Red uction and Ox idatio - PowerPoint Presentation

Slide1

Reduction and Oxidation

(

redox

chemistry)

Chapter 16

Slide2

What is a redox reaction?

Oxidation

is the addition of oxygen to a substance and

Reduction is the removal of oxygen from a substance.

lead oxide

+

carbon

lead

carbon monoxide



+

oxygen removed

reduction

oxygen added

oxidation

Reduction and oxidation always take place together. Why is this type of reaction called a

redox

reaction?

redox

=

red

uction and

oxidation

Which substances are oxidized and reduced in this reaction?

Slide3

Redox reactants – oxidized or reduced?

Slide4

Redox and electrons

Magnesium burns in oxygen to form magnesium oxide.

A redox reaction can also be explained in terms of the gain or loss of electrons.

magnesium

+

oxygen magnesium oxide



2Mg

(s)

O

2

(g)

2MgO(s)

+



What happens to the atoms and electrons in this reaction?

It is obvious that the magnesium has been oxidized, but what has happened to the oxygen?

Slide5

Oxidation and electron loss

When magnesium burns in oxygen to form magnesium oxide, what happens to magnesium and its electrons?

Oxidation is the loss of electrons.

oxidized

(electrons lost)



Mg

Mg

2+

O

2-

O

+

The magnesium has been oxidized.

The Mg atom has lost 2 electrons to form a Mg

2+

ion.

Slide6

Oxidation and electron gain

When magnesium burns in oxygen to form magnesium oxide, what happens to oxygen and its electrons?

Reduction is the loss of electrons.

reduced

(electrons gained)



Mg

Mg

2+

O

2-

O

+

The oxygen has been reduced.

The O atom has gained 2 electrons to form a O

2-

ion.

Slide7

Redox and OILRIG

An easy way to remember what happens to the electrons during oxidation and reduction is to think… OILRIG!

Slide8

Using OILRIG

What does OILRIG stand for in terms of redox reactions?

O

xidation

I

s

L

oss of electrons

R

eduction

I

s

G

ain of electrons

Slide9

What is a half-equation?

Redox reactions involve the transfer of electrons.

oxidation

:

Mg



Mg2+ + 2e-

magnesium

+

oxygen magnesium oxide



2Mg (s)

O

2

(g)

2MgO (s)+



reduction:

O

2 + 4e-  2O2-

Equations written to show what happens to the electrons during oxidation and reduction are called half-equations.

What are the half-equations for the oxidation and reduction processes in this reaction?

Slide10

What does each half-equation show?

Slide11

Redox reactions – summary

Slide12

Redox half equations

Slide13

Balancing half equations

Slide14

Balancing more challenging examplesBalance all atoms (but not O or H yet)

Eg

Cr

2O72-  2Cr3+Balance O by adding H2O to the other side

Eg

Cr2O72-  2Cr3+

+ 7H2OBalance H by adding H+ to the other side

Eg 14H+ + Cr2O72-

 2Cr3+ + 7H2OBalance charge by adding electrons (e-)

Eg 14H+ + 6e- + Cr2O72-  2Cr

3+ + 7H2O

Slide15

Combining half equations

Slide16

What are oxidation and reduction?

Oxidation

Reduction

addition of oxygen

e.g. 2Mg + O

2

®

2MgO

loss of oxygen

e.g. 2CuO + C

®

2Cu + CO

2

loss of hydrogen

e.g. CH

3

OH

® CH2O + H2

addition of hydrogene.g. C

2H4 + H

2  C2H6

loss of electrons

e.g. Al ® Al3+ + 3e-

gain of electrons

e.g. F

2

+ 2e

-

 2F

-

Oxidation and reduction can be used to describe any of the following processes:

Slide17

What are oxidizing and reducing agents?

Oxidizing

agents…

Reducing agents…

…oxidize other species

…are themselves reduced

…accept electrons

…reduce other species

…are themselves oxidized

…donate electrons

2NaCl + F

2

®

2NaF + Cl

2

For example, in the reaction below:

Fluorine:

oxidizes Cl

-

(to chlorine gas)

is reduced (to F

-

)

accepts electrons (from Cl

-

)

is an oxidizing agent

Slide18

Oxidizing and reducing agents

Common

oxidizing agents

:

Common

reducing agents:

sodium tetrahydrido-borate(III) (NaBH

4

)

lithium tetrahydrido-aluminate(III), (LiAlH

4

)

carbon monoxide (CO)

carbon (C)

zinc (Zn)

hydrogen (H

2

)

chlorine (Cl

2

)

manganese(IV) oxide

(MnO

2

)

potassium dichromate(VI)

(K

2

Cr

2

O

7

)

potassium manganate(VII)

(KMnO

4

)

concentrated sulfuric acid

(H

2

SO

4

)

hydrogen peroxide (H

2

O

2

)

Slide19

Spot the agent

Slide20

Oxidation numbers

Slide21

Working out oxidation numbers

Slide22

Changes in oxidation number

Oxidation numbers can be used to define the processes of oxidation and reduction.

oxidation

number

0

+1

H

2

®

2H

+

+ 2e-

During

oxidation, the oxidation number increases:

oxidation

number

+3

+2

Fe

3+

+ e- ®

Fe2+During reduction, the oxidation number decreases:

Slide23

Oxidation numbers in names

Oxidation numbers can be used in the names of compounds to indicate which oxidation state a particular element in the compound is in.

