Oxidation – Reductiona.k.a.

REDOXTextbook Sections:

4.4-4.6 and 20.1-20.2

Good website:http://www.wfu.edu/~ylwong/redox/

• Oxidation – the loss of one or more electrons by a substance (element, compound, ion)

• Reduction – the gain of one or more electrons by a substance (element, compound, ion)

O. I. L. R. I. G. Oxidation Is Loss, Reduction Is Gain

(of electrons)L.E.O. the lion says G.E.R.

Loss of Electrons is Oxidation Gain of Electrons is Reduction

Oxidation --------->A-2 A-1 + electronA-1 A + electronA A+1 + electronA+1 A+2 + electron

<--------- Reduction

• Redox reaction – a process where electrons are transferred from one substance to another

• How can you tell when a redox reaction is taking place?

1) Assign oxidation numbers to atoms in substances2) Compare oxidation numbers before and after reaction to determine if atom has lost or gained electrons

Rules For Assigning Oxidation NumbersUse the Rules in Order

1.The oxidation number of an atom in an element is 0.Examples: Na, H2, Br2, S8, NeOx. # 0 0 0 0 0

2.The oxidation number of a monatomic ion is the same as its charge.Examples: Na+1, Ca+2, Al+3, Cl-1, O-2

Ox # +1 +2 +3 -1 -2

3. The sum of the oxidation numbers of all atoms in a neutral compound is zero.

4. The sum of the oxidation numbers of all atoms in an ion is equal to the charge on the ion.

5. In compounds, fluorine is always assigned an oxidation number of -1 (the most electronegative element in a compound always has a negative oxidation number.)

6. Hydrogen’s oxidation number will be - +1 when bonded to a nonmetal (HCl)- -1 when bonded to a metal (NaH)Examples: NaH CaH2 HCl H2S

Na—H H—Ca—H H—Cl H—S—H

+1 -1 -1 +2 -1 +1 -1 +1 -2 +1

7. Oxygen usually has an oxidation number of -2. Exceptions: in peroxides, oxygen will be -1 and when combined with only F, it will be +2 and in O2 it will be 0Examples:

H2O CaO H2O2

H — O — H Ca—O H—O—O—H

O2-2 OF2

[O—O]-2 F—O—F

+1 -2 +1 +2 -2 +1 -1 -1 +1

-1 -1 -1 +2 -1

8. Halogens usually have an oxidation number of -1.

Examples: NaCl MgI2 OCl2

HOBrNa—Cl, I—Mg—I, Cl—O—Cl, H—O—

Br +1 -1 -1 +2 -1 +1 -2 +1 +1 -2 +1

** If none of the above rules help you get started…look for a atom

with a known charge and use that charge as its oxidation number

CdS: Cd-S

+2 -2

Use algebra to determine oxidation numbers of "difficult" atoms.

Example: H2SO4

H is +1 * 2 = +2 O is -2 * 4 = -8

2 + S + -8 = 0 S is +6

Example: ClO4-1

O is -2 * 4 = -8 -8 + Cl = -1

Cl is +7

Example: NH4+1

H is +1* 4 = +4

4 + N = 1

N is -3

FeSO4

O is -2 *4 = -8Recognize this as an ionic compound – sulfate has a -2

chargeFor sulfate: x + -8 = -2S = +6Then look at the compound as a wholeFe + 6 + -8 = 0Fe = +2

C3H8

H is +1 * 8 = 83C + 8 = 0 3C = -8 C = - 8/3oxidation numbers do NOT have to be integers

Once oxidation numbers have been assigned, compare them before and after the reaction.4 Fe (s) + 3 O2 (g) 2 Fe2O3 (s) 0 0 +3 -2

Fe is oxidized, going from 0 to +3

O is reduced, going from 0 to -2

Notice that a total of 12 electrons were lost and 12 electrons were gained

2 Fe2O3(s) + 3 C(s) 4 Fe(s) + 3 CO2(g) +3 -2 0 0 +4 -2

Fe is reduced, going from +3 to 0

C is oxidized, going from 0 to +4

O undergoes no change

12 electrons lost and 12 electrons gained

As seen in the above examples, oxidation and reduction ALWAYS occur together

Reducing agent (reductant)• causes reduction• loses electrons• undergoes oxidation• oxidation number increases

Oxidizing agent (oxidant)• causes oxidation• gains electrons• undergoes reduction• oxidation number decreases

Assign oxidation numbers, indicate what is oxidized and reduced, indicate what is the oxidizing agent and reducing agent Ca(s) + 2 H+1(aq) Ca+2(aq) + H2(g)

