Top 50 Industrial Chemicals (by weight) · Chem 471 Part 1: Primary Industrial Inorganic Chemicals 1.1 water, purification and treatment 1.2 methane (organic, but obtained from natural

1. sulfuric acid

2. nitrogen

3. oxygen

4. ethylene

5. calcium oxide (“lime”)

6. ammonia

7. phosphoric acid

8. sodium hydroxide

9. propylene

10. chlorine

11. sodium carbonate

12. methyl tert-butyl ether

13. ethylene dichloride

14. nitric acid

15. ammonium nitrate

16. benzene

17. urea

18. vinyl chloride

19. ethylbenzene

20. styrene

21. methanol

22. carbon dioxide

23. m-xylene

24. formaldehyde

25. terephthalic acid

26. benzene

27. hydrochloric acid

28. toluene

29. p-xylene

30. cumene

31. ammonium sulfate

32. ethylene glycol

33. acetic acid

34. phenol

35. propylene oxide

36. butadiene

37. carbon black

38. iso-butylene

39. potassium chloride

40. acrylonitrile

41. vinyl acetate

42. titanium dioxide

43. acetone

44. butyraldehyde

45. aluminum sulfate

46. sodium silicate

47. cyclohexane

48. adipic acid

49. nitrobenzene

50. bisphenol A

Top 50 Industrial Chemicals (by weight)_______________________________________________________________

_______________________________________________________________

31 are organic, but more inorganic chemicals are near the top of the list

(8 out of the top 10). Inorganics dominate by weight and are considered first.

Chem 471 Part 1: Primary Industrial Inorganic Chemicals

1.1 water, purification and treatment

1.2 methane (organic, but obtained from natural gas, main source of H2, CO, S)

1.3 hydrogen

1.4 carbon dioxide

1.5 oxygen major industrial gases

1.6 nitrogen

1.7 ammonia

1.8 nitric acid

1.9 sulfuric acid major industrial acids

1.10 phosphoric acid

1.11 calcium oxide (lime)

1.12 sodium carbonate limestone chemicals

1.13 concrete

1.14 iron and steel the industrial age begins

1.1 WATER ! Purification and Treatment

First up: water, the #1 industrial chemical used by weight

But wait!

Water is not listed in the “top 50” industrial chemicals.

• Why?

• most industrial water used to make other chemicals is

captive [used on site, not sold, shipped or taxed]

• large amounts of water are used to cool industrial

process equipment, with little or no chemical treatment

1.1.1 Water Purification by Distillation

• feed water is heated to the boiling point

• the water vapor produced is cooled and condensed

• non-volatile impurities (such as suspended particles,

dissolved salts, etc.) remain in the boiler

• provides high-purity water

• simple

• reliable

looks good, but …

Water Purification by Distillation

The diagram below shows how distillation is used to purify

water (and other liquids) in university labs.

Lab equipment is

unsuitable for

industrial water

purification

Why?

5 main reasons

Our first industrial

chemistry challenge.

boiler with a

220 V electric heater

in a quartz glass tube

• why 220 V ?

• why quartz glass ?distilled water

reservoir

condenser

cooled with

tap water


Lab water still in PSC 3020


Why is this equipment unsuitable

for industrial water purification ?

1. small scale, only a

few liters per hour

2. expensive electricity

is used to supply heat

3. batch operation

(continuous flow would

be more efficient − less

labor and downtime)

4. large amounts of

cooling water used

5. heat is lost down the

drain as warm tap water

Better: Multistage Flash Distillation (MSFD)

MSFD improves the economics of water distillation for

industrial production by using:

• natural gas or cogeneration* as the heat source

• large-scale production (up to 10,000 m3 per day)

• continuous flow operation

• heat recovery

hot water vapor pre-heats liquid water

for the next evaporation stage

cooling water not required

*such as waste heat from an electricity generating station

Multistage Flash Distillation (MSFD)

MSFD plant for the desalination of seawater at Jeleb Ali, Dubai

Seawater flowing into a series of distillation vessels at reduced

pressure “flashes” into vapor. Heat recovered from warm vapor

using heat exchangers pre-heats the seawater for evaporation.


A hot steam in B seawater in C pure water out D waste brine out

E steam out F heat exchanger G condensate collector H seawater heater

higher T,plower T,p

Ingenious Design: multistage distillation combined with

counter-current heat exchange/recovery

each stage: • an evaporation vessel

• a pure-water condensate collector

• a heat exchanger

Warm vapor at each stage preheats the seawater feed. Result:

seawater entering the heater is almost at the heater temperature,

reducing the heat input to about 90 MJ per m3 of product water.

10% to 20% of the water entering the plants is evaporated.

• Why multiple stages? If only one stage is used, the seawater

would only be heated about half way to the heater temperature.

• Why reduced pressure? Lowers the boiling point of water.


Typical plants produce 5,000 to 10,000 m3 drinking water per day,

providing about 60 % of global desalination capacity. MSFD

plants have larger capacity than reverse-osmosis plants. (up next)

• Why “flash” distillation? The vaporization of water under

reduced pressure is fast. Liquid converts to vapor “in a flash”.

• Why reduced pressure (no air)? The boiling temperature of

water is reduced. Because the pressure is determined by the

temperature at each stage, liquid water at each stage

spontaneously flows to the lower temperature stage, reducing

pumping costs. Also, no air is present to reduce the efficiency of

heat exchange between warm water vapor and cooler seawater.

Small amounts of acid (H2SO4 or H3PO4) are added to the

seawater feed. • Why?


