Chapter 15. Oxidation-Reduction Reactions

Oxidation-Reduction Reactions 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)

Oxidation-Reduction Reactions Corrosion: 2 Mg(s) + O2(g)  2 MgO(g) 4 Fe(s) + 3 O2(g)  2 Fe2O3(s) Rust!

Oxidation-Reduction Reactions 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 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-Reduction Reactions 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-Reduction Reactions 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-Reduction Reactions 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 CrO42 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 CrO42 Chromate ClO31 Chlorate Cr2O72 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) + HCCH(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 Reactions Who 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 Cells Voltages 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) 3.04 Al3+(aq) + 3 e1  Al(s) 1.66 Zn2+(aq) + 2 e1  Zn(s) 0.76 2 H 1+(aq) + 2 e1  H2(g) 0.00 Cu2+(aq) + 2 e1  Cu(s) 0.34 Ag1+(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 Cells A chlor-alkali cell

Electrolytic Cells A Hall-Heroult Cell for Aluminum

Electrolytic Cells An electrorefining cell for purifying copper

Electrolytic Cells An electroplating cell