chem 471 part 2: industrial electrochemistry

Chem 471 Part 2: Industrial Electrochemistry


2.1 Electrochemical Cells Background

2.2 Batteries

2.3 Fuel Cells

2.4 Corrosion

2.5 Sodium Chloride

2.6 Chloralkali Products (Cl2, NaOH, H2 from aqueous NaCl electrolysis)

2.7 Aluminum Production

2.8 Copper Refining

2.9 Other Electrochemical Industries


2.1 Electrochemical Cells - Background

why a whole section on electrochemistry?

what’s so special about electrochemistry?

advantages and disadvantages?

There are different ways to run chemical reactions.

Plan A. Direct Reaction (traditional “shake-and-bake” chemistry)

The obvious way to do chemistry is to bring the reactants into direct

contact. For example, mix hydrogen and oxygen

H2 + ½ O2 H2O

The spontaneous reaction to form water proceeds. O atoms (more

electronegative) gain electron density, and H atoms lose electron

density. Oxygen is reduced and hydrogen is oxidized.

Pt catalyst

Plan B. Electrochemical Reactions

A less obvious but important way to do chemistry:

oxidation and reduction reactions occur at different locations

electrons are transferred from the chemical being oxidized to the

chemical being reduced through an external circuit, usually a

metal wire (an electronic conductor)

the reactants are separated by an electrolyte solution (an ionic

conductor, but not an electronic conductor)

Industrial Electrochemistry:

use spontaneous electrochemical reactions to produce electric

current do electrical work (batteries and fuel cells, corrosion)

apply an electric current to drive electrons in the nonspontaneous

direction, to force chemical reactions that are impossible by

simply mixing the reactants (electrolysis cells)

Electrochemical Reaction of Oxygen and Hydrogen (spontaneous)

Bubble hydrogen gas over a Pt electrode (why platinum?) and bubble

oxygen over another Pt electrode. Dip the electrodes in an aqueous

hydrochloric acid solution (the electrolyte, an ionic conductor with

mobile H+(aq) and Cl-(aq) ions).Connect the electrodes with a metal

wire to conduct electrons externally to generate electric current.

anode (oxidation) cathode (reduction) H2(g) 2H+(aq) + 2e 2H+(aq) + ½ O2(g) + 2e H2O(l)

Overall: H2(g) + ½ O2(g) H2O(l) (n = 2 moles of electrons)

Electrolyse Water to Make Hydrogen and Oxygen (nonspontaneous)

Apply an electric current (using a battery or a dc power supply) to

drive the nonspontaneous reaction, stripping electrons from water and

“forcing” them onto hydrogen ions.

cathode (reduction) anode (oxidation) 2H+(aq) + 2e H2(g) H2O(l) 2H+(aq) + ½ O2(g) + 2e

Overall: H2O(l) H2(g) + ½ O2(g)

Examples of Anodic (Oxidation) Reactions

Ce3+(aq) Ce4+(aq) + 2e- simple electron transfer

Fe(s) Fe2+(aq) + 2e- anodic dissolution

2Cl- (aq) Cl2(g) + 2e- gas evolution

Pb(s) + SO42-(aq) PbSO4(s) + 2e- phase conversion

2Al(s) + 3H2O(l) Al2O3(s) + 6H+(aq) + 6e- oxide formation

CH3OH(l) + H2O(l) CO2(g) + 6H+(aq) + 6e- fuel oxidation

(“cold” combustion)

Examples of Cathodic (Reduction) Reactions

Fe3+(aq) + e- Fe2+(aq) electron transfer

Cu2+(aq) + 2e- Cu(s) metal deposition

2H2O(l) + 2e- H2(g) + 2OH-(aq) gas evolution

O2(g) + 4H+(aq) + 4e- 2H2O(l) gas depletion

PbO2 (s) +4H+(aq) + SO4

2-(aq) PbSO4(s) + 2H2O(l) phase

conversion

2CH2=CHCN + 2H2O(l) + 2e- (CH2CH2CN)2 + 2OH-(aq)

dimerization

Classifying EChem Cells as Spontaneous or Nonspontaneous (remember Chem 231, 232)

1. Using Changes in the Gibbs Free Energy ( G)

From thermodynamics, a chemical reaction is spontaneous at a given

temperature and pressure if it decreases the Gibbs free energy.

GT,p < 0 (spontaneous)

GT,p > 0 (nonspontaneous)

Example H2(g) + ½ O2(g) H2O(l) (at 25 oC, 1 bar)

Go = Gfo(products) Gf

o(reactants)

= Gfo(H2O(l)) Gf

o(H2(g)) 0.5 Gfo(O2(g))

= 237.13 kJ mol-1 0 0 = 237.13 kJ mol-1

Shows the conversion of pure hydrogen and pure oxygen to liquid

water is spontaneous at 25 oC and 1 bar.

Classifying EChem Cells as Spontaneous or Nonspontaneous (remember Chem 231, 232)

1. Using Changes in the Gibbs Free Energy ( G)

Voltage of an Electrochemical Cell Under reversible conditions

(fast electrode reactions, no side reactions) the cell voltage is

Eo = Go/nF = (237.13 kJ mol 1)/(2 (96485 C mol 1))

= 1.229 Volt (positive cell voltage spontaneous reaction)

Electrical Work Also under reversible conditions, the conversion

of one mole of H2(g) and one half mole of O2(g) can be used to

produce the electrical work

we = Go = nFEo = 237.13 kJ mol1

(negative electrical work work done on the surroundings)

Classifying EChem Cells as Spontaneous or Nonspontaneous (remember Chem 232)

1. Using Tables of Standard Reduction Potentials

A standard reduction potential is a voltage measuring the relative ease

of reducing (adding electrons) to molecules or ions in their standard

states (all gases at 1 bar and all dissolved ions at unit activity).

Examples

Fluorine is very easily reduced (F2 is a good oxidizer):

1/2 F2(g) + e- F-(aq) ERo = +2.87 Volt (25 oC)

Lithium ions are difficult to reduce (Li metal is good reducing agent):

Li+(aq) + e- Li(s) ERo = -3.045 Volt (25 oC)

Hydrogen is assigned zero standard reduction potential:

H+(aq) + e- 1/2 H2(g) ERo =0 Volt (definition)


Standard Reduction Potentials at 25 oC

Li+(aq) + e- → Li(s) - 3.04 V “difficult”

Rb+(aq) + e- → Rb(s) - 2.95

K+(aq) + e- → K(s) - 2.92

Ca2+(aq) + 2e- → Ca(s) - 2.76

Na+(aq) + e- → Na(s) - 2.71

Mg2+(aq) + 2e- → Mg(s) - 2.38

Al3+(aq) + 3e- → Al(s) - 1.71

Zn2+(aq) + 2e- → Zn(s) - 0.76

Fe2+(aq) + 2e- → Fe(s) - 0.41

Cd2+(aq) + 2e- → Cd(s) - 0.40

Ni2+(aq) + 2e- → Ni(s) - 0.23

Pb2+(aq) + 2e- → Pb(s) - 0.13

2H+(aq) + 2e- → H2(g) 0

Cu2+(aq) + 2e- → Cu(s) 0.34

Ag+(aq) + e- → Ag(s) 0.80 “easy”


1. Using Tables of Standard Reduction Potentials

Calculate Eo for the cell reaction Li(s) + ½ F2(g) = Li+(aq) + F-(aq)

½ F2(g) + e- F-(aq) ERedo = +2.87 V

Li(s) Li+(aq) + e- EOxo = -ER

o = +3.04 V ____________________________________________________________________________________________ _____________________________________________________________

Li(s) + ½ F2(g) Li+(aq) + F-(aq) Eo = +2.87 + 3.04 V

= +5.91 V

Warning! The actual cell voltage will be different if:

a) the chemicals are not in their standard states

b) the electrode reactions are slow

c) there are side reactions, such as

Li(s) + H2O(l) Li+(aq) + OH-(aq) + ½ H2(g)

2.2 BATTERIES

Electric power is generated by a spontaneous chemical reactions. In

a fuel cell, fuels such as hydrogen, methane, methanol, etc. are

reacted electrochemically with oxygen.

Advantages

• chemical energy is converted directly into electric current,

without electrical generators or moving mechanical parts

• from thermodynamic considerations, the full free energy change

of the chemical reaction can be converted into electrical work

(no heat engine Carnot limitation)

• batteries are the only practical power sources for small portable

electronic devices (radios, laptops, flashlights, cell phones, etc.).

BATTERIES

Disadvantages

In principle, any spontaneous chemical reaction can be used to

directly generate electric current electrochemically. In practice:

• Unwanted side reactions can occur. Aluminum electrodes, for

example, are coated with a thin insulating layer of Al2O3.

• The electrode reactions may be too slow to produce adequate

power (energy per unit time) for some applications.

• In competition with combustion engines in cars, trucks,

electricity power generating stations, etc., batteries and fuel

cells are still too expensive for large-scale use.

Important Battery Systems

Lead-Acid Cell [E = 2.05 V, secondary cell (rechargeable)]

cathode PbO2(s) + 4H+ + SO42- + 2e- 2H2O + PbSO4(s) 1.93 V

anode Pb(s) + SO42- PbSO4(s) + 2e- 0.36 V

overall PbO2(s) + Pb(s) + 4H+ + 2SO42- 2PbSO4(s) + 2H2O 2.29 V (Eo)

electrolyte aqueous H2SO4

Uses: car and truck engine ignition

emergency power supplies

electric vehicles (e.g., fork lifts)

About 150 million lead-acid batteries sold per year, worth about

$15 billion (45 % of total battery production).

(why not 2.05 V?)

Lead-Acid Cell (secondary cell, rechargeable)

E = 2.05 V The why are car lead-acid batteries rated at 12 V?

Pb PbO2

Lead-Acid Cell (secondary cell, rechargeable)

Good: inexpensive

rugged and reliable (proven 150-year-old technology)

rechargeable

can deliver large pulse currents (400 A) to start engines

Bad: toxic electrodes and electrolyte (lead and sulfuric acid)

heavy (low energy density), limiting the use of lead-acid

batteries in electric cars and portable electronic devices)

__________________________________________________________________________________________________________________________________________________________________

a) What have lead-acid battery manufacturers done to make the their

product more acceptable to consumers? b) How can lead acid batteries be

modified to avoid H2SO4 spills or leaks? c) Give 3 reasons why an

aluminum-acid battery (if developed!) would be better.

Zinc-Carbon Dry Cell (1.5 V, primary cell, not rechargeable)

cathode 2MnO2(s) + 2NH4+(aq) + 2e! Mn2O3(s) + 2NH3(aq) + H2O 0.6 V

anode Zn(s) Zn2+ + 2e! 0.76 V

overall 2MnO2(s) + Zn(s) + H2O Mn2O3(s) + Zn2+ + 2OH! 1.4 V (Eo)

electrolyte moistened NH4Cl/ZnCl2/MnO2/graphite powder

Uses: Small (but not miniature) portable

low-drain electronic devices such as

radios, clocks, flashlights, ...

Most popular battery ranked by number

produced: 10 billion per year.

First commercial dry battery. ______________________________________________________________________________

Why are zinc-carbon batteries called dry cells? Why is graphite powder used in

the electrolyte? Why is the terminology “carbon-zinc battery” inaccurate?

complicated!

Zinc-Carbon Dry Cell

Good:

very cheap

nontoxic

rugged

works in any orientation

Bad:

disposable

treated as hazardous waste in some jurisdictions

short shelf-life (zinc anode attacked by aqueous NH4Cl) __________________________________________________________________________________________________________

Why are gaskets and seals (often overlooked) very important in the construction

of dry cells? Why is carbon (not a metal) used as the cathode?

Development of a Better Dry Cell

A case study illustrating how industrial chemists produce new and

useful products (and big profits for chemical companies).

The short lifespan of dry cell batteries was bad for sales, consumer

confidence and company shareholders.

In the 1950’s a research team led by Lewis Urry at the Eveready

Battery Company improved the dry cell, replacing the NH4Cl/ZnCl2

electrolyte with aqueous KOH. The lifetime of the new battery was

longer, but the current was too low.

This problem was solved by using

powdered zinc/KOH gel as the anode

(why powdered zinc?), leading to the

development of the alkaline battery.

Alkaline Dry Cell (1.5 V, primary cell, not rechargeable)

cathode 2MnO2(s) + H2O(l) + 2e! Mn2O3(s) + 2OH-(aq) 1.28 V

anode Zn(s) + 2OH-(aq) ZnO(s) + H2O(l) + 2e! 0.15 V

overall 2MnO2(s) + Zn(s) Mn2O3(s) + ZnO(s) 1.43 V (Eo)

electrolyte aqueous KOH/zinc powder gel

Uses: Similar to those for

zinc-carbon cells

Almost 10 billion sold per year.

More expensive to produce than zinc-carbon cells, but can be

sold at higher prices.

Alkaline Dry Cell

Advantages over

zinc-carbon cells:

longer shelf life

(up to 10 yr)

about three times

the capacity for the

same size battery

Disadvantages:

corrosive KOH electrolyte

higher manufacturing costs

older models contained mercury to limit cathode side reactions ________________________________________________________________________________________________________________

Why is it important that zinc-carbon and alkaline cell voltages are the same (1.5 V)?

Lithium Ion Cell (3 to 4 V, secondary cell, rechargeable)

cathode CoO2(s) + Li+(nonaq) + e- LiCoO2(s)

anode Li (adsorbed on graphite) Li+(nonaq) + e-

overall Li(adsorbed on C) + CoO2(2) LiCoO2(s) 3 to 4 V

electrolyte LiClO4, LiBF4 or LiPF6 in an organic solvent such as

dimethylcarbonate or diethylcarbonate

Uses: laptops, notebooks, cellphones, cameras, electric power tools,

electric cars and trucks (all requiring higher energy densities than

can be provided by lead-acid or dry cells)

About 2 billion lithium ion cells are made per year. __________________________________________________________________________________________________________________________________

Questions. a) Why use expensive LiClO4, LiBF4 or LiPF6 salts for the

electrolyte, instead of much cheaper LiCl? b) Why use expensive organic

solvents such as alkylcarbonates, instead of cheaper organic solvents, such

as benzene or methanol? c) Why not use water as the solvent?

Lithium Ion Cell

Lithium metal is very good for high energy density batteries

because lithium is light (density 0.53 g cm-3, about half the

density of liquid water!) and strongly electropositive

Li(s) → Li+(aq) + e− Eo = 3.04 V

But aqueous electrolytes can’t be used. In contact with water,

lithium metal spontaneously reduces water:

Li(s) → Li+(aq) + e− 3.04 V

e− + H2O(l) → ½ H2(g) + OH− −0.83 V

Li(s) + H2O(l) → Li+(aq) + ½ H2(g) + OH−(aq) 2.21 V

Lithium Ion Cell

Lithium Ion Cell

Good: high energy density (500 to 1000 kJ per kg cell) compared

to other cells (about 100 kJ per kg for lead-acid cells)

rechargeable (secondary cell)

work in any orientation

Bad: expensive compared to lead-acid and dry cells

can catch fire* (Why? Li/organic solvent) if overheated or

if leaks develop

___________________________________________________________________________________________________________________________________

*cautionary tales about pushing battery technology to its limits:

Galaxy 7 smartphone recall due to lithium ion cells catching fire

cost Samsung about $10 billion

Boeing 787 Dreamliner fleet grounded for months due to lithium ion

battery fires and failures cost Boeing about $1 billion

Lithium Ion Cell

Lithium ion cell technology is actively investigated in university*,

government and industrial labs. Improved electrodes, electrolytes

(e.g., solid polymers) cell design are under development. Goals:

higher energy density (especially for electric vehicle applications)

faster recharge

higher reliability (less degradation, safer to operate)

*Prof. Jeff Dahn

Dal’s “Battery Man”

and

Herzberg Medalist

(Canada’s highest science award)

Less Commonly Used (but important) Battery Systems

Lithium Dry Cells (primary, about 3 V)

Li(s) + MnO2(2) LiMnO2(s)

Also called lithium metal cells. (Different from lithium ion cells.)