Advantages: using any feedwater, even brine or

seawater, very pure water is produced on a

large scale for drinking or crop irrigation

using low-maintenance plants

Main Disadvantage: even with heat recovery using efficient heat

exchangers, about twice as much energy is

required relative to reverse osmosis (up next)

Multistage vacuum distillation is widely used:

• in dry areas where pure water is valuable

• where low-cost energy is available, such

as local natural gas or waste heat from

electricity generation (cogeneration)


Shell and Tube Heat Exchanger

Bundles of small metal tubes (large surface → faster heat transfer)

carrying seawater absorb heat in a shell filled with warm vapor.

Used for heat recovery in MSFD and many other applications.

cool

seawater

IN

warm

seawater

OUT

warm

vapor

IN

cool

vapor OUT

Heat / Exchangers

Used in many applications,

not just industrial chemistry.

Can you spot the heat exchangers?

(4 !)

Heat Exchange: an example of a Unit Operation

Unit operations: the basic steps used for industrial

chemistry processes

examples of unit operations:

• heat exchange

• evaporation

• liquefaction

• crystallization

• gas absorption

• adsorption

• distillation

• filtration

• crushing

• drying

• extraction

• screening and sieving

• mixing

• chemical reaction

1.1.2 Water Purification by Reverse Osmosis

Dissolved salts and other impurities can be removed by pumping

water through semipermeable membranes (usually acetylcellulose

or polyamide) that are permeable to water molecules but almost

impermeable to ions and other solutes.

At constant pressure, water spontaneously moves from the pure water

side to the solution side of the membrane. To reverse this flow to

purify water, the pressure applied to the solution must exceed the

osmotic pressure [ about 3.5 bar (50 psi) for seawater ].

For economical production rates, higher pressures (50 to 100 bar)

and thin membranes (0.1 mm) with large surface areas are used.

Uses about 50 % less energy than MSFD.


Thermodynamics of Osmosis from Chem 232, recall:

and Osmotic Pressure = g h CtotRT

= solution density

g = gravitational acceleration

h = osmotic height

Ctot = total concentration of dissolved

solute particles

R = gas constant

T = temperature


Pure water at pressure p is in equilibrium

with water in a solution at pressure p + .

To produce pure water by

reverse osmosis, a pressure

higher than p + must be

applied to the solution,

forcing water from the solution

into the pure water on the other

side of the membrane.


reverse osmosis membrane element

For increased production:

• thin membranes (< 0.1 mm)

• large surface area

Perfectly semipermeable membranes are impossible in practice, so

multistage reverse osmosis units are used: the purified water from

one fiber bundle is used as feed water for the next fiber bundle.

Careful pretreatment of the raw water is essential to remove

suspended material and microbes, or else the membranes will clog.


reverse osmosis system

• the flow rate of water

through each membrane

element is about

400 m3 per day

• about 10 % is recovered as

purified water

• inlet pressure 40 to 70 bar

(the osmotic pressure of

seawater is 3.50 bar)


main advantage over MSFD:

• 50 % lower energy costs (water evaporation not required)

main disadvantages:

• higher installation and maintenance costs (filter element

monitoring and replacement) per tonne of water

• more extensive water pretreatment required to reduce

membrane clogging

• higher impurity levels than obtained by MSFD (the membranes are

not completely impermeable to Na+, Cl− and other small ions or molecules)

1.1.3 Production of Municipal Water for Cities and Towns

Question: What’s the easiest way to make a million dollars?

Answer: Start with about $950,000.

By analogy, what’s the easiest way to produce pure water?

Start with “almost pure” water from lakes, rivers or drilled wells.

Using this approach, about 1 km3 (1 billion tons) of municipal

water is produced per day in North America. A huge industry.

Halifax Water, for example, has about 450 employees.

The Brierly Brook Water Treatment Plant (8,000 m3 per day)

serves Antigonish.

Municipal water plants are very highly regulated. • Why?


Depending on the quality of the raw feedwater, some or all of the

following steps are used to produce safe, drinking-quality water:

• removal of coarse suspended material using screens or filters

• initial chlorination

• flocculation and sedimentation

• filtration through beds of sand

• treatment with activated carbon to adsorb impurities (optional)

• safety chlorination for treated water piped to consumers

• pH adjustment


Water Treatment Flow Diagram


Initial Chlorination

For strongly polluted water, the first treatment step is the addition

of sufficient chlorine to give 0.1 to 0.5 ppm free Cl2.

Advantage: Chlorine is a very strong oxidizing agent that

kills bacteria and viruses.

Disadvantage: Chlorination of dissolved organics (e.g., humic acid)

produces chemicals that taste bad and may be

carcinogenic.

Ozonization, an alternative to chlorination, avoids these problems, but is more

expensive. Treatment with high-intensity ultraviolet light is another option,

but used less frequently.


Flocculation and Sedimentation

If the raw water is highly turbid due to suspended particles and

colloids, aluminum sulfate and/or iron(III) sulfate is added

(typically 0.02 to 0.05 g/dm3), forming very small particles of

hydroxide precipitates with huge surface areas.

Al2(SO4)3 + 6H2O → 2Al(OH)3 (ppt) + 3H2SO4

Fe2(SO4)3 + 6H2O → 2Fe(OH)3 (ppt) + 3H2SO4

Suspended material adsorbed on the “sticky” gelatinous hydroxide

particles settles out as sludge or is removed by flotation.

Filtration can also remove suspended solids, but for highly turbid water

the frequent back-flushing needed to de-clog the filters is not economical.


Filtration

After flocculation, the water is filtered through a 1- to 2-meter thick

beds of sand. When the sand becomes covered with impurities,

it is cleaned by back-flushing. • What is used for back-flushing?

Activated Charcoal Treatment

Dissolved organic impurities (such as humic acid) and some metal

metal ions (such as iron) are removed by adsorption on activated

carbon. Carbon powder can be added during the flocculation step,

or the water can flow through a bed of granular charcoal, which

can be regenerated by heating to drive off adsorbed molecules.

Safety Chlorination

The final step in the production of drinking-quality water is the

addition of 0.1 to 0.2 ppm chlorine to prevent microbial reinfection

of the water during distribution or storage.

1.1.4 Water Softening

Untreated water (especially well water) often contains significant

amounts of calcium ions and bicarbonate ions.

Water contaminated with calcium ions and bicarbonate ions

produces a “hard” scale of precipitated calcium carbonate

Ca2+ + 2HCO3- → CaCO3 (ppt) + CO2 + H2O

Ca2+ + CO32- → CaCO3 (ppt)

Calcium carbonate scale deposits can partially block pipes and reduce

heat transfer in boilers and heat exchangers.

Calcium ions also interfere with the cleaning action of soaps by

forming insoluble salts (e.g., unsightly ring around the tub).


In some industrial applications (such as MSFD) scale formation

is reduced by adding an acid, usually H2SO4.

Added acid forms the more soluble calcium sulfate salt.

Adding too much acid can lead to corrosion problems.

Calcium hydroxide [Ca(OH)2 slaked lime] can be added to soften

municipal water supplies by precipitating calcium carbonate

Ca(HCO3)2 + Ca(OH)2 → 2CaCO3 (ppt) + 2H2O

during the flocculation or filtration steps.


Alternatively, Ca2+ and Mg2+ ions (and other metal ions) can be

removed by pumping the water through beds of cation-change resin

beads. Sodium polystyrene (PS) sulfonate is frequently used:

PS-SO3Na + 0.5 Ca2+(aq) → (PS-SO3Ca0.5) + Na+(aq)

When the cation exchanger is fully loaded with Ca2+ (or Mg2+) ions,

it is regenerated by reversing the above reaction by flushing the

ion exchanger with concentrated (5 to 10 %) aqueous NaCl.

PS-SO3Ca0.5 + Na+(aq) → PS-SO3Na + 0.5 Ca2+(aq)

1.1.5 Deionized Water

Ion-free water for steam turbines, pharmaceuticals and scientific

applications is prepared by using a cation exchanger (in H+ form)

and an anion exchanger (in the OH − form) in series:

Na+(aq) + Cl-(aq) + PS-SO3H + PS-N(CH3)3OH

(cation exchanger) (anion exchanger)

→ H+(aq) + OH -(aq) + PS-SO3Na + PS-N(CH3)4Cl

Unwanted ions (such as Na+ and Cl −) are exchanged for H+ and OH −,

which react form water

H+(aq) + OH -(aq) → H2O

• But ion exchange is unsuitable for the desalination of seawater. Why?

1.1.5 Deionized Water

Mixed Bed Water

Deionizing Cartridge

conductivity meter to

monitor purity of the

deionized water

quaternary ammonium

anion exchange beads

sulfonate cation

exchange beads

reverse osmosis

water inlet

deionized water outlet

drain

The operational life of ion

exchangers used in the

Chemistry Department is

extended by using feedwater

purified by reverse osmosis.

Next up: 1.2 METHANE (CH4)

Methane is the #2 industrial chemical by weight.

Natural gas is the main source of methane.

Methane, like water, is not included in many lists of industrial

chemicals, for similar reasons:

• most of the methane used to make other industrial

chemicals is consumed on site, not sold or shipped

• methane is obtained from natural gas, with only minor

chemical processing

• about 75 % of total methane is used for fuel

Based on cost, pollution and CO2 emission considerations, natural gas

is the preferred industrial fuel. The good, the bad and the ugly fossil

fuels are respectively natural gas, petroleum and coal.

Natural Gas

The composition of natural gas is highly variable.

“Typical” composition: 85 % methane

9 % ethane

3 % propane

1 % butanes

1 % nitrogen

Other constituents: heavier hydrocarbons (“condensate”), water,

hydrogen sulfide, carbon dioxide, helium

Natural gas is found in porous rock reservoirs, either associated

with crude oil deposits [ called associated natural gas, long

considered a waste product and was flared (burned) at oil wells for

safety reasons ] or in deposits without crude oil present [called

non-associated natural gas ].

Dry natural gas contains only small amounts (a few per cent) of

ethane, propane, butanes and heavier hydrocarbons that liquefy with

compression at room temperature.

Wet natural gas contains significant amounts of ethane, propane,

butanes and heavier hydrocarbons. (“Wet” does not refer to water

content in this context, but to liquid formation under compression.)

Sour natural gas contains significant amounts of hydrogen sulfide,

up to 35% H2S (!) for some deposits in Southern Alberta.

Sweet natural gas contains only trace amounts of hydrogen sulfide.

(Note: “sweet” does not refer to taste in this context.)

• Why are sour gas leaks very dangerous?

• Why is sweet gas much preferred over sour gas?

