Lecture 33 Highlights…
aNew Subject Today: The spellbinding and
electrifying world of electrochemistry
aOxidation and Reduction
aRedox Reactions
aLot’s of Definitions !!
Redox Reactions,
Electrochemistry and Batteries
Chemistry
The Science in Context
Chapters 5.5 and 17
1
Redox Reactions
Definition: Reactions where there is a simultaneous
transfer of electrons from one chemical species to
another….
9Oxidation: Loss of electrons
9Reduction: Gain of electrons
9Half Reactions: Two reactions
which combine to give an
overall “Redox” reaction
Oxidation
9Defined as:
• Loss of electrons (LEO)
• Gain of Oxygen
• Loss of Hydrogen
The last two definitions are useful in
organic chemistry and combustion
reactions.
Oxidation – Loss of Electrons
9 For example, the reaction of sodium metal
with chlorine gas to form sodium chloride
(NaCl). Sodium loses an electron as follows:
Na(s) Æ Na+ + eCl2 (g) + 2e- → 2 Cl-
When sodium metal loses the electron,
chemists say that the sodium metal has been
oxidized to the sodium cation (positively
charged ion.)
2
Oxidation – Gain of Oxygen
9 Sometimes, it is more obvious that oxygen
has been gained instead of electrons in going
from reactants to products. Such is the case
of combustion/burning of coal and rusting of
a metal:
C(s) + O2(g) Æ CO2(g)
2Fe(s) + 3O2(g) Æ 2Fe2O3(s)
Oxidation – Loss of Hydrogen
9In other reactions, oxidation can best be seen
as the loss of hydrogen:
CH3OH(l) Æ CH2O(l) + H2(g)
In going from methanol to formaldehyde, the
compound went from having four hydrogen
atoms to having two hydrogen atoms.
Reduction
9Defined as:
• Gain of electrons (GER)
• Loss of Oxygen
• Gain of Hydrogen
3
Reduction – Gain of Electrons
9Reduction is often seen as the gain of
electrons:
Ag+ + e- Æ Ag(s)
When the silver cation gains the
electron, chemists say that it has been
reduced to silver metal.
Reduction – Loss of Oxygen
9Sometimes, it is more obvious that oxygen
has been lost in going from reactants to
products:
Fe2O3(s) + 3CO(g) Æ 2Fe(s) + 3CO2(g)
The iron lost oxygen, so chemists say that the
iron ion has been reduced.
Reduction – Gain of Hydrogen
9In other reactions, oxidation can best be seen
as the gain of hydrogen:
CO(g) + 2H2(g) Æ CH3OH(l)
In going from carbon monoxide to methanol,
the compound gained four hydrogen atoms.
4
Oxidation - Reduction
9Neither can take place without the other.
Consider the reaction between zinc metal and
an aqueous copper (II) sulfate solution:
Zn(s) + Cu2+ Æ Zn2+ + Cu(s)
9The overall reaction may be decomposed into
two half-reactions:
Zn(s) Æ Zn2+ + 2e- (oxidation – loss of electrons)
Cu2+ + 2e- Æ Cu(s) (reduction – gain of electrons)
REDOX Agents
9The Oxidizing Agent is Reduced.
9The Reducing Agent is Oxidized.
Zn(s) Æ Zn2+ + 2e- (reducing agent – it is oxidized)
Cu2+ + 2e- Æ Cu(s) (oxidizing agent – it is reduced)
Keeping Track of Electrons ….
Bookkeeping by Oxidation Numbers
9The oxidation number of any element is
zero including diatomic elements (H2).
9In a simple compound the oxidation
number of an atom is the same as its
charge.
ƒ NaCl: Na oxidation number = +1, -1 for Cl
ƒ In H2S: H = +1, S = -2
ƒ In general, alkaline metals are +1 (IA family),
alkaline earths are +2 (IIA family)
9The sum of the oxidation numbers in a
neutral compound = zero.
5
Oxidation Numbers-II
9 O is usually -2 except “peroxides” (O is -1 in
H2O2, for example). H is +1 bonded to
nonmetals, -1 bonded to metals (in metal
hydrides, NaH).
9 In a covalent compound the oxidation
number of an atom indicates the number of
electrons partially lost or gained.
9 The sum of the oxidation numbers in a
polyatomic ion = its charge.
ƒ Ex. chromate: CrO42- : Cr = +6, O = 4 x -2; Net = -2
ƒ Or, H3O+ : H = 3 x +1, O = -2; Net =+1
9 Cl, Br and I are usually -1 (F is always -1), unless
they are combined with oxygen (Cl is +1 in ClO-).
Oxidation Numbers-III
9So, in an electrochemical reaction, if an
atom’s oxidation number increases (say
from 0 to +2 or from +2 to +3), then it
has been oxidized.
9If an atom’s oxidation number decreases
after a reaction, then it has been
reduced.
• Example: Mg(s) + FeCl2 (aq) → MgCl2(aq) + Fe (s)
• Ox. Num. = 0
+2, -2
+2, -2
0
• Mg is oxidized, Fe is reduced and chlorine is a socalled spectator ion !
• Why not the reverse reaction ?
Thanks..
Brown,
LeMay,
Bursten
6
Whiteboard Problems
Balancing Redox Equations
Al(s) + HCl(aq) Æ AlCl3(aq) + H2(g)
1.
2.
3.
Find and write the half-reactions:
Al Æ Al3+ (it is oxidized – loses electrons)
H+ Æ H2 (it is reduced – gains electrons)
Balance the atoms in the half-reactions:
Al Æ Al3+
2H+ Æ H2
Balance the charges in the half-reactions by adding
electrons:
Al Æ Al3+ + 3e2H+ + 2e- Æ H2
Balancing Redox Equations
4.
5.
6.
Make the number of electrons lost = the number
gained by multiplying one or both half-reactions by a
stoichiometric factor:
2 x (Al Æ Al3+ + 3e-)
3 x (2H+ + 2e- Æ H2)
Add the half-reactions to arrive at the balanced net
ionic equation
2Al Æ 2Al3+ + 6e6H+ + 6e- Æ 3H2
6H+ + 2Al Æ 2Al3+ + 3H2
Add the spectator ion back in and associate the ion
to arrive at the balanced molecular equation
6H+ + 2Al Æ 2Al3+ + 3H2
6ClÆ 6Cl2Al(s) + 6HCl(aq) Æ 2AlCl3(aq) + 3H2(g)
7