WEEK 7

Welcome Back!!

Lecture 19Monday: Simple Cells

Lecture 20Wednesday: StandardReduction Potentials

Lecture 21Friday: )G and E: NernstEquation and uses

19-1

Electron Transfer

19-2

• important in everyday life•corrosion breathing O2 CO2

•(and much more)

We start with electrochemical cells

Galvanic

(batteries)Electrolytic cells

Electricity is supplied to cause a reaction

e.g., electroplating

Silver plating

19-3

Galvanic cells: first example

Copper sulphate solution

Zinc rod

If you dip a zinc bar into a solution of

cooper sulphate something strange occurs.

The zinc electrode erodes -- and copper metal

forms in the solution.

Why?

As we will see soon Zn(s) has a greater tendency

to lose electrons than copper (so it is more easily

oxidized)

Cu2+ has a greater tendency to gain electrons

If we place copper and zinc in different solutions we can “watch” the

electrons transfer -- and make a battery.

19-4

Salt bridge

If the zinc and copper - containing portions of this redox system are

separated then electrons transfer via an external wire

19-5

So:

Zinc-Copper Cell:

1. An observation: If you place Zn metal in a CuSO4 solution, then you see that

Cu2+ ions plate out on the Zn surface; at the same time the Zn metal dissolves to

yield Zn2+ ions, and the solution also becomes warmer. Blue colour fades.

What is happening?

Zn(s) 6 Zn2+(aq) + 2e (oxidation)

2e + Cu2+(aq) 6 Cu(s) (reduction)

Zn(s) + Cu2+(aq) 6 Zn2+(aq) + Cu(s)

and if we arrange things properly we can generate electrical energy.

Because --electrons flow in wires connecting

electrodes

19-6

1. Zn(s) has a greater tendency to lose electrons (i.e.

be oxidized) than Cu(s)

2. Cu2+(aq) has a greater tendency to gain electrons

(i.e. be reduced) than Zn2+(aq).

3. Zn(s) is a stronger reducing agent than Cu(s).

4. Cu2+(aq) is a stronger oxidizing agent than

Zn2+(aq).

19-7

Cell notation:

anode on left (a comes before c) and

cathode on right

Zn(s)|Zn2+(aq)2Cu2+(aq)|Cu(s)

anode cathode

oxidation reduction

Reduction on the Right

As we will see the way the cell is written

implies the direction of the reaction

R O O R

19-8

Things to NOTE:

We have a Zn/Zn2+(aq) half-cell and a Cu/Cu2+(aq) half-cell.

The metal electrodes are connected via a wire and a voltmeter.

A Salt Bridge links the solutions.

Convention: Always write the cell diagram so that the ANODE is at the Left-

Hand Side,

CATHODE at the Right-Hand Side

An Ox Red Cat

Oxidation occurs at the Anode

Reduction occurs at the Cathode} always

19-9

Zn2+(aq) produced in the anode compartment

Cu2+(aq) consumed in the cathode compartment

To maintain electroneutrality, we need to get anions into the anode and cations

into the cathode.

This is where the SALT BRIDGE comes in.

It completes the electrical circuit by allowing ion migration; maintains

electroneutrality.

It prevents DIRECT reaction between Cu2+(aq) and Zn(s).

At the ANODE, Zn(s) 6 Zn2+(aq) + 2 e

The Zn(s) goes into solution as Zn2+

2e go along the wire to the Cu plate

CR-(aq) ions pop out of the SB to maintain neutrality.

At the CATHODE, 2e + Cu2+(aq) 6 Cu(s)

The 2 electrons entice a Cu2+ ion to plate out of the solution

K+(aq) ions pop out of the SB to maintain neutrality.

19-10

Other points:

» If we have 1.0 M Cu2+(aq) and 1.0 M Zn2+(aq), we get a voltage of exactly 1.10

volts, no matter how big or small the cell is. (we will see why soon)

» The "products" (Zn2+(aq) and Cu(s)) are not required to start with; they will be

made as soon as the current flows.

» The concentrations of BOTH cations [Cu2+] and [Zn2+] DO affect the cell

voltage - more on this later.

Cell voltage depends on )G !! (see later)

19-11

Summary:

often called galvanic cell

19-12

Other voltaic (galvanic) cells Zinc - Hydrogen Cell

Cathode: An inert Platinum surface in H+(aq) with H2(g) bubbling past it

What is happening?

Zn(s) 6 Zn2+(aq) + 2e (oxidation)

2e + 2H+(aq) 6 H2(g) (reduction)

Zn(s) + 2H+(aq) 6 Zn2+(aq) + H2(g)

If we have 1.0 M solutions and H2(g) at 1 atm, we get a cell

voltage of exactly 0.76 volts

19-13

Other types of electrodes and half-cells:

Consider the reaction:

Sn2+(aq) + 2Fe3+(aq) 6 Sn4+(aq) + 2Fe2+(aq)

All ions in solution - no metals for electrodes.

What to do:

use an INERT electrode.

Dip a solid conductor into the solution,

e.g. graphite, platinum, gold, and let the

ions in solution transfer electrons to or

from the surface of the inert electrode.

At the anode, Sn2+(aq) 6 Sn4+(aq) + 2e-

At the cathode, 1e- + Fe3+(aq) 6 Fe2+(aq)

If all ion concentrations are exactly 1.0 M

then we get a cell voltage of exactly 0.62

volts.

19-14

Recall:

For the cell on the previous page:

Pt (s) | Sn4+(aq) , Sn2+(aq) || Fe3+(aq) , Fe2+ (aq) | Pt(s)

19-15

shorthand version for coated electrodes:

silver silver chloride: Ag (s) | AgCl (s) | Cl- (aq) ||

half reaction (as a reduction) is AgCl(s) + e Ag(s) + Cl-(aq)

calomel: || Cl-(aq) | Hg2Cl2(s) | Hg(l)

half reaction: Hg2Cl2 + 2e 2Hg(l) + 2Cl-(aq)

as an anode

as a cathode

Standard reduction potentials:

Each “half reaction” (in anode and cathode) is given a voltage

Calomel electrodes sometimes used

Hg2Cl2 + 2e 2Hg(l) + 2Cl- (aq)

E = + .244V

Sometimes used as a reference electrode

19-16