Chapter 15.

Oxidation-ReductionReactions

Oxidation-ReductionReactions

Oxidation was originally understood as reaction of a substance with oxygen.

Combustion:

CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g)

+ Energy!

Metabolism:

C6H12O6(s) + 6 O2(g) 6 CO2(g) + 6 H2O(g)

+ Energy!

Oxidation-ReductionReactions

Corrosion:

2 Mg(s) + O2(g) 2 MgO(g)

4 Fe(s) + 3 O2(g) 2 Fe2O3(s)Rust!

Oxidation-ReductionReactions

Reduction was originally understood as the loss of mass of metal ores in smelting.

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

Oxidation is a process in which a substance loses electrons.

Reduction is a process in which a substance gains electrons.

4 Fe(s) + 3 O2(g) 2 Fe2O3(s)

4 (Fe0 Fe3+ + 3e1) Iron is oxidized

4 Fe0 4 Fe3+ + 12 e1 forms cation

3 (O2 + 4 e1 2 O2) Oxygen is reduced

3 O2 + 12 e1 6 O2 forms anions

Oxidation-ReductionReactions

A half-reaction is a chemical equation that shows either the oxidation or reduction part of an oxidation-reduction reaction.

Electrons appear as products in oxidations.

Zn(s) Zn2+(aq) + 2 e1

Electrons appear as reactants in reductions.

I2(aq) + 2 e1 2 I1 (aq)

L E O the lion says G E R!

Lose Electrons Oxidation

Gain Electrons Reduction

Oxidation-ReductionReactions

Examples:

Which of these half-reactions show oxidations? Which ones show reductions?

Zn(s) Zn2+(aq) + 2 e1

Cu2+(aq) + 2 e1 Cu(s)

Ag(s) Ag1+(aq) + e1

Cr2O72 + 14H1+ + 6 e1 2 Cr3+ + 7 H2O

2 F1(aq) F2(g) + 2 e1

Oxidation-ReductionReactions

In ionic compounds, we can look at changes in charge to figure out what's reduced and what's oxidized.

What happens with molecular compounds?

CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g)

Use Oxidation Numbers

Determining Oxidation Numbers

An oxidation number (or oxidation state) is the charge that an atom appears to have when the electrons in each bond in which it is participating are assigned to the more electronegative of the two atoms participating in the bond.

Determining Oxidation Numbers

Draw Lewis Structure of molecule or ion.

Use Electronegativities (p 281, Fig 7.12) to determine which atom "owns" electrons.

Lone pairs belong to the atom.

If two atoms of the same element are bonded together, 1 electron to each.

ONAtom = VEAtom - TEAtom

Determining Oxidation Numbers

The oxidation number of a main-group element can be anything from Column # (loses all e1) to (Column # 8) (fills valence shell).

The sum of all oxidation numbers in a species must equal its charge.

The oxidation number of an atom in its elemental form is zero.

If a Lewis structure is complex, just work with the atom of interest and the atoms bonded to it.

Determining Oxidation Numbers

Examples:

Determine the oxidation number of:

N in N2 H in H2 S in S8

C and H in CH4 N and O in N2O

S in H2SO4 Cl in ClO41

C in Urea, CH4N2O

Determining Oxidation Numbers

Polyatomic ions in which the central atom is a transition metal:

Oxygen's ON is 2. Hydrogen's ON is +1.

The sum of all oxidation numbers in a species must equal its charge.

What are the oxidation states of the metals?

MnO41 CrO4

2 Cr2O72

Vocabulary

A substance is oxidized if:

it loses electrons

it loses hydrogen atoms

it gains oxygen atoms

its charge or oxidation number increases (becomes more

positive).

Vocabulary

A substance is reduced if:

it gains electrons

it gains hydrogen atoms

it loses oxygen atoms

its charge or oxidation number decreases (becomes less

positive).

Vocabulary

An agent causes something to happen.

An oxidizing agent causes oxidation.

It does this by being reduced (gaining electrons).

A reducing agent causes reduction.

It does this by being oxidized (losing electrons).

