Analytical Chemistry-Exam 3- Electrochemistry

Electroanalytical Technique

Group of analytical methods based up on electrical properties of analytes when part of an electrochemical cell

General Advantages

Selective for particular redox state of a species

Cost- $4,000-25,000 for a good instrument compared to $10,000-$50,000-$250,000 for a good spectrometer

Measures activities (not concentration)

Activity usually more physiological importance

Fast

In situ

Info about:

Oxidation states

Stoichiometry

Rates

Charge transfer

Equilibrium constants

Electrochemical Cell: Basic Cell Set Up

Two electrodes

Electrolytes solution

External connection between electrodes (wire)

Internal connection via contact with a common solution or by different solutions connect by a salt bridge

Salt bridge

Acts to isolate two halves of electrochemical cell while allowing migration of ions and current flow

Usually consists of a tube filled w/ KCl

Separate species to prevent direct chemical reactions

Flow of Current (charge) in Cell

Electrons within wires between two electrodes

Ions within solution of each ½ cell (anions and cations) and through salt bridge

Electrochemical reaction at electrode

Reduction

Gain of e- net decrease in charge of species

Oxidation

Loss of e, net increase in charge of species

Net Reaction in Cell

Sum of reactions occurring in the two half cells

potential of overall cell=measure of the tendency of this reaction to proceed to equilibrium

@ equilibrium: potential=0

Larger the potential, the further the reaction is from equilibrium and the greater the driving force that exists

Types of Cells: Galvanic Cells

Reaction occurs naturally

Positive potential

Exothermic-produces energy

Types of Cells: Electrolytic Cell

Reaction does not occur naturally, requires external stimulus (energy) to occur

negative potential

Endothermic-requires energy

Types of Cells: Chemically Reversible Cell

A cell in which reversing the direction of the current simply reverses the chemical reaction

Electrodes

Cathode-electrode where reduction occurs

Anode-electrode where oxidation occurs

e supplied by electrical current via electrode

Species can both be in soultion, solids, or coated electrodes or combination

e is taken up by electrode into electrical circuit

Liquid Junctions

Interface between two solutions with different components of concentration

Small potentials may develop at a junction that affect overall cell potential

Galvanic cell w/o liquid junction

Two species have high potential for reaction but the reaction is slow

Mix two species directly into common solution

Not common

Electrode Potentials

For convenience, represent overall reaction in cell as two ½ reactions

i. one at anode and other at cathode

ii. each ½ reaction has certain potential associated with it

iii. by convention, write both ½ reactions as reduction

iv. potential of cell is defined as:

E_cell=E_cathode-E_anode

Electrode Potentials-Problems

Can not measure potential of just one electrode

i. need to compare to another electrode

ii. determine potential of all ½ cell reactions vs. a common reference electrode

iii. reference electrode-standard hydrogen electrode (SHE)

note: potential affected by pH, [H⁺], used as an early pH indicator, also dependent on P_H2

Standard Electrode Potential (E^o)

Measured E_cell when all species in solution or gas has an activity of 1.))

Activity (a)

proportional to molar concentration

a_x=y_x[X]

y_x is the activity coefficient of solute X

[X] is the molar concentration of solute X

If E^o is “+”

Indicates that the reaction is favored or spontaneous

If E^o is “-”

Indicates that the reaction is not favored or spontaneous and requires energy to proceed

Activity vs. Conecntration

Activity is ALWAYS lower than concentration

Activity coefficients are ALWAYS lower than or equal to 1

H2O forms a hydration sphere around ion

Hydration sphere hinders access by e^- - lower activity

Other ions (ionic strength) make the hydration sphere larger

As E^o increases

oxidizing ability of ½ cell reaction increases

Nernst Equation

Used for non-standard conditions of species

Activity Coefficients

Experimental determination of individual activity coefficients appears to be impossible

Can determine mean activity coefficient

Use Debye-Huckel equation; Note: At ionic strengths > 0.1, Debye-Huckel Equation Fails

Limitations in the Use of Standard Electrode Potentials (E^o)

E^o based on unit activities not concentrations

-activity= concentration only in dilute solutions

-at higher concentrations need to determine and use activity

Formal Potential (E^f or E^o):

used to compensate for problems with E^o in using activity and with side reactions

based on conditions of 1M concentration with all speceis being specified

gives better agreement than E^o with experimental data and Nernst Equation: conditions need to be similar to conditions where E^o was measured

Reaction Rates

some E^o ½ reactions listed in tables have been determined by calculations from equilibrium measurements rather than actual measurements of the ½ cell in an electrode system

Problem:

reaction is slow and difficult to see in practice

thermodynamics vs. kinetics

no suitable electrode

potentially usefully for computational purposes

Liquid Junction Potential

Potential that develops whenever two electrolytes of different ionic compositions come into contact

Due to the unequal distribution of cations and anions across a boundary as a result of the differences in rates at which ions migrate

Results in separation of "+” and “-” charges and creation of potential

Note: Equilibrium condition soon develops

ex.

Both H+ and Cl- move from high to low concentration

H+ smaller and more mobile relative to Cl- moves more quickly

Junction potential can be > 30 mV

-For simple system can calculate if know mobility and concentration of all ions present

-can decrease the junction potential by using salt bridge containing concentrated electrolyte

Currents in Electrochemical Cells

Ohm’s Law:

E=IR

Where

E=potential (V, voltage)

I=current (amps)

R=resistance (ohms)

R depends on concentration and types of ions in solution

Mass Transport Resulting From Current in Bulk Solution

-currents in solution are carried by movement of ions

-again, small ions (H⁺) move faster and carry mroe current than larger ions (Cl)

-species reacting at electrode don’t have to be the only species carrying current

Currents at Electrode Surfaces: Faradic

transfer of electron to/from electrode by redox reactions

governed by Faraday’s Law: amount of current is proportional to amount of species oxidized or reduced

Currents at Electrode Surfaces: Non-Faradic Current

due to processes other than redox reactions at electrodes

ex- charging current

-when first apply potential to electrode, get redistribution of ions near its surface to counter charge on electrode

movement of ions = current

-as system approaches equilibrium- get decrease in ion movement and current

Result of charging electrode is electric double layer by electrode surfaces. Electrode at this point is polarized

Effect of Current on Cell Potential

Potentials listed as E^o and E^o’ in Tables are Thermodynamic values; at equilibrium, no current

In practice, some current is always present

current causes: decrease in measured potential (E) for galvanic cell

increase in potential (EP needed to drive electrolytic cell

Ohmic Potential (IR drop)

Flow of ions (current) through solution (resistance, R) gives potential across cell according to Ohm’s law

Need to subtract from E_cell calculation to get “true” potential of the cell

Polarization Effects

Many electrochemical methods use current vs. potential curves

Polarization effects contribute to the non-linear regions of curve

Note: at high or low cell potential, get less “+” or “-” current than expected.

Due to polarization:

Solution or reaction can not keep up with changes in potential of system

Limits the rates of the overall reaction

Types of Polarization: Slow Mass Tranfer

Concnetration polarization

diffusion- concentration gradient

migration- ions move in potential

convection- mechanical stirring

Types of Polarization: Slow Intermediate Reactions

Reaction Polarization

Types of Polarization: Slow Transfer of Electron Between Electrode and Species

Charge-Transfer Polarization

Overvoltage or overpotential

degree of polarization of an electrode

differences between actual electrode potential (E) and equilibrium potential (E_eq): n=E-E_eq where E < E_eq

Polarization always reduces the electrode potential

N is always negative

Overvoltage is sometimes useful:

high overvoltage associated with the formation of H₂ and O₂ from H₂O

high n means takes much higher E than E^o to occur on many electrodes

can deposit metal without H₂ formation and interfering with electrodeposition process