Download File

Electrochemistry
Electron Transfer Reactions



Reactions where electrons flow
between atoms
Only reactions which produce
electricity
Only valence electrons are transferred
Oxidation State



Used to determine if electrons are
transferred in a reaction
Artificial numbers which DO NOT
INDICATE the atom’s ionic CHARGE
ELECTRON BOOKKEEPING
Oxidation Numbers


charge an atom would have if electrons
in its bonds belonged completely to the
more electronegative atom
oxidation numbers are written as +1
vs. 1+ to distinguish them from ionic
charges.

different than formal charge (charge an
atom would possess if all atoms had
same electronegativity)
Assigning Oxidation States


Rules on p.137
Differ some depending on resource
Oxidation States Shortcut Method
RULE 1
The oxidation state of an atom in the
most abundant, naturally occurring form
of an element is typically zero.
EX: diatomic
noble gas
metal
others
H2
Ne
Fe
P4, S8
Oxidation States Shortcut Method
Rule 2
The oxidation state of any Group 1A
atom in a compound is always +1.
Rule 2
The oxidation state of any Group 2A
atom in a compound is always +2.
Oxidation States Shortcut Method
Rule 2
The oxidation state of a monatomic ion
(one atom) is equal to the charge of the
ion.
Ex: Al3+, N3-, S2-
Oxidation States Shortcut Method
Rule 3a
The oxidation state of oxygen in a
compound is almost always -2.
Exceptions:
 peroxides O22- are -1
 when oxygen is bound to a more
electronegative atom (fluorine)
Oxidation States Shortcut Method
Rule 3b
The oxidation state of hydrogen in a
compound is almost always +1.
Exceptions:
 when hydrogen is bound to a less
electronegative atom (hydrides)
Ex: AsH3, SiH4, GeH4
Oxidation States Shortcut Method
Rule 3c
If a halogen is in a compound, it will
most likely have a -1 oxidation number.
The oxidation number of fluorine is
always -1.
Exception:
 Br, Cl or I is bonded to oxygen or fluorine
Oxidation States Shortcut Method
Rule 4
The sum of the oxidation states for
every atom in a formula must add up to
the overall charge of the formula.
Ex: H2O, (SO4)2-, (NH4)1+
Another Version of Rules
1.
2.
3.
4.
5.
6.
7.
Any element by itself is 0 (O in O2 is zero)
H is +1 except in metal hydrides (-1)
O is –2 except in peroxides (-1) or combined with F
(+2)
Metal in ionic compound equals its positive ionic
charge
In covalent compound the most electronegative
atom gets the negative charge.
Sum of compound’s oxidation numbers must equal
zero.
Sum of polyatomic ion’s oxidation numbers must
equal the charge of the ion.
What is REDOX?




REDOX stands for
REDuction/OXidation
Oxidation is often thought of as a
combination of a substance with
oxygen (rusting, burning)
Oxidation refers to a loss of electrons
Reduction refers to a gain of electrons
OIL RIG






Oxidation
Is
Loss of electrons
Reduction
Is
Gain of electrons
Increase in
oxidation
number
Oxidizing & Reducing Agents
REDUCING AGENT
 substance in a REDOX reaction that
donates an electron
 Contains element that is oxidized
 organic chemistry: oxidations remove
protons or add oxygen to the organic
molecule
Oxidizing & Reducing Agents
OXIDIZING AGENT
 Substance that readily transfers
oxygen atoms or that gains electrons
in a redox reaction
 Contains element that is reduced
 organic chemistry: reductions adds
protons or remove oxygen from an
organic molecule.
Applications of Redox







Batteries
Electroplating (mirrors, jewelry)
Fuel cell
Bleaching
Hardening of arteries
Cancer
arthritis
Using Organic Redox
majority of the energy generated in our
bodies comes from aerobic metabolism
C6H12O6(aq) + 6O2(g)  6CO2(g) + 6H2O(l)
 oxidation of compounds such as
carbohydrates (using the O2 that we
breathe) in redox reactions produces
ATP, our “energy currency”
Pyruvate + NADH + H+  lactate + NAD+

