Mr. Bouyer
Day 1 | Day 2 - 3 | Day 4 - 5 | Lab | Skills Test

  • Identify elements that change oxidation state from one side of a chemical equation to the other.
  • Identify reducing agents and oxidizing agents in a chemical equation.
  • Use the half-reaction method to balance redox reactions.

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

There will be a formal test before Lab Day. You will be given three unbalanced redox equations and asked to balance two of these using the half-reaction method. You may use a pencil, your paper periodic table, and a calculator on the test, but nothing else.

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 gaining 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

 
How do you know this
is a "Redox" equation?

Answer

hydrogen plus oxygen yields water

 

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 to be 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.
  5. The sum of the oxidation numbers in a polyatomic ion = charge on the ion.
  6. Elements in group 1, 2, and aluminum are always as indicated on the periodic table.

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

The oxidation number of elements not covered by these rules must be "calculated" using the known oxidation numbers in a compound.

Example #1: K2CO3 by rule,
K is + 1

by rule,
O is - 2

To calculate C

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

by rule,
O is - 2

To calculate S

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

Practice Problems:

Use the rules above to determine the oxidation number of the element indicated in each formula.
1. Sb in Sb2O5
2. N in Al(NO3)3
3. P in Mg3(PO4)2
4. S in (NH4)2SO4
5. Cr in CrO4-2 		6. Cl in ClO4-
7. B in NaBO3
8. Si in MgSiF6
9. I in IO3-
10. N in (NH4)2S 		11. Mn in MnO4 -
12. Br in BrO3 -
13. Cl in ClO -
14. Cr in Cr2O7 -2
15. Se in H2SeO3

Day 2 - 3

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

The top arrow indicates the Reduction half-reaction. The bottom arrow indicated the Oxidation half-reaction.

Practice Problems: 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. For each reaction below, draw arrows and show electron numbers as in the example here. The top arrow indicates the element that gains electrons, reduction, and the bottom arrow indicates the element that looses electrons, oxidation. An arrow shows what one atom of each of these elements gaines or looses.

This is the first thing that must be done in balancing a Redox reaction. Learn to do it well.
1.  Mg + O2 MgO
2.  Cl2 + I - Cl - + I2
3.  MnO4 - + C2O4 -2 Mn+2 + CO2
4.  Cr + NO2 - CrO2 - + N2O2 -2
5.  BrO3 - + MnO2 Br - + MnO4 - 		6.  Fe+2 + MnO4 - Mn+2 + Fe+3
7.  Cr + Sn+4 Cr+3 + Sn+2
8.  NO3 - + S NO2 + H2SO4
9.  IO4- + I - I2
10.  NO2 + ClO - NO3 - + Cl -

 

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.
Check to see that electrons, elements, and total charge are balanced.
  • 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.

Practice Problems:
  1. Identify the oxidizing agent and reducing agent in each equation:

    1. H2SO4 + 8HI H2S + 4I2 + 4H2O
    2. CaC2 + 2H2O C2H2 + Ca(OH)2
    3. Au2S3 + 3H2 2Au + 3H2S
    4. Zn + 2HCl H2 + ZnCl2

  2. To make working with redox equations easier, we will omit all physical state symbols. However, remember that they should be there.
    An unbalanced redox equation looks like this:

    MnO4- + H2SO3 + H +    Mn+2 + HSO4- + H2O

    Study how this equation is balanced using the half-reaction method.link to a local webpage
    It is important that you understand what happens in each step.
    Be prepared to ask questions about this process in class tomorrow.

Day 4 - 5 Explore this Internet Website about Chemistry


The next two days are your opportunity of gain experience in balancing redox reactions. There are many little things that happen in these equations that you must experience to understand. Your science facilitator expects you to ask questions when you are not sure what to do.

Practice Problems: Balance these Redox equations using the half-reaction method. The balanced equation and the detailed half-reaction solution are provided for each equation. Use these helps only after doing your best to balance the equation without help.
It is to your advantage to work all these problems to gain experience
with different situations that arise when working with redox equations.
link to a local webpage with useful information

  1. HNO3 + H3PO3     NO + H3PO4 + H2O
  2. Cr2O7-2 + H + + I -     Cr+3 + I2 + H2O
  3. As2O3 + H + + NO3- + H2O     H3AsO4 + NO
  4. CuS + NO3-     Cu+2 + NO2 + S
  5. H2SeO3 + Br -     Se + Br2
  6. Fe+2 + Cr2O7-2 Fe+3 + Cr+3
  7. HS - + IO3-     I - + S
  8. CrO4-2 + I -     Cr+3 + I2
  9. IO4- + I -     I2
  10. MnO4- + H2O2 Mn+2 + O2
  11. H3AsO4 + Zn     AsH3 + Zn+2

    See the balanced equations and half-reaction solutions.link to a local webpage

There will be a formal test tomorrow. Use the skills test for a review.

Research Links:

Chemistry Class

 

 

 

 

 

 

 

The oxidation numbers of elements change from one side to the other.

Hydrogen's oxidation number is 0 on the left and +1 on the right.
Oxygen's oxidation number is 0 on the left and -2 on the right.

 

 

 

 

 

 

 

 

 

oxidation number answers:
1. Sb + 5
2. N + 5
3. P + 5
4. S + 6
5. Cr + 6 		6. Cl + 7
7. B + 5
8. Si + 4
9. I + 5
10. N - 3 		11. Mn + 7
12. Br + 5
13. Cl + 1
14. Cr + 6
15. Se + 4

 

 

 

 

 

 

oxidizing and reducing agents:

  1. oxidizing agent - H2SO4 || reducing agent - HI
  2. oxidizing agent - H2O || reducing agent - CaC2
  3. oxidizing agent - Au2S3 || reducing agent - H2
  4. oxidizing agent - HCl || reducing agent - Zn

 

 

 

 

 

 

 

 

 

 

Balanced Redox Equations:

1.   2HNO3 + 3H3PO3 2NO + 3H3PO4 + H2O . . . . . . . . . . . . . solutionlink to a local picture

2.   14H+ + Cr2O7-2 + 6I - 2Cr+3 + 3I2 + 7H2O . . . . . . . . . . . . . solutionlink to a local picture

3.   3As2O3 + 4H+ + 4NO3- + 7H2O 6H3AsO4 + 4NO . . . . . . . solutionlink to a local picture

4.   CuS + 2NO3- + 4H+ Cu+2 + 2NO2 + S + 2H2O . . . . . . . . . solutionlink to a local picture

5.   H2SeO3 + 4Br - + 4H+ Se + 2Br2 + 3H2O . . . . . . . . . . . . . . solutionlink to a local picture

6.   6Fe+2 + Cr2O7-2 + 14H+ 6Fe+3 + 2Cr+3 + 7H2O . . . . . . . . . solutionlink to a local picture

7.   3HS - + IO3- + 3H+ I - + 3S + 3H2O . . . . . . . . . . . . . . . . . . . solutionlink to a local picture

8.   16H+ + 2CrO4-2 + 6I - 2Cr+3 + 3I2 + 8H2O . . . . . . . . . . . . . solutionlink to a local picture

9.   8H+ + IO4- + 7I - 4I2 + 4H2O . . . . . . . . . . . . . . . . . . . . . . . . solutionlink to a local picture

10.   6H+ + 2MnO4- + 5H2O2 2Mn+2 + 5O2 + 8H2O . . . . . . . . . . solutionlink to a local picture

11.   8H+ + H3AsO4 + 4Zn AsH3 + 4Zn+2 + 4H2O . . . . . . . . . . . . solutionlink to a local picture

More Redox Equations: