Chemistry 30 Chapter 14

1. Chemistry 30 Chapter 14 Electrochemical Cells pg. 610-669

2. 14.1 Technology of Cells and Batteries Pg. 612-621

3. Introduction to Electrochemistry • An electric cell converts chemical energy into electrical energy • Alessandro Volta invented the first electric cell but got his inspiration from Luigi Galvani. Galvani’s crucial observation was that two different metals could make the muscles of a frog’s legs twitch. Unfortunately, Galvani thought this was due to some mysterious “animal electricity”. It was Volta who recognized this experiment’s potential. • An electric cell produces very little electricity, so Volta came up with a better design: • A battery is defined as two or more electric cells connected in series to produce a steady flow of current • Volta’s first battery consisted of several bowls of brine (NaCl(aq)) connected by metals that dipped from one bowl to another • His revised design, consisted of a sandwich of two metals separated by paper soaked in salt water (an electrolyte).

4. Cells and Batteries • Alessandro Volta’s invention was an immediate technological success because it produced electric current more simply and reliably than methods that depended on static electricity. • It also produced a steady electric current –something no other device could do.

5. Cells and Batteries • Electric cells are composed of two electrodes (solid electrical conductors) and at least one electrolyte (aqueous electrical conductor) • In current cells, the electrolyte is often a moist paste (just enough water is added so that the ions can move). Sometimes one electrode is the cell container. • In an electric cell or battery, the electrode that is defined as the cathode is marked positiveand the electrode that is defined as the anode is marked negative. • Electricity is the flow of elections. • The electrons flow through the external circuit from the anode to the cathode. • To test the voltage of a battery, the red(+) lead is connected to the cathode (+ electrode), and the black(-) lead is connected to the anode (- electrode)

6. Spontaneous: electrons flow from anode (-) to cathode (+) • Nonspontaneous: electrons flow from anode (+) to cathode (-) • Electrical energy is supplied to force this reaction to occur. (p.640)

7. Basic Cell Design and Properties • A voltmeter is a device that measures the energy difference, per unit charge, between any two points in an electric circuit (called electric potential difference or the voltage (V)) • I.e. A 9V battery releases 6* as much energy compared with the electrons from a 1.5V battery. • The voltage of a cell, depends mainly on the chemical composition of the reactants in the cell, as it is a ratio of energy to charge and is not dependent on size. (eg. AA, B, C or D size batteries) • An ammeter is a device that measures the rate of flow of charge past a point in an electrical circuit (called electric current) • The larger the electric cell, the greater current (amperes (A)) that can be produced. Electric potential difference is like the potential energy difference between 1kg of water at the top of the dam and 1kg of water at the bottom of the dam. Electric current (or flow of electrons) is like the flow of the water. A larger drain would be like a larger electric cell, allowing more water (or electrons) to flow.

8. Basic Cell Design and Properties • The power of a battery is the rate at which it produces electrical energy. Power is measured in watts (W). Calculated using P = IV (Power = current x potential difference) • The energy density is a measure of the quantity of energy stored or supplied per unit mass. Measured in J/kg. • Table 1 summarizes some important electrical quantities and their units

9. Consumer, Commercial, and Industrial Cells • Technology proceeded the understanding of cells and batteries. Prediction-procedure-evidence-analysis cycle was used later on to understand why it worked the way it did. • Zinc chloride cell, dry cell, invented in 1865, Mercury and alkaline dry cell were more advanced dry cells • Problem is that the chemicals eventually become depleted and irreversible reactions… • See Table 2p. 616 Primary cells cannot be recharged, but are relatively inexpensive

10. Consumer, Commercial, and Industrial Cells Secondary cells can be recharged using electricity, but are expensive • E.g. Ni-Cd cells, lead-acid batteries, lithium-ion cell • They discharge and are charged

11. Consumer, Commercial, and Industrial Cells Fuel cells produce electricity by the reaction of a fuel that is continuously supplied. More efficient, and used for NASA vehicles, but still too expensive for general or commercial applications • More: Aluminum-Air cell – developed for possible use in electric cars

13. Homework… • Textbook: • pg. 614 #1-5 • Pg. 619 # 8, 10, 13-15 • Pg. 621 # 1, 5, 8, 12

14. 14.2a Voltaic Cells Pg. 622-626

15. 14.2 – Voltaic Cells • Questions? • Why is it that electrodes and electrolytes determine a cell’s characteristics such as voltage and current? • What happens in different parts of the cell? • To answer these questions, chemists use a cell that has two electrodes and their electrolytes separated.

