Galvanic Cell

Galvanic cell•A galvanic cell is a device in which chemical energy is changed to electrical energy.•A galvanic cell uses a spontaneous redox reaction to produce an electric current that can be used to do work. The system does work on the surroundings.

•A redox reaction involves the transfer of electrons from the reducing agent to the oxidizing agent.

Oxidation v reduction• Oxidation.• * Involves a loss of

electrons.• * Increase in

oxidation number.• * “To get more

positive.”• * Occurs at the anode of a galvanic cell.

• Reduction.• * Involves a gain of

electrons.• * Decrease in

oxidation number.• * “To get more

negative.”• * Occurs at the cathode of a galvanic cell.

Production of Current• Oxidation Reactions involve a transfer of

electrons.• Electric current is a movement of electrons.• In order to produce a usable current, the

electrons must be forced across a set path (circuit).• In order to accomplish this, an oxidizing agent

and something to oxidize must be separated from a reducing agent with something to reduce.

Pictures

• Oxidation reduction reaction in the same container will have electrons transferring, but we can’t harness them.• Separating the oxidation

reaction from the reduction reaction, but connecting them by a wire would allow only electrons to flow.

Oxidation Reduction

oxidation

reduction

Closer look

• We now have excess electrons being formed in the oxidizing solution and a need for electrons in the reducing solution with a path for them to flow through.• However, if electrons did flow through the wire it would

cause a negative and positive solution to form.

Oxidation Reduction

X → X+ + e- X+ + e- → X

That’s not possible• Or at least it would require a lot of energy.• A negative solution would theoretically be formed by

adding electrons, and a positive one by removing electrons.• The negative solution would then repel the electrons

and stop them from flowing in, and a positive solution would attract the electrons pulling them back where they came from.• Making it so the charged solutions wouldn’t form.• In order for this to work, I would need a way for ions to

flow back and forth but keeping the solutions mostly separated.

The Salt Bridge or the Porous Disk.•These devices allow ion flow to occur (circuit completion) without mixing the solutions. They are typically made of sodium sulfate or potassium nitrate

Closer look

• Now electrons can flow across the wire from the oxidation reaction to the reduction reaction.• As the oxidation reaction becomes positive, it removes

negative ions and adds positive ions to the salt bridge. • The reduction reaction does the reverse.

Oxidation Reduction

X → X+ + e- X+ + e- → X

Salt Bridge

e-e-

e-

Closer look

• Zooming in on the oxidizing side• This would make the salt bridge positive…

Oxidation side e-

Salt Bridge

e-

e-

+ ion+ ion

+ ion+ ion

- ion

- ion

- ion

- ion

Closer look

• (Zooming in on the reducing side)• if the reverse wasn’t happening on this side.

Reducing side

e-

Salt Bridge

e-

e-

+ ion+ ion

+ ion

+ ion- ion

- ion

- ion

- ion

Close up of salt bridge

• The ions keep flowing in the salt bridge to keep everything neutral.• Electrons do also travel across the salt bridge.• This decreases the cell’s effectiveness.

+ ion

+ ion

+ ion- ion

- ion

- ion

+ ion

+ ion

+ ion- ion

- ion- ion

Electrochemical cell

• This is the basic unit of a battery.• It is also called a galvanic cell, most

commercial batteries have several galvanic cells linked together.• Batteries always have two terminals.• The terminal where oxidation occurs is called

the anode.• The terminal where reduction occurs is

called the cathode.

Galvanic Cell

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Voltaic Cell: Cathode Reaction

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Voltaic Cell: Anode Reaction

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Electrochemical Half-Reactions in a Galvanic Cell

Cell Potential (Ecell)•Cell potential (electromotive force, emf) is the driving force in a galvanic cell that pulls electrons from the reducing agent in one compartment to the oxidizing agent in the other.•The volt (V) is the unit of electrical potential.•Electrical charge is measured in coulombs (C).• A volt is 1 joule of work per coulomb of charge transferred: 1 V = 1 J/C.•A voltmeter is a device which measures cell potential.

