Leo

(1)

(2)

(3)

LEO

says

(4)

LEO

says

L

oss of Electrons = Oxidation

Gain of Electrons =

R

eduction

(5)

Oxidation Numbers

• Oxidation is the loss of electrons;

Reduction is the gain of electrons

• Oxidation and reduction go together.

Whenever a substance loses electrons

and another substance gains electrons

• Oxidation Numbers are a system that

(6)

Oxidation Numbers

Oxidation numbers always refer to

single atoms

The oxidation number of an

uncombined element is always 0

O

₂

, H

₂

, Ne

Zn

The oxidation number of Hydrogen

is usually +1

Hydrides are an

exception They are -1

HCl, H

₂

SO

₄

The oxidation number of Oxygen is

usually -2

Peroxides are an exception

They are –1 In OF

₂

oxygen is a +2

H

₂

O, NO

2

_{, et}

Oxidation numbers of monatomic

ions follow the charge of the ion

O

2-

_{, Zn}

2+

The sum of oxidation numbers is

zero for a neutral compound. It is

the charge on a polyatomic ion

(7)

2-Practice Assigning

Oxidation Numbers

NO

₂

N

₂

O

₅

HClO

₃

HNO

₃

Ca(NO

₃

)

₂

(8)

Practice Assigning

Oxidation Numbers

Fe(OH)

₃

K

₂

Cr

₂

O

₇

CO

₃

CN

(9)

Practice Assigning

Oxidation Numbers

NO

₂

N= +4, O = -2

N

₂

O

₅

N = +5, O = -2

HClO

₃

H=+1, Cl=+5, O = -2

HNO

₃

H=+1, N = +5, O = -2

Ca(NO

₃

)

₂

Ca=+2, N =+5, O= -2

(10)

Practice Assigning

Oxidation Numbers

Fe(OH)

₃

Fe =+3, O=-2, H=+1

K

₂

Cr

₂

O

₇

K=+1, Cr=+6, O=-2

CO

₃2-

C=+4, O =-2

CN

-

C=+4, N=-5

(11)

Using Oxidation Numbers

• Careful examination of the oxidation

numbers of atoms in an equation allows

us to determine what is oxidized and

(12)

Using Oxidation Numbers

• An

increase

in the oxidation number indicates

that an atom has

lost electrons

and therefore

oxidized.

• A

decrease

in the oxidation number indicates

that an atom has

gained electrons

and therefore

reduced

• Example

Zn + CuSO

₄



ZnSO

₄

+ Cu

0 +2

+6-2

+2

+6-2

0 Zn:

0 

+ 2



Oxidized

(13)

Exercise

For each of the following reactions find

the element oxidized and the element

reduced

Cl

₂

+ KBr



KCl + Br

₂

Cu + HNO

₃



Cu(NO

₃

)

₂

+ NO

2

+ H

2

O

(14)

Exercise

For each of the following reactions find

the element oxidized and the element

reduced

Cl

₂

+ KBr



KCl + Br

₂

0 +1

-1

+1

-1

0 Br increases from –1 to 0 -- oxidized

Cl decreases from 0 to –1 -- Reduced

(15)

Exercise

For each of the following reactions find

the element oxidized and the element

reduced

Cu + HNO

₃



Cu(NO

₃

)

₂

+ NO

2 + H

2 O

0 +1

+5

-2

+2

+5-2

+4

–2 +1-2

• Cu increases from 0 to +2. It is oxidized

• Only part of the N in nitric acid changes from +5

to +4. It is reduced

• The nitrogen that ends up in copper nitrate

(16)

Exercise

For each of the following reactions

find the element oxidized and the

element reduced

HNO

₃

+ I

₂



HIO

₃

+ NO

2

1 +5

-2 0

+1

+5

-2

+4

-2

• N is reduced from +5 to +4. It is reduced

• I is increased from 0 to +5 It is oxidized

(17)

Oxidation-Reduction

Reactions

• _{All oxidation reduction reactions}

have

one element oxidized

and one

element reduced

• Occasionally the same element may

undergo both oxidation and

reduction. This is known as an

(18)

Balancing Redox

Reactions

• Many chemical reactions involving

oxidations and reductions are complex

and very difficult to balance by the

“guess and check” methods we learned

earlier.

(19)

Balancing Redox

Reactions

There are several basic steps

1. Assign oxidation numbers to the species in the

reaction

2. Find the substance oxidized and the substance

reduced

3. Write half reactions for the oxidation and reduction

4. Balance the atoms that change in the half reaction

5. Determine the electrons transferred and balance the

electrons between the half reactions

6. Combine the half reactions and balance the remaining

atoms

(20)

Balancing Redox Equations 1

Cu + HNO

₃



Cu(NO

₃

)

₂

+ NO

+ H

₂

O

1. Assign oxidation numbers

to the species in the reaction

2. Find the substance

oxidized and the substance reduced

3. Write half reactions for the oxidation and reduction 4. Balance the atoms that

change in the half reaction 5. Determine the electrons

transferred and balance the electrons between the half reactions

6. Combine the half reactions and balance the remaining atoms

(21)

Balancing Redox Equations 2

HNO

₃

+ I

₂



HIO

₃

+ NO

₂

+ H

₂

O

1. Assign oxidation numbers

to the species in the

reaction

2. Find the substance oxidized

and the substance reduced

3. Write half reactions for the

oxidation and reduction

4. Balance the atoms that

change in the half reaction

5. Determine the electrons

transferred and balance the

electrons between the half

reactions

6. Combine the half reactions

and balance the remaining

atoms

(22)

Balancing Ionic Redox Equations 3

Fe

2+

+MnO

4-



Mn

2+

+Fe

3+ (acidic) 1. Assign oxidation numbers

6. Combine the half reactions and balance the remaining atoms. You may need to add H+ or OH- and H₂O in ionic equations

(23)

Balancing Ionic Redox Equations 4

Br

₂



BrO

₃-

+ Br

-

_(basic)

1. Assign oxidation numbers

6. Combine the half reactions and balance the remaining atoms. You may need to add H+ or OH- and H₂O in ionic equations

(24)

Balancing Ionic Redox Equations 5

VO

₂+

+ Zn



VO

2+

+ Zn

2+

(Acidic)

1. Assign oxidation numbers to the species in the

reaction

6. Combine the half reactions and balance the remaining atoms. You may need to add H+_{or OH}-_{and H}

2O in ionic equations

(25)

(26)