Chemistry The Final Conflict… Chemistry The Final Conflict… Chapter 18: Electrochemistry and Electric Vehicles.

ChemistryChemistryThe Final Conflict…The Final Conflict…

Chapter 18: Electrochemistryand Electric Vehicles

Electrochemistry: The area of chemistry that examines the transformations between chemical and electrical energy.

Oxidation-reduction (redox) reactions involve an exchange of electrons between reacting species

(see chapter 4, section 4.8).

You may need to review the following terminology: oxidation, reduction, oxidizing agent, reducing agent and half-reaction .

In the following reaction, what is being oxidized, reduced? What is the oxidizing agent, the reducing agent?

(aq)H4(aq)SO4(aq)Fe2O(l)2H(g)O7FeS2 24

2222 pyrite)(s,

What are the half-reactions?

What is oxidized? reduced?

In this example, there is an exchange of electrons between the oxidized and reduced species that is thermodynamically favored (exergonic). The goal of using an electrochemical cell is to extract usable work from this electron transfer.

Fig. 18.1: A piece of zinc is immersed in a copper(II) sulfate A piece of zinc is immersed in a copper(II) sulfate solution. The Cu(II) is spontaneously converted to elemental Cu solution. The Cu(II) is spontaneously converted to elemental Cu and the solid Zn dissolves as Znand the solid Zn dissolves as Zn2+2+ ions. ions.

Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq)

Problem 12. Identify which elements (if any) undergo oxidation or reduction.

4 ClO3-(aq) → Cl- + 3 ClO4

-

ProblemProblem. Sometimes the cell reaction of NiCd . Sometimes the cell reaction of NiCd batteries is written with Cd metal as the anode and batteries is written with Cd metal as the anode and solid NiOsolid NiO22 as the cathode. Assuming the products of as the cathode. Assuming the products of the electrode reactions are solid hydroxides of Cd(II) the electrode reactions are solid hydroxides of Cd(II) and Ni(II), respectively, write a balanced chemical and Ni(II), respectively, write a balanced chemical equation for the cell reaction.equation for the cell reaction.

Fig. 18.2: An electrochemical electrochemical cellcell is a reaction system in which oxidation and reduction reactions occur in separate compartments (or cells) either consume or produce electrical energy. The cells are separated by a salt bridge or semi-permeable membrane that allows ions to migrate from one cell to the other. Electrons move from anode (oxd) to cathode (red).

Voltaic or Galvanic cell – chemical energy is used to produce electrical energy (G<0) (i.e. a battery).Electrolytic cell – an external source of electrical energy is used to do work on a chemical system (i.e. charging a car battery with the alternator after the car has started).

Alessandro Volta is credited with building the first battery…which was built by alternating layer of zinc and silver with paper-soaked brine between the metals. Volta coined the term “galvanism” to distinguishthe animal electricity” observed by his adversary,

Luigi Galvani

Cathode = reductionAnode = oxidation

Write the balanced redox reaction for this electrochemical cell.Which direction will the nitrate ions flow in the salt bridge?(see sample exercise 18.2)

Ni(NO3)2(aq) AgNO3(aq)

Cathode = reductionAnode = oxidation

Ni(NO3)2(aq)AgNO3(aq)

Cell potential (Ecell) or Electromotive Force (emf) is the voltage between the electrodes of a voltaic cell.

A Faraday (F) is the electrical charge on one mole of electrons or 9.65E4 Coulombs (C)/ mol e.

(The charge on one electron is -1.602E-19 C).

The quantity of charge flowing through an electrical circuit is: C = nF

Electrical work is done when a charge moves through an electrical potential,

welec = CEcell = –nFEcell

Since free energy is that available to do work. this equation becomes:

G = –nFEcell

[ 1 C۰V = 1 Joule (J)]

Fuel cellsFuel cells use a controlled electron transfer between use a controlled electron transfer between hydrogen and oxygen to produce electrical energy.hydrogen and oxygen to produce electrical energy.

HH22(g) + ½ O(g) + ½ O22(g) (g) H H22O(O(ll) ) GG°° = = 237 kJ237 kJ

What is the electromotive potential that can be What is the electromotive potential that can be produced by this cell under standard conditions (produced by this cell under standard conditions (EEoo)?)?

A standard potential (Eo) is the electromotive force of a half-reaction written as a reduction reaction in which all reactants and products are in their standard states (see Table A5.4 or en.wikipedia.org/wiki/Table_of_standard_electrode_potentials).

The standard cell potential (Eºcell) is the potential of a cell when all reactants and products are in their standard states, i.e. the pressure of all gases are 1 bar, and the concentration of dissolved species are 1 molar.

Eºcell = Eºcathode – Eºanode or Eºcell = Eºreduction – Eºoxidation

Fig. 17.7: The Standard Hydrogen Electrode (SHE) has a solution of 1 M HCl and is bathed in a stream of H2 gas at 1 bar (~1 atm) pressure.

This half-reaction has a defined potential of 0.00 V as either a reduction or oxidation reaction and is used to reference the potentials of other half-reactions.

Also see Table A6.1 in your text.

