1. Redox Reactions

2. Oxidation-Reduction Reactions (Redox)
Chemical reactions where reactants exchange electrons.
Eg: rusting and corrosion; all types of batteries, alkaline, NiCad, car; metabolism
OXIDATION: a process in which a substance loses electrons resulting in an increase in oxidation number.
REDUCTION: a process in which a substance gains electrons resulting in a lower oxidation number.
OXIDATION NUMBER: a positive or negative # assigned to an atom, ion or element. It indicates how electrons are shared. See rules for determining numbers.

3. Redox reactions are chemical reactions in which 2 or more atoms undergo a change in oxidation number.
Some redox reactions are spontaneous and result in a decrease in potential energy
They liberate energy which can be converted into electrochemical energy battery.
other redox reactions are not spontaneous and require energy to proceed.
Electrolytic cells hydrolysis of water
recharging a battery

4. LEO says GER Loss of electrons; oxidation / Gain of electrons; reduction)
Decreased
oxidation
number
REDUCTION
OXIDATION
increased
oxidation
number

5. Typical redox reactions
Magnesium burns in air to produce a bright light
2Mg(s) O2 (g) MgO(s)
Mg is oxidized by O2 ;O2 is called the oxidizing agent.
2Mg Mg e-
Each Mg loses 2 electrons (4 e- lost)
O2 is reduced by Mg; Mg is called the reducing agent.
O e O2-
Each O in O2 gains 2 electrons (4 e- gained)
IF ONE STUBSTANCE IS OXIDIZED, ANOTHER IN THE SAME REACTION MUST BE REDUCED.

6. Copper wire in silver nitrate solution
Cu(s) + 2AgNO3 (aq) Ag(s) + Cu(NO3)2(aq)
Net ionic equation:
Cu(s) + 2Ag+ (aq) Ag(s) + Cu2+(aq)
Note: spectator ions NO3- are not shown in equation because they do not react chemically.
Separately:
2Ag+(aq) e Ag(s)
gain of electrons  reduction copper wire is the reducing agent
Cu(s) Cu2+(aq) + 2e-
loss of electrons  oxidation silver ion is the oxidizing agent

7. Copper penny in nitric acid
Cu(s) + 4 HNO3 (aq) NO2 (g) + Cu(NO3)2(aq) + 2H2O(l)
Net ionic equation:
Cu(s) + 2NO3- (aq) + 4H+ (aq) NO2 (g) + Cu2+(aq) + 2H2O(l)
Separately:
Cu(s) Cu2+(aq) + 2e-
Cu loses of electrons  oxidation
2NO3- (aq) + 4H+ (aq) + 2e NO2 (g) + 2H2O(l)
How do you know N is reduced?
N gains of electrons  reduction

8. Oxidation States
A way of keeping track of the electrons.
need the rules for assigning (memorize).
The oxidation state of elements in their standard states is zero. Cu(s), Na(s), O2(g)
Oxidation state for monoatomic ions are the same as their charge. Cu2+, I-, N3-
Oxygen is assigned an oxidation state of -2 in its covalent compounds except as a peroxide. H2O, NO2 (ox. # -2) H2O2, Na2O2 (ox. # -1)
In compounds with nonmetals hydrogen is assigned the oxidation state +1. H2SO4, HCl, NaHCO3

9. In its compounds fluorine is always –1. SF6
The sum of the oxidation states must be zero in compounds.
PbCl4 NaBrO3
+4 + 4(-1) = (-2) = 0
7. The sum of the oxidation states must be equal the charge of the ion.
NO Cr2O7 2-
(-2) = (+6) + 7(-2) = -2
8. Group IA elements are always CsF, NaCl
Hydrogen is always +1 with the exception of hydrides.
(H is -1). LiH
10.Group IIA elements are always CaF2 , BaCl2

10. Oxidation States
Assign the oxidation states to each element in the following.
CO2
NO3-
H2SO4
Fe2O3
Na2Cr2O7
C = +4
O = -2
O = -2
N = +5
S = +6
O = -2
H = +1
Fe = +3
O = -2
Cr = +6
O = -2
Na = +1

11. Oxidation-Reduction summary
Oxidation means an increase in oxidation state - lose electrons.
Reduction means a decrease in oxidation state - gain electrons.
The substance that is oxidized is called the reducing agent.
The substance that is reduced is called the oxidizing agent.