More expensive than other dry cells, but last longer

Silver-Zinc “Button” Cells (primary or secondary, 1.8 V)

Ag2O(s) + Zn(s) 2Ag(s) + ZnO(s)

Stable voltage during discharge. Good for miniature devices, such as

watches, calculators and heart pacemakers.

Nickel-Cadmium Cells (secondary, about 1.5 V)

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

Low maintenance. Used for emergency power supplies. _________________________________________________________________________________________________________________________________________________________________

Note: Hundreds of other battery systems have been developed.

Energy and Power Density of Batteries

What’s the difference between energy and power?

Energy density an important design considerations for batteries. Why?

A battery can have a high energy density, but a low power density.

How is this possible? Why can this kind of battery still be useful?

Energy Densities

lead-acid batteries 0.12 MJ/kg

advanced Li ion batteries 1.0 MJ/kg

but…

gasoline 50 MJ/kg why so high?

liquid hydrogen 150 MJ/kg why even higher? _____________________________________________________________________________________________________________________________________________________________________

Is there some way to combine the high energy density of liquid fuels and

the advantages of electrochemical cells .… ?

.… 2.3 FUEL CELLS ( “Flow” Batteries)

batteries must be replaced (primary cells) or recharged

(secondary cells) when the reactants in the anode or

cathode compartments are used up

fuel cell: fresh reactants are pumped into the cell and

reaction products are pumped out

fuel cells can therefore operate continuously

chemical energy is converted directly into electrical work

fuel-cells are more efficient than heat engines because their

performance is not subject to the Carnot heat-engine limitation: :

w/qH = 1 (TC /TH)

Hydrogen/Oxygen Fuel Cells

cathode ½ O2 + H2O + 2e- 2OH- 0.40 V

anode H2 + 2OH- 2H2O + 2e- 0.83 V

overall H2(g) + ½ (g) H2O(l) Eo = 1.23 V

electrolyte aqueous KOH

Uses: limited by high cost and slow cathode reactions, but

promising and under active research and development

Other Fuel Cells The hydrogen/air cell operates at about 200 oC (to

speed up the electrode kinetics) using phosphoric acid

as the electrolyte and hydrogen from steam-reformed

natural gas. Cells operating at 600 oC using molten

carbonate electrolytes and methane fuel have also

been developed. Fuel cells are promising, but still too

costly to compete with other power supplies.


(PEM)


Application: Space Missions

cost not important here! (why?)

Apollo Service Modules carried

three H2/O2 fuel cells, each with

31 pairs of electrodes in series

30 V, power 500 W to 2000 W

used liquid H2 and O2 onboard for

the Service Module main engine

fuel cell waste (H2O) provided

drinking water for the crew!


H2(g) + ½ O2(g) H2O(l) 25 oC, 1 bar

Go = Gfo(products) Gf

o(reactants)

= Gfo(H2O(l)) Gf

o(H2(g)) ½ Gfo(O2(g))

= 237 kJ mol1 0 0 = 237 kJ mol1

Eo = Go/nF = (237 kJ mol1)/(2 96485 C mol1)

= 1.23 Volt

Electrical Work Under standard conditions, the reaction

of one mole H2 and one half mole O2 can

produce the electrical work

we = Go = nFEo = 237 kJ mol1

OK, but why is this “promising”?

Comparison with a Mechanical Heat Engine

Instead of a fuel cell, burn H2 in direct contact with O2 to produce

heat and high-pressure steam to run a turbine heat engine operating

with steam at TH = 500 K and the surroundings at TC = 300 K .

Using standard enthalpies of formation

H2(g) + ½ O2(g) H2O(l) 25 oC, 1 bar

heat released qH = Ho = Hfo(products) Hf

o(reactants)

= Hfo(H2O(l)) Hf

o(H2(g)) ½ Hfo(O2(g))

= 286 kJ mol1 0 0 = 286 kJ mol1

But from the Second Law of thermodynamics and the Carnot

limitation, the maximum mechanical work obtained is

wmech = Ho [1 (TC /TH )] = 114 kJ mol1

Important Nearly twice as much work can be produced electrically

in a fuel cell than in a heat engine (237 kJ vs. 114 kJ).

2.4 CORROSION

The spontaneous electrochemical oxidation of metals.

Example:

Rusting of iron and steel

to form iron oxides

corrosion can decide the lifetime of oil rigs, pipelines, reactors, …

costs about 1 trillion dollars per year

many industrial processes employ highly corrosive chemicals

equipment failures caused by corrosion can be dangerous

if corrosion can be understood, then steps can be taken to fight it

CORROSION

Corrosion reactions do not occur directly.

Instead, oxidation half-reactions such as

Fe(s) Fe2+ + 2e 0.44 V (Eo)

and reduction half-reactions such as

½ O2 + 2H+ + 2e H2O (acidic solutions) 1.23 V

2H+ + 2e H2 (acidic solutions) 0.00 V

½ O2 + H2O + 2e 2OH (alkaline solutions) 0.40 V

occur electrochemically at different locations on the metal.

Galvanic Corrosion When two different metals are brought into

contact in the presence of air and moisture, the metal with the less

positive reduction potential (easier to oxidize) will act as the anode

and dissolve, while the other metal will act as the cathode. Galvanic

corrosion can be very rapid.

