Transcript 4.4 Redox Reactions

Acid-Base titration
Objective: Determine the equivalence point.
Equivalence point nOH- = nH+
If 25.00mL of 0.0800M NaOH is needed to react
with 10.00 mL of HCl. What is the molarity of
HCl?
1.
2.
Write the reaction
Use M1V1=M2V2
In this case the formula could be written MOH- VOH- = MH+VH+
4.4 Redox
Reactions
Tracking electrons in oxidation
reduction reactions.
Oxidation-Reduction Reactions
(Red-Ox)
A type of electrochemistry with reactions that
involve a transfer of electrons between atoms.
 The charge or oxidation number changes
Ex: Fe2+  Fe 3+ + 1 eIn this red-ox ½ reaction the Iron (II) lost an
electron to become Iron (III).
 The electron does not just “go into space” so
there must be a companion ½ reaction that
gains an electron to make the reaction balance.

Oxidation Number
The oxidation number of an element in a
molecule is the charge that it would have if
all shared electrons were assigned to the
more electronegative elements in their
bonds.
 Draw the Lewis diagram for water
 Assign the bonding electrons to the more
En atoms.
 Determine what the charge would be.

Rules for assigning oxidation numbers.
1.
The oxidation state of an element is zero.

2.
The oxidation state of a monatomic ion is the same as its charge.

3.
4.
7.
Exceptions to this rule include peroxides (compounds containing the group),
where each oxygen is assigned an oxidation state of -1, as in hydrogen peroxide
(H2O2).
Hydrogen is assigned an oxidation state of +1.

6.
Fe2+ oxidation number is 2+
In compounds, Group 1 is +1, Group 2 is +2, Fluorine is -1
Oxygen is usually assigned an oxidation state of -2 in its covalent
compounds.

5.
including all elemental forms of the elements (N2, P4, S8, O2, O3).
Metal hydrides are an exception In LiH, hydrogen has an oxidation state of -1.
The sum of the oxidation states must be zero for an electrically neutral
compound.
For a polyatomic ion, the sum of the oxidation states must equal the charge
of the ion.
Assigning Oxidation Numbers
Cl2
 Sodium Metal
 Lead IV
 NaF
 CaCl2
 H2SO4
 CrO42 NO3
Red-ox reactions
A spontaneous red-ox reaction can be used
to perform electrical work.
 Fuel
Cells
 Primary voltaic cells (alkaline batteries)
 Storage Cells (car batteries)
Balancing Redox equations
ClO3- (aq) + I-
(aq)

Cl-
(aq)
+ I2 (s)
The reaction occurs in an acidic solution.
Split in to 2 half-reactions. An oxidation and a
reduction.
Balancing ½ reactions
1.
2.
3.
Balance the atoms of the element being
oxidized or reduced.
Balance the oxidation number by balancing
electrons.
Balance the charge. (add H+ if in acidic solution; add OH- if in
basic solution)
4.
5.
Balance Hydrogen by adding H2O molecules.
Check to make sure that oxygen is balanced.
I-
(aq)
 I2
(s)
(in acidic solution)
1.
2.
Balance the atoms of the
element being oxidized or
reduced.
Balance the oxidation
number by balancing
electrons. Oxidation ½ add electrons
to the right side
3.
Balance the charge. (add H+ if in
acidic solution; add OH- if in basic
solution)
4.
5.
Balance Hydrogen by adding
H2O molecules.
Check to make sure that
oxygen is balanced
ClO3-
(aq)
 Cl-
(aq)
1.
2.
(in acidic solution)
Balance the atoms of the
element being oxidized or
reduced.
Balance the oxidation
number by balancing
electrons. Reduction ½ adds e- to the
left side.
3.
Balance the charge. (add H+ if in
acidic solution; add OH- if in basic
solution)
4.
5.
Balance Hydrogen by adding
H2O molecules.
Check to make sure that
oxygen is balanced
Combining ½ reactions

Combine so that the number of e- cancel.


Multiply by coefficients to accomplish this.
Add together and eliminate what appears
on both sides.
The following reaction occurs in acidic solution:
Fe2+ (aq) + MnO4- (aq)  Fe3+ (aq) +
Mn2+ (aq)
The following reaction occurs in basic solution:
Cl2 (g) + Cr(OH)3 (s)  Cl- (aq) +
CrO42- (aq)