AP CHEMISTRY NOTES 12-1 ELECTROCHEMISTRY: ELECTROCHEMICAL CELLS Review: OXIDATION-REDUCTION REACTIONS – the changes that occur when electrons are transferred between reactants (also known as a redox reaction) OIL RIG Oxidation Is Loss of electrons . Reduction Is Gain of electrons Oxidation and reduction ALWAYS occur simultaneously Oxidation – the loss of electrons by a substance (the oxidation number rises, becoming more positive) Reduction – the gain of electrons by a substance (the oxidation number lowers, becoming more negative) Oxidizing Agent – the substance in a redox reaction which causes another substance to be oxidized (it itself is reduced) Reducing Agent – the substance in a redox reaction which causes another substance to be reduced (it itself is oxidized) Example: Balance the following redox reaction: Al + Sn4+ → Al3+ + Sn ELECTROCHEMICAL CELLS There are two main types of electrochemical cells: Galvanic (Voltaic) Cells – these cells are spontaneous and react with no outside energy input *These cells convert chemical energy into electrical energy (ie. batteries) Electrolytic Cells – these cells are non-spontaneous and require energy input in order for the reaction to occur *This is the type of cell used to electroplate metals onto the surface of other metals In an electrochemical cell, oxidation and reduction occur simultaneously. A typical galvanic (voltaic) cell: anode – the electrode where oxidation occurs (AN OX) *This is the negative electrode cathode – the electrode where reduction occurs (RED CAT) *This is the positive electrode Electron flow always occurs through the wire from the anode to the cathode (FAT CAT) salt bridge – a device used to maintain electrical neutrality in the cell by supplying ions to replace the charges; generally consists of a neutral salt *anions move to the anode side, cations move to the cathode side voltmeter – measures the cell potential (emf), usually in volts Example: In the previous diagram, what reactions occur at the anode and the cathode? In a standard electrochemical cell, substances are always in their standard states, and solutions are 1.0 M. To show the contents of an electrochemical cell, we use Standard Cell Notation: Reaction: M + N+ → N + M+ Standard Cell Notation: M | M+ || N+ | N (anode) (cathode) *If the “electrode” is an ion or a gas, then an “inert electrode” must be used (graphite or platinum are common) EXAMPLE: Draw the diagram of an electrochemical cell based on the following cell notation. Use potassium nitrate as the salt bridge. Be sure to clearly label the electrodes and solutions involved, as well as the direction of electron flow. In addition, show a balanced chemical equation for the reaction. Ni | Ni2+ || Ag+ | Ag AP CHEMISTRY NOTES 12-2 ELECTROCHEMISTRY – CALCULATING CELL POTENTIAL CELL POTENTIAL Electromotive Force (“emf” or “E” or “ε”) – the force which moves electrons from the anode to the cathode In order to determine whether a process is spontaneous, the cell potential must be calculated. The cell potential refers to a measure of electromotive force (emf) for a cell. This value can be determined by combining the potentials of both the oxidation and reduction half-reactions. To read a reduction potential chart: *The half-reactions are written as reductions. *The elements with more positive reduction potentials are more easily reduced. *The elements with more negative reduction potentials are more easily oxidized. EXAMPLE: Will a reaction occur if a piece of tin is dipped into a solution of copper(II) nitrate? EXAMPLE: What will happen if a silver bracelet is dropped into a solution of copper(II) nitrate? EXAMPLE: Which of the following is the stronger oxidizing agent? bromine sodium Which of the following is the stronger reducing agent? chromium mercury o Calculation of Cell Potential (E cell) To calculate cell potential: 1. Write oxidation and reduction half-reactions (use the reduction potential chart to determine which substance is oxidized and which is reduced if you are not given this information). 2. Use the reduction potential chart to find voltages for each half-reaction. 3. Change the sign of the voltage for the oxidation reaction (and re-write as oxidation if necessary) 4. Add the two voltages together (but do not multiply the voltages by the integers used to balance the half-reactions!) o o o E cell = E ox + E red 5. Determine the spontaneity of the cell: o *Positive E cell – spontaneous (galvanic or voltaic cell) o *Negative E cell – non-spontaneous (electrolytic cell) EXAMPLE: Calculate the cell voltage for the following reaction. In addition, determine the type of cell that would form as well as a balanced reaction for the process. Cr3+ + Cu → Cr + Cu2+ EXAMPLE: Calculate the voltage for the following reaction. In addition, determine the type of cell this would form as well as a balanced reaction for the process. Fe + Cu2+ → Cu + Fe2+ Recall that an electrolytic cell uses electrical energy to produce a chemical reaction (ie. electroplating, electrolysis, etc.) The only difference between electrolytic and galvanic cells is that in electrolytic cells: *the anode is positive (often an inert electrode is used) *the cathode is negative (often an inert electrode is used) *a battery (or some source of emf) must be used to “push” the reaction AP CHEMISTRY NOTES 12-3 APPLICATIONS IN ELECTROCHEMISTRY Faraday’s Law – the amount of a substance being oxidized or reduced at each electrode during electrolysis is directly proportional to the amount of electricity that passes through the cell Calculations involving electrolytic cells are best done by simply using dimensional analysis and a few key conversions factors: 1 Volt (V) = 1 Joule/Coulomb 1 Amp (A) = 1 Coulomb/Second 1 Faraday (F) = 96,500 Coulombs/Mole of electrons EXAMPLE: What mass of cobalt metal is plated when a current of 15 amps is passed for 60. minutes through a solution containing cobalt(II) ions? EXAMPLE: How long will it take to plate out one kilogram of aluminum from aluiminum ions using a current of 100 amps? SPONTANEITY OF CELLS Thermodynamically Favored (Spontaneous): ∆H < 0 ∆S > 0 K > 1 ∆G < 0 Ecell > 0 The relationship between free energy and the spontaneity of an electrochemical cell can be shown using the following equations: ∆Go = - n F Eo ∆Go = - RT ln K Where G = free energy n = # mol of electrons transferred F = Faraday’s Constant (96,500 C/mol e-) Eo = cell potential R = 8.3145 J/K.mol T = temperature (K) K = equilibrium constant o EXAMPLE: Calculate the Ecell and ∆G at 25 C for the following reaction. In addition, determine the equilibrium constant for the reaction. Cu2+ + Fe → Cu + Fe2+ Determination of Cell Potential at Other Than Standard Conditions To calculate the potential of a cell in which some or all components are not in standard states, the Nernst Equation is used: o EXAMPLE: Calculate the voltage of the following reaction (at 25 C) with the following conditions: 2+ 3+ o 2Al + 3Mn → 2Al + 3Mn E cell = 0.48 volts [Mn2+] = 0.50 M [Al3+] = 1.5 M .