Crevice Corrosion If a crack or a

crevice develops in a metal object, the

outer part of the crevice (exposed to air)

acts as the cathode and the inner part

acts as the anode (metal oxidation and

dissolution). Once crevice corrosion

begins, the depth of the crevice will increase

which can result is holes and structural failure.

Corrosion Caused by Paint Scratches

Paint or other surface coatings can protect metals from corrosion. But

if the coating is scratched, the exposed metal acts as a cathode

example: 1/2 O2 + 2H+ + 2e H2O

and the nearby painted metal acts as the anode and dissolves

example: Fe(s) Fe2+ + 2e

Causes paint to lift and peel

along the scratch, leading

to further corrosion.

Fix paint scratches!

Corrosion Control

avoid air, water, and other corrosive chemicals (usually impractical!)

coat metal with a protective layer (paint, a more resistant metal, oil)

repair holes or scratches in the protective layer

avoid contact with less electropositive metals (galvanic corrosion)

coat or connect the metal to be protected to a more electropositive

metal that acts as sacrificial anode which corrodes first

example: galvanized steel

(protective zinc coating)

Zn Zn2+ + 2e (anode) 0.76 V

1/2 O2 + 2H+ + 2e H2O (cathode) 1.23 V

Up next: Electrolytic Chemical Industries

Electrolysis Use an applied voltage to pump electrons into the

cathode and pull them out of the anode:

forcing a nonspontaneous electrochemical reaction to occur

Three most important electrolytic industries:

1. Brine (Aqueous NaCl) Electrolysis (chloralkali industry)

2Na+(aq) + 2Cl (aq) + 2H2O 2Na+(aq) + 2OH (aq) + Cl2 + H2

2. Aluminum Production

2Al2O3 + 3C 4Al + 3CO2

3. Copper and Zinc Electro-Refining

Cu2+(aq) + SO42(aq) + 2e Cu(s) + SO4

2(aq)

Zn2+(aq) + SO42(aq) + 2e Zn(s) + SO4

2(aq)

2.5 SODIUM CHLORIDE (“Salt”)

NaCl is not included in most “top 50" lists of industrial chemicals.

Why? Most industrial NaCl is “captive” (produced and used by the

same manufacturer to produce other chemicals, not sold). Also, NaCl

occurs naturally in very pure form (rock salt, typically 98 to 99% pure

NaCl), so little or no chemistry is involved in salt production.

But in terms of tonnage, NaCl ranks near sulfuric acid (#1 on the lists).

Uses of NaCl 45% chloralkali production (Cl2, NaOH, H2)

20% other industrial chemical manufacturing

25% ice control on roads

5% food products

5% miscellaneous

SODIUM CHLORIDE (“Salt”)

Vast amounts of NaCl are available in seawater and beds of rock salt.

NaCl production

1. from brine

(50 %)

2. underground mining

(30 %)

3. seawater evaporation

(20 %)

Sifto Salt Mine

under Lake Huron, near

Goderich, Ontario

1. Brine Water is pumped into drilled salt beds. A saturated solution

containing ~25 % NaCl (and other dissolved salts) is pumped out and

treated with Na2CO3 to precipitate Ca2+, Mg2+ and Fe3+. The purified

brine is used directly or evaporated to precipitate 99.8% pure NaCl.

2. Mining Mines are dug into deposits of rock salt (typically 98 to

99% pure NaCl). Clay, sand and other solid impurities are removed by

sieving or gravitational separation. This product is suitable for salting

roads. Further purification required for NaCl used in chemical

production is carried out by dissolving the solid NaCl in water and

using Ca(OH)2 or Na2CO3 to precipitate Ca2+, Mg2+ and Fe3+.

3. Seawater Evaporation First, Mg2+ and most Ca2+ are removed

as described above, then water is removed by evaporation, usually in

large, shallow solar ponds. Precipitated NaCl (typically 99.8% NaCl)

suitable for electrolysis or food products.

NaCl Production

2.6 CHLORALKALI PROCESSES

Aqueous NaCl solutions are electrolyzed to make Cl2, NaOH and H2.

Significance

70% of all industrial chemical products use Cl2 and/or NaOH

in one or more synthesis steps

this is the largest electrochemical industry (75 million tonnes Cl2,

80 million tonnes of NaOH per year)

illustrating the competitive nature of process economics, pollution

control and safety considerations, three different chloralkali

processes are used: 1. diaphragm cells

2. membrane cells

3. mercury cells

Chloralkali Industry

Aqueous NaCl solutions are electrolyzed to make Cl2, NaOH and H2.

why not use natural deposits of Cl2 and NaOH?

why are Cl2 and NaOH electrochemically synthesized, not

prepared by direct chemical reactions, such as

Ca(OH)2(aq) + Na2CO3(s) → CaCO3(s) + 2 NaOH(aq)

HCl(aq) + MnO2(s) → MnCl2(aq) + 2 H2O(l) + Cl2(g)

why is Cl so stable compared to Cl2?

why are mercury chloralkali cells banned in many countries

1. DIAPHRAGM Chloralkali Cells

simple but effective chlor-alkali technology

Anode plates are mounted vertically and parallel to one another. Flat,

hollow steel mesh cathodes fit between the anode plates in a “toast-

rack” arrangement. A diaphragm consisting of a mat of asbestos fibers

outside the cathodes provides a physical barrier between the anode

(Cl2 produced) and the cathode (H2 and OH produced) solutions.

