Section 14.2 Voltaic Cells p. 622 - 638

Voltaic cells

• Voltaic cells convert chemical energy to electrical energy.

• In redox reactions, oxidizing agents gain electrons and reducing agents lose electrons, but if electrons can be detoured through a wire and both agents can be separated in their own half cells, yet allowing for movement of positive and negative ions in solution through some pathway, an electrochemical cell can be made.

Classic Example: Zinc-Copper Battery (Voltaic Cell )

Reduction ½ Reaction: Cu 2+(aq) + 2 e – Cu(s)

• In the half-cell containing the strip of copper immersed in its own solution, copper ions gain electrons to form solid copper on the strip already present.

• The electrons travel through the wire from the other half-cell.

• In the solution, copper ions (+ive charge) leave the solution to become solid copper.

• The region immediately surrounding the copper strip lacks positive ions but negative ions, which do not react, remain.

• The solution immediately surrounding the strip is therefore negatively charged.

Oxidation ½ Reaction: Zn(s) Zn 2+(aq) + 2 e –

• In the half-cell containing the strip of zinc immersed in

its own solution, solid zinc atoms decompose into aqueous zinc ions, losing electrons in the process.

• These electrons travel through the wire to the other half-cell.

• In the solution, zinc ions (+ive charge) flood the solution immediately around the solid zinc strip making that region very positively charged.

• All positively charged ions (cations) from both half-cells are attracted to the negative region created in the copper half-cell.

• All negatively charged ions (anions) from both half-cells are attracted to the positive region created in the zinc half-cell.

• To maintain electrical neutrality, these ions in

solution migrate through a salt bridge or a porous cup.

Another Variation of an Voltaic Cell

Example #1.

• All batteries or electrochemical cells have a positive post and a negative post.

• These “posts” are the electrodes. • A positive electrode is called the cathode and a negative

electrode is called the anode. • Reduction ½ Reaction: Cu 2+

(aq) + 2 e – Cu(s)

(cathode) (G.E.R.P.C.)

• Oxidation ½ Reaction: Zn(s) Zn 2+(aq) + 2 e –

(anode) (L.E.O.N.A.)

Voltaic Cells with Inert Electrodes• Cells containing metals and metal ions, the electrodes are

usually metals, and the half-reactions occur on the surface of the metals.

• If an oxidizing or reducing agent is not a metal ion, then electrodes other than the metal ones have to be used.

• Acidic dichromate solution, for example, will spontaneously react with copper, and therefore copper electrodes would not be a good choice.

• In this case inert electrodes (unreactive electrodes), are usually used.

• Inert electrodes provide a location to connect a wire and a surface on which a half-reaction can occur.

• Common inert electrodes used are carbon (graphite) rod and platinum foil.

Read pgs. 622 – 626

pg. 626 Practice #’s 1 – 7

Standard Reduction Potentials - "E" Values

• Comparison of a gravitational system to an electrochemical cell:

Drop of gravitational energy determines water "pressure".

Water e –

V

Drop of electron's potential energy measured by voltmeter.

SRA

SOA

• "Electrical potential drop", commonly called "voltage", measures the tendency for electrons to change energy.

• The tendency for electrons to go from SRA to SOA in an electrochemical cell is measured by a voltmeter in the circuit.

• To compare the tendencies for gaining electrons by oxidizing agents, or losing electrons by reducing agents, each OA's or RA's half-cell must be measured against a standard half-cell.

· The cell chosen was the H2(g) , H +(aq) / Pt(s) half-cell,

arbitrarily assigned a value of 0.00 V.

Porous Barrier

V

X(s)

1 .0 X +(aq) 1 .0 mol/L

H +(aq)

Pt(s)

electrode

H2(g) at STP

H2(g)

Standard ½ Cell

from 1.0 HCl(aq)

Using the Hydrogen Half Cell As a Standard

· Each OA/RA combination is set up at "X(s) / X n+(aq)"

and the voltage measured against the standard H2(g), H +

(aq) / Pt(s) cell.

a) Cu(s) / Cu2+(aq) // H2(g) , H+

(aq) / Pt(s)

· Reads +0.34 V. Since the voltage of the standard cell is 0.00 V, then for:

· Cu2+(aq) + 2 e – Cu(s) E = + 0.34 V

· The "+" voltage means Cu2+(aq) has a greater

tendency to gain electrons than H+(aq).

· This means that Cu2+(aq) is a stronger OA than H+

(aq).

b) Zn(s) / Zn2+(aq) // H2(g) , H+

(aq) / Pt(s)

· Reads – 0.76 V, so:· Zn2+

(aq) + 2 e – Zn(s) E = – 0.76 V· The "–" voltage means Zn(s) has a greater tendency

to lose electrons than H2(g).· Zn(s) is a stronger RA than H2(g).

