1. Chapter 13 Oxidation-Reduction and Electrochemistry

2. Oxidation States The oxidation state of an atom indicates the number of electrons that it gains or loses when it forms a bond. Three important things to remember… –The oxidation state of an element is zero (even in the diatomic state). –The oxidation numbers for all the atoms in a molecule must add up to zero. –The oxidation numbers for all the atoms in a polyatomic ion must add up to its charge.

3. Most elements have different oxidation numbers that can vary depending on the molecule of which they are part. ElementOxidation # Alkali metals+1 Alkaline earths+2 Al, B+3 Oxygen-2 Halogens Hydrogen+1 This chart shows some elements that consistently take the same oxidation numbers. Remember that the transition metals can have varied oxidation numbers. The roman numeral will tell you the charge. Learn your polyatomic ions…given to you the first week of class

4. Examples…find oxidation #s H 2 SO 4 H 2 SO 3 Cr 2 O 7 2- CrO 4 2- N 2 O 5 P 4 O 10

5. Oxidation-Reduction Reactions In a redox reaction, electrons are exchanged by the reactants. –When an atom gains electrons, its oxidation # decreases, and it is said to have been reduced. –When an atom loses electrons, its oxidation # increases, and it is said to have been oxidized. Remember…LEO the lion goes GER –LEO…lose electrons oxidation –GER…gain electrons reduction

6. Oxidation and reduction go hand in hand. If one atom is losing electrons, another atom must be gaining them. –If an atom is losing electrons and being oxidized, it must be giving the electrons to another atom, which is being reduced. So if a reactant contains an atom that is being oxidized, it is a reducing agent. –If an atom is taking electrons and being reduced, it must be taking electrons away from another atom, which is being oxidized. So if a reactant contains an atom that is being reduced, it is an oxidizing agent.

7. Half-reactions An oxidation-reduction reaction can be written as two half-reactions: one for the reduction and one for the oxidation. Ex.Fe + 2HCl → FeCl 2 + H 2

8. Reduction Potentials Every half-reaction has a potential, or voltage, associated with it. The potentials are given as reduction half- reactions, but you can read them in reverse. –If the half-reaction is used in reverse, change the sign on the voltage. –If the half-reaction is doubled or halved, DO NOT change the voltage.

9. Determining E o cell 1.Determine each half-reaction 2.Find the voltage for each half-reaction 3.E o cell = E oxidation + E reduction IF E o cell > 0, the reaction takes place producing the voltage. IF E o cell < 0, the reaction will not take place, but the reverse reaction will take place.

10. Let’s go back to Fe + 2HCl → FeCl 2 + H 2 Determine the E o cell

11. Voltage and Spontaneity A redox reaction will occur spontaneously if its potential has a positive value. We also know from thermodynamics, that a reaction occurs spontaneously when the ΔG < 0. The relationship between reaction potential and free energy is given below. ΔG o = Standard Gibbs free energy change (kJ/mol) n = the number of moles of electrons exchanged in the reaction (mol) F = Faraday’s constant, 96,500 coulombs/mole E o = standard reaction potential (V)

12. Two important things… 1.If E o is positive, ΔG o is negative and the reaction is spontaneous. 2.If E o is negative, ΔG o is posative and the reaction is nonspontaneous.

13. Voltage and Equilibrium The standard reaction potential is related to the equilibrium constant for a reaction by … E o = standard reduction potential R = the gas constant, 8.31 (volt-coulomb)/(mol-K) T = absolute temperature n = the number of moles of electrons exchanged in the reaction F = Faraday’s constant, 96,500 coulombs/mol K = equilibrium constant

14. Voltage & Equilibrium, cont. At 25 o C, this equation simplifies to –If E o is positive, then K > 1 and the forward reaction is favored. –If E o is negative, then K < 1 and the reverse reaction is favored.

15. Galvanic Cells In a galvanic cell (AKA voltaic cell), a spontaneous redox reaction is used to generate a flow of current. (see picture next slide) In a galvanic cell, the two half-reactions take place in separate chambers, and the electrons that are released by the oxidation reaction pass through a wire to the chamber where they are consumed in the reduction reaction. That’s how current in created. Current is defined as the flow of positive charge, so current is always in the opposite direction from the flow of electrons.

17. Galvanic cells, cont. In any electric cell … – oxidation takes place at the anode – reduction takes place at the cathode Remember….AN OX, RED CAT –anode, oxidation –Reduction, cathode

18. Galvanic cells, cont. The salt bridge maintains electrical neutrality in the system by providing enough negative ions to equal the positive ions being created at the anode (during oxidation) and providing positive ions to replace the positive ions being used up at the cathode (during reduction). The salt bridge can be an actual salt, or it can be a slim passage that allows ions to move between the two chambers.

19. Under standard conditions (all concentrations are 1M) the voltage of the cell is the same as the total voltage of the redox reaction. Under nonstandard conditions, the cell voltage can be computed by using the Nernst Equation. E cell = cell potential under nonstandard conditions (V) E o cell = cell potential under standard conditions (V) R = the gas constant, 8.31 (volt-coulomb)/(mol-K) T = Kelvin temperature n = number of moles of electrons exchanged in the reaction (mol) F = Faraday’s constant, 96,500 coulombs/mole Q = the reaction quotient

20. At 25 o C, the Nernst equation reduces to The most important thing you should understand from the Nernst Equation –As the concentration of the products of a redox reaction increases, the voltage decreases. –As the concentration of the reactants in a redox reaction increases, the voltage increases.

21. Electrolytic Cells In an electrolytic cell, an outside source of voltage is used to force a nonspontaneous redox reaction to take place. Electrolytic cells are used for electroplating Use the following formula for determining any missing variable when electroplating.

23. What do you have to know for test? Write half-reactions Determine who is oxidized, who is reduced, who is the oxidizing agent, and who is the reducing agent. Determine E o cell Relate E o cell to equilibrium Relate E o cell to free energy