This source: http://www.jesuitnola.org/upload/clark/refs/red_pot.htm

Fig. 17.8: The SHE can be used in a cell with either oxidation or reduction reactions to determine the standard potential of that half-reaction. Use the above figures and the following cell potential equation to find the standard potentials for the Cu and Zn reactions.

Eºcell = Eºcathode – Eºanode

SHE: 2H+(aq) + 2e- → H2(g) E0= 0.0000

Zn2+(aq) + 2e- → Zn(s) E0= -0.7618

Cu2+(aq) + 2e- → Cu(s) E0= 0.52

One of the first batteries built was that of Allesandro Volta (1745-1827). Calculate the standard cell potential (Eºcell ) for this battery that had a Ag/Ag+ cell connected to a Zn/Zn2+ cell by a salt bridge (a blotter soaked with a NaNO3 solution).

Eºcell = Eºcathode – Eºanode

From Table A6.1:

Ag+ + 1 e Ag(s) E = 0.7996 V

Zn2+ + 2 e Zn(s) E = -0.7618 V

Compare two reduction half-reactions; the more positive of the two will occur as the cathode reaction (for a voltaic cell). The other reaction will be the anode.

See sample and practice exercise 18.4.

Problem. A Voltaic cell based on the following pair of half-reactions is constructed. Write a balanced equation for the overall cell reaction, and identify which half-reaction takes place at the anode and cathode.

AgBr(s) + e- → Ag(s) + Br-(aq) E° = 0.095V

MnO2(s) + 4 H+ + 2e- → Mn2+(aq) + 2 H2O(l) E° = 1.23V

The Nernst Equation (Walther Nernst, 1864-1941) can be used to calculate the cell potential for non-standard conditions usually when the conc. of dissolved species ≠ 1 .

Where R is the gas constant, T is the temperature in Kelvin, n is the number of moles of electrons transferred between the oxidation and reduction reactions, F is the Faraday constant and Q is the reaction quotient for the system.

What is the value of Q when all the reaction species are at standard conditions?

lnQnF

RTcell

ocell EE

Qlogn

V 0.0592cell

ocell EE Nernst Eq. at 25ºC.

Fig. 18.5: The standard lead-acid storage battery found in your car uses the following reaction.

Pb(s) + PbO2(s) + 2 H2SO4(aq) 2 PbSO4(s) + 2 H2O(l)

With an electrolyte of 4.5 M H2SO4 each cell produces ~2.0 V.

Fig. 17.10: The cell potential will decrease as reactants are converted to products. This is an example of the cell potential in a standard lead-acid car battery as the sulfuric acid is used up while the battery discharges.


Problem 18.53. Permanganate ion can oxidize sulfite to sulfate in basic solution as follows:

2MnO4–(aq) + 3SO3

2–(aq) + H2O( ) → 2MnO2(s) + 3SO4

2–(aq) + 2OH–(aq)

Determine the Standard Potential for the reaction at 298 K and when the concentrations of the reactants and products are as follows:

[MnO4–] = 0.150 M, [SO3

2–] = 0.256 M, [SO42–] = 0.178 M,

and [OH–] = 0.0100 M.

Will the value of Ern increase or decrease as the reaction proceeds?

Also see Table A6.1 in your text.

This source: http://www.jesuitnola.org/upload/clark/refs/red_pot.htm

The Nernst Equation can be used to predict the cell potentials under non-standard conditions. In this example (A) the two identical cells with differing concentrations of dissolved silver will produce a cell potential (see p. 867).

At equilibrium the Nernst Equation:

changes since at equilibrium Ecell = 0 and Q = K, so the equation can be rearranged as follows:

Qlogn

0.0592cell

ocell EE

0.0592

nKlog

Klogn

0.05920

cello

cello

E

E

Using standard reduction potentials from the table in Appendix , calculate the value of the equilibrium constant for the following reaction at 298 K.

5 Fe2+(aq) + MnO4–(aq) + 8 H+(aq)

5 Fe3+(aq) + Mn2+(aq) + 4 H2O(l)

Fig. 17.1: the predicted effect of temperature on the cell potential of a lead-acid battery using the Nernst Eqn.

Other T-dependent factors have a greater influence on battery performance.

Differences between an electrolytic cell and a voltaic cell.

From a practical stand point one of the important characteristics of a battery is its ability to do work.

wcell = C·Ecell

Battery capacity can be expressed in coulombs x volts (= joules). Other common definitions of battery capacity are useful, for example:

1 ampere (amp) = 1 coulomb(C)/sec or 1 C = 1 amp·sec

The Faraday constant can be written: F = 9.65E4 A۰s/mole

1 watt = 1 volt·amp = 1 J/s

The watt (James Watt, 1736-1819) is a widely used unit electrical power.

Consequently, a cell producing 1 volt of potential and 1 amp of current will produce 1 watt of power.

Problem 78. In the electrolysis of water, how long will it take to produce 100.L of H2(g) at STP using an electrolysis cell through which flows a current of 50.mA?

Fig. 17.13: The discharge-charging cycle of a lead-acid battery.