Standard Reduction Potentials•The measured potential of a voltaic cell is affected by changed in concentration of the reactants as the reaction proceeds and by energy losses due to heating of the cell and external circuit.• In order to compare the output of different cells, the standard cell potential (Eo

cell) is obtained at 298 K, 1 atm for gases, 1 M for solutions, and the pure solid for electrodes.

•The Standard Hydrogen Electrode is considered the reference half-cell electrode, with a potential equal to 0.00 V.• It is obtained when platinum is immersed in 1 M H+(aq), through which H2(g) is bubbled.

The Standard Electrode (Half-Cell) Potential (Ehalf-cell)

•A standard electrode potential always refers to the half-reaction written as a reduction.•Oxidized form + n e- reduced form Eo

half-cell

•If you need the oxidation, you will have to reverse the reaction•Reversing a reaction changes the sign of the potential.

•Eocell = Eo

reduction + Eooxidation

Spontaneous reactions•As the potential increases in value (more positive), the reaction is more likely to occur (spontaneity occurs). •Eo

cell must be positive for the cell to produce electricity.•A substance will have a spontaneous reaction another substance with a lower Eo

cell.•Although some half-reactions must be manipulated with coefficients, NEVER MULTIPLY THE Eo

cell BY THE COEFFICIENT!!!

Galvanic Cell Problems•Consider a galvanic cell based on the following reactions. For each give the balanced half cell reactions and calculate Eo

cell.

•Cu2+(aq) + Zn(s) Zn2+

(aq) + Cu(s)

•Fe3+(aq) + Cu(s) Fe2+

(aq) + Cu2+(aq)

•Al (s) + Cd2+ (aq) Cd(s) + Al3+(aq)

Calculating an Unknown Eo

half-cell from Eocell.

•A voltaic cell based on the reaction between aqueous Br2 and vanadium (III) ions has •Eo

cell = 1.39 V:

•Br2(aq) +2V3+(aq) + 2H2O(l) 2VO2+(aq) + 4H+(aq) +2Br-

(aq)• •What is the standard electrode potential for the reduction of VO2+ to V3+?

Line Notations.•The components of the anode compartment are written to the left of the cathode compartment.•Double vertical lines separate the half-cells and represents the wire and salt bridge.• Within each half-cell, a single vertical line represents a phase boundary. A comma separates half-cell components in the same phase.•Half-cell components appear in the same order as in the half-reaction, while electrodes appear at the extreme left and right of the notation.•Mg(s) Mg2+

(aq) Al3+(aq) Al(s)

• Fe(s) Fe2+(aq) H+, MnO4

-(aq), Mn2+

(aq) Pt(s)

Describing a Galvanic Cell.•A cell will always run spontaneously in the direction that produces a positive cell potential.•A complete description of a galvanic cell always includes four items:

•1. The cell potential (always positive for a galvanic cell) and the balanced cell reaction.•2. The direction of electron flow, obtained by inspecting the half-reactions and using the direction that gives a positive Eo

cell.

•3. Designation of the anode and cathode.•4. The nature of each electrode and the ions present in each compartment. A chemically inert conductor is required if none of the substances participating in the half-reaction is a conducting solid.

Diagramming Voltaic Cells

•In one compartment of a voltaic cell, a graphite rod dips into an acidic solution of K2Cr2O7 and Cr(NO3)3; in the other, a tin bar dips into a Sn(NO3)2 solution. A KNO3 salt bridge joins the half-cells. The tin electrode is negative relative to the graphite. •Diagram the cell, show balanced equations, and write the cell notation.

Description of a Galvanic Cell.

•Describe completely the galvanic cell based on the following half-reactions under standard conditions:• •Ag+ + e- Ag Eo

cell = 0.80 V (1)•Fe3+ + e- Fe2+ Eo

cell = 0.77 V (2)• •In addition, draw the cell and write the line notation.