Vocabulary

Typical Oxidizing Agents:

O2 Oxygen MnO41 Permanganate

ClO41 Perchlorate CrO4

2 Chromate

ClO31 Chlorate Cr2O7

2 Dichromate

NO31 Nitrate Cl2, other halogens

O3 Ozone H2O2 Peroxides

Vocabulary

Typical Reducing Agents:

Hydrogen sources:

H2 Hydrogen

NaBH4 Sodium borohydride

LiAlH4 Lithium aluminum hydride

Active metals (1A and 2A metals, zinc)

Reduced Carbon (C, CO)

Examples:

Identify the substance being oxidized, the sub-stance being reduced, the oxidizing agent, and the reducing agent:

2 H2(g) + HCCH(g) C2H6(g)

2 HCl(aq) + Zn(s) ZnCl2(aq) + H2(g)

4 C3H6O(l) + NaBH4(s) + 4 H2O(l) 4 C3H8O(l) + NaB(OH)4(aq)

Examples:

Identify the substance being oxidized, the sub-stance being reduced, the oxidizing agent, and the reducing agent:

2 Na(s) + Cl2(g) 2 NaCl(s)

C2H4O(l) + H2(g) C2H6O(l)

2 KMnO4 + 5 H2O2 + 3 H2SO4 2 MnSO4 + K2SO4 + 5 O2(g) + 8 H2O

Predicting ReactionsWho does what to whom?

Cu2+(aq) + 2 Ag(s) Cu(s) + 2 Ag1+(aq)

Cu(s) + 2 Ag1+(aq) Cu2+(aq) + 2 Ag(s)

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

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

2 Ag1+(aq) + Zn(s) Zn2+(aq) + 2 Ag(s)

2 Ag(s) + Zn2+(aq) 2 Ag1+(aq) + Zn(s)

A Daniell Cell

Daniell CellsVoltages from Daniell Cells provide information

about energy changes:

Cu(s) + 2 Ag1+(aq) Cu2+(aq) + 2 Ag(s)

0.34 V

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

1.10 V

2 Ag1+(aq) + Zn(s) Zn2+(aq) + 2 Ag(s)

1.56 V

An Energy Scale for Redox RXN’s

Standard Reduction Potentials are potentials measured against a standard hydrogen elec-trode, with all solutions at 1.0 M and gas pressures at 1.0 atm.

A Standard Hydrogen Electrode is an elec-trode in which the half-reaction is:

2 H1+ + 2 e1 H2(g)

Materials are at standard conditions, and the half-cell potential is 0.00 V.

An Energy Scale for Redox RXN’s

From chemguide.co.uk

An Energy Scale for Redox RXN’s

Some Standard Reduction Potentials:Reaction E0, V

Li1+(aq) + e1 Li(s) 3.04Al3+(aq) + 3 e1 Al(s) 1.66Zn2+(aq) + 2 e1 Zn(s) 0.76

2 H 1+(aq) + 2 e1 H2(g) 0.00Cu2+(aq) + 2 e1 Cu(s) 0.34Ag1+(aq) + e1 Ag(s) 0.80

Cr2O7 + 14H1+ + 6 e1 2 Cr3+ + 7 H2O 1.33

F2(g) + 2 e1 2 F1(aq) 2.87

An Energy Scale for Redox RXN’s

Strongest oxidizing agents have highest standard reduction potentials (E0's, in volts). Strong oxidizing agents have low-energy vacant orbitals.

Strongest reducing agents have lowest (often negative) standard reduction potentials. Strong reducing agents have electrons in high-energy orbitals.

Predicting Reactions

Balancing equations for Redox Reactions

Use Table of Standard Reduction Potentials:

Find half-reaction for reduction, write it as given.

Find half-reaction for oxidation, write it as reverse of what is given

(electrons are products).

Predicting Reactions

Balance half-reactions so number of electrons is equal for each.

Add half-reactions, cancelling out electrons and any other species that appear on both sides of equation.

Predicting Reactions

Will the Reaction be Spontaneous?

Calculate E0 for RXN: E0 reduction

- E0 oxidation

E0 RXN

If E0 RXN is positive, the reaction is exo- thermic and will proceed spontaneously.