anode

2Cl(aq) Cl2(g) + 2e Diaphragm Cell

(side view)

cathode

2 H2O(l) + 2e H2(g) + 2OH(aq)

DIAPHRAGM Chloralkali Cells

typical diaphragm cell:

dimensions 3 m by 2 m by 2 m

1,200 amp current

voltage 3.2 V to 3.8 V

3.5 tonnes Cl2 per day


Advantages:

simple low-cost cell design

cells are easy to operate

Disadvantages:

chloride contamination Ions are free to diffuse through the pores

of the diaphragm, a simple physical barrier, so the NaOH solution is

contaminated with Cl ions from the brine on the other side of the

diaphragm. The solution leaving the cathode chambers is 15% NaCl

and 12% NaOH. The NaCl content must be reduced.

low NaOH concentration After leaving the cells, evaporation is

used to concentrate the NaOH from to 50% (the usual commercial

product), adding to the cost of the diaphragm process. NaCl has a low

solubility in concentrated NaOH solutions, so most of the NaCl

precipitates out, reducing the NaCl concentration to about 1%.


Disadvantages:

electrical resistance of the diaphragm The diaphragm partially

blocks the ionic current flowing between the anode and cathode. To

overcome this resistance for useful production rates, a voltage

difference of 3.2 to 3.8 V is required, considerably higher than 2.2 V

required for a cell with no internal resistance. And the brine feed must

be purified to avoid precipitation of Mg(OH)2 and Ca(OH)2 which

would clog the diaphragms and further raise their resistance.

diaphragm lifetime Due to clogging, asbestos diaphragms

must be replaced every few months. Requires cell dismantling.

asbestos Requires special handling and disposal procedures.

2. MEMBRANE Chloralkali Cells

ingenious materials science – membranes that selectively

transport Na+ ions !

The problems with asbestos diaphragms prompted research to develop

membranes that transport Na+ ions, but block the undesirable flow

of Cl and OH anions between cathode and anode compartments.

Membrane Chloralkali Cell

Diagram not to scale!

In practice, the membranes are

very thin (about 0.2 mm) and

the electrodes are very close

together (4 mm apart). Why?

MEMBRANE Chloralkali Cells

Cell Room:


ingenious materials science – membranes that selectively

transport Na+ ions !

Chloralkali cell membranes are thin sheets of perfluorinated

polyethylene with side chains terminating in sulfonate groups

(tradename “Nafion”). The membranes contain microscopic “pockets”

(about 2 nm diameter) connected and lined with sulfonate groups. Na+

ions are transported through the membrane by being passed from one

sulfonate group to the next, making the membrane a cation conductor.


Advantages:

very pure NaOH solutions are produced (< 50 ppm Cl)

less power is required than for the mercury process

no mercury or asbestos is used.

Disadvantages:

highly purified brine is required to avoid fouling the membranes

with precipitated Mg(OH)2) and other impurities

because the membranes are not perfectly cation selective (small

amounts of OH leakage), the NaOH concentration is limited to a

maximum of about 30 %, requiring some water evaporation

the Cl2 produced is contaminated with some oxygen.

3. MERCURY Chloralkali Process

Ingenious technology!

Anode: 2Cl (aq) Cl2(g) + 2e

Cathode: Na+(aq) + Hg(l) + e NaHg(l amalgam)

Liquid mercury serves as the cathode, also absorbs sodium metal,

forming a liquid metal sodium + mercury solution (an amalgam).

The Na-Hg amalgam (typically 0.5% Na by weight), after leaving the

cell, is treated with water. Produces concentrated (~ 50 %) aqueous

NaOH of very high purity (< 30 ppm chloride).

2NaHg(l) + 2H2O(l) 2Hg (l) + 2Na+(aq) + 2OH (aq) + H2(g)

The mercury is recycled to the cell.

MERCURY Chloralkali Cells

A typical cell consists of thin (about 3-mm thick) layer of mercury in

the bottom of a shallow steel trough which slopes slightly to promote

the flow of mercury. Horizontal anodes (adjustable in height) are fitted

in the cell lid together with slits through which Cl2 gas is drawn off.

Mercury chloralkali plants

have cell rooms larger

than a football field, use

100 MW of power and

250,000 amperes of current

to produce 250,000 tons

of chlorine per year.


Graphite Anodes were used for many years. But these anodes are not

completely inert and the electrode kinetics are slow, requiring a

0.5 V overpotential to speed up the oxidation of Cl ions.

Dimensionally Stable Anodes (DSA)

After considerable R&D, the anodes are now

titanium mesh coated with RuO2 and smaller

amounts of other metal oxides (Co3O2 or PdO2)

to act as catalysts. These anodes are called

DSA’s because they are resistant to wet chlorine

and last 5 to 10 years. Important:

The catalysts reduce overpotentials to < 0.04 volt.

DSA’s are also used in diaphragm and membrane cells.


Cell Voltage

The standard voltage for the mercury cell (3.25 V)

2Cl(aq) Cl2(g) + 2e 1.36 V

Na+(aq) + Hg(R) + e NaHg(l amalgam) 1.89 V

3.25 V

is less favorable (more negative than for the simple cell (2.20 V)

2Cl (aq) Cl2(g) + 2e 1.36 V

2H2O(l) + 2e H2(g) + 2OH (aq) 0.84 V

2.20 V

As a result, mercury chloralkali cells require a larger voltage and

energy costs are higher. This disadvantage is offset by the high purity

of the products and the high concentration of aqueous NaOH produced

(~50 weight %, suitable for direct sale). For adequate production rates,

about 4.5 volts is applied to each mercury cell.

