Oxidation States and Oxidation-Reduction Reactions

• We determine whether a given chemical reaction is an oxidation–reduction reaction by keeping track of the oxidation numbers (oxidation states) of the elements involved in the reaction.

• Oxidizing Agent — The substance that oxidizes another substance; also known as oxidant.

• Reducing Agent —  A substance that gives up electrons, thereby causing another substance to be reduced; also known as reductant.

Balancing Redox Equations

• Law of Conservation of Mass: The amount of each element must be the same on both sides of the equation.

• Half Reactions — Equations that show either oxidation or reduction alone.

• In the overall redox reaction, the number of electrons lost in the oxidation half-reaction must equal the number of electrons gained in the reduction half-reaction.

• When this condition is met and each half-reaction is balanced, the electrons on the two sides cancel when the two half-reactions are added to give the balanced oxidation–reduction equation.

Balancing Equations by the Method of half-Reactions

1. Divide the equation into one oxidation half-reaction and one reduction half-reaction.

2. Balance each half-reaction.

   1. First, balance elements other than H and O.

   2. Next, balance O atoms by adding H2O as needed.

   3. Then balance H atoms by adding H+ as needed.

   4. Finally, balance charge by adding e^– as needed.

• This specific sequence (a) — (d) is important, and it is summarized in the diagram in the margin. At this point, you can check whether the number of electrons in each half-reaction corresponds to the changes in the oxidation state.

3. Multiply half-reactions by integers as needed to make the number of electrons lost in the oxidation half-reaction equal the number of electrons gained in the reduction half-reaction.

4. Add half-reactions and, if possible, simplify by canceling species appearing on both sides of the combined equation.

5. Check to make sure that atoms and charges are balanced.
   

Balancing Equations for Reactions Occurring in Basic Solution

1. Balance the half-reactions as if they occurred in acidic solution.

2. Count the number of H+ in each half-reaction, and then add the same number of OH– to each side of the half-reaction.

Voltaic Cells

• Voltaic Cell — A device in which the transfer of electrons takes place through an external pathway rather than directly between reactants present in the same reaction vessel.

• Electrodes — The two solid metals connected by the external circuit.

• Anode — The electrode at which oxidation occurs.

• Cathode —  The electrode at which reduction occurs.

• Half-Cell — Each compartment of voltaic cell.

• One half-cell is the site of the oxidation half-reaction, and the other is the site of the reduction half-reaction.

• Whichever device is used to allow ions to migrate between half-cells, anions always migrate toward the anode and cations toward the cathode.

Cell Potentials under Standard Conditions

• Potential Difference — The difference in potential energy per electrical charge between two electrodes is measured in volts.

• Cell Potential — The potential difference between the two electrodes of a voltaic cell.

• Electromotive Force (EMF) — The potential difference provides the driving force that pushes electrons through the external circuit.

• Standard Cell Potential (Standard EMF) — The cell potential under standard conditions.

• Standard Reduction Potentials — Standard half-cell potentials are tabulated for the reduction reaction.

• Standard Hydrogen Electrode — An electrode designed to produce this half-reaction.

  • Whenever we assign an electrical potential to a half-reaction, we write the reaction as a reduction.

• Half-Cell Potential — standard reduction potentials.

• Because electrical potential measures potential energy per electrical charge, standard reduction potentials are intensive properties.

• Changing the stoichiometric coefficient in a half-reaction does not affect the value of the standard reduction potential.

Free Energy and Redox Reactions

• The change in the Gibbs free energy, ∆G, is a measure of the spontaneity of a process that occurs at constant temperature and pressure.

• The emf, E, of a redox reaction also indicates whether the reaction is spontaneous.

• The relationship between emf and the free-energy change is ∆G = –nFE

• Faraday constant is the quantity of electrical charge on 1 mol of electrons.

• A positive value of E and a negative value of ∆G both indicate a spontaneous reaction.

• When the reactants and products are all in their standard states, can be modified to relate ∆G° and E°.

• ∆G° = –nFE°

Cell Potentials under Nonstandard Conditions

The Nernst Equation

• The effect of concentration on cell emf can be obtained from the effect of concentration on free-energy change.