• The E values for oxidizing agents are listed in the table of reduction half reactions.

• Note: If the RA for a half reaction is not a metal an inert electrode such as C(s) or Pt(s) is used.

C(s) / Cr2O7 2–

(aq) , H +(aq) // H2(g) , H +

(aq) / Pt(s)

Measuring Standard Reduction Potentials• Example 1:Find the net reaction and Ecell.

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

(aq) / Cu(s) SRA OA SOA RA

Measuring Standard Reduction Potentials• Example 1:Find the net reaction and Ecell.

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

(aq) / Cu(s) SRA OA SOA RA

Cathode : Cu2+(aq) + 2 e – Cu(s)

Anode: Zn(s) Zn2+(aq) + 2 e –

Net: Cu2+(aq) + Zn(s) Zn2+

(aq) + Cu(s)

E°cell = E°r - E°r cathode anode

= +0.34 - (-0.76)= + 1.10 V

• Example 2:Find the net reaction and Enet.

Mg(s) / Mg2+(aq) // Cr2O7

2–(aq) , H+

(aq) / C(s) SRA OA SOA OA

• Example 2:Find the net reaction and Enet.

Mg(s) / Mg2+(aq) // Cr2O7

2–(aq) , H+

(aq) / C(s) SRA OA SOA OA

Cathode: Cr2O7 2–

(aq) + 14 H +(aq) + 6 e – 2 Cr 3+

(aq) + 7 H2O(l) Anode: 3 (Mg(s) Mg 2+

(aq) + 2 e – )Net:Cr2O7

2–(aq) + 14 H +

(aq)+ 3 Mg(s) 2 Cr 3+(aq) + 7 H2O(l) + 3 Mg2+

(aq)

E°cell = E°r - E°r cathode anode

= +1.23 - (-2.37)= +3.60 V

• Even though we balance the electrons by an appropriate factor, we do not change the reduction potentials with those factors.

• If the electrical potential of a cell is positive, then the reaction is spontaneous – a requirement for all voltaic cells.

Summary

• L.E.O.N.A. and G.E.R.P.C.• Electrons flow from the negative anode electrode to the

positive cathode electrode through a conducting wire.• Half-cells contain species involved in separate half

reactions.• Electrical neutrality is maintained via a “salt bridge” or

porous cup, allowing anions (–ive ions) to migrate to the anode electrode and cations (+ive ions) to migrate to the cathode electrode.

• E°cell = E°r - E°r cathode anode

Read pgs. 627 – 632

pg. 621 Practice #’s 10 – 16

CorrosionCorrosion is an electrochemical process in which a metal is oxidized by substances in the environment, returning the metal to an ore-like state.

In general, any metal appearing below the various oxygen half-reactions in a redox table will be oxidized in our environment.

Rusting of Iron

O2(g) + 2 H2O(l) + 4 e– → 4 OH – (aq)

2 [ Fe(s) → Fe2+(aq) + 2 e– ]

2 Fe(s) + O2(g) + 2 H2O(l) → Fe(OH)2(s)

The iron(II) hydroxide precipitate is further oxidized to eventually form rust: Fe2O3·xH2O(s).

cathode:

anode:

net:

The rusting of iron requires the presence of oxygen and water and is accelerated by the presence of acidic solutions, electrolytes, mechanical stresses, and contact with less active metals.

Corrosion Prevention

Protective Coatings Paint and other metals, such as tin and zinc are most commonly used.

Tin adheres well to the surface of iron and the outer surface of the tin coating has a thin, strongly adhering layer of tin oxide that protects the tin.

Iron has been galvanized when it has been coated with a layer of zinc.

Zinc is a stronger reducing agent than iron, thus more easily oxidized.

Fe2+(aq) + 2 e–→ Fe(s) E°r = –0.45 V

Zn2+(aq) + 2 e–→ Zn(s) E°r = –0.76 V

Zinc is oxidized instead of the iron it is protecting.

Cathodic Protection

Recall that oxidation (loss of electrons) takes place at the anode of a voltaic cell.

Cathodic protection is when the iron is forced to become the cathode by supplying the iron with electrons

Impressed Current Cathodic Protection

An impressed current is an electric current forced to flow toward an iron object by an external potential difference which is provided by a constant power supply.

This is commonly used for pipelines.

Sacrificial Anode

A sacrificial anode is a metal that is more easily oxidized than iron and connected to the iron object to be protected.

The water pipe is turned into the cathode and magnesium is used as the sacrificial anode.

Read pgs. 634 – 636

pg. 637 Practice #’s 17 – 24

pgs. 637 – 638 Section 14.2 Questions #’s 1 – 12