ElectrochemistryChapter 17

Electrochemistry

• The branch of chemistry that links chemical reactions to the production or consumption of electrical energy.

• In chemistry, electrical energy is stored in electrons. – In other words, electrochemistry is based upon

the principles oxidation-reduction reactions.

Redox Revisited

• Earlier this year, we examined Redox reactions and how to write acid-base half reactions

• To review principles of redox reactions:– A redox reaction is the sum of two half reactions, the

reduction and oxidation reactions. – Reduction and oxidation reactions happen

simultaneously, so the number of electrons gained and lost must match exactly

– Oxidation=loss of electrons– Reduction=gain of electrons.

Redox revisited

• Lets look at a common reaction of Zinc metal immersed in a copper sulfate solution.

The entire reaction is represented as a single replacement reaction where the blue CuSO4 solution becomes clear as the Zinc replaces the copper:

Zn + CuSO4 → Cu + ZnSO4

Redox Revisited

Zn + CuSO4 → Cu + ZnSO4

The Redox half reactions are then represented as:

Zn(s) → Zn+2(aq) + 2e-

Cu+2(aq) + 2e- → Cu(s)

Zn(s) + Cu+2(aq) + 2e- → Cu(s) + Zn+2

(aq) + 2e-

So:Zn(s) + Cu+2

(aq) → Cu(s) + Zn+2(aq)

Electrochemical (Galvanic) cells• An apparatus that

converts chemical energy into electrical work, or vice versa– Contains two

compartments– Bridge that allows flow

of energy (electrons)• “salt bridge” • Usually a piece of tubing

filled with an electrolyte

Electrochemical (Galvanic) cells

• Anode-compartment in which the oxidation half reaction takes place

• Cathode-compartment which the reduction half reaction takes place. – We represent the reactions that take place using

cell diagrams• Cell diagrams are symbols that show how the

components of an electrochemical cell are connected.

Electrochemical (Galvanic) cellsSalt bridge replaced with porous disk to allow ion flow and minimum mixing of solutions. Oxidizing agent “pulls” electrons through wire from reducing agent. Let there be light.

Homework

• Pg 879-880 # 13-19.

Cell Potential

• The “pull” on the electrons is called the cell potential (E °cell), or “electromotive force” (emf), is measured in volts.– Volt: 1 joule of work per coulomb of charge

transferred. 1 J/C– coulomb: defined as the charge transported by a

constant current of one ampere in one second:

Standard Reduction Potentials

• E °--standard reduction potentials in Volts. – E °cell= E °cathode- E °anode

• Pay close attention to sign of E for certain reactions• If the reaction is in reverse, change the sign for the

reduction potential. • See Table 17.1 on page 843 in your book. • All reduction potentials are given with all solutes at

1M and all gases at 1 atm pressure. • E °cell is always positive

Standard Reduction Potentials

Zn + CuSO4 → Cu + ZnSO4

The Redox half reactions are then represented as:Zn(s) → Zn+2

(aq) + 2e- -E °anode=.76

Cu+2(aq) + 2e- → Cu(s) E °cathode=.34

Zn(s) + Cu+2(aq) → Cu(s) + Zn+2

(aq) E °cell = 1.1V• This cell produces 1.1 volts of electrical energy

(work).

Standard Reduction Potentials

• Another example, consider the galvanic cell based on the reaction:

Al+3(aq) + Mg(s) → Al(s) + Mg2+

(aq)

Give the redox half reactions, make sure to balance the reactions (equal # of electrons), and calculate E °cell (E ° for half reactions on table 17.1)

Standard Reduction Potentials

Al+3(aq) + Mg(s) → Al(s) + Mg2+

(aq)

The Redox half reactions are then represented as:3(Mg(s) → Mg+2

(aq) + 2e)- -E °anode =2.37V

2(Al+3(aq) + 3e- → Al(s)) E °cathode=-1.66V

2Al+3(aq) + 3Mg(s) → 2Al(s) + 3Mg+2

(aq) E °cell = .71V• This cell produces .71 volts of electrical energy

(work).

Standard Reduction Potentials

• Sometimes there are multiple possibilities for redox potentials, in this case, pay close attention to what the equation they give states the cell is based on.

• Example: a galvanic cell is based on the reaction:– MnO4

-(aq) + H+

(aq) + ClO3-(aq) → ClO4

-(aq) + Mn+2

(aq) + H2O(l)

• There are multiple oxidation reactions for MnO4-(aq) , so you must

consult table 17.1 and match the reactants and products.

Representing Cells with Line Notations

• Consider key components of this galvanic cell:

Zinc solid electrode

Zinc ions in solution

Spectator ions

Spectator ions

Copper ions in solution

Copper solid electrode

Salt bridge

Steps to Representing Cells with Line Notations

• Rule #1…list everything• Separate the cathode/anode with the double line

notation (II) that represents the salt bridge• Separate substances in different states of matter in

the same compartment with a single line (I), separate substances in the same state in the same compartment with a comma.

• Dispense spectator ions (usually water)• When there is no solid electrode listed, assume

there is Platinum (Pt) in both.

Representing Cells with Line Notations

• Consider key components of this galvanic cell:

Zinc solid electrode

Zinc ions in solution

Spectator ions

Spectator ions

Copper ions in solution

Copper solid electrode

Salt bridge

homework

• Pg 880, #’s 25-35 odd

Determining Spontaneity in Galvanic cells

• 1. Any reduction reaction is spontaneous when paired with the reverse of any reaction below it on Table 17.1.

• 2.If the Cell potential calculated is negative, the reaction is not spontaneous (yes, I know I told you that the cell potential is always positive…in a cell that works).

Cell Potential, Electrical Work, And Free Energy (∆G)

• Potential Difference (V)= work (J)/charge©– When a cell produces a current (V), the cell

potential is positive, and the current can be used to do work.

– Since the work does not stay in the system (the cell), the sign for work is negative. So:

q=charge in coulombs transferred from anode to cathode

q

w E

qEw

qEw

Cell Potential, Electrical Work, And Free Energy (∆G)

qEw Work measured in Joules

Coulombs based on Faraday Constant

Cell potential difference in V or J/C

The Faraday Constant: the charge on one mole of electrons is 96,485 Coulombs of charge . When 1.33 moles is transferred:

nFq = 1.33 mol e- x 96,485 C/mol e-

This equation calculates amount of work done. HOWEVER, since some work and energy is always lost to the surroundings as heat, there is a way to calculate Maximum work

Cell Potential, Electrical Work, And Free Energy (∆G)

• All galvanic cells have a maximum potential that they never reach because of energy lost as heat. To calculate maximum work, use the maximum potential in your equation. And then the Maximum work equals ∆G.

maxmaxmax nFEqEGw

Dependence of a Cell on Concentration

• Simply put…if the product concentration is raised above 1.0M, E °cell will be less than what is listed 17.1

• If the reactant concentration is above 1.0M, E °cell will be greater than listed in table 17.1

Dependence of a Cell on Concentration

• The dependence of cell potential on concentration relies directly on the dependence of free energy on concentration. Remember that

• And since• And • Then:

)ln(QRTGG cellcell nFEqEG max

nFEG

)ln(QnF

RTEE “Nerst Equation”

Nerst Equation

At 25°C:

E=E°-

n= moles of electrons

homework

• Pg 881, #’s 39-59 odd.