Mercury Chloralkali Cells

Advantages:

very pure Cl2 gas NaOH solutions (< 50 ppm Cl) are produced

50 % NaOH solution produced directly (no H2O evap. needed)

simple cell design, low maintenance

brine purification to remove Mg2+ and Ca2+ ions less rigorous than

for diaphragm and membrane cells

Disadvantages:

higher voltage and energy cost compared to diaphragm and

membrane chloralkali cells

high costs of mercury release avoidance and monitoring

mercury cells are banned in Japan and Europe

Chloralkali Industry

Three competing technologies are available.

1. diaphragm cells

2. membrane cells

3. mercury cells

Consensus?

The overall economics of mercury and diaphragm cells are similar.

But membrane cells are about 10% less costly to operate and are

gradually becoming the choralkali technology of choice.

2.7 ALUMINUM

2nd-largest electrolytic industry (after chloralkali production)

about 60 million tonnes of aluminum produced per year

aluminum metal is widely used because it is remarkably

light

strong

malleable (can be drawn and stamped without cracking)

corrosion-resistant ( why?)

essential for aircraft construction

main alternative to copper for electrically conducting wire

aluminum is the most important nonferrous metal (more important

than copper, lead or zinc)

ALUMINUM

Aluminum Ore

The most important aluminum ore is bauxite, a sedimentary rock

containing hydrated aluminum oxides, such as Al(OH)3 and AlO(OH),

together with iron oxides, silicates, clay and other impurities.

Bayer Process for Alumina (Al2O3) Production from Bauxite

Crushed bauxite ore and aqueous NaOH are heated in a pressure vessel

to about 150 oC and 20 bar ( why use high pressure?). Al3+ ions

dissolve, but iron oxide, silicates and other impurities are insoluble.

After filtration and cooling, pure Al(OH)3 (gibbsite) is precipitated.

Gibbsite is heated to 1200 oC to drive off water, producing high purity

aluminum oxide: Al(OH)3(s) = Al2O3(s) + H2O(g)

More than 90% of Al2O3 production is used to make aluminum. Most

of the rest is used to make refractory bricks, glass and abrasives.

ALUMINUM

Aluminum Production

Direct Chemical Reaction?

No. In principle, aluminum can be obtained from the oxide by reaction

with carbon (as used for iron production):

2Al2O3 + 3C 4Al + 3CO2

But this reaction is thermodynamically unfavorable. Aluminum “likes”

oxygen too much to give it up to carbon.

Small amounts of aluminum were first prepared by reacting anhydrous

aluminum chloride with potassium (a very strong reducing agent):

AlCl3 + 3K = Al + KC1

Due to the high cost of the reagents and the low yields, aluminum

produced by this reaction was more expensive than gold or platinum!

ALUMINUM

Aluminum Production

Electrolysis of aqueous Al3+ solutions?

No. Electrochemical reduction of aqueous Al3+ ions might seem to be a

possibility. But water is reduced instead:

Al3+(aq) + 3e Al(s) 1.71 V

2H2O(l) + 2e H2(g) + 2OH(aq) 0.83 V

(more favorable)

Electrolysis of Al3+ ions in nonaqueous molten Al2O3?

No. The melt is nonconducting nonionic liquid. Also, alumina is a

refractory material used to make firebricks to line furnaces. It melts at

2020 oC, a temperature too high for economical industrial processing.

ALUMINUM

Aluminum Production

Hall-Heroult Process – finally the breakthrough

Electrolysis of Al2O3 dissolved in molten cryolite (Na3AlF6). The

melts contain 7 to 12% aluminum oxide, near the aluminum oxide +

cryolite eutectic (minimum melting point) at 10.5% Al2O3 and 960 oC.

Al2O3 + NaAlF3 melts are ionic conductors

eutectic depression of the freezing point (from 2020 oC for pure

Al2O3 to 1000 oC) makes the process economical

Cryolite (sodium hexfluoroaluminate) is a rare mineral. There is only one mine with

commercial deposits (Ivittut, Greenland). Today most cryolite is synthesized using

6NaOH + 2Al2O3 + 6HF = 2Na3AlF6 + 6H2O

ALUMINUM

Aluminum Production

Hall-Heroult Electrolysis Cells

The cells are constructed using large steel tanks (9 m 3 m, 1 m deep)

open at the top and lined with alumina refractory bricks then carbon.

The base of the tank is lined with carbon blocks inlaid with steel bars

to improve the electrical conductivity.

Molten aluminum, which is slightly denser than the alumina + cryolite

melt, collects at the bottom of the tank and acts as the cathode.

The anodes, also carbon, are lowered into the tank from above at a rate

of about 2 cm per day to compensate for the carbon lost by reaction.

To reduce heat loss, a crust of solid alumina + cryolite is allowed to

form at the top of the exposed melt. .

ALUMINUM

Aluminum Production

Hall-Heroult Electrolysis Cells

About 4.5 volt is applied to each cell,

which is significantly higher than the

standard cell voltage owing to slow

anode kinetics and the electrical resistance

of the melt and carbon electrodes.

To reduce pollution, the

fluoride-containing gases leaving

the top of the cells pass over beds

of powdered alumina adsorbents.