Problems. If it takes 6.0 seconds of discharge for a car battery to start the engine and the starter drew a current of 230 A, what mass of Pb will be oxidized to PbSO4 in this time?

How long will it take to re-charge the battery with an alternator current of 30.0 A?


See sample & practice exercise 18.7 and assume 100% efficiency in these reactions.

a. What will be the oxidation reaction that occurs in the above voltaic cell?b. What will be the reduction reaction that occurs in the above cell?c. What will be the standard cell potential for the above cell at 25 C?d. On the diagram, indicate the direction of electron flow through the external wire.e. On the diagram, indicate the anode and cathode compartments.f. On the diagram, indicate the direction of migration of sulfate through the

salt bridge?g. Calculate the cell potential when the concentration of Cd2+ is 0.050 M and the

concentration of Mg2+ is 0.0025 M (at 298 K).

Mg Cd

MgSO4(1 M, aq) CdSO4(1 M, aq)Salt bridge

Fig. 17.15: Thin coatings of metals can be applied using electrolytic reactions.

How many grams of silver can be plated from a silver nitrate solution using a 20. mA current for 15 minutes?

Fuel cells are a voltaic device in which there is a flow of reactants to the anode and cathode.

The hydrogen fuel cell uses streams of H2 and O2 gases that diffuse to the following anode and cathode reactions, respectively.

H2(g) 2 H+(aq) + 2 e– Eº = 0.000 V

O2(g) + 4 H+(aq) + 4 e– 2 H2O(l) Eº = 1.229 V

The overall reaction is:

2 H2(g) + O2(g) 2 H2O(l) Eºcell = 1.229 V

17_04.jpg

Fig. 17.4: The “disposable” zinc-acid (dry cell) battery.

Zn + 2 NH4Cl + 2 MnO2 2 Zn(NH3)2Cl2 + Mn2O3 + H2O, Ecell=1.5 V

17_05.jpg

Fig. 17.5: The alkaline cell has the same potential as the classic dry cell but the zinc anode is oxidized to Zn(OH)2.

17_06.jpg

Fig. 17.6: The Ni-Cd or nickel-cadmium battery can be recharged because the products adhere to the respective electrodes and the reaction can be readily reversed.

Cd(s) + NiO(OH)(s) + 2 H2O(l) Cd(OH)2(s) + 2 Ni(OH)2(s)

17_09.jpg

Fig. 17.9: In the nickel-metal hydride battery (NiMH) , hydrogen atoms are stored in the interstitial spaces of the metal matrix. They can migrate from these spaces to participate in the anode half-reaction. NiO(OH) is simultaneously reduced at the cathode.

In the lithium-ion battery, Li+ is stored in pure graphite anodes (see section 10.8; each six carbon ring of graphite stores 1 lithium ion). The cathodes are made of porous transition metal oxides ( i.e. MnO2) which can form highly stable complexes with Li+ ions. In a fully charge battery there is a concentration gradient between the anode and the cathode. The lithium ions migrate down the gradient and at the same time electrons flow in the external circuit to balance the charge.

The electrodes in the Li-ion cell react with oxygen and water so they must be either entirely solid-state or use non-aqueous solvents.

See pages 862-863 for description of the Li-ion cell.

Some common types of batteries and uses.

This tutorial explores the concept of cell potential (Ecell)

as a measure of how much electrical energy is stored in an electrochemical cell. Includes practice exercises.

»PC version

»Mac version

Cell Potential Tutorial

You should view these tutorials on your own. They are found on your CD or the publishers website.

17_16.jpg

Fig. 17.16:

17_18.jpg

Fig. 17.18: Photochemical cells can be used to directly convert sunlight energy to electricity. This cells uses sunlight to catalyze the reduction of water to H2 at the cathode and oxidation to form O2 at the anode.

17_19.jpg

Fig. 17.19: “Biological” batteries can use bacteria to catalyze redox reactions and harness the electron flow to do useful work.

Learn how fuel cells use a redox reaction between hydrogen and oxygen to produce electrical energy. Includes practice exercises.

»PC version

»Mac version

Fuel Cell Tutorial


W. W. Norton & CompanyIndependent and Employee-Owned

This concludes the Norton Media Libraryslide set for chapter 17

ChemistryChemistryThe Science in Context

byThomas Gilbert,Rein V. Kirss, &Geoffrey Davies

This tutorial illustrates the reactions that occur at the electrodes of a typical zinc copper battery, and explores how the energy released by a voltatic cell is used to do work on the surroundings. Includes practice exercises.

»PC version

»Mac version

Zinc-Copper Cell Tutorial


Learn how the potential of an electrochemical cell can be used to determine the free energy available to do work, and explore the relationships between free energy, cell potential and equilibrium constant. Includes practice exercises.

»PC version

»Mac version

Free Energy Tutorial


This tutorial explores the oxidation-reduction reactions that power rechargeable Ni-Cd batteries and describes the changes in reaction quotient as a battery loses its charge or is recharged. Includes practice exercises.

»PC version

»Mac version

Ni-Cd Battery Tutorial


Chemistry The Final Conflict… Chemistry The Final Conflict… Chapter 18: Electrochemistry and Electric Vehicles.

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