Electrochemistry

1. Electrochemistry Chapter 17 AP Chemistry

2. Decoding the Past – The Real Dr. Frankenstein • As you watch the following video, answer the questions on the sheet provided: Decoding the Past - The Real Dr. Frankenstein

3. Review of Terms • Oxidation-reduction reaction (Redox) • Involves a transfer of electrons from the reducing agent to the oxidizing agent • Oxidation • Loss of electrons • LEO • OIL • Reduction • Gain of electrons • GER • RIG • Oxidation number • Oxidizing agent • Reducing agent • Half-reactions • Overall reaction is split into two half-reactions, one involving oxidation and one reduction Work on the sample Review Problem

4. Review Problem • Balance the following oxidation-reduction reaction IN BASIC SOLUTION using the half-reaction method. Be sure to identify the oxidizing agent and reducing agent. HXeO4-(s)  XeO64- + Xe (g)

5. Forefathers of Electrochemistry • Luigi Galvani • Italian physician who observed a frog’s leg twitch when it was touched with two different metals • In attempting to explain what happened, Galvani thought that the animal tissue in the frog’s leg was the source of electricity • Alessandro Volta • Italian physicist who disputed Galvani’s hypothesis • Resulting controversy resulted in discovery that electric currents could be produced by chemical reactions • Volta used this discovery to create the first chemical battery

6. What is Electrochemistry? • In general, all chemistry is electrical in the sense that it involves the behavior of electrons and other charged particles • The term electrochemistry is reserved specifically for the study of the interchange of chemical and electrical energy

7. Today’s Technology and Electrochemistry • All of the following involve the principles of electrochemistry: • Remote controls for TVs, DVD players, CD players, stereos • Itty bitty teeny tiny batteries • Calculators • Silverware • Metal-plated jewelry

8. The Galvanic Cell (aka Voltaic cells)A Type of Electrochemical Cell • Defined as a device in which chemical energy is changed to electrical energy • Examples – batteries and fuel cells • Name comes from the work of Volta and Galvani • Uses a spontaneous redox reaction to produce a current that can be used to do electrical work

9. Galvanic Cells and Their Construction • Electrons flow from one terminal to the other when the terminals are connected by an external circuit • Terminals are called electrodes • Anode • Oxidation occurs here • Electrons leave the cell here • Cathode • Reduction occurs here • Electrons are accepted by the species being reduced and enter the cell here • Electrodes are submerged in an electrolyte • A salt solution that contains ions • Electrolyte may be involved in the reaction or the ions may be used to carry the charge • Can contain a salt bridge or a porous-disk connection in order to neutralize charge buildups in electrode compartments • Completes the circuit!

10. Electrode Compartments in Galvanic Cells • Anode • Oxidation occurs here • Cathode • Reduction Occurs here

11. Galvanic Cells and Their Construction Salt Bridge Porous Disk

12. A Closer Look at a Galvanic (Voltaic) Cell Galvanic Cells

13. Cell Potentials and Electrochemistry

14. Helpful Neumonic Devices • AN OX • Oxidation occurs at the anode • RED CAT • Reduction occurs at the cathode • FAT CAT • The electrons in a galvanic (voltaic) cell always flow From the Anode To the CATode

15. Cell Potential • Recall that a galvanic cell consists of an oxidizing agent in one compartment that “pulls” electrons through a wire from a reducing agent in the other compartment • The “pull” or driving force on the electrons is called the CELL POTENTIAL or electromotive force (emf) • If pull occurs spontaneously, cell is a good battery!

16. Measuring Cell Potential • A VOLTMETER is used to measure cell potential • The unit of electrical potential is the volt (V) • Defined as 1 joule of work per coulomb of charge transferred

17. How to Describe a Galvanic Cell • The cell potential (always positive for a galvanic cell) and the balanced cell reaction is written somewhere on the diagram • The direction of electron flow is given • Obtained by inspecting the half-reactions and using the direction that gives a positive E0cell • The anode and cathode are designated • The nature of each electrode and the ions present in each compartment are labeled • A chemically inert conductor such as Pt is required if none of the substances participating in the half-reaction is a conducting solid • Example – Fe2+ and Fe3+

18. Standard Reduction Potentials • Reaction in a galvanic cell is always an oxidation-reduction reaction that can be broken down into two half-reactions • Each half-reaction has a cell potential • We can obtain the overall cell potential by summing the half-cell potentials! • A cell will always run spontaneously in the direction that produces a POSITIVEcell potential • Each potential is measured against a standard called the STANDARD HYDROGEN ELECTRODE (SHE) • SHE consists of a piece of inert Platinum that is bathed by hydrogen gas at 1 atm • SHE is assigned a potential of ZERO volts

