Redox & Oxidation number Term 1: Week 4

The bronze statue of Great Buddha in Kamakura, Kanagawa, Japan

Outline

• Oxidation and reduction

• Strength of oxidizing and reducing agents

• Ionic equations

• Oxidation number

Oxidation and Reduction Reactions in Daily Life

Corrosion

Oxidation with hydrogen peroxide Batteries

Oxidation of apple

Combustion

Oxidation and reduction in biology

2n CO2 + 2n DH2 + photons → 2(CH2O)n + 2n DO Carbon dioxide + electron donor + light energy → carbohydrate + oxidized electron donor

Oxidation and reduction in medicine

• Oxidation is linked with the effects of aging of humans.

• To forestall the effects of oxidation, doctors recommend antioxidants - natural reducing agents such as vitamin C and vitamin E.

Jacob Tsiperovich

Oxidation and Reduction

• Originally oxidation meant “addition of oxygen” and reduction meant “removal of oxygen”.

• REDOX – a contraction of reduction-oxidation

It implies that two processes always act together

Fe2O3 loses oxygen

and is reduced

Fe2O3 (s) + 3 CO (g) 2Fe (s) + 3 CO2 (g)

CO gains oxygen and is oxidized

Oxidation and Reduction

• Oxidation is the loss of electrons.

• Reduction is the gain of electrons.

2Na0 + Cl20 2Na+Cl-

2Na 2Na+ + 2ē

Cl2 + 2ē 2Cl-

Memory aid

Oxidizing and reducing agents

2 C(s) + O2(g) CO(g)

Oxidizing and reducing agents

Sodium is oxidized – it is the reducing agent

Chlorine is reduced – it is the oxidizing agent

2Na0 + Cl20 2Na+Cl-

2Na 2Na+ + 2ē

Cl2 + 2ē 2Cl-

Active metals: Lose electrons easily Are easily oxidized Are strong reducing agents

Active nonmetals: Gain electrons easily Are easily reduced Are strong oxidizing agents

Oxidizing and reducing agents

1. Free (elemental) nonmetals become negative ions:

F2 + 2ē 2F-

O2 + 4ē 2O2-

2. Positive (usually metal) ions become neutral:

Ag+ + ē Ag

3. Higher oxidation states become lower:

8H+ + MnO4

- + 5ē Mn2+ + H2O

Fe3+ + ē Fe2+

Oxidizing agents

Strong Oxidizing Agents

• Metal Oxyacids

– Chromium Reagents (H2CrO4; K2Cr2O7 + H2SO4;

CrO3 + H2SO4)

– Manganese reagents (KMnO4)

– Osmium Tetroxide (OsO4)

• Nitric Acid and Nitrous Acid

– (HNO3, HNO2)

• Halogens

– (F2 > Cl2 > Br2 > I2)

• Forms of Oxygen and Peroxides

– (O3, H2O2)

HCl H+(aq) + Cl-(aq)

Acids as Oxidizing Agents

Hydrogen ion in hydrochloric acid can be an oxidizing agent because it can be reduced to H2.

2H+(aq) + Zn(s) Zn2+

(aq) + H2(g)

In nitric acid solution, the nitrate ion is a more powerful oxidizing agent than the hydrogen ion.

HNO3 H+(aq) + NO3

-(aq)

Cu(s)+4H+(aq)+2NO3

-(aq) Cu2+

(aq)+2NO2(g)+2H2O(l)

1. Active metals forms ions plus electrons:

Zn Zn2+ + 2ē

Na Na+ + ē

2. Nonmetals combine with other nonmetals, such as F and O, which they take from compounds with metals:

C + [O2-] CO + 2ē

3C + Fe2O3 3 CO + 2 Fe

3. Lower oxidation states become higher:

NO + 2 H2O NO3- + 4H+ + 3ē

Reducing agents

Activity Series

F2(g) + 2ē 2 F- (aq)

Fe3+ (aq) + ē Fe2+(aq)

Cu2+(aq) + 2ē Cu(s)

2H+(aq) + 2ē H2(g)

Ni2+(aq) + 2ē Ni(s)

Fe2+(aq) + 2ē Fe(s)

Zn2+(aq) + 2ē Zn(s)

Al3+(aq) + 3ē Al(s)

Li+(aq) + ē Li(s)