12. Agents Oxidizing agents Reducing agents gets reduced gains electrons.
More negative oxidation state.
Reducing agents
gets oxidized.
Loses electrons.
More positive oxidation state.

13. Identify the… Oxidizing agent; Reducing agent
Substance oxidized; Substance reduced in the following reactions
Fe (s) + O2(g) ® Fe2O3(s)
Fe2O3(s)+ 3 CO(g) ® 2 Fe(l) + 3 CO2(g)
c) SO3-2 + H+ + MnO4- ® SO4-2 + H2O + Mn+2

14. homework
P. 653 #2
P.657 # 9-11
P. 659 #12, 13
P.662 #19

15. Half-Reactions
All redox reactions can be thought of as happening in two halves.
One produces electrons - Oxidation half.
The other requires electrons - Reduction half.
Write the half reactions for the following.
Na + Cl2 ® Na+ + Cl-
SO3-2 + H+ + MnO4- ® SO4-2 + H2O + Mn+2

16. Steps to Balancing Redox Equations
In aqueous solutions the key is the number of electrons produced must be the same as those required.
For reactions in acidic solution an 8 step procedure.
For reactions in basic solutions, one more step is required.

17. Acidic Solution Write separate half reactions
For each half reaction balance all reactants except H and O
Balance O using H2O
Balance H using H+
Balance charge using e-
Multiply equations to make electrons equal
Add equations and cancel identical species
Check that charges and elements are balanced.

18. BrO3- + I- Br- + I2 in an acidic solution
Write separate half reactions
BrO Br I I2
For each half reaction balance all reactants except H and O
BrO Br I I2
done
Balance O using H2O
BrO Br H2O I I2
Balance H using H+
BrO H Br H2O I I2

19. BrO3- + 6H+ + 6e- Br- + 3 H2O 2 I- I2 + 2e-
Balance charge using e-
BrO3- + 6H e Br H2O 2 I I e-
Multiply equations to make electrons equal
BrO3- + 6H+ + 6e Br- + 3 H2O [2 I I e- ]
Add equations and cancel identical species
BrO3- + 6H I Br I H2O
8. Check that charges and elements are balanced.

20. Practice
The following reactions occur in acidic solutions. Balance them.
a) MnO4- + H2O2 ® Mn+2 + O2
b) I- + NO3- ® I2 + NO(g)
c) Cr2O Fe2+ ® Cr3+ + Fe3+
d) Mn+2 + BiO3- ® Bi+3 + MnO4-

21. Basic Solution
Do everything you would with acid, but add one more step.
Add enough OH- to both sides to neutralize the H+

22. Al + NO3- Al+3 + NH3 in a basic solution
Write separate half reactions
Al Al NO NH3
For each half reaction balance all reactants except H and O
done done
Balance O using H2O
Al Al NO NH H2O
done
Balance H using H+
Al Al NO H NH H2O

23. Balance charge using e-
Al Al e NO3- + 7H+ + 6e NH3 + H2O
Add OH- to both sides to balance H+ . Create H2O.
Al Al e NO H e OH- NH H2O + 7OH-
Al Al e- NO3- + 7H2O + 6e NH H2O + 7OH-
Multiply equations to make electrons equal
2[Al Al3+ + 3e- ] NO H2O e NH H2O + 7OH-
Add equations and cancel identical species
2Al + NO H2O Al NH OH-
9. Check that charges and elements are balanced.

24. Practice
The following reactions occur in basic solutions. Balance them.
Cr(OH)3 + OCl- + ® CrO4-2 + Cl- + H2O
MnO4- + Fe+2 ® Mn+2 + Fe+3
Fe(OH)2 + H2O2 ® Fe(OH)4 -
d) S2O OCl- ® SO Cl-

25. Homework:p. 671 # 5 p673 #6-8, 3,4,7,