The oxidation state is usually put in brackets in roman numerals after the name of the element in question.

For example:

iron(III) hydroxide

Fe(OH)

3

iron is in oxidation state +3

iron(II) hydroxide

Fe(OH)

2

iron is in oxidation state +2

Slide24

Oxidation or reduction?

Slide25

Half equations

Chloride ions can be oxidised to produce chlorine. The

half equation

for this reaction is:

All redox reactions can be illustrated using half equations. Half equations can be combined to give the equation for the overall redox process.

Half equations are used to show the loss or gain of electrons when a species undergoes oxidation or reduction.

2Cl

–

®

Cl2

+ 2e-

-1

0

One element in a half equation changes oxidation state. Here chlorine has changed its oxidation state from

-1 to 0.

Slide26

Combining half equations

To combine half equations:

Step 1:

Write the half equations. (You may need to work these out if complex ions and other species such as H

+

are involved.)

Step 2:

Make sure that the number of electrons in each half equation is the same, so that the electrons cancel out. Do this by multiplying one or both equations to make the number of electrons the same in each case.

Step 3:

Add the half equations and cancel the electrons. It may be possible to cancel other species that appear on both sides – often H

+

or H

2O.

Slide27

Combining half equations – example

Chlorine oxidizes iron(II) to iron(III) and is itself reduced to chloride ions. Write a balanced equation for this reaction.

Step 3:

Add the half equations and cancel the electrons.

Cl

2(aq)

+ 2Fe

2+

(aq)

® 2Cl-(aq) + 2Fe3+(aq)

2Fe

2+(aq) ® 2Fe3+

(aq) + 2e-

Step 2:

Eqn.

A involves 2 electrons and Eqn. B involves 1 electron, so multiply both sides of Eqn. B by two.

chlorine has been reduced

Step 1:

Write the half equations.

Eqn. A

Eqn. B

Cl

2(aq)

+ 2e

-

®

2Cl

-

(aq)

Fe

2+

(aq)

®

Fe

3+

(aq)

+ e

-

iron(II) has been oxidized

Slide28

Combining half equations

Slide29

Comparing reactivity

The orders of reactivity of metals with water, oxygen and air can be compared.

with water

lithium

copper

gold

sodium

magnesium

silver

potassium

with oxygen

zinc

lead

magnesium

iron

copper

calcium

with acid

aluminium

iron

copper

magnesium

zinc

lead

calcium

What patterns can you see in these lists?

Slide30

The reactivity series

Combining the information from all the reactions gives an overall order of reactivity called the

reactivity series

.

One way to remember this order is to learn this silly sentence:

p

lease

sendcharlie’smonkeys

andzebrasinlargecages

securelyguarded!

Slide31

What is the order of reactivity?

Slide32

Using the reactivity series

The reactivity series can be used to make predictions about the reactions of metals.

Predictions can be made about simple reactions of metals with oxygen, water and acids.

Predictions can also be made about more complex reactions where one metal is competing with another.

potassium

sodium

calcium

magnesium

aluminiumzinc

ironleadcoppersilvergold

increasing reactivity

Slide33

Predicting simple reactions

Use the reactivity series to predict if a reaction will take place and how intense the reaction will be:

acid

gold

metal

reacts with

calcium

sodium

oxygen

oxygen

oxygen

silver

prediction

zinc

water

no reaction

fizzing

burns vigorously

very slow reaction

burns moderately

Slide34

Corrosion of metals

Metals corrode when they are left in contact with air or water.

Which metals corrode quickly and which corrode slowly?

When iron corrodes, it is called rusting.

Rusting is the oxidation reaction of iron with oxygen in the presence of water.

oxygen

iron oxide

iron

+



3O

2

Fe

2

O

3

2Fe

+



Slide35

Does salt make iron rust faster?

People who live by the seaside often claim that their cars go rusty faster.

Does salt speed up the rate of the rusting reaction?

Slide36

What is needed for iron to rust?

Slide37

Preventing rust

Rusting destroys a huge amount of iron and steel every day.

People spend a lot of money making sure that their iron and steel buildings, engines, lorries and ships do not rust.

What methods could you use to prevent things rusting?

painting

plastic coating

oiling

galvanising

Slide38

Rusting: sacrificial protection

Sacrificial protection is another way of preventing rust.

This involves attaching big blocks of magnesium or zinc to the iron hull of a ship or water pipe.

Because magnesium is more reactive than iron, it corrodes first, leaving the iron intact.

Eventually the magnesium blocks have to be replaced because they have corroded completely away.

iron

reactive metal

Slide39

Rusting: true or false?

Slide40

What to do now…..Read Ch

16 and jot down key points

Complete the following questions

16.1 TY 16.1.1, 16.1.2, 16.1.3 Q1-6 p411 Harder examples TY16.1.4, Q1,2 p41316.2 TY 16.2.1 Q1-3 p41816.3 Q1-3 p422Chapter review p423-424