Production of Methane from Crude Natural Gas

Four main steps:

1. dehydration (water removal) by bubbling the gas through a

liquid glycol in which water (but not methane) is very soluble

2. chemical absorption to remove H2S and CO2 acid gases using

aqueous amine (such as mono- or diethanolamine) solutions

H2S(g) + HO-(CH2)2-NH2(aq) → HS-(aq) + HO-(CH2)2-N+H3(aq)

CO2(g) + H2O(l) + HO-(CH2)2-NH2(aq) → CO32-(aq) + O-(CH2)2-N

+H3(aq)

3. removal of higher-boiling gases (ethane, propane, butanes, …)

by dissolution (physical absorption) in a liquid, such as hexane

4. removal of mercury and other heavy metals using activated

carbon or zeolite molecular sieves (physical adsorption)

Natural Gas Processingprovides chemical feedstocks and

a pollution-free clean-burning fuel

The glycol absorbent is regenerated by heating to vaporize the dissolved water.

The amine absorbent is regenerated by heating to drive off H2S, from which sulfur

is obtained by the Claus process (details later). Fractional distillation of the oil

absorbent yields valuable ethane, propane and butane “natural gas liquids” (NGLs).

Scrubbing Towers

Water, hydrogen sulfide

and natural gas liquids

are removed (“scrubbed

out”) from raw natural

gas in towers using

counter-current flows

of liquid glycol, aqueous

amine, or hexane,

respectively.

The towers are filled with

packed beds of ceramic

spheres or rings. • Why?

another unit operation

Located in southwest Alberta, near

Pincher Creek. Operating since 1962.

products

• natural gas (3 million m3 per day)

• ethane (petrochemical feedstock)

• propane (fuel)

• butane (petroleum refining)

• condensate (gasoline)

• sulfur (shipped to sulfuric acid plants)

Shell Waterton Gas Complex

• Associated natural gas always wet. Why?

• Why is it important to sweeten sour natural gas?

• No olefins [e.g., ethylene (H2C=CH2)] are in natural gas. Why?

• Water, but not methane, is very soluble in glycol absorbents. Why?

• Ethane, propane, and butanes are much more soluble in liquid

hexane absorbents than methane. Why?

• NGLs (acronym for Natural Gas Liquids such as ethane and propane) are

valuable natural gas byproducts. Why valuable?

• Chemical absorption (e.g., absorption of H2S gas by aqueous amine solutions)

more efficient than physical absorption (e.g., the absorption of propane by

liquid hexane). Why?

• Describe a chemical gas absorption process that keeps us alive.

• After extensive purification, small amounts of ethanethiol (CH3CH2SH) and

other organic sulfur compounds are often added to natural gas! Why?

1.3 HYDROGEN

Hydrogen: • #1 element on the chemistry periodic table

• H is the most abundant element in the universe

• H2 is most abundant molecule in the universe

But no free hydrogen exists on Earth. • Why?

Many scientists, including chemists, believe hydrogen is produced

industrially by the electrolysis of water:

2H2O(l) = 2H2(g) + O2(g)

Wrong! Electrolysis is too expensive for large-scale hydrogen

production due to the high cost of electricity.

Industrial production of hydrogen (about 50 million tonnes/year):

Steam Reforming and Shift Conversion of Hydrocarbons

Methane (from natural gas) is the cheapest source of hydrogen.

A two-step process is used:

(synthesis gas)

steam reforming CH4(g) + H2O(g) → CO(g) + 3H2(g)

water gas shift CO(g) + H2O(g) → CO2(g) + H2(g)

(WGS) conversion

___________________________________

overall CH4(g) + 3H2O(g) → CO2(g) + 4H2(g)

4 moles of H2 are produced per mole of CH4 (2 moles of H2 from

methane and another 2 moles from the steam).

Ethane, propane, butane and heavier alkanes (“naphtha”) are also

used to produce hydrogen, but at higher cost.

(synthesis gas)

steam reforming CnH2n+2 + nH2O → nCO + (2n+1)H2

shift conversion nCO + nH2O → nCO2 + nH2

________________________________

overall CnH2n+2 + 2nH2O → nCO2 + (3n+1)H2

__________

In the petrochemical industry, reforming refers to a process that changes the structure of

hydrocarbons to modify their properties.

• Why is nCO + nH2O → nCO2 + nH2 called a water gas shift (WGS) reaction?

• Why is it cheaper to make H2 from methane than from ethane, propane, butanes, …?

essential step before reforming: Desulfurization

Reforming reactions use nickel oxide catalysts supported on Al2O3

(alumina) or MgO-Al2O3. These catalysts are susceptible to

poisoning by sulfur compounds, so the natural gas or other feed

must be thoroughly desulfurized (sweetened) to < 10 ppb sulfur.

This is usually done by adding hydrogen to the untreated feed,

heating the mixture to 350 to 450 oC, and reducing the S-containing

compounds to H2S (e.g., R1-S-R2 + 2H2 = R1-H + R2-H + H2S)

using cobalt oxide or mixed nickel-molybdenum catalysts.

The H2S is removed by adsorption on zinc oxide or scrubbed out

using an alkaline absorbent, such as aqueous solutions of organic

amines or sodium hydroxide.

H2S recovered from desulfurization as a byproduct is used to make

sulfur and sulfuric acid, the most important industrial acid.

Steam Reforming CnH2n+2 + nH2O → nCO + (2n+1)H2(reduction of water using hydrocarbons)

After desulfurization of the methane or other hydrocarbons, steam

is added and the mixture is preheated to about 500 oC by

combustion of natural gas or another fuel.

The reforming reactions take place at about 1200 oC in steel tubes

packed with the NiO catalyst. Despite the relatively high

temperatures, a catalyst is still required owing to the high stability

of methane molecules.

