Chapter 7

Electrochemistry and corrosion

A-Electrochemistry

Electric cells are composed of two electrodes–solid electrical

conductors and at least one electrolyte (aqueous electrical

conductor).

In current cells, the electrolyte is often a moist paste (just enough

water is added so that the ions can move). Sometimes one electrode

is the cell container.

The positive electrode is defined as the cathode and the negative

electrode is defined as the anode.

Electrochemical Reactions

In electrochemical reactions, electrons are transferred from one species to another.

Oxidation Numbers

In order to keep track of what loses electrons and what gains them, we assign

oxidation numbers.

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Oxidation and Reduction

• A species is oxidized when it loses electrons.

Here, zinc loses two electrons to go from neutral zinc metal to the Zn2+ ion.

A species is reduced when it gains electrons.

Here, each of the H+ gains an electron and they combine to form H2.

H+ oxidizes Zn by taking electrons from it.

Zn reduces H+ by giving it electrons.

Assigning Oxidation Numbers

1. Elements in their elemental form have an oxidation number of 0.

2. The oxidation number of a monatomic ion is the same as its charge.

3. Nonmetals tend to have negative oxidation numbers, although some are

positive in certain compounds or ions.

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Oxygen has an oxidation number of −2, except in the peroxide ion in which it

has an oxidation number of −1.

Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal.

Fluorine always has an oxidation number of −1.

The other halogens have an oxidation number of −1 when they are negative;

they can have positive oxidation numbers, however, most notably in

oxyanions.

4. The sum of the oxidation numbers in a neutral compound is 0.

5. The sum of the oxidation numbers in a polyatomic ion is the charge on the

ion.

Galvanic Cells:

Produces electrical current spontaneous by chemical reactions à Battery

Cu2+ + 2e- D Cu E0 = +0.34 V

Zn2+ + 2e- D Zn E0 = −0.76V

Cu2+ + Zn D Cu + Zn2+ E0 = 0.34 – (-0.76) = 1.10 V

Parts of the voltaic or galvanic cell…

Anode à the electrode where oxidation occurs

After a period of time, the anode may appear to become smaller as it falls into

solution.

Cathode à the electrode where reduction occurs

After a period of time it may appear larger, due to ions from solution plating onto

it.

Salt Bridge à a device used to maintain electrical neutrality in a galvanic cell.

This may be filled with neutral salt or it may be replaced with a porous barrier.

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Electron Flow à always from anode to cathode (through the wire).

Anode Cathode

The diagram to the right illustrates what really happens when a Galvanic cell is

constructed from zinc sulfate and copper (II) sulfate using the respective metals as

electrodes.

Zn/Zn+2 is the anode

Zn Zn+2 + 2e- E° = +0.76 V

Cu/Cu+2 is the cathode

Cu+2 + 2e- Cu E° = 0.34 V

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Zn Zn+2 + 2e- E° = +0.76 V

Cu+2 + 2e- Cu E° = 0.34 V

Cu+2 + Zn Cu + Zn+2 E°cell = 1.10 V

Shorthand Notation

Electromotive Force (emf)

• Water only spontaneously flows one way in a waterfall.

• Likewise, electrons only spontaneously flow one way in a redox reaction—

from higher to lower potential energy.

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• The potential difference between the anode and cathode in a cell is called the

electromotive force (emf).

• It is also called the cell potential, and is designated Ecell.

Batteries

A battery is a group of galvanic cells connected in series

The potentials of the individual cells add to give the total battery potential

Secondary cells can be recharged by adding electricity

Figure 1: One of the Cells in a 12V Lead Storage Battery

Lead –acid battery:

If a number of cells are connected in series, the arrangement is called a battery.

The lead storage battery is one of the most common batteries that is used in the

automobiles. A 12V lead storage battery is generally used, which consists of six

cells each providing 2V. Each cell consists of a lead anode and a grid of lead

packed with lead oxide as the cathode. These electrodes are arranged alternately,

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separated by a thin wooden piece and suspended in dil.H2SO4 (38%), which acts as

an electrolyte (Fig. 1).Hence it is called Lead-acid battery.

