Electrochemical Methods: Intro Electrochemistry Basics Electrochemical Cells The Nernst Equation Activity Reference Electrodes (S.H.E) Standard Potentials Note: I think you should read Chapter 14 of Harris FIRST, before reading the Skoog book.

Review of the Basics Oxidation –Loss of electrons –Always occurs at the anode –Happens because of the action of a reducing agent Reduction –Gain of electrons (charge is reduced) –Always occurs at the cathode –Happens because of the action of the oxidizing agent

Electric Charge (in Coulombs) and Work The charge in coulombs (q) is equal to the number of moles of electrons (n) times the Faraday Constant (F) Voltage represents electrical potential (potential to do work) If some total charge in coulombs (q) is moved through some electrical potential (E, in volts V) then work is done!

Ohm’s Law and Power Power is the work done in some unit time (e.g. joules of work per second) The units of Power are Watts (W) Ohm’s law and power are related! Ohm’s law relates electrical resistance, current and potential!

Let’s Work Some Problems A 6.00V battery is connected across a 2.00 K  resistor, how many electrons flow through the circuit per second? How many joules of heat (heat is work) are produced per electron? What voltage would the battery need to be to deliver a power at Watts?

Electrochemical Cells A complete cell contains: –anode –cathode –completed circuit (for electrons to flow) –a salt bridge (usually!) –an electrolyte solution –chemical species that undergo reaction. There are two basic electrochemical cells: –A GALVANIC cell uses spontaneous chemical reactions to generate electricity –In a “electrolytic” cell, an electrical potential is applied to the cell to drive some reaction.

What is happening at the electrode(s) and how do we describe the cell?

The Standard Hydrogen Electrode (SHE) The basis by which all other measurements are made. Assigned a potential of zero by definition! Not practical for regular use

Standard Potentials Standardized potentials (E o ), listed as reductions, for all half-reactions Measured versus the S.H.E (0) Used in predicting the action in either a galvanic cell or how much energy would be needed to force a specific reaction in a non-spontaneous cell Assumes an activity of one for the species of interest (usually a fair approximation) at a known temperature in a cell with the S.H.E. Assumes that the cell of interest is connected to the (+) terminal of the potentiometer (voltmeter) and the S.H.E. is connected to the (-) terminal

Better Oxidizing Agents in upper left hand corner. Better Reducing Agents in lower Right hand corner

A cell is constructed (all activities equal 1), in which iron (III) is oxidizes copper (II). Draw a picture of the cell, write the shorthand diagram, and calculate the voltage of the cell from Standard Potentials!

Cells and the Nernst Equation For any cell, the measured potential between the anode and the cathode can be calculated:

Using E cell Write the calculation as described earlier Calculate E + and E - Calculate E cell Write a balanced cell reaction, by adding the two half-reactions –Write out the right cell half reaction –Write out the left cell half reaction and reverse it –Add the two reactions together to get a net, balanced cell reaction. If, you use the conventions described here,then: –If E cell >0, the reaction is spontaneous to the right –If E cell <0, the reaction is spontaneous to the left

Nernst Equation Accounts for potentials of cells where the reagents are not at an activity of 1 –Remember that standard potentials are at A=1 Accounts for the number of electrons transferred in a reaction, the temperature of the reaction, LeChatelier’s Principle and a variety of other factors Used to calculate E + and E - under non-standard conditions –Most real cases!

Arsenic solid is reduced to arsine (AsH 3(g) ). Write the Nernst Equation for the reaction. Find E when pH = 3.00 and the partial pressure of arsine is 1.00 torr.

The Nernst Equation for Complete Reactions….. Setup a series of two Nernst equations –One for E + –One for E - Solve each Nernst equation to get either E + or E - Add together to get E cell

A galvanic cell is assembled in which the left cell is the anode and a cadmium metal electrode is oxidized to cadmium ion in cadmium nitrate. In the right cell, the cathode, silver ion is reduced to silver metal on a silver metal electrode in 0.50M silver nitrate. 1. Draw the cell (both a picture and a schematic diagram) 2. Write the half and net cell reactions 3. Calculate the net cell voltage 4. Indicate in which direction the cell is spontaneous

Using the Nernst Equation to Solve for a Cell as a Chemical Probe! Best done by example. Solve for some unknown quantity! The cell is described as follows. Solve for the concentration of chloride if the measured cell voltage is V

“Complications” in Cells A variety of complicating factors can influence the ability of a cell to have the ideal potential value Some are always present, some can be corrected for or evaluated Their influence is usually less significant than the potential or current imparted by the two half-reactions

Junction Potentials Arise because of the differing mobility of different ions in solution A charge (potential) develops when one ion moves to an electrode more rapidly than the other ion Counteracted by the use of strong electrolytes in solution and/or in the salt-bridge –The strong electrolyte helps overcome mobility-induced differences in charge between two cells Rarely significant in electroanalytical chemistry

Concentration Polarization Arises because of differing analyte concentrations in solution when a reaction is initiated An electrical layer is created at the electrode surface This layer resists the flow of electrical charge unless fresh ions are brought to the electrode –Stirring –Diffusion

Ohmic Potential (IR Drop) Some electrical energy must be used to get ions moving in solution This energy is not registered on the potentiometer, because it is used to impart kinetic energy to the ions So, in reality the calculation of Ecell is: