Chemistry For Engineers Module 1 Electrochemistry PDF

CHEMISTRY FOR ENGINEERS MODULE 1 ELECTROCHEMISTRY ================================================= - Describe at least three types of corrosion and identify chemical reactions responsible for corrosion. - Write and balance half-reactions for simple redox processes. - Use standard reduc...

CHEMISTRY FOR ENGINEERS MODULE 1 ELECTROCHEMISTRY ================================================= - Describe at least three types of corrosion and identify chemical reactions responsible for corrosion. - Write and balance half-reactions for simple redox processes. - Use standard reduction potentials to calculate cell potentials under both standard and nonstandard conditions. - Use standard reduction potentials to predict the spontaneous direction of a redox reaction. - Calculate the amount of metal plated, and the amount of current needed, or the time required for an electrolysis process. INTRODUCTION ============ - Uniform corrosion - occurs evenly over a large portion of the surface area of a metal. - Galvanic corrosion - occurs when two different metals contact each other in the presence of an appropriate electrolyte. - Crevice corrosion - occurs when two pieces of metal touch each other, leaving a small gap or crevice between the metals. Oxidation-Reduction Reactions and Galvanic Cells ================================================ - Oxidation is the loss of electrons from some chemical species. - Reduction is the gain of electrons to some chemical species. Oxidation-Reduction and Half-Reactions ====================================== 𝐴𝑔^+^ + 1𝑒^−^ → 𝐴𝑔~(𝑠)~ 2𝐴𝑔^+^ + 2𝑒^−^ → 2𝐴𝑔~(𝑠)~ +-----------------------------------------------------------------------+ | 𝐶𝑢~(𝑠)~ → 𝐶𝑢^2+^ + 2𝑒^−^ | | | | (𝑎𝑞) | +=======================================================================+ | 2𝐴𝑔^+^ + 2𝑒^−^ → 2𝐴𝑔~(𝑠)~ | +-----------------------------------------------------------------------+ | 𝐶𝑢~(𝑠)~ + 2𝐴𝑔^+^ → 𝐶𝑢^2+^ + 2𝐴𝑔~(𝑠)~ | +-----------------------------------------------------------------------+ - The Cu was oxidized and is the reducing agent. - The Cu facilitated the reduction of Ag^+^ by losing electrons. - The Ag^+^ was reduced and is the oxidizing agent. - The Ag^+^ facilitated the oxidation of Cu by gaining electrons. Building a Galvanic Cell ========================  Figure 3: Galvanic Cell ======================= - NH4+ will flow into the Ag^+^ beaker to offset the removal of Ag^+^ from solution. - Cl^--^ will flow into the Cu^2+^ beaker to offset the production of Cu^2+^ in solution. Terminologies for Galvanic Cells ================================ - The electrode where oxidation occurs is the anode. - The electrode where reduction occurs is the cathode. - Cell notation lists the metals and ions involved in the reaction. - A vertical line, \|, denotes a phase boundary. - A double vertical line, \|\|, denotes a salt bridge. - The anode is written on the left, the cathode on the right. General form of cell notation 𝐶𝑢~(𝑠)~ \| 𝐶𝑢^2+^ (1 𝑀) \|\| 𝐴𝑔^+^ (1 𝑀) \| 𝐴𝑔~(𝑠)~ (𝑎𝑞) (𝑎𝑞) Atomic Perspective on Galvanic Cells ==================================== Galvanic Corrosion and Uniform Corrosion ========================================  Cell Potentials =============== Measuring Cell Potential ======================== - The behavior of cell potentials is akin to state functions. - If a specific standard electrode is chosen, comparison to all other electrodes will result in a practical system for determining cell potential.  - The half-cell notation is: Pt(s) \| H2 (g, 1 atm) \| H^+^ (1 M). - The half-cell is assigned a potential of exactly zero volts. - The cell potential is attributed to the other half-reaction. Figure 8: SHE Cell set-up ========================= Standard Reduction Potentials ============================= Table 1: Standard Reduction Potentials for Several Half-Reactions Involved in the Cells =======================================================================================  𝐸^°^ = 𝐸^°^ − 𝐸^°^ 𝑐𝑒𝑙𝑙 Problem 1 ========= 𝐶𝑢~(𝑠)~ \| 𝐶𝑢^2+^ (1 𝑀) \|\| 𝐴𝑔^+^ (1 𝑀) \| 𝐴𝑔~(𝑠)~ (𝑎𝑞) (𝑎𝑞) Nonstandard Conditions ====================== *F* = 96,485 J V^-1^ mol^-1^ or 96,485 C mol^-1^ 𝑄 = Problem 2 ========= [ 𝐽] ) 𝑉𝑚𝑜𝑙 Cell Potentials and Free Energy =============================== Problem 3 ========= 𝑉𝑚𝑜𝑙 Equilibrium Constants ===================== Figure 11: Cell Potential vs. logK ================================== ∆G^0^ = -RT ln K Batteries ========= Primary Cells ============= 2𝑀𝑛𝑂~2(𝑠)~ + 𝐻~2~𝑂~(𝑙)~ + 2𝑒^−^ → 𝑀𝑛~2~𝑂~2(𝑠)~ + 2𝑂𝐻^−^  - Lithium is the anode. - Manganese(IV) oxide is the cathode as in the alkaline battery, but in this case the MnO2 reacts with lithium ions. - Zinc is the anode. - Oxygen reacts at the cathode. - In a zinc-air battery, one of the reactants is oxygen from the surrounding air. As a result, these batteries can offer a very attractive energy density.  Secondary Cells =============== - Nickel-metal-hydride batteries are an example of secondary cells. The anode reaction is The complex cathode reaction can be represented as Nickel-metal-hydride batteries have become popular as rechargeable cells. ========================================================================= The lead-acid storage battery found in cars is a secondary cell. - The anode for a lead-acid battery is lead metal. - The cathode for a lead-acid battery is lead oxide.  Fuel Cells ========== Limitations of Batteries ======================== 𝐻~2~ → 2𝐻^+^ + 2𝑒^−^ Electrolysis ============ - Passive electrolysis: the electrodes are chemically inert materials that simply provide a path for electrons. - Active electrolysis: the electrodes are part of the electrolytic reaction. Electrolysis and Polarity =========================  Passive Electrolysis in Refining Aluminum ========================================= Active Electrolysis and Electroplating ====================================== 𝐴𝑔~(𝑠)~ + 2𝐶𝑁^−^ → 𝐴𝑔𝐶𝑁^−^ + 𝑒^−^ (𝑎𝑞) 𝐴𝑔(𝐶𝑁)^−^ + 𝑒^−^ → 𝐴𝑔~(𝑠)~ + 2𝐶𝑁^−^ Barrel plating is often used to apply coatings to small parts Electrolysis and Stoichiometry ============================== Current and Charge ================== Problem 4 ========= 96,485 𝐶 Problem 5 ========= 3.6𝑜 𝑥 10^6^𝐽 Calculations Using Masses of Substances in Electrolysis ======================================================= Problem 6 ========= 𝑚𝑜𝑙 Problem 7 ========= 𝐼 Batteries in Engineering Design ===============================  Figure 14: Lithium-ion Batteries ================================ Figure 13: Schematic of Lithium-ion Battery =========================================== - Li^+^/Li has one of the largest standard reduction potentials - Both lithium and carbon are relatively light 1. For each of the following oxidation-reduction reactions, identify the half reactions and label them as oxidation or reduction a. Cu (s) + Ni^2+^ (aq) Ni (s) + Cu^2+^ (aq) b. 2 Fe^3+^ (aq) + 3 Ba^2+^ (aq) = 2 Fe(s) 2. Write a balanced chemical equation for the overall cell reaction in each of the following galvanic cells. a. Ag(s) \| Ag^+^(aq) ║ Sn^4+^(aq), Sn^2+^(aq) \| Pt(s) b. Al(s) \| Al^3+^(aq)║ Cu^2+^(aq) \| Cu(s) 3. Based on the cell potential measured for the cells 4. Using values from the tables of standard reduction potentials, calculate the cell potentials of the following cells: a. Fe(s) \| Fe^2+^(aq) ║ Hg^2+^(aq) \| Hg(ℓ) b. Pt(s) \| Fe^2+^(aq), Fe^3+^(aq) ║ MnO ^--^(aq), H^+^(aq), Mn^2+^(aq) \| Pt(s) c. Pt(s) \| Cl2(g) \| Cl^--^(aq) ║ Au^+^(aq) \| Au(s) 5. Suppose that you cannot find a table of standard reduction potentials. You remember that the standard reduction potential of Cu^2+^ + 2 e^--^ → Cu(s) is 0.337 V. Given that ∆*G*f°(Cu^2+^) = 65.49 kJ mol^--1^ and that 6. Assume the specifications of a Ni-Cd voltaic cell include delivery of 0.25 A of current for 1.00 h. What is the minimum mass of the cadmium that must be used to make the anode in this cell? Ans. 0.52 grams. 7. If a current of 15A is run through an electrolysis cell for 2.0 hours, how many moles of electrons have moved? 8. In a copper plating experiment in which copper metal is deposited from copper (II) ion solution, the system is run for 2.6 hours at a current of 12.0A. What mass of copper is deposited?