Redox reactionsRedox reactions are everywhere! Our body uses redox reactions to convert

food and oxygen to energy plus water and CO2 which we then exhale. The

electronics batteries also rely on redox reactions.

An example and important termsRedox reactions have some terms you should be familiar with.

We will go over these terms using the following example reaction:

2Fe2 O3(s) + 3C(s) → 4Fe(s)+3CO2 (g)

Questions to answer:

1. Is this reaction a redox reaction, and how do we know?

2. If this is a redox reaction, what is being reduced or oxidized?

Remember OIL RIG

3. What is the reducing agent in this reaction?

4. What is the oxidizing agent in this reaction?

5

Question 1:

Yes, if there is transfer of electrons which can be checked to know whether

change in oxidation number occurred from reactant to products.

↓ ↓ ↓ ↓ ↓↓

+3,−2 0 0 +4,−2 (Oxidation numbers)

The oxidation numbers can be use to answer question 1, because we can

show that the oxidation numbers for carbon and iron changed during the

reaction by transfer of electrons.

2Fe2O3(s ) + 3C(s) → 4Fe(s)+3CO2 (g)

Question 2

Carbon being oxidized because it is losing electrons as the oxidation number

increases from 0 to +4. Iron is being reduced because it is gaining electrons

when the oxidation number decreases from +3 to 0.

Question 3

The reducing agent is the reactant that is being oxidized (and thus causing

something else to be reduced), so C is the reducing agent.

Question 4:

The oxidizing agent is reactant that is being reduced (and thus causing

something else to be oxidized), so Fe2 O3(s) is the oxidizing agent.

Common types of redox reactions

1. A combustion reaction is a redox reaction between a compound and

molecular oxygen (O2). A hydrocarbon burns in oxygen (reactants) to give the

(products) carbon dioxide and water.

Examples of combustion reactions are as follows: Octane, a hydrocarbon is a

component of gasoline, and this combustion reaction occurs in the engine of

most cars: 2C8H18+25O2 → 16CO2 (g)+18H2O

Combustion of octane

C3H8+5O2 → 3CO2+4H2O C5H12+8O2 → 5CO2+6H2O

Combustion of propane Combustion of pentane

2. Disproportionation reactions

A disproportionation reaction (or auto-oxidation reaction) is a reaction in

which a single reactant is both oxidized and reduced. The following reaction is

for the disproportionation of hypochlorite, ClO−

3ClO−(aq)→ClO3−(aq)+2Cl−(aq)

When analyzed, it is seen that the oxidation numbers for chlorine, the reactant

ClO− is being oxidized to ClO3− (increase in oxidation number +1 to +5. At the

same time, the chlorine in some other molecules of ClO− are being reduced to

Cl− (decrease in oxidation number +1 to -1). Oxygen has an oxidation number

of −2 in both ClO− and ClO3− so it does not get oxidized or reduced in the

reaction.

3. Single replacement reactions

A single replacement reaction (or single displacement reaction) involves

two elements trading places within a compound. For example, many metals

react with dilute acids to form salts and hydrogen gas. The following

reaction shows zinc replacing hydrogen in the single replacement reaction

between zinc metal and aqueous hydrochloric acid:

Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2 (g)

Oxidation-reduction of Organic Compounds

An organic compound commonly is said to be "reduced" if reaction leads to

an increase in its hydrogen content or a decrease in its oxygen content. The

compound would be "oxidized" if the reverse change takes place:

This is a very unsatisfactory definition because many oxidation- reduction

or redox reactions do not involve changes in hydrogen or oxygen content, as

the following example illustrates:

CH3Cl + Mg CH3-Mg-Cl (carbon is reduced; magnesium is oxidized)

Redox reactions are better defined in terms of the concept of electron

transfer. Thus an atom is said to be oxidized if, as the result of a reaction, it

experiences a net loss of electrons; and is reduced if it experiences a net gain

of electrons.

This simple definition can be used to identify oxidation or reduction processes at

carbon in terms of a scale of oxidation states for carbon based on the

electronegativities of the atoms attached to it. The idea is to find out whether in a

given reaction carbon becomes more, or less, electron-rich. We will use the

following somewhat arbitrary rules:

1. Elementary carbon is assigned the zero oxidation state.

2. The oxidation state of any chemically bonded carbon may be assigned

by adding -1 for each more electropositive atom and +1 for each more

electronegative atom, and 0 for each carbon atom bonded directly to the

carbon of interest. That is,

-1 for electropositive atoms, H, B, Na, Li, Mg

+1 for electronegative atoms, halogens, 0, N, S

0 for carbon.

The rationale for this mode of operation can be seen if we look more closely at the example of CH3Cl + Mg CH3-Mg-Cl

Chlorine is more electronegative than either carbon or magnesium. Carbon is more electronegative than magnesium.

δ+ δ-

Thus CH3C1 is written properly with a polar bond as CH3----C1,

δ- δ+

whereas the C-Mg bond is oppositely polarized, CH3----Mg. If all of the

bonds were ionized completely, we could write

and it would be completely clear that carbon gains two electrons (is reduced),

while magnesium loses two electrons (is oxidized). But because covalent, or

at most polar, bonds actually are involved, it is much more difficult to determine

whether oxidation or reduction occurs.

3. In compounds with multiple bonds ( C=O, -C=N), the attached

heteroatom is counted twice or three times, depending on whether the bond is

double or triple.

4. A formal positive charge on carbon changes the oxidation state by + 1, and

a formal negative charge by -1 ; an odd electron on carbon leaves the oxidation

state unchanged.

To illustrate, the oxidation state of carbon in four representative examples

is determined as follows:

Any reaction that increases the degree of oxidation of carbon corresponds to

a loss of electrons (oxidation), and a reaction that decreases the oxidation

level corresponds to a gain of electrons (reduction). Two examples follow:

The terminology "redox" should not be confused with the mechanism of a

reaction, as there is no connection between them. A moment's reflection also

will show that virtually all reactions theoretically can be regarded as redox

reactions, because in almost every reaction the reacting atoms experience some

change in their electronic environments.

Traditionally, however, reactions are described as redox reactions of carbon only

when there is a net change in the oxidation state of the carbon atoms involved.

An indication of just how arbitrary this is can be seen by the example of addition

of water to ethene.

This reaction usually is not regarded as an oxidation-reduction reaction

because there is no net change in the oxidation state of the ethene carbons,

despite the fact that, by our rules, one carbon is oxidized and the other reduced:

Apart from indicating when oxidation or reduction occurs, the oxidation

scale is useful in balancing redox equations. For example, consider the

following oxidation of ethenylbenzene (styrene) with potassium

permanganate:

To determine how many moles of permanganate ion are required to oxidize

one mole of styrene in this reaction, first determine the net change in

oxidation state of the reacting carbons:

Second, determine the net change in oxidation state of manganese for

MnO4 MnO2:

Mn(VI1) Mn(1V) net change = 3

Therefore 3 moles of styrene is needed for every 8 moles of permanganate:

To get the overall atom and electrical balance for Equation 1, the

requisite amounts of H20 must be added, but the 3:8 ratio will remain

unchanged:

Because KOH reacts in a nonoxidative way with carboxylic acids to form

carboxylate salts (RC02H + KOH RC02K + H2O), the final

equation is

AssignmentFor each of the following reactions determine the oxidation state of the

carbons in the reactants and products and decide whether the overall

changes involve oxidation, reduction, or neither.

1. CH4 + CI2 CH3CI + HCI

2. CH3CH=CH2 + HCI CH3CHCH3

Cl3. (CH3)2C=CH2 + (CH3)3CH (CH3)2CHCH2C(CH3)3

4. CH3OH CH2=O + H2

References

Basic Principles of Organic Chemistry. 2nd Ed. by

John D. Roberts and Marjorie C. Caserio

https://www2.estrellamountain.edu/faculty/farabee/biobk/BioBookEnzym.html

https://www.onlinemathlearning.com/redox-reaction.html

https://www.docsity.com/en/redox-reaction-lecture-notes/1981176/

https://www.onlinemathlearning.com/redox-reaction.html