Using Organic Redox
BREATHALYZER
 Ethanol is oxidized to acetic acid
 Oxygen is reduced to water
CH3CH2OH(l) + O2(g) → CH3COOH(l) + H2O(l)

Electric current produced is measured
and displayed as blood alcohol content
Practice
For each reaction determine the:
1.
oxidation numbers of each atom
2.
the oxidized element
3.
the reduced element
4.
Compound containing the reduced element (oxidizing agent)
5.
Compound containing the oxidized element (reducing agent)
1.
2.
3.
4.
2Na + Cl2  2NaCl
2H2 + O2  H2O
4Fe + 3O2 → 2Fe2O3
C6H12O6 + 6O2  6CO2 + 6H2O
Balancing Redox Reactions
1. Write separate equations (halfreactions) for oxidation and reduction
2. For each half-reaction:
Balance elements involved in e- transfer
 Balance number e- lost and gained by
multiplying half reactions by whole numbers
3. Add half-reactions/cancel like terms (e-)

Acidic and Basic
4. Acidic conditions:


Balance oxygen using H2O
Balance hydrogen using H+
5. Basic conditions:
 Balance oxygen using OH Balance hydrogen using H2O
Examples

Acidic conditions:
MnO

CN
4(aq)
2
(aq)
 Fe

 Mn
acid
2
(aq)
3
(aq)
 Fe

(aq)
 MnO 2(s)
Basic conditions:

(aq)
 MnO
4 (aq)

 CNO
base
Voltaic Cells


converts chemical energy into
electrical energy
perform electrical work
battery - two or more voltaic cells
connected in series to produce a steady
flow of current
Voltaic Cells
Cathode – reduction half reaction;
positive
Anode – oxidation half reaction; negative
Electrons flow from anode  cathode
(ac)
Voltaic Cells
Half cell – one compartment of voltaic
cell
Salt bridge – electrolyte solution or
membrane which allows ions to flow to
balance overall charge of half cell
Voltaic Cells
Zn  Zn2+ + 2eoxidation
anode (-)
Cu2+ + 2e-  Cu
reduction
cathode (+)
Cell EMF
Electromotive force (EMF)
Difference in potential that pushes ethrough the external circuit
Volt (V) =
1 Joule (J)
1 Coulomb (C)
Standard Cell
Each half cell is at standard conditions:
1atm, 1M, 25°C
Cell EMF
Standard Reduction Potential (E0r)
Standard half cell potential of a reduction
reaction
 Measured in Volts relative to standard
hydrogen electrode (0 V)
 Appendix E (p.1117)
Cell EMF
Standard Cell Potential (EMF)
E0 cell = E0r cathode – E0r anode
 Intensive property – doesn’t change
with amount (like density, boiling point)
 + value spontaneous
Cell EMF
Standard Cell Potential (EMF)
 Cathode is half reaction with more
positive E0r
 More positive E0r means more likely to
be reduced (oxidizing agent)


Halogens, oxygen
Very negative E0r (reducing agent)
Cell EMF
Al
Ag1+
NO31-
E0r = -1.66 V
E0r = +0.80 V
E0r = +0.96 V
NO31- is strongest oxidizing agent
Al is strongest reducing agent
Cell EMF – calculate E0cell
Example
Cd2+ + 2e-  Cd
Sn2+ + 2e-  Sn
E0 cell
E0r = -0.403 V
E0r = -0.136 V (cathode)
= E0r cathode – E0r anode
= (-0.136) – (-0.403)
= 0.267 V (spontaneous)
Cell EMF - calculate E0r
Example
In1+  2e- + In3+
Br2 + 2e-  2 Br1E0 cell = 1.46 V
E0r = ?
E0r = +1.06 V