Best oxidizing agent

Worst oxidizing agent Best reducing agent

Worst reducing agent

Comparing oxidizing and reducing strength

𝐬𝐭𝐫𝐨𝐧𝐠𝐞𝐫𝐨𝐱𝐢𝐝𝐢𝐳𝐢𝐧𝐠

𝐚𝐠𝐞𝐧𝐭 +

𝐬𝐭𝐫𝐨𝐧𝐠𝐞𝐫𝐫𝐞𝐝𝐮𝐜𝐢𝐧𝐠

𝐚𝐠𝐞𝐧𝐭

𝐬𝐩𝐨𝐧𝐭𝐚𝐧𝐞𝐨𝐮𝐬

𝐰𝐞𝐚𝐤𝐞𝐫𝐫𝐞𝐝𝐮𝐜𝐢𝐧𝐠

𝐚𝐠𝐞𝐧𝐭 +

𝐰𝐞𝐚𝐤𝐞𝐫𝐨𝐱𝐢𝐝𝐢𝐳𝐢𝐧𝐠

𝐚𝐠𝐞𝐧𝐭

Fe2+(aq) + Cu(s) Fe(s) + Cu2+

(aq)

In which direction will reaction go spontaneously? 18

The formation of hydrogen gas in the reaction of a metal with an acid is a special case of a more general phenomenon – one element displacing (pushing out) another element from a compound by means of a redox reaction.

2H+(aq) + Zn(s) Zn2+

(aq) + H2(g)

Fe is a stronger reducing agent than Cu.

Fe + Cu2+ Fe2+ + Cu

Activity Series

F2(g) + 2ē 2 F- (aq)

Cu2+(aq) + 2ē Cu(s)

2H+(aq) + 2ē H2(g)

Fe2+(aq) + 2ē Fe(s)

Zn2+(aq) + 2ē Zn(s)

Li+(aq) + ē Li(s)

Best oxidizing agent

Worst oxidizing agent Best reducing agent

Worst reducing agent

Fe2+(aq) + Cu(s) Fe(s) + Cu2+

(aq)

Fe(s) + Cu2+(aq) Fe2+

(aq) + Cu(s)

Electron transfer

Zn

Cu

CuSO4(aq) ZnSO4(aq)

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Zn(s) Zn2+(aq) + 2ē

Cu2+(aq) + 2ē Cu(s)

Example of oxidation: corrosion

The rusting of iron is an electrochemical process that begins

with the transfer of electrons from iron to oxygen.

Fe → Fe2+ + 2 e−

Redox reaction in the presence of water:

4 Fe2+ + O2 → 4 Fe3+ + 2 O2−

Rust formation:

Fe2+ + 2 H2O ⇌ Fe(OH)2 + 2 H+

Fe3+ + 3 H2O ⇌ Fe(OH)3 + 3 H+

Dehydration:

Fe(OH)2 ⇌ FeO + H2O

Fe(OH)3 ⇌ FeO(OH) + H2O

2FeO(OH) ⇌ Fe2O3 + H2O

Corrosion

Corrosion is even faster in the

presence of salts and acids,

because these materials make

electrically conductive solutions that

make electron transfer easy.

Zinc is added to any metal

that will be submerged in

water and exposed to stray

currents to provide protection

against galvanic corrosion.

Fe2+ + Zn Fe + Zn2+

Corrosion

Gold and platinum are called

noble metals because they

are resistant to losing their

electrons by corrosion.

Other metals may lose their

electrons easily, but are

protected from corrosion by

the oxide coating on their

surface, such as aluminum

oxide.

Ionic equations

To make the essential processes of redox

reactions clearer ionic equations are employed.

• ions in solution are written separately

• only species that change are shown

(not spectator ions)

SnCl2(aq) + Fe2(SO4)3(aq) FeSO4(aq) + SnCl4(aq) Sn2+(aq) + Fe3+(aq) Fe2+(aq) + Sn4+(aq)

Half Equations

Ag+ (aq) + Cu(s) Ag(s) + Cu2+ (aq)

Ag+ (aq) + ē Ag(s)

Ag+ gains electrons, is reduced, and is the oxidizing agent.

Cu(s) Cu2+ (aq) + 2ē

Cu loses electrons, is oxidized, and is the reducing agent.

Balancing Half Equations

Oxidation: Cu(s) Cu2+ (aq) + 2ē

Reduction: Ag+ (aq) + ē Ag(s) ×2

2Ag+ (aq) + 2ē 2Ag(s)

Net: Cu(s) + 2Ag+ (aq) 2Ag(s) + Cu2+

(aq)

What if there is NO complete

electron transfer from one

substance to another?

C(s) + O2(g) CO2(g)

To overcome this problem, the

concept of OXIDATION NUMBER (ON)

was introduced.

