Reference: NC Science Olympiad Page, Philip Dail…NC StateUniv.

REDOX REACTIONS AND BALANCING EQUATIONS

The Process:

  • Redox reactions: the transfer of electrons
  • Acid/base reactions: the transfer of protons.
  • Precipitation reactions: not a transfer as such but the combination of existing ions in solution such that an insoluble compound is formed.

Examples of Redox Reactions:

  • Production of ATP in the mitochondria
  • Corrosion of metals…ex. rusting of iron
  • Batteries

LEO the lion goes GER: lose electrons oxidation, gain electrons reduction!

OILRIG: oxidation is losing, reduction is gaining

e– + 2H+ + N+3O2– --N+2O + H2O Reduction

2I– -- I20 + 2e– Oxidation

  • BIG GOAL IN REDOX . . .Keeping up with where the electrons go!!!!

OXIDATION NUMBERS

1.All free pure elements have an oxidation number of 0 since they are only made of neutral atoms.

2.For all molecules and formulas for ionic compounds, the sum of the oxidation numbers in them equals zero.

  1. For polyatomic ions, the sum of the oxidation numbers equals the charge you have learned (or are supposed to have learned!!!!) on the polyatomic ion.
  1. Elements not combined with other elements have an oxidation number of zero, e.g. O2, P4, S8
  1. Oxygen when combined always has an oxidation number of –2 except in peroxides, H2O2.
  1. Hydrogen when combined always has an oxidation number of +1 except in certain metal hydrides, NaH, when it is –1.

7.The "charge" of a single atom that has become an ion, ex. Na+, is its oxidation number. As in the case of the sodium ion here, the oxidation number is +1.

Consider several examples to show how to assign oxidation numbers.

CrO42– NH3 N2H4 N2

NONO2NO3-

II.NOW WHAT TO DO WITH THESE NUMBERS??? One use is to answer this question.

Is the following equation for a redox reaction?

CH4 + O2 -- CO2 + H2O

Which of the following equations represent oxidation-reduction reactions?

1.CH4 + O2 -- 2H2O + CO2

2.Zn + 2HCl -- ZnCl2 + H2

3.2Na + 2H2O -- 2NaOH + H2

4.MnO2 + 4HCl -- Cl2 + 2H2O + MnCl2

a. 1 only b. 2 and 3 only c. 1 and 4 only d. 1, 3, and 4 only e. All are oxidation-reduction reactions.

A helpful hint: any time an element exists on one side of the arrow by itself but is combined with other elements on the other side of the arrow, that element is undergoing either reduction or oxidation, and the reaction is redox.

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BALANCING REDOX REACTIONS IN AQUEOUS SOLUTIONS

IN ACID

NO2– + I– -- NO(g) + I2(s)

Step One: Take the equation and split it into 2 "logical" half reactions.

NO2– -- NO

I– -- I2

Step Two: Balance all elements by using coefficients EXCEPT hydrogen and oxygen.

NO2– -- NO

2I– -- I2

Step Three: Balance the O's by adding H2O's wherever needed.

NO2– -- NO + H2O

2I– -- I2

Step Four: Balance H's by adding H+'s wherever needed.

2H+ + NO2 -- NO + H2O Note, adding the 2 H+'s on the left took care of the H's.

2I– -- I2 Again, no H's to worry about.

The steps up to this point balance all the elements BUT NOT THE CHARGE!!!!

Step Five: Add electrons in order to balance the charge.

e– + 2H+ + NO2– --NO + H2O

2I– -- I2 + 2e–

Step Six: Multiply the two equation by a set of numbers so that e lost equals e gained.

2X (e– + 2H+ + NO2 --NO + H2O)

2I– -- I2 + 2e–

Now add the two halves together since the electrons are equal.

2I– + 4H+ + 2NO2– -- 2NO + 2H2O + I2

In Basic Solution:

Use the same steps 1-6. YOU MUST LEARN THESE STEPS. After doing the first 6 steps, you simply convert the H+'s to OH–'s.

Example: The following reaction occurs in a basic solution.

Cr(OH)3 + ClO3– --CrO42– + Cl–

H2O + Cr(OH)3 -- CrO42– + 5H+ + 3e-

6e– + 6H+ + ClO3– -- Cl– + 3H2O

2H2O + 2Cr(OH)3 --2CrO42– + 1OH+ + 6e–

Now add the two half reactions and will have the following.

6H+ + ClO3– + 2H2O + 2Cr(OH)3 -- Cl– + 3H2O + 2CrO42– + 1OH+

"Collect" all the waters and hydrogen ions on one side

ClO3– + 2H2O + 2Cr(OH)3 -- Cl– + H2O + 2CrO42– + 4H+

This is not basic yet since the H+'s are clearly in the equation.

To accomplish this conversion, add 4 OH–'s to each side.

4OH– + ClO3– + 2Cr(OH)3-- Cl– + H2O +2CrO42– + 4H+ + 4OH–

4OH– + ClO3– + 2H2O + 2Cr(OH)3 -- Cl– + 5H2O + 2CrO42–

OXIDIZING AND REDUCING AGENTS

  • The reducing agent is the substance oxidized.
  • The oxidizing agent is the substance reduced.

One common confusion about redox is how to identify the oxidizing and reducing agents in the reaction. First, the two agents are always reactants in the original equation. They are not the H+'s and H2O's added to balance H's and O's.

TYPES OF REDOX REACTIONS
  1. SIMPLE REDOX
  1. Single replacement
  2. Combustion
  3. Binary compound synthesis
  1. REACTIONS INVOLVING OXOANIONS: ex. Cr2O72-
  1. ATYPICAL REDOX REACTIONS (memorize for AP)
  1. Hydrogen reacts with a hot metallic oxide to produce the elemental metal and water.
  2. Metal sulfide reacts with oxygen to produce the metallic oxide and sulfur dioxide.
  3. Chlorine gas reacts with dilute sodium hydroxide to produce sodium hypochlorite, sodium chloride and water.
  4. Copper reacts with concentrated sulfuric acid to produce copper(II) sulfate, sulfur dioxide and water.
  5. Copper reacts with dilute nitric acid to produce copper(II) nitrate, nitrogen monoxide and water.
  6. Copper reacts with concentrated nitric acid to produce copper(II) nitrate, nitrogen dioxide and water.

Remember: All single replacement, all combustion, some synthesis, some decomposition and atypical reactions are redox. Acid/Base and Precipitation ARE NOT REDOX.