1 Oxidation – Reduction Notes (G1) describe oxidation and reduction processes In this unit, we will be dealing with a branch of chemistry called electrochemistry. Electrochemistry is concerned with the conversion of chemical energy into electrical energy and vice versa. Consider the reaction: Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) This reaction can also be described by the set-up below called an electrochemical cell. Note: The salt bridge allows the flow of ions between the two sides of the electrochemical cell to maintain a balance in charge. define and identify o oxidation _________________________________________________________ examples: __________________________________________________ o reduction_________________________________________________________ examples: __________________________________________________ o oxidizing agent_____________________________________________________ examples: __________________________________________________ o reducing agent_____________________________________________________ examples: __________________________________________________ o half-reaction_______________________________________________________ examples: __________________________________________________ o redox reaction_____________________________________________________ examples: __________________________________________________ 2 determine the following: o the oxidation number of an atom in a chemical species Oxidation number is a number assigned to an element in a compound that represents the number of electrons lost or gained by that element. Here are some rules to follow when determining the oxidation number: 1. The oxidation number of an atom in its elemental form is zero 2. The oxidation number of monatomic ion equals the charge on the ion 3. The sum of the oxidation numbers of all the atoms in a compound equals zero 4. The sum of the oxidation numbers of all the atoms in a polyatomic ion equals the charge on the ions Some assumptions can be made about oxidation numbers: 1. Alkali metals usually have an oxidation number of +1 2. Alkaline earth metals usually have an oxidation number of +2 3. Oxygen usually has an oxidation number of -2 4. Halogens usually have an oxidations number -1 Find the oxidation number of the following species. 1. Na in NaCl 2. Cr in CrO3 3. Ga in Ga(OH)3 3 4. P in PO43- 5. Cr in CrO42- 6. Cl in ClF3 7. P in P4 8. P in H4P2O7 4 relate change in oxidation number to gain or loss of electrons determine the following: o the change in oxidation number an atom undergoes when it is oxidized or reduced o whether an atom has been oxidized or reduced by its change in oxidation number Oxidation is the loss of electrons, so the species undergoing oxidation experiences an increase in oxidation number. Reduction is the gain of electrons, so the species undergoing reduction experiences a decrease in oxidation number. In most redox reactions, the change in oxidation number is easy to spot Consider the reaction: Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Consider the reaction: 2 Ag+(aq) + Cu(s) → 2 Ag(s) + Cu2+(aq) However, there some more difficult ones to spot… ClO3- → ClO4- H2O2 → H2O Cr3+ → CrO42- NO2 → N2O3 5 (G2) analyse the relative strengths of reducing and oxidizing agents from data for a series of simple redox reactions, create a simple table of reduction halfreactions identify the relative strengths of oxidizing and reducing agents from their positions on a half-reaction table use the “Standard Reduction Potentials of Half-Cells” table to predict whether a spontaneous redox reaction will occur between any two species o Things to notice on the “Standard Reduction Potentials of Half-Cells” table All the half-reactions are written as reduction reactions. The oxidation reaction is the reverse reaction. In general, metals are found in the bottom half of the table. In general, halogens and oxyanions (anions containing oxygen) are found in the top half of the table. Some metals maybe found on both sides of the table because it has more than one common oxidation number. The table has two columns… On the left, species higher up are stronger oxidizing agents On the right, species farther down are stronger reducing agents Predict whether or not a reaction will occur when the following are mixed. a) Zn and Cu b) Zn2+ and Cu2+ 6 c) Zn2+ and Cu d) Cu2+ and Zn e) Cl2 and Br- f) Sn and Mn g) Ni2+ and Pb h) H+ and SO42- 7 (G3) balance equations for redox reactions balance the equation for a half-reaction in solutions that are acidic, basic, or neutral o A half-reaction must be balanced for both mass and charge. o Follow these 4 steps to balance a half-reaction: 1. Balance the major atoms by inspection (the major atoms are any atoms other than oxygen and hydrogen). 2. Balance the oxygen atoms by adding H2O molecules. 3. Balance the hydrogen atoms by adding H+ (the reactions are initially treated as if they occur in acidic solutions). 4. Balance the overall charge by adding electrons. Balance the half-reaction RuO2 ↔ Ru. The half-reaction occurs in acidic solution. Balance the half-reaction Cr2O72- ↔ Cr3+. The half-reaction occurs in acidic solution. Balance the half-reaction Pb ↔ HPbO2-. The half-reaction occurs in basic solution. Balance the following half-reaction which occurs in acidic solution: H2 → 8 write the equations for reduction and oxidation half-reactions, given a redox reaction balance the equation for a net ionic redox reaction in acidic or basic solution o There are two methods for balancing redox reactions: using half-reactions and using oxidation numbers. The method of half-reactions is easier to use with more complicated redox reactions while the method of oxidation numbers is quicker to use when balancing simple redox reactions. Method of half-reactions breaking the equation into separate reduction and oxidation half-reactions Balance the equation Os + IO3- → OsO4 + I2 (in acidic solution). 9 Balance the equation MnO4- + C2O42- → MnO2 + CO2 (in basic solution). Balance the equation ClO2- → ClO3- + Cl- (in basic solution). 10 Method of oxidation numbers An increase in oxidation number in one species must be balanced by the corresponding decrease in oxidation number in a second species The overall change in oxidation numbers must be zero Balance the equation U4+ + MnO4- → Mn2+ + UO22+ (in acidic solution). Balance the equation Zn + As2O3 → AsH3 + Zn2+ (in basic solution). Balance the equation P4 → H2PO2- + PH3 (in acidic solution). 11 (G4) determine the concentration of a species by performing a redox titration Redox Titrations o As with acid-base titrations, redox titrations allow us to determine the concentration of a substance which is capable of undergoing an oxidation or reduction reaction. Oxidizing Agents o One of the most useful oxidizing agents is KMnO4 (permanganate). MnO4- is a strong oxidizing agent that can oxidize a large number of other substances. K is a spectator ion. Consider the reaction between Fe2+ and MnO4-. Reducing Agent o NaI and KI are common reducing agents because I- can be oxidized by a large number of other substances. Consider the reaction between I- and OCl-. 12 When 25.00 mL of a solution containing an unknown concentration of Fe2+ is titrated to an endpoint with acidic KMnO4, the titration requires 17.25 mL of acidified 0.1000 M KMnO4. The equation for the reaction is: MnO4- + 8 H+ + 5 Fe2+ → Mn2+ + 4 H2O + 5 Fe3+ What is the concentration of Fe2+ in the solution? 13 A 25.00 mL sample of bleach is reacted with excess KI according to the equation: 2 H+ + OCl- + 2 I- → Cl- + H2O + I2 Using starch solution as an indicator, the I2 produced requires exactly 46.84 mL of 0.7500 M Na2S2O3 to bring the titration to the endpoint according to the equation: 2 S2O32- + I2 → S4O62- + 2 IWhat is the [OCl-] in the bleach? 14 demonstrate familiarity with at least two common reagents used in redox titrations (e.g., permanganate, dichromate, hydrogen peroxide) select a suitable reagent to be used in a redox titration, in order to determine the concentration of a species calculate the concentration of a species in a redox titration from data (e.g., grams, moles, molarity) (from G3) identify reactants and products for various redox reactions performed in a laboratory, and write balanced equations *These are addressed in labs and in previous notes and calculations.