1. ELECTROCHEMISTRY 9.1 and 9.2
To play the movies and simulations included, view the presentation in Slide Show Mode.

2. Electron Transfer Reactions
Electron transfer reactions are oxidation-reduction or redox reactions.
Eg: rusting and corrosion; all types of batteries, alkaline, NiCad, car; metabolism

3. Terminology for Redox Reactions
OXIDATION—loss of electron(s) by a species;
REDUCTION—gain of electron(s
OXIDIZING AGENT—electron acceptor; species is reduced.
REDUCING AGENT—electron donor; species is oxidized.

4. 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
LEO the lion says GER!
GER!

5. Another way to remember
OIL RIG

6. 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.

7. 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

8. 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

9. 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

10. Common oxidation numbersLi +1 Na Cs Rb K Fr Ba +2 Be Mg Sr Ca Ra B +3 C +4 -2 -4 N
+5 +4
+3 +2
+1 -3 O -1 -2 F -1 Ne H +1 He Al +3 Si +4 -4 P +5 +3 -3 S
+6 +4 +2 -2 Cl
+7 +5
+3 +1 -1 Ar Sc 3+ Ti +4 +3 +2 V
+5 +4 +3 +2 Cr +6 +3 +2 Mn
+7 +6
+4 +3 +2 Fe +3 +2 Co +3 +2 Ni +2 Cu +2 +1 Zn +2 Ga +3 Ge +4 -4 As 5+ 3+ 3- Se 6+ 4+ 2- Br +5 +1 -1 Kr +4 +2 Y +3 Zr +4 Nb +5 +4 Mo +6 +4 +3 Tc +7 +6 +4 Ru
+8 +6 +4 +3 Rh +4 +3 +2 Pd +4 +2 Ag +1 Cd +2 In +3 Sn +4 +2 Sb +5 +3 -3 Te +6 +4 -2 I
+7 +5 +1 -1 Xe +6 +4 +2 Lu +3 Hf +4 Ta +5 W +6 +4 Re +7 +6 +4 Os +8 +6 Ir +4 +3 Pt +4 +2 Au +3 +1 Hg +2 +1 Tl +3 +1 Pb +4 +2 Bi +5 +3 Po +2 At -1 Rn Lr +3
Common oxidation numbers

11. Oxidation numbers and the periodic table
Some observed trends in compounds.
Metals have positive oxidation numbers.
Transition metals typically have more than one oxidation number.
Nonmetals and semimetals have both positive and negative oxidation numbers.
No element exists in a compound with an oxidation number greater than +8.
The most negative oxidation numbers equals 8 - the group number

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

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

14. 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.

15. 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.

16. 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.

17. 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

18. 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.

19. 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-

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

21. 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

22. Balance charge using e- Al Al3+ + 3e- NO3- + 9H+ + 8e- NH3 + 3 H2O
Add OH- to both sides to balance H+ . Create H2O.
Al Al e- NO H e OH- NH H2O + 9 OH-
Al Al e- NO3- + 9H2O + 8e NH H2O + 9OH-
Multiply equations to make electrons equal
8[Al Al3+ + 3e- ] 3 (NO H2O e NH H2O + 9OH-)
Add equations and cancel identical species
8Al + 3NO H2O Al NH OH-
9. Check that charges and elements are balanced.

23. Practice
The following reactions occur in basic solutions. Balance them.
Cr(OH)3 + OCl- + ® CrO4-2 + Cl-
MnO4- + Fe+2 ® Mn+2 + Fe+3
Fe2+ + H2O2 ® Fe OH -
d) S2O OCl- ® SO Cl-