1. CHAPTER 20
Oxidation-Reduction Reactions
LEO SAYS GER

2. 20.2

3. Section 20.2 - Oxidation Numbers
Objectives
Determine the oxidation number of an atom of any element in a pure substance.
Define oxidation and reduction in terms of a change in oxidation number, and identify atoms being oxidized or reduced in redox reactions.

4. Remember Oxidation Numbers

5. Oxidation Numbers What is an oxidation number:
An oxidation number is a positive or negative number assigned to an atom to
keep track of electron transfers & electron sharing

6. Rules for Assigning Oxidation Numbers
The oxidation number of any uncombined element is zero.
2.The oxidation number of a monatomic ion equals its charge. +1 -1
2Na + Cl2 2NaCl

7. Rules for Assigning Oxidation Numbers
3. The oxidation number of O in compounds is -2,
Exceptions:
in hydrogen peroxide: H2O2 where it is -1,
in compounds with the more electronegative fluorine F2O +1 -2 +1 -1 -1 +2
F2O
H2O
H2O2
2(+1) + 1(-2) = 0
2(-1) + 1(+2) = 0
2(+2) + 2(-2) = 0

8. Rules for Assigning Oxidation Numbers
4. The oxidation number of H in compounds is +1, Exception: in metal hydrides it is -1. +1 -2 +1 -1
H2O
NaH
Warning!
Don't get too bogged down in these exceptions. In most of the cases you will come across, they don't apply!

9. Rules for Assigning Oxidation Numbers
5. For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal 0. +1 -2 +2 -2 +1
Ca(OH)2
H2O
(+2) + 2(-2) + 2(+1) = 0
Ca O H
2(+1) + (-2) = 0
H O

10. Rules for Assigning Oxidation Numbers
6. For a polyatomic ion, the sum of the oxidation numbers in the formula is equal to the ionic charge of the ion ? -2 ? -2
SO42-
NO3-
X + 3(-2) = -1
N O
X + 4(-2) = -2
S O
THUS X = +6
THUS X = +5

11. 20.1

12. Section 20.1 – The Meaning of Oxidation and Reduction
Objectives:
Define oxidation and reduction in terms of loss or gain of oxygen and the loss or gain of electrons
State the characteristics of a redox reaction
Describe what happens to iron when it corrodes

13. Oxidation and Reduction
We already learned why salt is spread on roads during winter months
Lowers the freezing point of water preventing the buildup of slippery ice
But…what about the affect of the salt that clings to the metallic parts of cars?
RUST
The corrosion of metal is one example of an oxidation-reduction reaction

14. Other Examples of Oxidation-Reduction Reactions
When methane burns in air, it oxidizes and forms oxides of carbon and hydrogen
One oxide of carbon is carbon dioxide
Elemental iron slowly oxidizes into compounds such as iron (III) oxide, or rust
Hydrogen peroxide also releases oxygen when it decomposes – when you pour it on an open cut
Bleaching stains in fabrics is an oxidation-reduction reaction

15. Oxidation and Reduction
What is an oxidation-reduction reaction?
Classic Definition: A reaction where oxygen is transferred from one substance to another.
In an oxidation-reduction reaction:
The substance gaining oxygen is oxidized
The substance losing oxygen is reduced.
Oxidation & reduction always occur simultaneously
An abbreviation for oxidation-reduction reactions is redox reactions

16. Redox Reactions
Modern concepts of redox reactions include reactions that do not even involve oxygen
Redox reactions are currently understood to involve the shift of electrons between reactants
Oxidation
Complete or partial loss of electrons, or
gain of oxygen
Reduction
Complete or partial gain of electrons, or
loss of oxygen

17. Why call it Oxidation?
With the exception of fluorine, oxygen is the most electronegative element
When oxygen bonds with atoms of different elements, the electrons from the other atom shift toward oxygen
When a substance gains electrons, it acting like oxygen

18. Oxidation and Reduction
Until now our view of a reaction:
Mg + S  MgS
Reaction viewed as Redox:
Mg S → Mg S2-
The magnesium atom changes to a more stable magnesium ion by losing 2 electrons, and is thus oxidized
The sulfur atom is changed to a more stable sulfide ion by gaining 2 electrons, and is thus reduced.