The hot H2 + CO gases exiting the reformer tubes and the flue

gases from the fuel combustion are cooled using heat exchangers

for the co-generation of steam._________________________________

Challenges The high temperatures of reforming reactions requires special materials for

reactor construction and significant amounts of fuel for heating. Sulfur can never be

completely eliminated from the feed, so the catalysts are eventually poisoned. Coking

[carbon deposition, e.g. CH4(g) = C(s) + 2H2(g) ] also limits catalyst life.

(synthesis gas)

Steam Reforming CnH2n+2 + nH2O → nCO + (2n+1)H2

?

?

The name synthesis gas for the

CO + H2 mixtures comes from its

use as a synthetic fuel gas and as a

chemical intermediate in the

synthesis of other chemicals,

mostly ammonia and methanol.

WGS Conversion Reactions CO + H2O → CO2 + H2

More steam is added to the H2 + CO mixture and the CO is

converted to CO2 using metal oxide catalysts. Fe2O3/Cr2O3

catalysts are used for high-temperature (HT) conversion (350 to

380 oC). If the sulfur content is very low (< 0.1 ppm), lower

temperature (LT) conversion is feasible on CuO/ZnO catalysts.

Removal of Carbon Dioxide Physical or chemical absorption is

used to remove CO2 from H2 + CO2 mixtures produced in shift

reactors. H2, a small nonpolar molecule, is relatively insoluble in

organic solvents. Physical absorption of CO2 can therefore be

accomplished by using organic solvents (e.g., methanol or glycols)

to preferentially dissolve the CO2. For chemical absorption of CO2,

an acid gas, alkaline solutions of amines or K2CO3 are be used. _________________________________

When the absorbents become saturated with CO2, they can be regenerated (• how?).

Pure CO2 can be recovered and sold as a byproduct.

H2 Production from Steam Reforming and Shift Conversion of Hydrocarbons

(CO, CO2, and unreacted H2O, CH4 removed by methanol and liquid N2 scrubbers)

Uses of Industrial Hydrogen

• ammonia production (about 45 %)

• petroleum refining processes (about 40 %)

hydrocracking of heavy oil to produce lighter product (“upgrading”)

e.g. R-CH2-CH2-R′ + H2 → R-CH3 + CH3-R′

hydrotreating to remove sulfur and nitrogen (as H2S and NH3)

e.g. R-CH2-NH2 + H2 → R-CH3 + NH3

• methanol production from syngas (about 10 %)

CO + 2H2 → CH3OH

• miscellaneous (oil and fat hydrogenation, metal refining, …)

The “Hydrogen Economy”

Burning fossil fuels (natural gas, petroleum, coal) produces

pollutants, such as sulfur and nitrogen oxides and mercury. Also

carbon dioxide. Rising atmospheric CO2 levels, combined with the

greenhouse effect, cause global warming. These problems could

be overcome by switching from fossil fuels to hydrogen:

H2 + ½ O2 → H2O

• What are the advantages/disadvantages of H2 as a fuel?

• If H2 replaces fossil fuels, where will the H2 come from?

• Prototype motor vehicles using electric motors powered by

electrochemical hydrogen fuel cells have been produced. The

fuel-cell reaction is H2 + ½ O2 → H2O, so no CO2 is produced

(“carbon neutral”) and no global warming. Is this correct?

Syngas (carbon monoxide + hydrogen) from Coal

For many years, syngas was produced mainly from coal, a complex

mixture of aromatics consisting of C, H, O, N, S and smaller

amounts of metals and other elements. For simplicity, using carbon

to represent coal, the main coal gasification reaction is

C(s) + H2O(g) → CO(g) + H2(g)

Steam is passed over white-hot coal heated (• how?) to 1000 oC.

The CO and H2 produced can be used to make chemicals (mainly

NH3 and CH3OH) or as a fuel gas, which is easier to transport

(• why?) and burns much more cleanly (• why?) than coal.

Coal gasification dwindled in the 1950s as cheaper petroleum and

natural gas became widely available for reforming. In the future,

when petroleum and natural gas become scarce, vast deposits of

coal will remain. Coal gasification might regain importance as a

major source of energy and essential industrial chemicals.

Alternative energy sources exist, such as nuclear reactors,

Coal in the 21st Century? Yes! And Industrially Important

Coal is dirty, but cheap and widely available. Currently provides

25 % of the primary energy and 40 % of the electricity generated

worldwide. “Clean” ways to use coal are actively investigated in

many research and development projects.

Coal and Nova Scotia

Approximately 60 % of the electricity generated in Nova Scotia is

produced by burning imported coal to run steam turbines.

Lingan Generating Station on Cape Breton

4,000 tons of coal

burned per day

produces 25 % of Nova

Scotia’s electricity

• Why built way up on

Cape Breton’s east coast?

• Is there something “fishy”

about this photo?

Electricity Generation in Nova Scotia

_________________________________________

Year 2007 2018

__________________________________________

coal 78 % 60 %

oil, natural gas 14 % 14 %

wind 1 % 17 %

hydro 7 % 7 %

biomass 1 % 2 %

__________________________________________

1.4 CARBON DIOXIDE

#22 on the top 50 list.

Most industrial CO2 is produced by steam reforming and shift

conversion of methane or light hydrocarbons:

CnH2n+2 + 2n H2O → n CO2 + (3n+1) H2

CO2 is removed from the gas mixture by physical or chemical

absorption (already described). Heating the absorbents releases

CO2 and regenerates the absorbents.

