aq)

Pb(s) + PbO2(s) + H+(aq) Pb2+(aq) + H2O(l)

18.3 Electrochemistry and Its Applications

1. To understand the concept of electrochemistry 2. To learn to identify the components of an

electrochemical (galvanic) cell 3. To learn about commonly used batteries 4. To understand corrosion and ways of preventing

it 5. To understand electrolysis 6. To learn about the commercial preparation of

aluminum

Electrochemistry: An introduction

• Electrochemistry – a study of the interchange of chemical and electrical energy

• Two types of processes:– Production of an electric current from a chemical

(redox) reaction– The use of an electrical current to produce a

chemical change

• Making an electrochemical cell

• If electrons flow through the wire charge builds up.

• Solutions must be connected to permit ions to flow to balance the charge.

Electrochemistry: An Introduction • A salt bridge or porous disk connects the half cells and allows

ions to flow, completing the circuit.

• Electrochemical battery (galvanic cell) – device powered by an oxidation-reduction reaction where the oxidizing agent is separated from the reducing agent so the electrons must travel through a wire from the reducing agent to the oxidizing agent

• Anode – electrode where the oxidation occurs

• Cathode – electrode where the reduction occurs

• Electrolysis – a process where electrical energy is used to produce a chemical change

• 2H2O(l) 2H2(g) + O2(g)

Batteries

• Lead Storage Battery – Anode reaction - oxidationPb + H2SO4 PbSO4 + 2H+ + 2e

– Cathode reaction - reduction

PbO2 + H2SO4 + 2e + 2H+ PbSO4 + 2H2O

– Overall reaction Pb + PbO2 + 2H2SO4 2PbSO4 + 2H2O

• Electrical Potential – the pressure on electrons to flow from one electrode to the other in a battery– Measured in volts

• Dry Cell Batteries – do not contain a liquid electrolyte – Acid version

• Anode reaction - oxidation

Zn Zn2+ + 2e • Cathode reaction – reduction

2NH4+ + 2MnO2 + 2e Mn2O3 + 2NH3 + 2H2O

– Alkaline version• Anode reaction - oxidation

Zn + 2OH ZnO + H2O + 2e • Cathode reaction – reduction

2MnO2 + H2O + 2e Mn2O3 + 2OH

– Other types

• Nickel-cadmium – rechargeable – products turned back into reactants by the use of external source of current

• Silver cell – Zn anode, Ag2O cathode

• Mercury cell – Zn anode, HgO cathode

Corrosion• Corrosion is the oxidation of metals to form mainly

oxides and sulfides.– Some metals, such as aluminum, protect

themselves with their oxide coating. – Corrosion of iron can be prevented by coatings,

by alloying and cathodic protection.

Cathodic protection of an underground pipe

Electrolysis

• Electrolysis – a process involving forcing a current through a cell to produce a chemical change that would not otherwise occur