Oxidation and Reduction (REDOX) reactions?
1. A reaction in which electrons are transferred from
one atom to another is called an oxidation and
reduction reaction or REDOX reaction.
Rules for Assigning Oxidation Numbers
The charge the atom would have in a molecule
(or an ionic compound) if electrons were completely transferred.
1. Free elements (uncombined state) have an oxidation
number of zero. All Diatomics have an oxidation state
of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
2. In monatomic ions, the oxidation number is equal to
the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
3. The oxidation number of oxygen is –2.
Exception: In H2O2 it is –1.
4. The oxidation number of hydrogen is +1 except when
it is bonded to metals in binary compounds. In these
cases, its oxidation number is –1.
5. Group IA metals are +1, IIA metals are +2 and fluorine is
always –1.
6. The sum of the oxidation numbers of all the atoms in a
molecule or ion is equal to the charge on the
molecule or ion.
HCO3Oxidation numbers of all
the atoms in HCO3- ?
How to write oxidation numbers using rules?:Biochemhelp
Find the oxidation state of Cr in
K2Cr2O7
Element
K
Cr
O
Total Charge of the
compound has to
equal ZERO
Subscript
Oxidation
State of 1
atom
Sum of
oxidation
states
0
Oxidation-Reduction Reactions
“REDOX” reactions
1. Determine the oxidation number of the
boldface underlined element in the following
formulas:
NaClO4
NH4+
HCO3-
Concept Check
1. Give the oxidation number for the following
atoms:
Cr3+ Cr = _____
O2 = _____
Cl-1 Cl = _____
Mg = _____
2. Give the oxidation number for the following
atoms:
AgNO3 N = _____
NH3
N = _____
Concept Check
Give the oxidation numbers for the atoms in
(NH4)2CrO4
NH4NO3
Redox Reactions
1. OXIDATION — loss of electron(s) by a substance.
Increase in oxidation number. (Mg  Mg2+ + 2e-)
2. REDUCTION — gain of electron(s) by a substance.
Decrease in oxidation number. (S+ 2e-  S2-)
3. OXIDIZING AGENT — electron acceptor and the
substance is reduced. (Non-metals, Cl atom gains 1 e-)
4. REDUCING AGENT — electron donor and the
substance is oxidized. (Metals, Na atom loses 1 e-)
How would you classify most metals and nonmetals?
You can’t have one… without the other!
1. Reduction (gaining electrons) can’t happen
without an oxidation to provide the electrons.
2. You can’t have 2 oxidations or 2 reductions in
the same equation. Reduction has to occur
when there is oxidation.
LEO the lion says GER!
o l x
s e i
e c d
t a
r t
o i
n o
s n
GER!
a l e
i e d
n c u
t c
r t
o i
n o
s n
Another way to remember
OIL
RIG
Oxidation is loss
(oxidation number increases)
Reduction is gain
(oxidation number decreases)
LEO
OXIDATION
1. Loss of e2. Increase in
oxidation number
3. Reducing Agent
GER
REDUCTION
1. Gain of e2. Decrease or
reduction in
oxidation number
3. Oxidizing agent
Reaction of zinc with hydrochloric acid - YouTube
Zn + 2HCl  ZnCl2 + H2
Examining Redox Reactions
1. After oxidation numbers have been
assigned, it can be determined that the
reaction may or may not have been redox.
2. Is this reaction a redox reaction? Explain why or
why not.
NaCl + AgNO3  AgCl + NaNO3
Redox Reactions
Oxidation-Reduction Basics - YouTube
4 Fe(s) + 3 O2(g)  2 Fe2O3(s)
1. Assign oxidation numbers to all atoms.
2. Identify what is oxidized and what is reduced.
3. Identify the oxidizing and reducing agent.
Concept Check
1. Give the oxidation number and identify what
is oxidized, reduced, and identify the
oxidizing and reducing agents.
Cu + 2HNO3  Cu(NO3)2 + H2
Oxidation and Reduction half-reactions
Half-reaction
1. Shows either the oxidation or reduction part of a
redox reaction.
2. Shows the electrons gained or lost.
3. Follow the law of conservation of matter:
a) Mass of atom(s) on reactant side = mass of
atom(s) on the product side.
b) Conservation of charge: net charge must be the
same on both sides of the equation, but does not
necessarily equal zero.
Half Reactions
Oxidation half-reaction
1. Shows an atom or
ion losing 1 or more
electrons while its
oxidation number
increases.
Fe(s)  Fe3+(aq) + 3e-
Reduction half-reaction
1. Shows an atom or
ion gaining 1 or
more electrons
while its oxidation
number decreases.
Fe3+(aq) + 3e-  Fe(s)
For oxidation, the e- is
on the product side.
For reduction, the e- is
on the reactant side.
Write the Half-reaction from an
equation
Cu(s) + AgNO3(aq)  Cu(NO3)2(aq) + Ag(s)
Step 1: Assign an oxidation number to each element.
Step 2: Write a partial half-reaction to show the change in
oxidation state. (no e- yet)
Step 3: Now show the number of e- needed to explain how the
oxidation number changed.
Step 4: Keep conservation of mass and charge for ½ reactions.
Step 5: There must be a balance between the number of
electrons lost and gained. Balance by multiplying through by
the lowest common denominator.
Step 6: Once the e- lost and gained are equal, we can cancel
the e- on both sides and add the 2 half reactions.
1. A half reaction does not occur by itself.
2. At least two such reactions must occur so that
the electron released by one reactant is accepted
by another in order to complete the reaction.
3. Thus, oxidation and reduction reactions must
take place simultaneously in a system.
2Mg (s) + O2 (g)
2Mg
O2 + 4e-
2MgO (s)
2Mg2+ + 4e- Oxidation half-reaction (lose e-)
2O2-
Reduction half-reaction (gain e-)
19.1
Half reactions
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
Zn(s)  Zn2+(aq) + 2e- OXIDATION ½ reaction
Cu2+(aq) + 2e-  Cu
REDUCTION ½ reaction
Concept Check
Balance the following reactions using the half-reaction
method.
Li(s) + CuSO4(aq)  LiSO4(aq) + Cu(s)
Ag(s) + S(s)  Ag2S(s)
AgS animation
The Galvanic or Voltaic Cell
Zn/Cu
Cu/Ag
Cell
Construction
Salt bridge –
KCl in agar
Provides conduction
between half-cells
Observe the
electrodes to see
what is occurring.
Cu
1.0 M CuSO4
Zn
1.0 M ZnSO4
What about half-cell
reactions?
What about the sign of the
electrodes?
-
+
cathode half-cell
Cu+2 + 2e-  Cu
Cu
plates out or
deposits on
electrode
Why?
anode half-cell
Zn  Zn+2 + 2e-
Cu
1.0 M CuSO4
What
happened at
each
electrode?
Zn electrode
erodes
or dissolves
Zn
1.0 M ZnSO4
Galvanic cell
• cathode half-cell (+)
REDUCTION
Cu+2 + 2e-  Cu
• anode half-cell (-)
OXIDATION
Zn  Zn+2 + 2e-
• overall cell reaction
Zn + Cu+2  Zn+2 + Cu
Spontaneous reaction that produces electrical current!
Now for a standard cell composed of
Cu/Cu+2 and Zn/Zn+2, what is the voltage
produced by the reaction at 25oC?
Standard Conditions
Temperature - 25oC
All solutions – 1.00 M
All gases – 1.00 atm
Now replace the light bulb with a volt meter.
+
-
1.1 volts
cathode half-cell
Cu+2 + 2e-  Cu
anode half-cell
Zn  Zn+2 + 2e-
Cu
1.0 M CuSO4
Zn
1.0 M ZnSO4
We need a standard electrode to
make measurements against!
The Standard Hydrogen Electrode (SHE)
25oC
1.00 M H+
1.00 atm H2
H2 input
1.00 atm
Half-cell
2H+ + 2e-  H2
EoSHE = 0.0 volts
Pt
inert
metal
1.00 M H+
Electrode - an
electrical conductor
which carries charge
to or from a liquid
undergoing
electrolysis.
Lead bromide
Electrolysis
Electrolysis is the
break-down of a
substance by
electricity
Electrolyte - a molten
or aqueous solution
through which an
electrical current can
flow.
Electrolysis experiments
•
•
•
Electrolysis only happens in:
- molten ionic liquids or
- aqueous solutions containing ions.
There must be a complete circuit.
A lamp or ammeter shows that
electricity is flowing around the circuit.
Electrolysis of zinc chloride
At the electrodes
Cathode (-)
(negative electrode)
• Positive ions go here
(cations).
• As metal ions are positive,
they go to the cathode.
• Ions gain electrons. They
are reduced and become
neutral atoms.
Anode (+)
(positive electrode)
• Negative ions go here
(anions).
• As non-metal ions are
negative, they go to the
anode.
• Ions lose electrons. They are
oxidised and become neutral
atoms
(which react together to form
molecules).