ALUMINUM

Aluminum Production

Hall-Heroult Process Electrode Reactions. Ideally:

cathode Al3+(in the melt) + 3e Al(liquid)

anode 2O2 O2 + 4e

But it’s impossible to find economical electrodes that resist attack by

oxygen at cell temperatures (~1000 oC). In practice, the cells operate

with consumable carbon anodes. The overall cell reaction is

2Al2O3 + 3C 4Al + 3CO2

Although the anodes must be replaced, formation of CO2 as a reaction

product (instead of O2) changes the applied cell voltage from about

5.5 V to 4.5 V, significantly reducing energy costs.

ALUMINUM

Aluminum production is energy-intensive, requiring 14 to 18 kW hr

per tonne of aluminum (compared to 3 kW hr per tonne of chlorine

produced by electrolysis). Aluminum production consumes about 5 %

if the electricity generated in North America. There are ten aluminum

production plants in Canada. Why are nine located in Quebec?

The formation of CO2 at the anodes significantly reduces the energy

costs for aluminum production. Why ?

Pure aluminum is generally too soft for many applications (vehicle and

machine parts, structural beams, cans, foil, …) How is it hardened?

Raw aluminum is manufactured in the

form of large multi-tonne billets.

How are aluminum billets economically

shaped to form consumer products?

2.8 COPPER

2nd-most important nonferrous metal (after aluminum)

about 20 million tonnes of copper are produced per year

excellent electronic and heat conductor

resists corrosion

main uses

building construction (pipe, wiring, roofing, etc.,)

electronics (wiring, circuit boards, electric motors, etc.,)

consumer products

machine parts

important copper alloys: brass (copper + zinc)

bronze (copper + tin)

COPPER

Copper resists corrosion, but architectural copper turns green. Why?

COPPER

Copper Ores

The main copper ores are the sulfide minerals chalcopyrite (CuFeS2),

bornite (Cu5FeS4), covellite (CuS) and chalcocite (Cu2S) obtained from

large open-pit mines. Chile, Peru and U.S. are the largest producers.

open pit copper mine

(Bingham, Utah)

challenges: 1) typical copper ores contains only 0.6 % Cu

2) copper is not easily extracted from sulfides

COPPER

Copper Extraction

Step 1 froth flotation to enrich Cu content from 0.6% to 20%

Finely-ground ore is mixed with water, oil and surfactants in

large tanks. Compressed air is injected to form streams of

bubbles. Oil-coated grains of copper ore preferentially stick to

the bubbles and are floated to the top of the tanks where they

are skimmed off and collected. Clay, sand and other impurities

sink to the bottom of the tanks as sludge.

Step 2 smelting to remove iron and silicate impurities as slag

Enriched copper ore is dried and melted together with

limestone (CaCO3) and silica (SiO2). Iron impurities dissolve in

the molten Ca/Si slag. Liquid Cu2S, insoluble in the slag and

denser, sinks to the bottom of the crucibles and is tapped off.

COPPER

Copper Extraction

Step 3 roasting Cu2S to produce Cu2O

Hot Cu2S from the smelter is reacted with oxygen from air

Cu2S + 3/2 O2 = Cu2O + SO2

to convert cuprous sulfide to cuprous oxide. Pollution control is

important here. Why? What byproducts can be obtained?

Step 4 heating Cu2O to produce blister copper (raw copper)

Cuprous oxide is easily decomposed by heating to form

“blister” copper (named after it’s appearance). But the purity of

the copper is too low for many applications.

Step 5 electrorefining of blister copper to produce 99.9% pure Cu

COPPER

Copper Electrorefining

To purify copper by electrorefining, impure blister copper is cast to

form thick anodes for oxidation in aqueous copper sulfate/sulfuric acid

solutions. The cathodes are thin sheets of previously purified copper.

anode Cu(s, impure anode) Cu2+(aq) Eo = 0.34 V

cathode Cu2+(aq) Cu(s, pure cathode) Eo = 0.34 V

_________________________________________

overall Cu(s, impure anode) Cu(s, pure cathode)

Impurities are “left behind” when the aqueous Cu2+ ions produced by

oxidation at the anode are reduced and deposit on the cathode.

COPPER

Copper Electrorefining Cells

The cell design is very simple: vertical parallel plate electrodes in

open tanks lined with rubber or plastic. The electrolyte is an aqueous

solution containing about 5% dissolved CuSO4 + 15% H2SO4.

The applied cell voltage is very low ( Why?), only about 0.4 V

COPPER

Copper Electrorefining Cells

Ingenious electrochemistry!

at the anode Copper and less-noble metals (such as iron, zinc,

nickel) are oxidized to form aqueous metal ions.

Nobler metals (such as silver, gold, selenium

tellurium) do not dissolve and sink to the bottom of the

tank as sludge.

cathode Dissolved Cu2+ ions migrate through the electrolyte and

are reduced at the cathode as 99.9 % pure copper. Less

noble metal ions (Fe2+, Zn2+, Ni2+) remain in solution.

The sludge is used to improve the overall economics. How?

2.9 Other Industrial Electrochemical Industries

hydrometallurgical production of copper and zinc by

electrochemical extraction from ore leachate (eliminates high

temperature smelting and roasting steps, no SO2 produced)

chlorate (e.g., NaClO3) and perchlorate (e.g., NaClO4)

F2 from electrolysis of molten KF + HF

water electrolysis (source of ultrapure H2 and O2)

magnesium and sodium from electrolysis of molten chloride salts

organic electrosyntheses (e.g., adiponitrile for nylon manufacture)

electroplating

electromachining (e.g., “drilling” square holes!)

chem 471 part 2: industrial electrochemistry

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