If E0 RXN is negative, the reaction is endo-thermic and will not proceed spontaneously.

Predicting Reactions

Balance the reactions.

Which of them are spontaneous ?

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

Zn(s) + Ag1+(aq) Zn2+(aq) + Ag(s)

Cu(s) + H1+(aq) Cu2+(aq) + H2(g)

Fe(s) + Al3+(aq) Fe2+(aq) + Al(s)

Cu(s) + Cr2O72–(aq) Cu2+(aq) + Cr3+(aq)

Batteries

Batteries are galvanic (a.k.a. voltaic) electro-chemical cells in which a spontaneous reac-tion is used to convert chemical energy to electrical energy.

The earliest batteries common galvanic cells were Daniell Cells, from 1836.

Batteries

Batteries usually contain an electrolyte, which is a solution that contains ions. Electric cur-rent can flow through the electrolyte.

The anode in a battery is the half-cell with the lower (more negative) E0. Oxidation occurs at the anode. Anions flow toward the anode.

The cathode in a battery is the half-cell with higher (more positive) E0. Reduction occurs at the cathode. Cations flow toward the cathode.

A Daniell Cell

Batteries Dry cell (ca. 1900, developed from Leclanché

cell, 1866. Alkaline cell, ca. 1950)

Batteries

Dry cell, acid form; Cell potential ~1.5 V

Anode half-reaction:

Zn(s) Zn2+(aq) + 2 e1–

Cathode half-reaction:

2 MnO2(s) + 2 NH4Cl(aq) + 2 e1– Mn2O3(s) + 2 NH3(aq) + H2O(l) + 2

Cl1–(aq)

Batteries

Dry cell, alkaline form; Cell potential ~1.5 V

Anode half-reaction:

Zn(s) + 2 OH1–(aq) ZnO(s) + H2O(l) + 2 e1–

Cathode half-reaction:

2 MnO2(s) + 2 H2O(l) + 2 e1– Mn(OH)2(s) + 2 OH1–(aq)

Batteries

“Button” battery, ca. 1940, HgO or Ag2O

Batteries

Button battery; Cell potential ~1.5 V

Anode half-reaction:

Zn(s) + 2 OH1–(aq) ZnO(s) + H2O(l) + 2 e1–

Cathode half-reaction:

Ag2O(s) + 2 H2O(l) + 2 e1– 2 Ag(s) + 2 OH1–(aq)

HgO(s) + H2O(l) + 2 e1– Hg22+(aq) + 2 OH1–(aq)

Batteries Lead-Acid Battery (ca. 1859)

Batteries Detail of Lead-Acid Battery: Six cells in series

produce about 12 volts.

Batteries

Lead-acid battery; Cell potential ~2.1 V

Anode half-reaction:

Pb(s) + SO42–(aq) PbSO4(s) + 2 e1–

–0.36 V

Cathode half-reaction:

PbO2 (s) + 4 H1+(aq) + SO42–(aq) + 2 e1–

PbSO4(s) + 2 H2O(l)

+1.69 V

Batteries

Nickel-Metal Hydride (Ni-MH) Battery; cell potential ~1.4V

Anode half-reaction:

MH(s) + OH1–(aq) M(s) + H2O(l) + e1–

Cathode half-reaction:

NiO(OH)(s) + H2O(l) + e1– Ni(OH)2(s) + OH1–(aq)

MH(s) is metal alloy, e.g. LaNi5, into which hydrogen is absorbed.

Batteries

Lithium ion battery; cell potential ~3.7V

Chemistry is complex because water can-not be used in the electrolyte.

Lithium embedded in graphite is the anode.

Electrolytic Cells

An electrolytic cell is an electrochemical cell in which a non-spontaneous chemical change is caused to occur by application of electrical energy.

Many important chemical processes are carried out in electrochemical cells.

Electrolytic CellsA chlor-alkali cell

Electrolytic Cells

A Hall-Heroult Cell for Aluminum

Electrolytic Cells

An electrorefining cell for purifying copper

Electrolytic Cells

An electroplating cell