Oxidation-Reduction Reactions

Oxidation and ReductionOxidation-reduction reactions always occur

simultaneoulsy.Redox Reactions

OxidationLoss of electrons Gain oxygen

2Fe2O3 + 3C2 4Fe + 3CO2

Reduced Oxizided

ReductionGain ElectronsLoss of Oxygen

Redox Reactions that Form IonsBetween metal and nonmetals, electrons are

transferred from the metal to the nonmetal.Increases stabilityMg + S Mg2+ + S2-

Oxidation: Mg Mg2+ + 2e- (loss of electrons)Reduction: S + 2e- S2- (gain of electrons)

Oxidizing and Reducing AgentsReducing Agent: substance that loses

electrons.Oxidizing Agent: substance that accepts the

electrons is the oxidizing agent.Mg + S Mg2+ + S2-

Mg: reducing agent, oxidizedS: oxidizing agent, reduced

LEO the lion goes GERLEO: Losing Electrons is OxidationGER: Gaining Electrons is Reduction

Redox with Covalent CompoundsIn covalent compounds, polar molecules

involve unequal sharing of electronsThe shift in electrons is redox for it is the

partial gain and loss of electronsH2OOxygen: electrons shift toward

Reduced, oxidizing agentHydrogen: electrons shift away

Oxidized, reducing agent

Processes Leading to RedoxOxidation

Complete loss of electrons (ionic reactions)

Shift of electrons away from an atom in covalent bond

Gain of OxygenLoss of Hydrogen by a

covalent compoundIncrease in oxidation

number

ReductionComplete gain of

electrons (ionic reactions)

Shift of electrons toward from an atom in covalent bond

Loss of OxygenGain of Hydrogen by a

covalent compoundDecrease in oxidation

number

CorrosionIron, corrodes by being oxidized to ions of

iorn by oxgyen2Fe +O2 + H2O 2Fe(OH)2

To protect iron, a piece of magnesium is placed in electrical contact.

When oxygen or water attack the iron object, iron lose electrons.

Because Mg is more easily oxidized, the Mg immediately transfers electrons to the iron, preventing their oxidation to iron ions.

Oxidation Numbers

Oxidation NumbersA positive or negative number assigned to an

atom to indicate its degree of oxidation or reduction.

Rule of Thumb: when bonded, the oxidation number is the same as its ionic charge.

In a chemical reaction:Increase in oxidation number oxidationDecrease in oxidation number reduction

Rules for Assigning Oxidation NumbersMonatomic ions equal ionic charge; Br1-: -1H: compounds is +1; metal hydrides is -1

H2O: +1, NaH: -1

O: compounds is -2; peroxides is -1, or in compounds with F it’s +H2O: -2, H2O2: -1

Atoms in elemental form or diatomic is 0.S: 0, H2: 0

For compound, the sum of oxidation numbers must equal 0.H2O H(+1), O(-2) 2(+1) + 1(-2) =0

For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion.NO3

2- : N(4), O(-2) 1(+4) + 3(-2) = 2-

Oxidation Number PracticeNaCl

(+1,-1)H2O

(+1,-2)SO2

(+4, -2)CO3

2-

(+4,-2)Na2SO4

(+1,+6,-2)

Oxidation-Number Changes in Chemical Reactions +1 +5 -2 0 +2 +5 -2 02AgNO3 +Cu Cu(NO3)2 + 2Ag

Ag: reducedCu: oxidized

Let’s Try These:Cl2 + 2HBr 2HCl +Br

2H2 + O2 2H2O

2KNO3 2KNO2 + O2

AnswersLet’s Try These:0 +1 -1 +1 -1 0 Cl2 + 2HBr 2HCl +Br

0 0 +1 -22H2 + O2 2H2O

+1 +5 -2 +1 +3 -2 02KNO3 2KNO2 + O2

Balancing Redox Reactions

How to tell if it’s a redox rxnIf the oxidation number of an element in a

reacting species changes 0 0 +2 -2N2 + O2 2NO

Balancing by Oxidation No.1) Assign oxidation numbers to all the atoms

+3 -2 +2 -2 0 +4 -2

Fe2O3 + CO Fe + CO2

2) Identify which atoms are oxidized and reduced.

3) Use brackets to connect that atoms undergoing oxidation, and other set for those reduced.

+2 Oxidation +3 -2 +2 -2 0 +4 -2

Fe2O3 + CO Fe + CO2

-3 reduction

Balancing by Oxidation No.Make the total increase in oxidation number

equal the total decrease using coefficients (+2)x3=6

+3 -2 +2 -2 0 +4 -2

Fe2O3 + 3CO 2Fe + 3CO2

(-3)x2=6

Make sure the equation is balanced for both atoms and charge

Let’s PracticeKClO3 KCl + O2

HNO2 + HI NO + I2 + H2O

Bi2S3 + HNO3 Bi(NO3) 3 + NO + S + H2O

SbCl5 + KI SbCl3 +KCl + I2

Balancing Half-ReactionsRedox Reactions

Half-ReactionsEquation showing just the oxidation or

reduction portion of the redox reaction.S + HNO3 SO2 + NO +H2O 0 +4 -2

Oxidation Half: S SO2

+5 -2 +2 -2

Reduction Half: NO3- NO

Balancing Half-ReactionsTo balance:

Write the unbalanced ionic equationWrite separate half reactions for oxidation &

reductionBalance atoms in each half-reactionAdd enough electrons to one side of each half-

reaction to balance the chargesMultiply each half-reaction by an appropriate

number to make the numbers of electrons equal

Add the half reaction to show the overall equation

Add the spectator ions and balance the equation

Half-ReactionsS + HNO3 SO2 + NO +H2O

Ionic Form: S + H+ + NO3- SO2 + NO +H2O

0 +4 -2

Oxidation Half: S SO2

+5 -2 +2 -2

Reduction Half: NO3- NO

Balancing Atoms in Half –Reactions2H2O + S SO2 + 4H+ 4H+ + NO3

- NO + 2H2O

Half-ReactionsAdd e- to each side of half reactions to

balance chargesOxidation: 2H2O + S SO2 + 4H+ + 4e-

Reduction: 4H+ + NO3- + 3e- NO + 2H2O

Multiply each half reaction by an appropriate number to make the numbers of electrons equalOxidation: 6H2O + 3S 3SO2 + 12H+ + 12e-

Reduction: 16H+ + 4NO3- + 12e- 4NO +

8H2O

Subtract the terms that appear on both sides and add in the spectator ions6H2O + 3S + 16H+ + 4NO3

- + 12e- 3SO2 + 12H+ + 12e-+ 4NO + 8H2O

3S + 4HNO3 3SO2 + 4NO + 2H2O

Half-ReactionsLet’s Practice:KMnO4 + HCl MnCl2 + Cl2 + H2O + KCl