1. Oxidation-Reduction Reactions
Chapter 19
Oxidation-Reduction Reactions

2. Sect. 19-1: Oxidation and Reduction
Rules for assigning oxidation states
Element by itself, 0
Monatomic ion equals the charge on the ion
More electronegative element in a compound is assigned number it would have as an ion
Oxidation # of fluorine is always -1
Oxygen is always -2, unless with fluorine when it is +2 or when peroxide it is a -1

3. Rules con’t
Hydrogen in most compounds is +1, unless with a metal then it’s a -1
Group 1 element = +1, group 2 = +2, Al= +3
Sum of oxidation numbers in a compound is 0
Sum of oxidation numbers in a polyatomic ion is equal to charge on the ion

4. Reduction – a decrease in oxidation state due to a gain of electrons
Oxidation – an increase in oxidation state during a reaction by losing electrons
Na  Na+ + e-, sodium is oxidized because its oxidation # increases from 0 to +1
Reduction – a decrease in oxidation state due to a gain of electrons
Cl2 + 2e-  2Cl-, chlorine is reduced because its oxidation number decreased from 0 to -1
OIL RIG (oxidation is loss, reduction is gain of electrons)

5. Oxidation-reduction reaction (redox reaction) – any chemical process in which elements undergo changes in oxidation number
Half-reaction – the part of the reaction involving oxidation or reduction alone

6. Redox reactions are not limited to ionic compounds
In the case of covalent compounds, the electron is not actually transferred, but rather partially shared in a manner different than before the reaction

7. Sect. 19-2: Balancing Redox reactions
Steps for balancing:
1. Write formula & ionic equations for the reaction.
2. Assign oxidation numbers & delete substances that do no show a change in oxidation #.
3. Write oxidation half reaction; balance atoms & charge.

8. 4. Write reduction half-reaction; balance atoms & charge.
5. Multiply coefficients of the half-reactions so that electrons lost = electrons gained.
6. Combine half-reactions & cancel anything common to both reactants & products.
7. Combine ions to form the original compounds. Check to be sure all ions balance.

9. Method 2 for balancing
Identify the change in oxidation number for the substance being oxidized and the substance being reduced
Add coefficients as needed to make the numbers match for electrons being lost/gained

10. Sect. 19-3: Oxidizing and Reducing agents
Reducing agent – substance that has the potential to cause another substance to be reduced (aka: substance that was oxidized)
Oxidizing agent – substance that has the potential to cause another substance to be oxidized (aka: substance that was reduced)

11. The more active an element is, the more likely it is to lose electrons and therefore is an excellent reducing agent (see pg. 603 chart for strength of reducing and oxidizing agents)
Fluorine is the most active oxidizing agent, but the weakest reducing agent due to its high electronegativity
The negative ion of a strong oxidizing agent is a weak reducing agent and vice versa

12. Autooxidation – a process in which a substance acts as both an oxidizing agent and a reducing agent
Example: H2O2  H2O + O2
hydrogen peroxide is both oxidized and reduced

13. Sect. 19-4: Electrochemistry
Electrochemistry – the branch of chemistry that deals with electricity-related applications of redox reactions

15. Electrochemical Cells (pg. 607)
Electrode – a conductor used to establish electrical contact with a nonmetallic part of a circuit, such as an electrolyte
Half-cell – a single electrode immersed in a solution of its ions
Anode – electrode where oxidation occurs
Cathode – electrode where reduction occurs
Electrochemical cell – a system of electrodes and electrolytes in which either chemical reactions produce electrical energy or an electric current produces chemical change

16. Voltaic cell – electrochemical cells in which the redox reaction occurs spontaneously and produces electrical energy
Zinc-carbon dry cells
Alkaline batteries
Mercury batteries

17. Electrolysis – the process in which an electric current is used to produce a redox reaction
Electrolytic cell – electrochemical cells in which electrical energy is required to produce a redox reaction and bring about a chemical change
See pg. 611 figure (& reading below) for difference between electrolytic and voltaic cells

18. Electroplating – an electrolytic process in which a metal ion is reduced and a solid metal is deposited on a surface
Object to be plated is the cathode and the anode is a piece of plating metal
Rechargeable cells act as voltaic cell when converting chemical energy to electrical energy and as an electrolytic cell when it is recharging (converting electrical to chemical energy)
Example: car battery

19. Reduction potential – the tendency for the half-reaction to occur as a reduction half-reaction in an electrochemical cell
Electrode potential – the difference in potential between an electrode and its solution
Standard electrode potential (E°)– a half-cell potential measured relative to a potential of zero for the standard hydrogen electrode
The half-reaction with the larger E° value will be the cathode
E° cell = E° cathode – E° anode