Oxidation number (ON)

• Oxidation number – is a number assigned

to atom or an ion to describe its relative

state of oxidation or reduction.

H2O ((H+)2O-2) ON of H in H2O = +1

ON of O in H2O = -2

HCl (H+Cl-) ON of H in HCl = +1

ON of Cl in HCl = -1

Don’t misunderstand!

• Oxidation number has no structural or physical significance. It is not a charge of atom! Oxidation number is relative value of oxidation, which can be equal to the charge.

Recommended name Common (trivial) name HClO chloric (I) acid hypochloric acid FeSO4 iron (II) sulphate Fe2(SO4)3 iron (III) sulphate

used in systematic nomenclature

Advantages of Oxidation Numbers

• useful in balancing equations Oxidation: Cu(s) Cu2+ (aq) + 2ē

Reduction: Ag+ (aq) + ē Ag(s) ×2

2Ag+ (aq) + 2ē 2Ag(s)

Cu(s) + 2Ag2+ (aq) 2Ag(s) + Cu2+ (aq)

Guidelines for Determining Oxidation Numbers

1. The algebraic sum of oxidation numbers

in neutral compound must be zero; in a

polyatomic ion, the sum must be equal to

the ion charge (Al2O3 or MnO4-)

Guidelines for Determining Oxidation Numbers

2. Each atom in a pure element has an

oxidation number of zero (Cu, I2 or S8)

3. Elements of Group 1A – 3A form monoatomic

ions with positive charge and the oxidation number is equal to the group number.

Guidelines for Determining Oxidation Numbers

Example

K: 1s2 2s2 2p6 3s2 3p6 4s1 K+: 1s2 2s2 2p6 3s2 3p6 4s0

K+: [Ar] – octet stability

Guidelines for Determining Oxidation Numbers

4. The oxidation number of H is +1 and

fluorine is always -1 in compounds with

other elements.

Exceptions:

When H forms a binary

compound with a metal, the

metal forms positive ion and H

becomes a hydride ion H-

Guidelines for Determining Oxidation Numbers

5. The oxidation number of O is -2 in

most compounds

Exceptions:

Oxygen can have an oxidation

number -1 in a class of compounds called peroxides

Guidelines for Determining Oxidation Numbers

6. Cl, Br and I are always -1 in

compounds except when combined with

oxygen and fluorine

( Cl has an oxidation number -1 in NaCl, but in the ion

ClO- has an oxidation number +1)

Guidelines for Determining Oxidation Numbers

7. When there is a conflict between two of these

rules or an ambiguity in assigning an oxidation

number, apply the rule with the lower number

and ignore the conflicting rule. 1. The algebraic sum of oxidation numbers in neutral compound must

be zero; in a polyatomic ion, the sum must be equal to the ion

charge

2. Each atom in a pure element has an oxidation number of zero

3. Elements of Group 1A – 3A form monoatomic ions with positive

charge and the oxidation number is equal to the group number

4. The oxidation number of H is +1 and fluorine is always -1 in

compounds with other elements

5. The oxidation number of O is -2 in most compounds

6. Cl, Br and I are always -1 in compounds except when combined with oxygen and fluorine

1. The algebraic sum of oxidation numbers in

neutral compound must be zero; in a

polyatomic ion, the sum must be equal to

the ion charge

2. Each atom in a pure element has an

oxidation number of zero

3. Elements of Group 1A – 3A form

monoatomic ions with positive charge and

the oxidation number is equal to the group

number

4. The oxidation number of H is +1 and

fluorine is always -1 in compounds with

other elements

5. The oxidation number of O is -2 in most

compounds

6. Cl, Br and I are always -1 in compounds

except when combined with oxygen and fluorine

Example

Mg(NO3)2

(+2)+2×(ON(N)+3×(-2))=0

ON(N) = +5

Exercise

• H3PO4 ON (P) =

• Cr2O72- ON (Cr) =

• H2C2O4 ON (C) =

• NaClO3 ON (Cl) =

+5

+6

+3

+5

(+1)×3 + ON(P) + (-2)×4 = 0

2×ON(Cr) + (-2)×7 = -2

(+1)×2 + 2×ON(C) + (-2)×4 = 0

(+1) + ON(Cl) + (-2)×3 = 0

WELL DONE!!!

October's MONTHLY QUIZ is at 8 a.m. Wednesday. Groups F and G in 2/302 D, E, H, I in 5/103 A, B, C in 3/143 Turn up in good time with calculator, pencil and ruler. Periodic tables will be provided.

When the student is ready, the master appears.

Buddhist Proverb