16. Voltaic Cells (aka Galvanic Cell) • A device that spontaneously produces electricity by redox • Uses chemical substances that will participate in a spontaneousredox reaction. • The reduction half-reaction (SOA) will be above the oxidation half-reaction (SRA) in the activity series to ensure a spontaneous reaction. • Composed of two half-cells; which each consist of a metal rod or strip immersed in a solution of its own ions or an inert electrolyte. • Electrodes: solid conductors connecting the cell to an external circuit • Anode: electrode where oxidation occurs (-) • Cathode: electrode where reduction occurs (+) • The electrons flowfrom the anode to the cathode (“a before c”) through an electrical circuit rather than passing directly from one substance to another • A porous boundary separates the two electrolytes while still allowing ions to flow to maintain cell neutrality • Often the porous boundary is a salt bridge, containing an inert aqueous electrolyte (such as Na2SO4(aq) or KNO3(aq)), • Or you can use a porous cup containing one electrolyte which sits in a container of a second electrolyte.

17. A Salt Bridge can be used to separate the two electrolytes while still permitting ions to move between the two solutions • A porous boundary (unglazed ceramic) can be used as well to separates the two electrolytes while still permitting ions to move between the two solutions

18. Using this design, a cell can be split into two parts connected by a porous boundary • Each part, called a half-cell, consists of one electrode and one electrolyte. • A single line (I) indicates a phase boundary (such as the interface of an electrode and an electrolyte in one half-cell) • A double line (II) represents a physical boundary (such as a porous boundary between half-cells)

19. Voltaic Cells (aka Galvanic Cells) The single line represents a phase boundary (electrode to electrolyte) and the double line represents a physical boundary (porous boundary) Memory device: “Anox ate a redcat” Anode oxidation; reduction cathode Voltaic cells can be represented using cell notation: The SOA present in the cell always undergoes reduction at the cathode The SRA present in the cell always undergoes oxidation at the anode

20. Match the cell notation to the descriptions • Sn(s) Sn4+(aq) Cu2+(aq)Cu(s) • Mg(s) MgCl2(aq) SnCl4(aq)Sn(s) • Sn(s) SnCl2(aq) CuCl2(aq)Cu(s) • Mg(s) Mg2+(aq) Cu2+(aq)Cu(s) • Mg(s) Mg2+(aq) Sn2+(aq)Sn(s) • Sn(s) SnCl2(aq) SnCl4(aq)Sn(s) • Copper placed in a solution of copper(II) chloride and tin metal placed in a solution of tin(II) ions • A copper-magnesium cell • Magnesium in a solution of magnesium chloride and tin in a solution of tin(II) chloride • A tin(IV) ion solution containing tin and a solution of magnesium ions containing magnesium • Two tin electrodes in solution of tin(II) chloride and tin (IV) chloride respectively • Copper place in a copper(II) solution and tin place in a tin(IV) solution • Note: • It is a common practice to write the cell notation with the anode on the left-hand side and the cathode on the right-hand side. *Diagrams or photos of cells can have any arrangement

21. A Voltaic Cell Demonstration • A voltaic cell is the arrangement of two half-cells separated by a porous boundary.

22. Voltaic Cells – What is going on? • Example: Silver-Copper Cell Cu(s) Cu2+(aq)Ag+(aq)Ag(s) • Use the activity series to determine which of the entities is the SOA. The SOA present in the cell always undergoes a reduction at the cathode. Write the reduction half reaction • Use the activity series to determine which of the entities is the SRA. The SRA present in the cell always undergoes an oxidation as the anode. Write the oxidation half-reaction • Balance the half-reactions and add together to create the net equation. The cathode is the electrode where the strongest oxidizing agent present in the cell reacts. The anode is the electrode where the strongest reducing agent present in the cell reacts. Memory device: “Anox ate a redcat” Anode oxidation; reduction cathode

23. Voltaic Cells – What is going on? • Example: Silver-Copper Cell • Silver ions are the strongest oxidizing agents in the cell, so they undergo a reduction half-reaction at the cathode, creating more Ag(s) • Copper atoms are the strongest reducing agents in the cell, so they give up electrons in an oxidation half-reaction and enter the solution (Cu2+ = blue ions) at the anode. • Electrons released by the oxidation of copper atoms at the anode travel through the connecting wire to the silver cathode. (Ag+(aq) win in the tug of war for e-’s over Cu2+(aq)) • Since the positive silver ions are being removed from solution, you would assume that the solution would become negatively charged. This does not happen. Why? • Cations (positively charged ions) move from the salt bridge into the solution in the cathode compartment to maintain an electrically neutral solution. • Anions (negatively charged ions) move from the salt bridge into the solution in the anode compartment to maintain an electrically neutral solution.

24. Voltaic Cell Animation

25. Voltaic Cell Summary • A voltaic cell consists of two-half cells separated by a porous boundary with solid electrodes connected by an external circuit • SOA undergoes reduction at the cathode (+ electrode) – cathode increases in mass • SRA undergoes oxidation at the anode (- electrode) – anode decreases in mass • Electrons always travel in the external circuit from anode to cathode • Internally, cations move toward the cathode, anions move toward the anode, keeping the solution neutral

26. Voltaic Cells with Inert Electrodes The copper electrode will decrease in mass and the blue colour of the electrolyte increases (Cu2+), which indicates oxidation at the anode. The carbon electrode remains unchanged, but the orange colour of the dichromate solution becomes less intense and changes to greenish-yellow (Cr3+), evidence that reduction is occurring in this half cell • Inert electrodes are needed when the SOA or SRA involved in the reaction is not solid. If this is the case, usually a graphite (C(s)) rod or platinum strip is used as the electrode. • Inert (unreactive) electrodes provide a location to connect a wire and a surface on which a half-reaction can occur.

29. Homework… • Textbook pg. 626 #2-8 • Study for the 6 week exam over break! • Enjoy the break! 

30. 14.2b Standard Cells and Cell Potential Pg. 627-633

31. Standard Cells and Cell Potentials • A standard cell is a voltaic cell in which each half-cell contains all entities shown in the half-reaction equation at SATP conditions, with a concentration of 1.0 mol/L for the aqueous solutions. • Standardizing makes comparisons and scientific study easier • Standard Cell Potential, E0 cell = the electric potential difference of the cell (voltage) operating under standard conditions. E°cell represents the energy difference (per unit charge) between the cathode and anode. E0 cell = E0rcathode – E0ranode • Where E0r is the standard reduction potential, and is a measure of a standard ½ cell’s ability to attract electrons, thus undergoing reduction. • The higher the E0r , the stronger the OA

32. The standard cell potential is the difference between the reduction potentials of the two standard half-cells: • All standard reduction potentials are based on the standard hydrogen ½ cell being 0.00V. This means that all standard reduction potentials that are positive are stronger OA’s than hydrogen ions and all standard reduction potentials that are negative are weaker. • If the E0cell is positive, the reaction occurring is spontaneous. • If the E0 cell is negative, the reaction occurring is non-spontaneous

33. Standard Hydrogen Half-Cell • The standard hydrogen half-cell is chosen as a reference and is therefore assigned an electrode potential of exactly zero volts. It is called a reference half-cell. • Standard reduction potentials for all other half-cells are measured relative to that of the standard hydrogen half cell. • The relative potential of the hydrogen ion reduction half-reaction is defined to be zero volts:

34. A reduction potential that has a positive value means that the oxidizing agent is a stronger oxidizing agent than hydrogen ions. • A negative reduction potential means that the oxidizing agent is a weaker oxidizing agent than hydrogen ions.

35. Rules for Analyzing Standard Cells • Determine which electrode is the cathode. The cathodes is the electrode where the strongest oxidizing agent present in the cell reacts. I.e. The OA that is closet to the top on the left side of the redox table = SOA If required, copy the reduction half-reaction for the strongest oxidizing agent and its reduction potential • Determine which electrode is the anode. The anode is the electrode where the strongest reducing agent present in the cell reacts. I.e. The RA that is closet to the bottom on the right side of the redox table = SRA If required, copy the oxidation half-reaction (reverse the half-reaction) • Determine the overall cell reaction. Balance the electrons for the two half reactions (but DO NOT change the E0r) and add the half-reaction equations. • Determine the standard cell potential, E0cell using the equation: E0 cell = E0rcathode – E0ranode

36. Measuring Standard Reduction Potentials • Both voltage and the direction of current must be known. • The magnitude of the voltage determines the numerical value of the half-cell potential, and the direction of the current determines the sign of the half-cell potential. • Half-reaction equations can be multiplied by appropriate factors to balance the electrons (see Text pg.629), but the reduction potentials are not altered by the factors used to balance the electrons. E°cell= E°r (cathode) - E°r (anode) • A positive cell potential (E°cell >0) indicates that the net reaction is spontaneous – a requirement for all voltaic cells.

37. Standard Cells and Cell Potentials #1a

38. Standard Cells and Cell Potentials #1a • Example: What is the standard potential of the cell represented below: • Determine the cathode and anode • Determine the overall cell reaction • Determine the standard cell potential

39. Standard Cells and Cell Potentials #1b

40. Standard Cells and Cell Potentials #1b

41. Standard Cells and Cell Potentials #2 • Example: What is the standard potential of an electrochemical cell made of a cadmium electrode in a 1.0 mol/L cadmium nitrate solution and chromium electrode in a 1.0 mol/L chromium(III) nitrate solution? Cd2+(aq)Cd(s) Cr2+(aq) Cr(s)H2O(l) E0 cell = E0rcathode – E0ranode = (-0.40V) - (-0.91V) = + 0.51V The E0cell is positive, therefore the reaction is spontaneous. SOA SRA cathode anode

42. Standard Cells and Cell Potentials #3 • Example: A standard lead-dichromate cell is constructed. Write the cell notation, label the electrodes, and calculate the standard cell potential. Pb(s) Pb2+(aq) Cr2O72-(aq) H+(aq) Cr3+(aq)C(s) E0 cell = E0rcathode – E0ranode = (+1.23V) - (-0.13V) = + 1.36V The E0cell is positive, therefore the reaction is spontaneous. SRA SOA anode cathode

43. Standard Cells and Cell Potentials #4 • Example: A standard scandium-copper cell is constructed and the cell potential is measured. The voltmeter indicates that copper the copper electrode is positive. Sc(s) Sc3+(aq) Cu2+(aq)Cu(s) E0 cell = +2.36V Write and label the half-reaction and net equations, and calculate the standard reduction potential of the scandium ion. E0 cell = E0rcathode - E0ranode 2.36V = (+0.34V) - (x) E0ranode = -2.02V anode cathode

44. Cell Potentials Under Nonstandard Conditions • The electric potential difference or voltage of a cell decreases slowly as the cell operates • If electrons flow, oxidation and reduction occur, which in turn, changes the concentration from the standard 1.0 mol/L • If electrons are allowed to flow, eventually an equilibrium will be reached when the flow ceases

45. Nonstandard conditions for concentrations, temperature, and pressure will cause differences between the cell potentials predicted from a standard redox table and the ones measured in the laboratory • These differences are generally small if the conditions are relatively close to standard values • Other more important reason for discrepancies are the purity of substances, the presence of oxide coatings on the metals, and the type and size of porous boundaries

47. Homework… • Textbook: • Pg.631 #10,11 • Pg. 633 #12-16

48. 14.2c Corrosion Pg. 634-635

49. Human History… • The Copper ages, Bronze ages, Steel ages… • Describe metals that were widely refined from ores and used for tools and weapons.

50. Corrosion • An electrochemical process in which a metal reacts with substances in the environment, returning the metal to an ore-like state. • Because we live in an oxidizing (oxygen) environment, oxidation (corrosion) of some metals occurs spontaneously. • As a metal is oxidized, metal atoms lose electrons to form positive ions. • Preventing corrosion and dealing with the effects of corrosion are major economical and technological problems for our society. Eg. A ship’s hull.