ELECTROCHEMISTRYPgs. 652 - 654

How does our lab from Friday link to corrosion? Corrosion is the process of

returning metals to their natural state It’s a REDOX reaction!!

Fe (s) + O2 (g) Fe2O3 (s)

LOTS of metals corrode, but not all of them corrode to the same extent: Ex Aluminum!!

Aluminum will be oxidized by the air Al (s) + O2 (g) Al2O3 (s)

A thin layer of Al2O3 will cover the metal and protect it from further corrosion

How can we protect these metals from corrosion? The Mg will react instead of the

iron…but why???

So what does this have to do with the lab?? It all comes down to HOW ACTIVE a

metal is!! What was the most active metal you saw

in the lab? What was the least active metal? What does it all mean???

Activity SeriesLi Rb

KCsBaSrCaNaMgAlZnCrFeNiSnPbCuHgAgAu

Most reactive

Least reactive

How does electronegativity

relate?

How does electronegativity relate?

Electronegativity = how much you “love” electrons

More electronegative = more you “love” electrons = more likely to ________

Activity SeriesLet’s look at Mg and Cu:

Mg + CuCl2 Cu + MgCl2

Cu + MgCl2 Mg + Cu Cl2

Li Rb

KCsBaSrCaNaMgAlZnCrFeNiSnPbCuHgAgAu

Most reactive

Least reactive

For a reaction to happen the solid metal must be above the aqueous metal in the activity series

Electrochemistry

The study of the interchange of chemical and electrical energy

Two types of processes in electrochemistry: The production of an electric current

from a chemical (redox) reaction The use of an electric current to produce

a chemical change

But first, a demo review from yesterday… When iron metal is dipped into an

aqueous solution of blue copper sulfate, the iron becomes copper plated Why?

The iron loses e- to the copper What type of reaction is this????

Fe(s) + Cu2+(aq) Fe2+

(aq) + Cu (s)

Copper Plating – An Example

CuSO4(aq)

Fe

FeSO4(aq)

Zn

Cu

Fe

Since the copper is plating the iron, the solution will get lighter as more copper is used.

Does the reverse happen?

Fe(s) + Cu2+(aq) Fe3+

(aq) + Cu (s)

Can we go backwards?….

Fe(s) + Cu2+(aq) Fe3+

(aq) + Cu (s)

•Some metals are better reducing agents than others

• (AKA: some metals lose e- easier than others.)

• The reaction is only spontaneous one way… the reverse reaction requires an outside source of energy to work.

What does all of this mean? To capture the electrical energy, the

two half-reactions must be physically separated Called electrochemical cells

Can create electricity or be used to create a chemical change!!

Galvanic Cells

Invented by Alessandro Volta in 1800

Galvanic cells: electrochemical cells used to convert chemical energy into electrical energy Examples alkaline batteries Made of half cells

One part of the galvanic cell where oxidation or reduction is occurring

Schematic for separating the oxidizing and reducing agents in a redox reaction.

Cu2+ + 2e- Cu Fe2+ Fe3+ + e-

Cu2+

Why won’t the reaction continue??

Cu2+ + 2e- Cu Fe2+ Fe3+ + e-

Build up of chargeswould require largeamounts of energy

Solutions mustbe connectedto allow ionsto flow!

Salt Bridge: contains a strong electrolyte held in place by gel

Porous Disk: allows ion flow without mixing

solutions

Allows ions to pass between solutions, but doesn’t allow the solutions to mix

Parts of a Galvanic Cell

Electrode: Conductor in a

circuit that carries electrons to a metal

Anode = oxidation Negatively charged

Cathode = reduction Positively charged

Steps of a Galvanic Cell

e- created at anode Shown in

oxidation half-reaction

e- leave zinc and pass through wire

e- enter cathode and cause reduction Shown in half-

reaction Positive and

negative ions pass through salt bridge to finish the circuit

Figure 21.5 A galvanic cell based on the zinc-copper reaction.

Oxidation half-reaction Zn(s) Zn2+(aq) + 2e-

Reduction half-reaction Cu2+(aq) + 2e-C Cu(s)

Overall (cell) reactionZn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Schematic of a battery.

Electron flow anode to cathode (- to +)

oxidation to reduction

reducing agent to oxidizing agent

Let’s practice drawing a Cu/Zn Galvanic Cell

Cu2+ + 2e- Cu Cathode/reduction Zn Zn2+ + 2e- Anode/oxidation Cu2+ + Zn Zn2+ + Cu

ZnCu

SO42-

Zn2

+

Voltaic Cell Shorthand

Oxidation half cell is listed first with reduced and oxidized species separated by a line.

Reduction is next in the opposite order.

Double line separates the two and represents a salt bridge and electron transfer:

Zn|Zn2+||Cu2+|Cu

Voltaic Cell Shorthand

Draw shorthand notation for a Mg-Pb cell where the nitrate ion is present. You might want to refer to the activity series to determine what is oxidized and what is reduced! Draw a diagram for this galvanic

cell on the scratch paper provided! Label: anode, cathode, direction of e-

flow Write-out the ½ rxns and combined

reaction.