19. Standard Reduction Potentials Using SHE Reduction Half-Reaction Standard Hydrogen Electrode

20. Table of Standard Reduction Potentials • Half-reaction cell potentials are listed in a convenient table! • The values in the table correspond to REDUCTION half-reactions with all solutions at 1 M, all gases at 1 atm, and 25°C (298K) for all Cu2+ + 2 e-→ Cu E0 = -0.34 V versus SHE SO42- + 4 H+ + 2 e-→ H2SO3 + H2O E0 = 0.20 V versus SHE Symbol for standard conditions!

21. How to Read Table of Standard Reduction Potentials • Elements that have the MOST POSITIVE reduction potentials are easily REDUCED • In general, non-metals • Elements that have the LEAST POSITIVE reduction potentials are easily OXIDIZED • In general, metals • Table can also be used to tell the strength of various oxidizing and reducing agents • Another form of the activity series • Metals having LESS POSITIVE reduction potentials are MORE active and will replace metals with more positive potentials

22. Calculating Overall Standard Cell Potential (E0cell) Given Standard Reduction Potentials • Decide which element is oxidized or reduced using the table of reduction potentials • THE MORE POSITIVE REDUCTION POTENTIAL GETS TO BE REDUCED • Write both equations AS IS from the chart with their voltages • REVERSE the equation that will be OXIDIZED and change the sign of the voltage! • This is now E0oxidation • Balance the two half-reactions using integers • Number of electrons lost must equal number gained • DO NOT MULTIPLY VOLTAGE VALUES • Add the two half reactions and the voltages together

23. Standard Reduction Potentials for Many Common Half-Reactions At 25ºC (298 K)

24. Practice! • Consider a galvanic cell based on the reaction: • Give the balanced cell reaction and calculate E0 for the cell

25. Very Handy Line Notation • Line notation can be thought of as an “Ion Sandwich” in alphabetical order Anode metal | Anode ion || Cathode ion | Cathode metal • “|” indicates phase boundary (solid → solution or gas or solution or gas → solid) • “||” indicates salt bridge • Phase (AND concentration if not 1M) specified in parentheses • ZnSO4 (aq, 0.5M) • A comma should be used to separate 2 components in the same phase

26. Practice! • Calculate the cell voltage for the following reaction. Draw a diagram of the galvanic cell for the reaction and label completely • See previous slide for requirements of a complete diagram Fe3+ (aq) + Cu (s) →Cu2+ (aq) + Fe2+ (aq)

27. Cell Potential, Electrical Work, and Free Energy Combining thermodynamics, electrochemistry, and not to mention a bit of physics!

28. emf and Work • The work that can be accomplished when electrons are transferred through a wire depends on the “push” or emf (electromotive force) • emf is defined in terms of potential difference (in volts) between two points in the circuit • Recall that a volt represents a joule of work per coulomb of charge transfered • Thus, one joule of work is produced (or required) when one coulomb of charge is transferred between two points in the circuit that differ by potential of one volt

29. emf and Work • Work is viewed from the point of view of the system • Therefore, if work flows OUT of the system, it is assigned a MINUS sign • When a cell produces a current (aka a battery), the cell potential is positive and the current can be used to do work (like running a motor) • Therefore, emf and work have opposite signs!

30. Whoa….what in the world is “q”? • q is the quantity of charge in coulombs transferred • The charge of 1 mole of electrons is a constant called the faraday (F) • Has the value of 96,485 coulombs of charge per mole of electrons • So,

31. Before Continuing…Some Insight into Spontaneity and Gibb’s Free Energy • In chemistry, we often refer to processes as either spontaneous or nonspontaneous • A spontaneous process is said to occur if it occurswithout outside intervention • Has nothing to do with the speed of the reaction • To explore the idea of spontaneity, consider the following physical and chemical processes: • A ball rolls down a hill but never spontaneously rolls back up the hill • If exposed to air and moisture, steel rusts spontaneously. However, the iron oxide in rust does not spontaneously change back to iron metal and oxygen gas • Heat flow always occurs from a hot object to a cooler one. The reverse process never occurs spontaneously