19. Oxidation and Reduction
Mg + S MgS
The over all process is represented as the two component processes below:
Mg Mg e-
S + 2e- S2-
Oxidation – loss of electrons
Reduction – gain of electrons

20. Example 2: Oxidation and Reduction
Each sodium atom loses one electron – oxidation
Each chlorine atom gains one electron – reduction

21. LEO says GER Gain Electrons = Reduction Lose Electrons = Oxidation
Sodium is oxidized
Gain Electrons = Reduction
Chlorine is reduced

22. Remember
Oxidation - reduction reactions involve a transfer of electrons.
LEO - GER
Lose Electrons Oxidation
Gain Electrons Reduction

23. Practice Problem
Determine what is oxidized and what is reduced in each reaction.
2Na + S Na2S
4Al + 3O2 2Al2O3
Na is oxidized
(Na1+ S2-)
S is reduced
Al is oxidized
(Al3+ O2-)
O is reduced

24. Practice Problem
Identify these processes as either oxidation or reduction.
S2- S + 2e-
Zn2+ + 2e- Zn
Oxidation – is loss
Reduction – is gain

25. Redox in Covalent Compounds
It is easy to see the loss and gain of electrons in ionic compounds, but what about covalent compounds? (complete electron transfer does not occur)
Remember when we did bonding, we looked at Electronegativity.
 Electronegativity = strong pull on the electron.
If the electrons are shared between two atoms, the atom with the higher Electronegativity has the electron most of the time.
So the oxidation number will show the electron belonging to the atom with the higher Electronegativity
(In reality the e- is not completely transferred, but it helps us keep track of where the e- is)

26. Redox in Covalent Compounds
It is easy to see the loss and gain of electrons in ionic compounds, but what about covalent compounds? (complete electron transfer does not occur)
Consider: 2H2 + O H2O
Oxygen is highly electronegative
In H there is a shift of bonding electrons away from H
Hydrogen is oxidized
In O there is a shift of electrons toward O
Oxygen is reduced
loss
gain

27. Oxidation and Reduction
In some reactions involving covalent reactants and products, the partial electron shifts are less obvious
A few general guidelines:
Gain of oxygen – oxidation
Loss of oxygen – reduction
Loss of hydrogen by a covalent compound – oxidation
Gain of hydrogen by a covalent compound – reduction
Increase in oxidation number – oxidation
Decrease in oxidation number – reduction

28. Terminology for Redox Oxidation Reduction Oxidizing Agent
Loss of electrons
Increase in oxidation number
Increase in oxygen
Oxidizing Agent
Electron acceptor
Species that is reduced
Reduction
- Gain of electrons
- Decrease in oxidation
number
- Decrease in oxygen
Reducing Agent
- Electron donor
- Species that is oxidized

29. Trends in Oxidation and Reduction
Active metals:
Lose electrons easily
Are easily oxidized
Are strong reducing agents
Active nonmetals:
Gain electrons easily
Are easily reduced
Are strong oxidizing agents

30. Oxidation – Number Changes in Chemical Reactions
An increase in oxidation number = oxidation
A decrease in oxidation number = reduction
Determining all of the oxidations in a chemical equation allows us to determine the oxidizing agent and reducing agent (which species is oxidized and which species is reduced)
Determine the oxidation numbers in the following chemical equation.
2AgNO3 + Cu Cu(NO3)2 + 2Ag +1 +5 -2 +2 +5 -2
Copper is oxidized
Silver is reduced

31. 20.3

32. Section 20.3 – Balancing Redox Equations
Objectives:
Describe how oxidation numbers are used to identify redox reactions.
Balance a redox equation using the oxidation-number-change method.
Balance a redox equation by breaking the equation into oxidation and reduction half-reactions, and then balance using the half-reaction method.

33. Identifying Redox Equations
In general, all chemical reactions can be assigned to one of two classes:
oxidation-reduction, in which electrons are transferred:
Single-replacement, combination, decomposition, and combustion
All other reactions…
No electron transfer
Double-replacement and acid-base reactions

34. Identifying Redox Equations
In an electrical storm, oxygen and nitrogen react to form nitrogen monoxide:
N2(g) + O2(g) → 2NO(g)
Is this a redox reaction?
If the oxidation number of an element in a reacting species changes, then that element has undergone either oxidation or reduction; therefore, the reaction as a whole must be a redox reaction.
YES!

