14.2b Standard Cells and Cell Potential

1. 14.2b Standard Cells and Cell Potential Pg. 627-633

2. 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

3. 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

4. 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:

5. 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.

6. 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

7. 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.

8. Standard Cells and Cell Potentials #1a

9. 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

10. Standard Cells and Cell Potentials #1b

11. Standard Cells and Cell Potentials #1b

12. 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

13. 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

14. 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

15. 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

16. 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

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