Anode: Pb

Cathode: PbO2

Electrolyte: H2SO4 (38%)

EMF=2V

To increase the current output of each cell, the cathode and the anode plates are

joined together, keeping them in alternate positions. The cells are connected

parallel to each other (anode to anode and cathode to cathode). The cell is

represented as

Pb | PbSO4 (s), H2SO4 (aq.) | PbSO4 (s), Pb

In the process of discharging, i.e. when battery produces current, the reactions at

the electrodes are as follows:

At anode:

Pb à Pb+2 + 2e-

Pb (s) + SO4 (aq.) à PbSO4 (s)

At cathode:

PbO2 (s) + SO4 (aq.) + 4H+ (aq.) + 2e– à PbSO4 (s) + 2H2O

Therefore, overall reaction is

Pb (s) + PbO2 (s) + 2H2SO4 (aq.) à 2PbSO4 (s) + 2H2O

During discharging the battery, H2SO4 is consumed, and as a result, the density of

H2SO4 falls; when it falls below 1.20 g/cm3, the battery needs recharging. In

Discharging, the cell acts as a voltoic cell where oxidation of lead occurs.

During recharging, the cell is operated like an electrolytic cell, i.e. electrical energy

is supplied to it from an external source. The electrode reactions are the reverse of

those that occur during discharge.

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PbSO4 (s) + 2e– àPb (s) + SO4– – (aq.)

PbSO4 (s) + 2H2O à PbO2 (s) + 2H2SO4 + 2e–

2PbSO4 (s) + 2H2O à Pb (s) + PbO2 (s) + 2H2SO4 (aq.)

During this process, lead is deposited at the cathode, PbO2 is formed at the anode

and H2SO4 is regenerated in the cell.

Advantages: Lead acid batteries are used for supplying current to railways, mines,

laboratories, hospitals, automobiles, power stations, telephone exchange, gas

engine ignition, Ups (stand-by supplies). Other advantages are its rechargeability,

portability and Its relatively constant potential & low cost.

Disadvantages: Use of Conc.H2SO4 is dangerous; Use of lead battery is fragile.

B. CORROSION

What is corrosion?

Corrosion is a natural event

It represents a return of metals to their more natural state as minerals (oxides).

Metal Corrosion

The destruction of a material by chemical or electrochemical reaction to its

environment”

Typically a transfer of electrons from one metal to another through an

Oxidation-Reduction Reaction.

Oxidation - Reduction Reaction

Anodic metal gives up electrons (oxidation)

Cathodic metal accepts electrons (reduction)

Or gases accept electrons (reduction)

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Al→ Al+ 3+3e−Fe→Fe+2+2e−

Cu2 ++2e−→Cu

2 H++2e−→H2( gas )

Acceleration of Corrosion

Physical Characteristics

time of exposure (more time, more corrosion)

exposed area (less, increases corrosion rate)

Environmental Characteristics

acidic environment

sulfur gas environment

temperature (high temps, more corrosion)

moisture (oxygenated moisture)

Passivation

Refers to a material becoming "passive" that is, being less affected by

environmental factors such as air and water. Passivation involves a shielding outer-

layer of base material,

Types of Passivation

A protective film in oxidizing atmospheres

Chromium, nickel, titanium, aluminum

Metal oxide layer adheres to parent metal

Barrier against further damage

Self-healing if scratched

Forms of Corrosion

Uniform corrosion of a single metal

Usually an electrochemical reaction.

Relatively slow and predictable.

Rusting of exposed steel, tarnished silver.

Easily corrected with coatings and regular maintenance.

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Galvanic Corrosion

2 dissimilar metals, electrolyte, electrical connection and oxygen.

Pitting Corrosion

Localized corrosion forming holes.

Difficult to initially detect.

Reinforcement Corrosion

Corrosion Products

Fe + 2OH = Fe(OH)2

Oxidation of Fe(OH)2

Fe(OH)3 (rust)

Corrosion of Metals in Concrete

Concrete is Normally Highly Alkaline

Protects Steel from Rusting if Properly Embedded

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If Corrosion Occurs, the Reaction Products are greater in volume than the

original steel.

Corrosion Initiation and Rate Depends On

Amount of Concrete Cover, Quality of Concrete

Details of Construction, & Exposure to Chlorides

Avoiding Corrosive Situations

Choose couple metals close on the galvanic series

Use large anode, and small cathode areas

Electrically insulate dissimilar metals

Connect a more anodic metal to the system

Corrosion Prevention

Coatings

Barrier films

Inhibitive Pigments

Sacrificial treatments

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Paint

Active Cathodic Protection

Cathodic protection is an electrochemical means of corrosion control in which the

oxidation reaction in a galvanic cell is concentrated at the anode and suppresses

corrosion of the cathode in the same cell. This is achieved by placing a more easily

corroded metal to act as the anode of the electrochemical cell in contact with the

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