Transcript Introduction to Redox

Introduction to Redox
Mrs. Kay
Chemistry 12
Chapter 18 Pages:713-729
Redox Reactions
• combustion of gas in
a car
• rusting of metals
• bleaching hair
• reactions in batteries
• Cut apples turning
brown
What do you think
Redox stands for?
Redox = oxidation and reduction
reactions
• Oxidation: loss of
electrons
• Reduction: gain of
electrons
Hint:
Leo the lion says Ger
• redox reactions are a family of reactions that are
concerned with the transfer of electrons between species
• Like acid-base reactions, redox reactions are a matched
set -- you don't have an oxidation reaction without a
reduction reaction happening at the same time
Single Displacement redox reaction
• Cu(s) + 2AgNO3(aq) ---> Cu(NO3)2(aq) + 2Ag(s)
• The silver nitrate solution is transparent, when
the copper wire is placed in it, the reaction
begins slowly. The wire is coated with silver,
while the copper is broken down into ions
Ex: Mg(s) + O2(g)  MgO(s)
• Mg loses 2 electrons to become Mg+2
• O2 gains 2 electrons to become O-2
• The total reaction = redox reaction
• The half reactions = oxidation and
reduction reactions
half reactions for the overall redox
reaction
• Mg  Mg+2 + 2e- (Magnesium loses
electrons, so it is oxidized)
• O2 + 4e-  2O2- (Oxygen gains electrons,
so it is reduced)
We’re Not Finished yet!
• When combining half reactions, you must make
sure that the electrons gained = the electrons
lost. This ensures balancing of the redox
reaction. (so we multiplied the first half reaction by
2, so that the 4e- balanced out)
• 2Mg  2Mg+2 + 4e• O2 + 4e-  2O2• Total redox reaction is: 2Mg + O2  2MgO
Let’s see that again…
• The unbalanced reaction
is as follows:
• Look at each half reaction
separately: aluminum
metal being oxidized to
form an aluminum ion
with a +3 charge and
oxygen being reduced to
form two (2) oxygen ions,
each with a charge of -2.
• If we combine those
two (2) half-reactions,
we must make the
number of electrons
equal on both sides.
The number 12 is a
common
Taking care of the number of
atoms, you should end up with:
Practice Writing the half reactions
and balance the following
equations.
1.
2.
3.
4.
Fe + Br2  FeBr3
Ni + HgCl2  Hg + NiCl2
Sn+4 + Cu  Sn+2 + Cu+2
CO+ I2O5  CO2 + I2 (this one’s tricky!)
Vocabulary:
• Examine : 2Mg + O2  2MgO
• Mg oxidation # increased from 0 to +2,
therefore it has oxidized.
• O2 oxidation # decreased from 0 to –2,
therefore it has reduced.
• Oxygen has been reduced because of
magnesium, so magnesium is the
reducing agent (the atom who gave its
electrons away)
• Magnesium has been oxidized because of
oxygen, so the oxygen is the oxidizing
agent (the atom who accepted electrons)
• In redox: the reducing agent is oxidized
and the oxidizing agent is reduced.
Assigning oxidation numbers
Pg 721-726
• Actual or hypothetical charges assigned
using a set of rules.
– Used to describe redox reactions, identify
redox reactions and to identify oxidizing
agents and reducing agents.
Oxidation Numbers from Lewis
structures
• Using electronegativities
• In molecules, the more electronegative atom will
have the electrons found closer to its nucleus.
• Molecules don’t have CHARGES, but you can
assign Ox # based on which atom “owns” the
electrons more
• So the Oxygen = -2 and the Hydrogen each = +1
Oxidation Numbers
•
•
•
Molecules with the same atoms will have
no difference in electronegativity, so they
share electrons evenly.
Ex: Cl2 occurs where Cl-Cl, then pair of
electrons are shared equally, so there is
NO OWNING of electrons.
Cl will have an oxidation number of zero.
Oxidation numbers for ionic
compounds
• The oxidation number for an ionic
compound is equal to its charge
• Example
– MgO is made of Mg 2+and O2– So, Magnesium has an ox# of +2 and oxygen
has an ox# of -2.
Oxidation Number Rules: Pg 724
Rules
Example
1. A pure element has
an ox # of zero.
Ne, Na, Cl2 = 0
2. A monatomic ion ox # Ca2+ = +2
equals its charge
3. Hydrogen with a
metal = -1, Hydrogen
with a non-metal =+1
Ex: H2S : H=+1
NaH: H=-1
Oxidation Number rules
Rules
Examples
4. Oxygen is usually -2,
unless its in a peroxide
(H2O2 or OF2)
Li2O : O = -2
H2O2: O= -1
OF2: O=+2
PCl3: Cl=-1, so P=+3
CS2: S=-2, so C=+4
5. In molecules, the
more electronegative
atom usually has the
same charge that it
would have had as an
ion
Oxidation number rules
Rules
Examples
6. The sum of all ox# in CF4: C=+4, because 4
a compound or molecule F=-1 each
equal zero
7. The sum of ox# in a NO3polyatomic ion equal the
charge of the ion
Exercise: For the following
reactions
1. draw a diagram showing the loss and gain of
electrons
2. identify the substance oxidized, the
substance reduced, the oxidizing agent and
the reducing agent.
3. Write the oxidation and reduction half
reactions
Practice together:
Loss of e-
• Example:
Substance oxidized
Reducing agent
0
1,2.
0
+2
2Mg + O2  2MgO
Gain of eSubstance
reduced
3.
-2
2Mg  2Mg+2 + 4eO2 + 4e-  2O2-
Oxidizing
agent
Work on the following in class,
finish for homework:
1. Ni + HgCl2  Hg + NiCl2
2. 2Na + 2H2O  2NaOH + H2
3. Cl2 +2NaBr  Br2 + 2NaCl
4. 4NH3 + 7O2  4NO2 + 6H2O
Practice:
•
•
•
•
Page 715 # 1-4
Page 716 # 6 & 7
Page 726 # 9-12
http://staff.prairiesouth.ca/~chemistry/che
m30/6_redox/redox1_2.htm