Smaller amounts of CO2 (< 10 % total production) are obtained from

fermentation, coal gasification, and lime (CaO) manufacturing:

C6H12O6 (and other sugars) → 2C2H5O6 + 2CO2

C + 2H2O → 2H2 + CO2

CaCO3 → CaO + CO2

Uses of Industrial Carbon Dioxide

Solid CO2 (dry ice) is an important food-industry refrigerant. It

sublimes (turns to vapor) at −78 oC at atmospheric pressure, rather

than melting to form a liquid which could damage food products.

Beverage carbonation is another food-industry application.

In recent years, CO2 is used increasingly as the “working fluid”

(refrigerant) is air conditioners and refrigerators by replacing

chlorofluorocarbons. (• Why being replaced?)

CO2 competes with nitrogen in enhanced petroleum recovery. In

this application, CO2 gas is pumped down oil wells together with

water and surfactants to force oil out of porous reservoir rocks.

CO2 is also used to make chemicals. For example, the synthesis of

urea (an important fertilizer and organic intermediate) with the

overall reaction CO2 + 2NH3 → NH2CONH2 + H2O.

1.5 NITROGEN and 1.6 OXYGEN

Nitrogen and oxygen are discussed together because they are produced

together on a large scale by the distillation of liquid air.

N2 #2 in the “top 50”

O2 #3

Major Uses

N2 • NH3 synthesis (subsequent HNO3 and fertilizer production)

• inert atmospheres to prevent fire and other oxidation processes

• advanced petroleum recovery

O2 • steel making (O2 + 2C → 2CO; Fe2O3 + 3CO → 2Fe + 3CO2)*

• production of organic chemicals by hydrocarbon oxidation

• oxy-acetylene welding and cutting torches* • Why not use air instead of O2?

NITROGEN and OXYGEN

Two main processes are used for air liquefaction:

a) Linde Process Compressed air is cooled by adiabatic

(heat q = 0) Joule-Thomson expansion in a throttle valve.

b) Claude Process (more efficient • why?) Compressed air

expands adiabatically in a turbine engine, loses energy by doing

useful work, and cools.

The process air is first filtered to remove particulate matter, then

compressed to about 50 bar. A platinum catalyst to convert traces

of hydrocarbons (such as compressor lubricant) to water and

carbon dioxide, then H2O and CO2 are removed using molecular sieves.

The purified air is cooled in a heat exchanger and then using a JT

throttle valve or a Claude expansion turbine.

Liquefaction of Air

From thermodynamics:

Expansion Through a Throttle Valve

Joule-Thomson coefficient (T/p)H > 0, expanding air cools

Adiabatic (heat q = 0) Expansion in a Claude Turbine Engine

First Law U = q + w = w < 0

Expanding air does work on the engine (w < 0), loses energy,

and cools

Nitrogen (normal boiling point 77 K) is more volatile than

oxygen (normal boiling point 90 K) and is collected as vapor

at the top of a distillation column.

Liquid oxygen is collected at the bottom end.

Cold nitrogen and oxygen leaving the column are used to precool

the incoming air in an efficient heat exchanger.

Argon (and smaller amounts of neon, krypton and xenon) can be

obtained by further distillation of the liquid oxygen.

Liquefaction of Air

Oxygen and Nitrogen from the Fractional Distillation of Liquid Air

• Before liquefaction, why is it important to remove traces of hydrocarbons, H2O and CO2?

• Why are special pumps, valves and fittings needed to handle oxygen?

• What are molecular sieves?

Oxygen and Nitrogen from the Fractional Distillation of Liquid Air

O2/N2/Ar plant near Leipzig, Germany

11 2

3

2

33

3

1st distillation tower for

O2 and N2 separation

2nd distillation tower for

O2 and Ar separation

molecular sieve units

for removing H2O, CO2

1.7 AMMONIA

Global production: 150 million tonnes/year

Major use: agricultural fertilizers (about 85% of total production),

ammonia and ammonium nitrate, phosphate and sulfate salts.

Ammonia is also a raw material for the synthesis of nitric acid, urea,

plastics and many other nitrogen-containing compounds.

Historical Background

For many years, nitrogen-containing fertilizers were obtained from natural sources,

such as saltpeter (KNO3), “Chile niter” (NaNO3) and guano (sea-bird poop!)

from South America. Growing world population made these naturally-occurring

supplies inadequate for agricultural needs in the early 1900's. World War I naval

blockades reduced overseas shipments and vastly increased demand for

nitrogen-containing explosives. The search was on for another source of ammonia.

Direct Synthesis of Ammonia from Nitrogen and Hydrogen ?

An obvious possibility: Use cheap N2 from the distillation of liquid air

and H2 from steam reforming (historically) coal to make ammonia.

Big problem: Unfortunately, the triply-bonded N2 molecule is very

stable and inert (difficult to react).

From lab experiments, it was known that N2 + H2 + NH3 mixtures

contain hardly any NH3 at atmospheric pressure.

Even with a favorable 1:3 N2:H2 ratio, the equilibrium NH3 mole

fraction is only 0.0001 at 1300 K (hot enough for fast reaction rates).

Many experts concluded that the production of ammonia using

atmospheric nitrogen was economically impossible.


Other experts were less pessimistic, based on the considerations:

a) The ammonia synthesis reaction

N2 + 3H2 → 2NH3 Ho = -91.4 kJ at 25o C

is strongly exothermic and therefore more favorable at lower

temperatures. Try low temperatures.

b) Four moles of reactant gases (N2 + 3H2) are converted to two

moles of product gas (2NH3). Due to the significant (50 %) reduction

in volume, ammonia synthesis is also favored at higher pressures.

