Redox Reactions Oxidation & Reduction Oxidation: ________________ electrons Reduction: ________________ electrons One can use the acronym LEO the lion says GER Losing electrons is Oxidation Gaining electrons is Reduction Oxidation Number Rules Sometimes the number represents an actual charge, but often does not. It is simply an electronic book-keeping system All of the oxidation numbers in a molecule must add up to the overall charge of that molecule 1. Free elements (including the HOFBrINCl’s) have ox. # = 0. Ex. K(s) = 0 Cl2(g) = 0 1. Single atom ions: ox. # = ionic charge. Na+ = +1 Cl- = -1 3. All ox. #s in a neutral compound must add up to 0. Al2O3: (2 x 3+) + (3 x 2-) = 0 4. All ox. #s in a polyatomic ion must add up to the charge of the ion. 5. Oxygen in a compound is nearly always -2. (Exceptions: peroxides O22- (with a column IA or IIA metal), O is -1, or in OF2, O is +2.) 6. Hydrogen in a compound is nearly always +1. (Exception: H with a column IA or IIA metal. H is the anion. Hydride = -1. Ex. NaH, CaH2). 7. In covalent compounds (without H or O) the oxidation number is the charge of the most EN element Examples: • K2O2 • NH4Cl • NaNO3 • Mg(NO3)2 Eg// Fe2(Cr2O7)3 The oxidation states of atoms in a polyatomic ion must add up to the overall charge of that ion. (Cr2O7- must add up to -1; but overall molecule is neutral) Balancing Redox Reactions Redox reactions occur when reactants gain or lose electrons A reactant can be considered an oxidizing agent or a reducing agent An oxidizing agent causes the oxidization of another element but it itself gets reduced A reducing agent cause the reduction of another element but it itself gets oxidized Methods of Balancing Redox Reactions 1. Write the complete equation and assign oxidation numbers to all elements. 2. Identify the elements oxidized (ox. # ↑) and reduced (ox. # ↓). 3. Use lines to connect the elements being oxidized and reduced. Show the change in oxidation number. 4. Make the total change in oxidation numbers equal. Electrons gained must equal electrons lost. 5. Balance any remaining atoms by inspection. 6. Balance the hydrogen atoms on each side by adding H+ if in acidic solution. (If the reaction is in base, next add enough OH- to each side to neutralize all the H+.) Fe2O3 + CO Ex// Fe + CO2 Half Reaction Method In this method we break down the reaction into 2 half-reactions; an oxidation reaction and a reduction reaction 1. Write the unbalanced ionic reaction. 2. Write the half reactions for the oxidation and reduction. 3. Balance the atoms being oxidized and reduced. (You may need to check the oxidation numbers to be sure which ones they are.) 4. Balance oxygen by adding water to the other side of the equation. 5. Balance the hydrogen atoms on each side by adding H+. (If the reaction is in base, next add enough OH- to each side to neutralize all the H+.) 6. Balance the charge on both sides of each half-reaction by adding electrons. The net charge on both sides of each equation should be equal. 7. If necessary, multiply one or both half-reactions by a factor that will make electrons lost = electrons gained. 8. Add the half-reactions together, cancel any common terms, including electrons, and make sure any remaining atoms are also balanced. Ex// S + NO3- SO2 + NO (in acid) Example: Try both methods to balance the following reaction Oxidation Numbers: Half-Reaction: Cr(OH)3(s) + ClO3-(aq) CrO42-(aq) + Cl-(aq) Cr(OH)3(s) + ClO3-(aq) CrO42-(aq) + Cl-(aq) Galvanic Cells • Galvanic cells convert the chemical energy of redox reactions in electrical energy to perform work • The key to gathering the electron flow is to separate the oxidation-reduction halfreactions, • Therefore forcing the electrons to travel through a wire, creating current. • Galvanic cells contain two metal electrodes: a cathode and an anode • At the anode: an oxidation reaction occurs where the metal electrode loses electrons to the solution housing it • At the cathode: a reduction reaction occurs where the metal electrode gains electrons transferred to the solution from the anode solution • The difference between the potential energy at the anode and cathode is called Cell Voltage • Standard cell potentials can be calculated using reduction potential tables. • Formula for standard potential:
