ElectrochemistryPrepared by

Odyssa Natividad RM. Molo

Electrochemistry

• Branch of chemistry that deals with electricity and its relation to chemical reaction

Electrochemical processes- Redox rxns in which the energy

released by a spontaneous rxn is converted to electricity or in which electricity is used to drive a nonspontaneous chemical rxn

Redox Reaction

• Oxidation-reduction reaction• Chemical reactions in which the

oxidation state of one or more substances changes

Oxidation state/number- a (+) or (-) whole # assigned to an

element in a molecule or ion on the basis of a set of formal rules; to some degree it reflects the (+) or (-) character of that atom

A Spontaneous Redox Rxn

Redox reaction

• Oxidation– Involves loss of electron

• Reduction– Involves gain of electrons

• Oxidizing agent/oxidant– Substance responsible for another substance

to be oxidized; one that undergoes reduction• Reducing agent/reductant

– Gives up electrons/undergoes oxidation; caused another substance to be reduced

How to determine if redox or not?

• Keep tract of the oxidation number of all the elements involved in the reaction.– Activity: Rules on assigning oxidation #s

(p44 LM)– Ex:

• Zn(s) + 2 H+(aq) Zn2+(aq) + H2(g)

• In any redox rxn, both oxidation & reduction must occur. If one substance is oxidized, another must be reduced.

Balancing Redox Reaction

• The law of conservation of mass must be observed.

• + the gains & losses of electrons must be balanced( Electrons are neither created nor destroyed in any chemical rxn)

• Two methods:– Oxidation number method (LM p 44)– Ion-electron method (LM p48)

ELECTROCHEMICAL CELL

Electrochemical cell

• Apparatus which converts chemical energy from spontaneous reaction to produce electricity

• Also called voltaic cell or galvanic cell• Transfer of electron takes place through

an external pathway rather than directly between reactants

• Based on the principle formulated by Alessandro Volta (Italian physicist) & Luigi Galvani (Italian physiologist)

Voltaic cell diagram

Components of Electrochemical Cell

• Electrodes– simply where redox occurs (anode: oxidation; cathode:

reduction)– Types: inert, membrane & metallic

• Charge carriers– Metal wire (usually wire with alligator clips)

• Where electrons pass through as current that be measured by a voltmeter

– Salt bridge• Where electrons travel through its ion to transfer from the

anode to cathode• Usually made up of U-tube with an electrolyte soln that will

not react with other specie in the cell• Maintains the electrical neutrality of the cell by transferring its

cations or anions

Types of Electrode

• Inert electrode– Usually made of C(graphite) or Pt (inert &

nonreactive to components that undergo redox)– Usually used when gases or liquids are formed as by-

products of redox since the charge carrier cannot be connected to these

• Membrane electrode– Specialized; also called ion-selective electrode– Detects electron transfer from or to a specific species

• Metallic electrode– Composed of metal strip & its metallic ion soln– Not only act as sites of rxn but also participate in

redox rxn– Also called active electrodes

Cell emf

• Q: Why does electrons flow spontaneously in the way that they do?

• A: There is a “driving force” that pushes the electrons through an external circuit in a voltaic cell

• Comparison: – Electron Flow to Flow of Water in

Waterfall

Electron Flow to Flow of Water in Waterfall

• Water flows over the waterfalls because of its PE is lower at the bottom of the falls that at the top. Likewise, if there is an electrical connection between the anode & cathode of a voltaic cell, electrons flow from anode to cathode to lower their PE

Electromotive force, emf

• Electromotive = “causing electron motion”

• Potential difference between two electrodes of a voltaic cell which provides the driving force that pushes electron through the external circuit

Cell emf/cell potential

• Standard electrochemical cell potential; aka cell voltage

• Measured in volts• E°cell = E°red(cathode) - E°red(anode)

• Standard condition: (with note sign)– 1M conc for R & P in soln & 1 atm

pressure for those that are gases at 25°C

Cell emf cont…

• Factors that cell emf depends on:– Specific rxns that occur at the cathode &

anode– Conc of reactants & products– Temp (assumed to 25°C unless noted)

• Standard Reduction Potential (SRP), E°red – Standard electrode potential for reduction half

rxn• Standard Hydrogen Electrode, SHE

– Reference half-rxn– 2 H+(aq, 1M) + 2e- H2(g, 1atm) E°red = 0V

SHE

SRP

Points to remember on SRP

• SRPs are intensive properties or has intrinsic properties. Thus changing the stoichiometric coefficient in a half-rxn does not affect the value of SRP– Ex: Zn2+(aq) + 2e- Zn(s) E°red = -0.76

2 Zn2+(aq) + 4e- 2 Zn(s) E°red = -0.76V

• The more positive the value of E°red, the greater the tendency for the reactant to be reduced, undergo reduction, & oxidize another specie

Practice Exercise

• Exercise 18.3 page 479 (refer to SRP page 478)– Compute corresponding E°cell

– Indicate which rxn occurs in the anode & cathode

– Which electrode is consumed?– Which is the OA? RA?– Is the rxn spontaneous or not?– Write the cell diagram

Spontaneity of Redox Rxn

• Determining std change in Gibb’s free energy of the rxn (G) makes it possible to determine the spontaneity of a rxn using this eqn:

G = -nFEcell where n = # of e- transferred during rxn

F = Faraday’s constant, 96500 C/mol

• G is (-) spontaneous; G is (+) nonspontaneous

Spontaneity cont…

• Voltaic cells use redox rxns that proceed spontaneously. Any rxn that can occur in a voltaic cell to produce a (+) emf must be spontaneous.

• General statement:– “A (+) value of E° indicates a

spontaneous process & a (-) value indicates a nonspontaneous one”

– E°= emf under std condition– E= emf under nonstandard condition

Cell Diagram

• Conventional notation that shows the components of an electrochemical cell

• Conventions: (Std rule: ABC)– Anode = left side– Cathode = right side– Boundary

• between diff phase (ex electrode & soln) = single vertical line (│)

• between half-cell compartment (salt bridge) = double vertical line (││)

• between diff species within the same soln = comma (,)

– Species in aqueous soln are placed on either side of the double vertical line

Practice Exercise

• Example 18.1 page 482• Exercise 18.4 page 483

Effect of Concentration on Cell emf

• As a voltaic cell is discharged, the Reactant/s of the rxn are consumed & the Product/s are generated so the conc of these substances change. The emf progressively drops until E=0, at which point, we say the cell is “dead”. At that point the conc of the R & P cease to change; they are at eqlbm.

• The emf generated under nonstandard conditions can be calculated using an eqn first derived by Walther Nernst, a German chemist who established many of the theoretical foundations of electrochemistry.

Nernst Equation

Practice Exercise

• Calculate the emf at 298K generated by the cell below:1) Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

where [Cu2+] = 5.0M & [Zn2+] = 0.050M

2) 2 Al(s) + 3 I2(s) 2Al3+(aq) + 6I-(aq) where [Al3+] = 4.0 x 10-3M & [I-]=0.010M

3) Zn(s) + Cd2+(aq) Zn2+(aq) + Cd(s) where [Cd2+] = 0.0750M & [Zn2+] = 0.950M

Concentration Cells• Cell based solely on the emf generated

because of a difference in a concentration.• more dilute soln = anode; more conc =

cathode.• Applications: pH meter; regulation of

heartbeat in mammals.

Cell emf & Chemical Equilibrium

• From Nernst eqn: As R are converted to P, the value of Q inc, so the value of E dec. The cell emf eventually reaches E = 0, where G = 0; system is at eqlbm; Q = K.

• Rearranging Nernst eqn (@ T = 298K):