0 +1 +2 0

Ca is oxidized – increasing from 0 to +2H+1 is reduced – decreasing from +1 to 0 Ca is the reducing agentH+1 is the oxidizing agent

2 Fe+2(aq) + Cl2(aq)2 Fe+3(aq) + 2 Cl-1(aq)

+2 0 +3 -1

Fe+2 is oxidized – increasing from +2 to +3

Cl2 is reduced – decreasing from 0 to -1

Fe+2 is the reducing agentCl2 is the oxidizing agent

• In general, • metals and anions act as reducing

agents (are oxidized) and • nonmetals and cations act a oxidizing

agents (are reduced).

• Periodic table – in general, metals on left of table are more active, metals become less active as you move to the right side of the table

Predicting Products of Redox Reactions

The simple ones you already know:• Decomposition (except of acids, bases,

carbonates & hydrates)• Composition (except of two oxides)• Combustion• Replacement

Replacement ReactionsReplacement Reactions are redox reactions.

General pattern: A + BX AX + B

Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g)The Mg is oxidized and the H+ is reduced.

Fe(s) + Ni(NO3)2(aq) Fe(NO3)2(aq) + Ni(s)

• The net ionic equation shows the redox chemistry well:

Fe(s) + Ni+2(aq) Fe+2(aq) + Ni(s)

Fe is oxidized to Fe+2 Ni+2 is reduced to Ni.

• Always keep in mind that whenever one substance is oxidized, some other substance must be reduced.

The Activity Series• The activity series is a list of metals in

order of decreasing ease of oxidation. • The metals at the top of the activity

series are called active metals and are easily oxidized (they WANT to be ions)

• The metals at the bottom of the activity series are called noble metals and NOT easily oxidized (they WANT to be atoms)

Oxidation of copper metal by silver ions.

(Dime Lab)

• A metal in the activity series can ONLY be oxidized by a metal ion below it (metal doing the replacing has to be above what it is replacing in the activity series)

• If we place Cu into a solution of Ag+ ions, then Cu+2 ions can be formed because Cu is above Ag in the activity series:

• Cu(s)+2AgNO3(aq)Cu(NO3)2(aq)+2Ag(s) or

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

• This is only part of the story – more later !

• Special Cases:1. Hydrogen reacts with a hot metallic oxide to produce the metal element and water.

Ex: H2 + MgO Mg + H2O

2. A metal sulfide reacts with oxygen to produce the metallic oxide and sulfur dioxide.

Ex: 2MgS + 3O2 2MgO + 2SO2

3. Chlorine gas reacts with dilute sodium hydroxide to produce sodium hypochlorite, sodium chloride, and water.Cl2 + 2NaOH NaClO + NaCl + H2O

4. Copper reacts with concentrated sulfuric acid to produce copper(II) sulfate, sulfur dioxide, and water.

Cu + 2H2SO4 CuSO4 + SO2 + 2H2O

5. Copper reacts with dilute nitric acid to produce copper(II) nitrate, nitrogen monoxide, and water.Cu + HNO3 Cu(NO3)2 + NO + H2O

6. Copper reacts with concentrated nitric acid to produce copper(II) nitrate, nitrogen dioxide, and water.Cu + HNO3 Cu(NO3)2 + NO2+ H2O

Reactants (oxidizing agents) Products

MnO4- in acidic solution Mn2+

MnO2 in acidic solution Mn2+

MnO4- in neutral or basic solution MnO2(s)

MnO4- in very basic solution MnO4

2-

Cr2O72- in acidic solution Cr3+

HNO3, concentrated NO2

HNO3, dilute NO

H2SO4, hot, concentrated SO2

Metallic ion (higher charge) Metallous ion (lower charge)

Elemental Halogen Halogen ion

Na2O2 NaOH

HClO4 Cl-

C2O42- CO2

H2O2 H2O

Memorize… These are reduction reactions (oxidation number decreases)

Reactants (Reducing Agents) Products

Halogen ions Halogen element

Metal element Metal ion

SO32- or SO2 SO4

2-

NO2- NO3

-

Halogen element, dilute basic solution

Hypo-halogen-ite ion(Ex: ClO-, BrO-)

Halogen element, concentrated basic solution

Halogen-ate ion(Ex: ClO3

-, BrO3-)

Metallous ion (lower charge) Metallic ion (higher charge)

H2O2 O2

Memorize… These are oxidation reactions (oxidation number increases)