• The free-energy change for any chemical reaction, ∆G, is related to the standard free-energy change for the reaction, ∆G°:

• ∆G = ∆G° + RT ln Q

• The quantity Q is the reaction quotient, which has the form of the equilibrium-constant expression except that the concentrations are those that exist in the reaction mixture at a given moment.

• Substituting ∆G = –nFE gives:

• –nFE = –nFE° + RT ln Q

• 
  Concentration cell — A cell based solely on the emf generated because of a difference in a concentration.

Batteries and Fuel Cells

• Battery — is a portable, self-contained electrochemical power source that consists of one or more voltaic cells.

• Primary Cells — Batteries that cannot be recharged and must be either discarded or recycled after the voltage drops to zero.

• Secondary Cells — Batteries that can be recharged from an external power source after its voltage has dropped.

• A 12-V lead–acid automotive battery consists of six voltaic cells in series, each producing 2 V. The cathode of each cell is lead dioxide 1PbO22 packed on a lead grid.  The anode of each cell is lead. Both electrodes are immersed in sulfuric acid.

• A major advantage of the lead–acid battery is that it can be recharged.

• Alkaline Battery: The anode is powdered zinc metal immobilized in a gel in contact with a concentrated solution of KOH. The cathode is a mixture of MnO2(s) and graphite, separated from the anode by a porous fabric.

• The battery is sealed in a steel can to reduce the risk of any of the concentrated KOH escaping.

• Nickel-Cadmium Battery: During discharge, cadmium metal is oxidized at the anode while nickel oxyhydroxide is reduced at the cathode.

• A single nicad voltaic cell has a voltage of 1.30 V. Nicad battery packs typically contain three or more cells in series to produce the higher voltages needed by most electronic devices.

• Because cadmium is toxic, these batteries must be recycled.

• Lithium-Ion Battery: Mostly used in portable electronic devices, including cellphones, and laptop computers.

• Because lithium is a very light element, Li-ion batteries achieve a greater specific energy density—the amount of energy stored per unit mass—than nickel-based batteries.

• A Li-ion battery produces a maximum voltage of 3.7 V per cell, nearly three times higher than the 1.3 V per cell that nickel–cadmium and nickel–metal hydride batteries generate.

Hydrogen Fuel Cells

• The thermal energy released by burning fuels can be converted to electrical energy.

• The thermal energy may convert water to steam, for instance, which drives a turbine that in turn drives an electrical generator.

• Fuel Cells — Voltaic cells that perform high rate of conversion of chemical energy to electricity using conventional fuels.

Corrosion

• Corrosion reactions — Are spontaneous redox reactions in which a metal is attacked by some substance in its environment and converted to an unwanted compound.

Corrosion of Iron

• Rusting of iron requires both oxygen and water, and the process can be accelerated by other factors such as pH, presence of salts, contact with metals more difficult to oxidize than iron, and stress on the iron.

• The corrosion process involves oxidation and reduction, and the metal conducts electricity.  Electrons can move through the metal from a region where oxidation occurs to a region where reduction occurs, as in voltaic cells.

Preventing Corrosion of Iron

• Objects made of iron are often covered with a coat of paint or another metal to protect against corrosion. Covering the surface with paint prevents oxygen and water from reaching the iron surface.

• If the coating is broken, however, and the iron is exposed to oxygen and water, corrosion begins as the iron is oxidized.

• With galvanized iron, which is iron coated with a thin layer of zinc, the iron is protected from corrosion even after the surface coat is broken.

• Cathodic Protection — Protecting metal from corrosion by making it the cathode in an electrochemical cell.

• Sacrificial Anode — The metal that is oxidized while protecting the cathode**.**

Electrolysis

• Electrolysis reactions — Processes driven by an outside source of electrical energy and takes place in electrolytic cells.

• An electrolytic cell consists of two electrodes immersed either in a molten salt or in a solution.

• Electroplating — Uses electrolysis to deposit a thin layer of one metal on another metal to improve beauty or resistance to corrosion.

• Electrometallurgy — Processes that used to produce or refine metals are based on electrolysis.

• The electrodes in an electrolytic cell can be inert or active, meaning that the electrode can be involved in the electrolysis reaction. Active electrodes are important in electroplating and in metallurgical processes.