35. Balancing Redox Equations
It is essential to write a correctly balanced equation that represents what happens in a chemical reaction
Many redox reactions are too complex to be balanced by trial and error
Fortunately, there are two systematic methods based on the fact that the total electrons gained in reduction equals the total lost in oxidation
The methods:
Using oxidation-number changes
Use half-reactions

36. Using Oxidation-Number Changes
Type of “chemical bookkeeping” to keep track of electron transfers
Balance a redox equation by comparing the increases and decreases in oxidation number
Start with the skeleton equation
Fe2O3 + CO Fe + CO2 (unbalanced)
Assign oxidation numbers to all the atoms in the equation
Fe2O3 + CO Fe + CO2 +3 -2 +2 -2 +4 -2

37. Using Oxidation-Number Changes
2. Identify which atoms are oxidized and which are reduced
Iron decreases in oxidation number from +3 to 0, a change of -3. Iron is reduced
Carbon increases in oxidation number from +2 to +4, a change of +2. Carbon is oxidized
3. Use brackets to connect the atoms that undergo oxidation and another to connect those that undergo reduction. Write the oxidation number change on the line.
Fe2O3 + CO Fe + CO2
+2 oxidation
-3 reduction
As the equation is written, the number of electrons transferred in oxidation, does not equal the number of electrons transferred in reduction

38. Using Oxidation-Number Changes
4. Make the total increase in oxidation number equal to the total decrease in oxidation number by using appropriate coefficients
Fe2O3 + CO Fe + CO2
Using the numbers you multiplied the oxidation number changes by, add appropriate coefficients to the equation
Fe2O3 + 3CO 2Fe + 3CO2
5. Finally, make sure that the equation is balanced for both atoms and charges. If necessary, finish balancing the equation by inspection
3 X (+2) = + 6
2 X (-3) = - 6

39. Using Oxidation-Number Changes
Balance this redox reaction by using the oxidation-number-change method
K2Cr2O7 + H2O + S KOH + Cr2O3 + SO2
3 X
+ 4 oxidation
= + 12 +6 +3 +4
4 X
- 3 reduction
= - 12 2
K2Cr2O7 + H2O + S KOH + Cr2O3 + SO2 2 3 4 2 3

40. Rules shown on page 651 – bottom
Using Half-Reactions
A half-reaction is an equation showing just the oxidation or just the reduction that takes place
They are then balanced separately, and finally combined
1. Write unbalanced equation in ionic form
2. Write separate half-reaction equations for oxidation and reduction
3. Balance the atoms in the half-reaction
4. Add enough electrons to one side of each half-reaction to balance the charges
5. Multiply each half-reaction by a number to make the electrons equal in both
6. Add the balanced half-reactions to show an overall equation
7. Add the spectator ions and balance the equation
Rules shown on page 651 – bottom

41. You can’t have one… without the other!
Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons.
You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation
GER!

42. Using Half-Reactions
Consider the reduction of Ag+ ions with copper metal.
Cu + Ag give--> Cu Ag

43. Using Half-Reactions
Step 1: Divide the reaction into half- reactions, one for oxidation and the other for reduction.
Ox Cu ---> Cu2+
Red Ag+ ---> Ag
Step 2: Balance each element for mass. Already done in this case.
Step 3: Balance each half-reaction for charge by adding electrons.
Ox Cu ---> Cu e-
Red Ag+ + e- ---> Ag

44. The equation is now balanced for both mass & charge
Using Half-Reactions
Step 4: Multiply each half-reaction by a factor so that the reducing agent supplies as many electrons as the oxidizing agent requires.
Reducing agent Cu ---> Cu e-
Oxidizing agent 2 Ag e- ---> 2 Ag
Step 5: Add half-reactions to give the overall equation.
Cu Ag > Cu Ag
The equation is now balanced for both mass & charge

45. Zn0 + H+1N+5O3-2 Zn+2(N+5O-23)3 + N+4O-22 + H2+1O-2
Practice Problem
Balance the following equation using the half reaction method
Zn + HNO3 Zn(NO3)3 + NO2 + H2O
Step 1: Identify the oxidized and reduced species by assigning oxidation numbers
Zn0 + H+1N+5O3-2 Zn+2(N+5O-23)3 + N+4O-22 + H2+1O-2
Step 2: Write the half reactions
Oxidation: Zn0 Zn+2 + 2e-
Reduction: N+5 + 1e- N+4
2( N+5 + 1e- N+4)
2N+5 + 2e- 2N+4
Step 3: Balance species oxidized and reduced in equation
Zn + 2HNO3 Zn(NO3)3 + 2NO2 + H2O
Step 4: Balance remaining atoms
Nitrogen is not balanced because some are not reduced, balance H and O
Zn + 4HNO3 Zn(NO3)3 + 2NO2 + 2H2O

46. Choosing a Balancing Method
If the oxidized and reduced species appear only once one ach side of the equation and no acids and bases are involved in the reaction – oxidation number change usually works
Sometimes the same element is both oxidized and reduced – reactions when this occurs are best balanced by the half-reaction method
Equations for reactions that take place in acidic or alkaline solution are best balanced by the half-reaction method