Try high pressures.


c) Further guidance from the relatively new science of chemical

thermodynamics indicated that the yield the ammonia could be

substantially increased in practice by removing ammonia product

and recycling unreacted hydrogen and nitrogen through the reactor.

This discovery had major implications for the efficient production of

ammonia and many other industrial chemicals.

d) At suitable temperatures and pressures, and by using recycling,

ammonia synthesis from nitrogen and hydrogen might be practical,

especially if suitable catalysts could be found to speed up the reaction.

Over 6000 possible catalysts were tested, mostly by trial and error.

Iron-oxide catalysts (magnetite (Fe2O3) with additives such as CaO

and K2CO3) were found to be effective and cheap.


e) Using heat exchangers, the heat released by the exothermic

ammonia synthesis reaction could be used to warm the N2 and H2

reactants entering the catalytic reactor to the operating temperature,

improving the process economics.

Haber Process

Combining these ideas (low temperatures, high pressure, product

removal, reactant recycle, catalysts, and heat exchange), led to the

development (1908 to 1913) of Haber process for the economical

and large-scale production of synthetic ammonia.

A triumph of industrial chemistry.

The Haber process is still used today, with only minor changes.

Haber Ammonia Synthesis Loop

Lots going on in this diagram!

Modern Ammonia Synthesis Plants

Haber Process The iron oxide-catalyzed reaction of N2 and H2.

Reactant Feed Ratio The 1:3 ratio of N2:H2 is obtained by mixing

purified N2 (from the distillation of liquid air) and H2 (usually

from the steam reforming and shift conversion of natural gas).

Low Temperature vs. High Pressure Optimization Favorable

thermodynamic conditions for NH3 production are low temperatures

and high pressures. But as the temperature is lowered, the reaction

rate becomes too slow, even with catalysts. But as the pressure

is increased, equipment costs for the additional compression and

stronger reactor vessels becomes prohibitive. As a comprimise, most

plants operate with reactant gases entering the catalyst bed at 400 oC

and 250 bar, leaving the reactor at 450 to 500 oC.

Modern Ammonia Synthesis Plants

Product Removal and Reactant Recycle The conversion of

N2 + 3H2 to 2NH3 per pass through the reactor is only 15 to 20 %.

To increase overall yield, NH3 is removed by condensation (< 0 oC)

from gas leaving the reactor. Unreacted N2 and H2 are compressed,

mixed with fresh make-up N2 and H2, and returned to the reactor.

Heat Recovery To save fuel costs, the reactants are preheated by

using a heat exchanger and the hot gases leaving the reactor.

Scrubbing Impurities CO, CO2, H2O and H2S impurities, which

can “poison” (deactivate) the catalysts, are removed by using a

condenser between the make-up gas supply and the reactor to

condense liquid ammonia which absorbs (scrubs out) impurities,

while letting the nitrogen and hydrogen pass.

Thermodynamics of Ammonia Synthesis

The equilibrium constant K for the ammonia synthesis reaction

N2(g) + 3H2(g) = 2NH3(g)

is given by the expression K = pNH32/pN2 pH2

3 = exp(-Go/RT).

pN2 , pH2 and pNH3 are the partial pressures of the gases and

Go is the standard Gibbs energy change for the reaction. ________________________________________________________________________________________________________________________________________________________________

t / oC T / K Go / kJ mol−1 K________________________________________________________________________________________________________________________________________________________________

25 298 -32.7 545,000

227 500 9.60 0.0992

477 750 66.7 0.0000225

727 1000 123.8 0.00000034_________________________________________________________________________________________________________________________________________________________________

Notice the huge (exponential) increase in K at low temperatures.


Starting with 1 mole of nitrogen and 3 moles of hydrogen (in the correct

stoichiometric ratio for ammonia synthesis) and letting x denote the extent of

reaction (the fraction of the initial nitrogen and hydrogen that reacts to form

ammonia) at equilibrium

N2(g) + 3H2(g) = 2NH3(g)

initial no. of moles: 1 3 0

initial no. of moles: 1 - x 3 - 3x 2x

mole fraction yi: x

x

24

1

−

−

x

x

24

33

−

−

x

x

24

2

−

The mole fraction the gases

yN2 = (1 - x)/(4 - 2x) yH2 = (3 - 3x)/(4 - 2x) yNH3 = 2x/(4 - 2x)

are obtained by dividing the number of moles of each gas by total number of

moles: 1 - x + 3 - 3x + 2x = 4 - 2x.


Using the fact that the partial pressure of a gas is its mole fraction times the total

pressure p = pN3 + pH2 + pNH3 gives

24

22

3

H2N2

2

NH3

3

H2N2

2

NH3 1

)1(

)2(

27

16

))((

)(

px

xx

pypy

py

pp

pK

−

−===

Note that x is zero (no ammonia synthesis) in the limit p 0.

Conversely, x → 1 (complete conversion to NH3) in the high-pressure limit p → 4.

To illustrate that effect of temperature and pressure on the yield of ammonia, solving for

x under different conditions gives

T = 500 K p = 1 bar x = 0.157

T = 750 K p = 1 bar x = 0.00307

T = 750 K p = 100 bar x = 0.213


High-pressure reactors for

ammonia synthesis required

the development of strong

corrosion-resistant steels and

special techniques for making

large thick-walled vessels.

Ammonia/Urea Fertilizer Plant near Belle Plaine, Saskatchewan

Converts natural gas, air and water to 1800 tonnes of ammonia

and 2800 tonnes of urea (2NH3 + CO2 = H2N-CO-NH2 +H2O) per day.