32. Spontaneity and Gibb’s Free Energy • The driving force for a spontaneous process is an increase in the entropy of the universe • Entropy can be viewed as a measure of molecular randomness or disorder • Natural progression of things is from order to disorder (from lower entropy to higher entropy) • Entropy is related to another thermodynamic quantity called Gibb’s Free Energy (symbolized by G) – more on this later! • A process at constant temperature and pressure is spontaneous in the direction in which the free energy decreases

33. More on Gibb’s Free Energy • Gibb’s free energy is qualitatively useful by telling us whether a process is spontaneous or not • It is quantitatively useful because it can tell us how much work can be done with a given process • Thus, G is defined as the energy available in a system that is available to do useful work • Maximum possible work obtainable from a process at constant temperature and pressure is equal to the change in free energy:

34. Gibb’s Free Energy and Work • ∆G for a spontaneous process represents the energy that is free to do useful work • ∆G for a nonspontaneous process represents the minimum amount of workthat must be expendedto make the process occur

35. So, How Does This Relate to Galvanic Cells? • Recall that for a galvanic cell: • And: • Since: • We can make some substitutions to come up with a relationship between Gibb’s Free Energy, work, and cell potential at constant temperature and pressure:

36. Summary of Free Energy and Cell Potential G = Gibb’s Free Energy n = number of moles of electrons F = Faraday constant = 96, 485 coulombs per mole of electrons • This relationship is important because it confirms that a galvanic cell will run in the direction that gives a positive value for E0 • +E0corresponds to a negative ∆G value (spontaneous) • -E0corresponds to a positive ∆G value (nonspontaneous)

37. Practice! Using the Table of Standard Reduction Potentials, calculate ∆G0 for the reaction: Is this reaction spontaneous?

38. Dependence of Cell Potential on Concentration

39. Concentration and Le Chȃtelier’s Principle • So far, we have described galvanic cells under standard conditions • All solutions at 1 M, all gases at 1 atm, and 25°C (298K) for all • What would happen to the cell potential if the solutions were not at 1M? • Can be answered qualitatively in terms of Le Chȃtelier’sPrinciple

40. A Quick Rundown of Equilibrium • Virtually everything you encounter, including your own bodily processes and senses, is the result of one or more equilibrium reactions • Chemical equilibrium can be defined as the condition where reactant and product concentrations remain constant • Occurs when a forward reaction and its reverse reaction proceed at the same rate • While concentrations do not change, products and reactants continue to interconvert at equal rates

41. A Quick Rundown of Le Chȃtelier’s Principle • If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance

42. A Table that Summarizes Le Chȃtelier’s Principle

43. Effect of Concentration on E0 • An increase in reactant concentration will favor the forward reaction and thus, increase the driving force on the electrons • Ewill increase • An increase in product concentration will oppose the forward reaction and thus, decrease the driving force on the electrons • Ewill decrease

44. Practice! • For the cell reaction: Predict whether Eis larger or smaller than E0 for the following cases: • [Al3+] = 2.0 M, [Mn2+] = 1.0 M • [Al3+] = 1.0 M, [Mn2+] = 3.0 M

45. Concentration Cells • Because cell potentials depend on concentration, we can construct galvanic cells where both compartments contain the same components but at different concentrations • An increase in concentration of reactant will increase cell potential • Increase in product concentration will decrease cell potential

46. For a More Quantitative Approach, Use the Nernst Equation! • This equation is used to calculate the potential of a cell in which some or all of the components are not in their standard states • Remember, standard states are 1M and gases at 1 atm

47. Oh My Goodness…There’s So Many Variables! • R = gas constant 8.315 J/K·mol • F = Faraday constant • 96, 485 colombs per mole of electrons • E = Energy produced by reaction • T = Temperature in Kelvin • n = number of electrons exchanged in BALANCED redox equation • Q (Reaction Quotient)

48. What is a Reaction Quotient (Q)? • A reaction quotient is an expression that is obtained by applying the law of mass action • The rate of a chemical reaction is directly proportional to the products of the reactants • Expression uses using initial concentrations of substances rather than equilibriumconcentrations • Always written like so:

49. More on The Nernst Equation • Potential calculated from the Nernst equation is the maximum potential before any current flow has occurred • As the cell discharges and current flows from anode to cathode, the concentration will change • As reactants are being converted to products, Ecell will decrease • Eventually, the cell potential reaches zero • Zero potential means reaction is at equilibrium (a dead battery) Also, Q = K (equilibrium constant) And ∆G = 0 as well

50. Practice!