CLASS CONCEPTS

1. Introduction to Chemistry

2. The Periodic Table

3. Quantum Numbers

4. Electron Configuration

5. Chemical Families

6. Oxidation Numbers

7. Chemical Formulas

8. Chemical Names

9. Formula Mass

10. Percentage Composition

11. Reaction Types

12. Balancing Equations

13. The Mole Concept

14. Solution Concentration

15. Stoichiometry

16. Kinetic Theory

17. The Gas Laws

18. Enthalpy & Heat

19. Reaction Rates

20. Acids & Bases

21. pH Scale

22. Salts

23. Net Ionic Equations

24. Redox Reactions

25. Organic Chemistry

26. Nuclear Chemistry

24. Oxidation - Reduction Reactions (REDOX)

 

Oxidation / Reduction - "Redox" - reactions in which atoms or ions undergo changes in electron structure.

Oxidation is defined as:
  • The loss of electrons from an atom or ion.
     
  • Also, the combination of oxygen with other substances.
Reduction is defined as:
  • The gain of electrons by an atom or ion.
     
  • Also, "reducing" a substance into its components.


LEO  the lion goes  GER

Loose Electrons, Oxidize         Gain Electrons, Reduce

Question: How do you know the equation below is a "Redox" equation?

2H2 + O2 → 2H2O

Answer the question.

 
Oxidation and Reduction must both occur in a Redox reaction. If one particle gains electrons in a reaction, some other particle must lose them.

You have learned to read the oxidation number of many elements from the Periodic Table. While that information is important, the following rules are your guide when working with Redox Equations.

Rules for assigning oxidation numbers:

  1. The oxidation number of a free element = 0.
     
  2. The oxidation number of a monatomic ion = charge on the ion.
     
  3. The oxidation number of hydrogen = +1 and rarely −1.
     
  4. The oxidation number of oxygen = −2 and in peroxides −1.

    Hydrogen peroxide, H2O2, is the only peroxide you are responsible for recognizing.

  5. The oxidation number of fluorine always = −1.
     
  6. The sum of the oxidation numbers in a polyatomic ion = charge on the ion.
     
  7. Elements in group 1, 2, and aluminum are always as indicated on the periodic table.
 
Oxidation numbers of elements not covered by the rules above must be "calculated" using known oxidation numbers in a compound.

Example #1: K2CO3

  • By rule, K is +1 and O is −2
  • The oxidation number for C must be calculated.
    • The sum of all the oxidation numbers in this formula equal 0.
    • Multiply the subscript by the oxidation number for each element.
      • K = (2) (+1) = +2
      • O = (3) (−2) = −6
      • therefore, C = (1) (+4) = +4
Example #2: HSO4−
  • By rule, H is +1 and O is −2
  • The oxidation number for S must be calculated.
    • The sum of all the oxidation numbers in this formula equal −1.
    • Multiply the subscript by the oxidation number for each element.
      • H = (1) (+1) = +1
      • O = (4) (−2) = −8
      • therefore, S = (1) (+6) = +6
 
Redox Problems #1

 
Reducing Agents and Oxidizing Agents:

Reducing agent - the reactant that gives up electrons.

  • The reducing agent contains the element that is oxidized (looses electrons).
  • If a substance gives up electrons easily, it is said to be a strong reducing agent.
Oxidizing agent - the reactant that gains electrons.
  • The oxidizing agent contains the element that is reduced (gains electrons).
  • If a substance gains electrons easily, it is said to be a strong oxidizing agent.
Example: Fe2O3(cr) + 3CO(g) → 2Fe(l) + 3CO2 (g)

Notice that the oxidation number of C goes from +2 on the left to +4 on the right. The reducing agent is CO, because it contains C, which loses e−.

Notice that the oxidation number of Fe goes from +3 on the left to 0 on the right. The oxidizing agent is Fe2O3, because it contains the Fe, which gains e−.

 
Charting Reducing Agents and Oxidizing Agents:

In any Redox equation, at least one particle will gain electrons and at least one particle will lose electrons. This is indicated by a change in the particle's oxidation number from one side of the equation to the other.

The change is shown in a chart like this:

  • The top arrow shows the element that gains electrons, (reduction).
     
  • The bottom arrow shows the element that looses electrons, (oxidation).
     
  • The arrows show the electrons gained or lost by one atom of each element.
 
Redox Problems #2

 
Balancing Redox Equations by the Half-reaction Method

1. Decide what is reduced (oxidizing agent) and what is oxidized (reducing agent).
  • Do this by drawing arrows as in the practice problems.
2. Write the reduction half-reaction.
  • The top arrow in step #1 indicates the reduction half-reaction.
  • Show the electrons gained on the reactant side.
  • Balance with respect to atoms / ions.
    • To balance oxygen, add H2O to the side with the least amount of oxygen.
    • THEN: add H + to the other side to balance hydrogen.

Remember that the arrow in step #1 indicates the number of electrons gained by one atom.

3. Write the oxidation half-reaction.

  • The bottom arrow in step #1 indicates the oxidation half-reaction.
  • Show the electrons lost on the product side.
  • Balance with respect to atoms / ions.
    • To balance oxygen, add H2O to the side with the least amount of oxygen.
    • THEN: add H + to the other side to balance hydrogen.

Remember that the arrow in step #1 indicates the number of electrons lost by one atom.

4. The number of electrons gained must equal the number of electrons lost.

  • Find the least common multiple of the electrons gained and lost.
  • In each half-reaction, multiply the electron coefficient by a number to reach the common multiple.
  • Multiply all of the coefficients in the half-reaction by this same number.

5. Add the two half-reactions.

  • Write one equation with all the reactants from the half-reactions on the left and all the products on the right.
  • The order in which you write the particles in the combined equation does not matter.

6. Simplify the equation.

  • Cancel things that are found on both sides of the equation as you did in net ionic equations.
  • Rewrite the final balanced equation.

7. Check the equation.

  • There should be no electrons in the equation at this time.
  • The number of each element should be the same on both sides.
  • It doesn't matter what the charge is as long as it is the same on both sides.
  • If any of these are not balanced, the equation is incorrect. The only thing to do is go back to step #1 and begin looking for your mistake.
Sample Problem:   An unbalanced redox equation looks like this -
Note: the equation is already in net ionic form.
MnO4− + H2SO3 + H + → Mn+2 + HSO4− + H2O

See how this equation is balanced using the half-reaction method.

 
Balancing Redox Equations

 
Redox Equations

 
Transformations of Copper