• Where does the CO2 required for urea production come from?

• Why is this plant better for Saskatchewan than importing fertilizer? (Win-win-win)

• Synthetic fertilizers (ammonia, ammonium nitrate, urea, etc.)

derived from the Haber process boost crop yields by up to 500 %.

• As a result of higher food production, global population has

increased from 1.6 billion in 1900 to 7.6 billion in 2018.

• 500 million tonnes of synthetic fertilizers are produced per year,

essential for modern agriculture.

• > 50 % of the nitrogen in humans is from Haber plants !

• Crops such as beans and clover produce nitrates and other fertilizers

naturally by “fixing” atmospheric nitrogen at ambient temperature

and pressure. • So why do we need the Haber process?

• Why were advances in steelmaking needed for the Haber process?

Haber Process: Important Considerations

The high-pressure ammonia

synthesis reactor used by

by Haber and his team more

than a century ago. On display

at the Karlsruhe Institute of

Technology, Germany.

• Ugly pieces of old steel?

or

• A monument to one of the

most important advances in

in the history of civilization?

8. NITRIC ACID

No. #14 in the “top 50” list. Applications:

• 60% for fertilizers (mostly ammonium nitrate)

• 15% for explosives (mostly ANFO, ammonium nitrate/fuel oil)

• 15% for adipic acid, precursor for nylon and other plastics

Historically, nitric acid was made from saltpeter and sulfuric acid.

NaNO3 + H2SO4 → HNO3 + NaHSO4

Today, nitric acid is produced by the 3-step Ostwald process for the

oxidation of ammonia. The overall reaction is deceptively simple:

NH3 + 2 O2 → HNO3 + H2O (strongly exothermic)

But in practice, HNO3 synthesis is tricky because the oxidation

2NH3 + 3/2 O2 → N2 + 3H2O is thermodynamically more favorable

due to the stability of triply-bonded N2 molecules.

Three-step HNO3 synthesis: • air oxidation of NH3 to NO

Ostwald Process • air oxidation of NO to NO2

(in production since 1906) • hydration of NO2 to HNO3

Step 1 Catalytic oxidation of NH3 to NO at 850 to 950 oC using

oxygen from air. The catalyst is a red-hot platinum alloy

fine-mesh gauze containing 5% rhodium

4NH3 + 5O2 → 4NO + 6H2O (exothermic)

Without the catalyst, there is a strong tendency to produce unwanted

nitrogen (4NH3 + 3O2 → 2N2 + 6H2O). The role of the catalyst

is to select the formation of the desired NO product.

The hot gas leaving the catalytic reactor is used to generate steam for

other applications to improve process economics.

Step 1 Catalytic oxidation (combustion) of NH3 to NO

• the critical step in nitric acid production

• the platinum catalyst (alloyed with rhodium for strength) gives

up to 99 % selective conversion of NH3 to the desired NO product

• a triumph of catalysis!

• one of the most efficient industrial chemistry catalytic processes

• problem: loss of the precious ($$$) metal catalyst, eroded from the

gauze as PtO2

• solution: use a downstream Au-Pd gauze or bed of marble chips

to adsorb the lost Pt

• 80% of the lost Pt can be recovered

Step 2 Air oxidation of NO to NO2 NO from the first step is

reacted with additional oxygen from air to give NO2.

Catalysts not required.

2NO + O2 → 2NO2

Some of the NO2 dimerizes: 2NO2 = N2O4.

Step 3 Absorption and hydration of NO2 in liquid water

3NO2 + H2O → 2HNO3 + NO

Aqueous HNO3 is produced together with some NO reactant from the

previous step. Because the oxidation of NO occurs in the gas phase,

mass transfer of NO from the aqueous solution to the gas phase plays a

major role in the overall reaction rate for Step 3.

Ostwald Process Flow Sheet for Nitric Acid Production

SCR: Selective Catalytic Reduction of NO and NO2 to H2O and N2 (for pollution control)

bleacher: hydrogen peroxide used to increase product purity (2NO2 + H2O2 = 2HNO3)

Chemical Reactor Design

• temperature ? • pressure ?

• size and shape? • flow rates (residence times) ?

• heat exchange ? • catalyst ?

• transport processes ? • reaction rates … ?

Weird Kinetics of the Gas-Phase Oxidation of NO

3rd-Order Reaction Mechanism? The reaction 2NO + O2 → 2NO2 is

famous because it appears to be “third-order” with observed rate law

suggesting a simultaneous NO + NO + O2 tri-molecular collision.

But in the gas phase, three-body collisions are non-existent or at least

extremely improbable.

Negative Activation Energy? Rates of chemical reactions almost

always increase with temperature. The oxidation of NO is highly

unusual because it slows down as the temperature increases, suggesting

the rate constant k for the reaction 2NO + O2 → 2NO2 decreases with

temperature: dk/dT < 0.

The Arrhenius law k = k0exp(-Eact/RT) for the temperature dependence

of rate constants gives a negative activation energy?

What’s going on? Three-body collisions? Negative Eact?

Actual Mechanism of the 2NO + O2 → 2NO2 Reaction

Using only bimolecular reaction steps:kf (fast)

Step I NO + O2 ↔ NO3 ( at equilibrium)kb (fast)

k′ (slow)

Step II NO3 + NO → 2NO2 (slow, rate-determining)

Actual Mechanism of the 2NO + O2 → 2NO2 Reaction

Step I: NO + O2 ↔ NO3

fast forward reaction and fast back reaction

forward reaction rate = backward reaction rate

kf [NO] [O2] = kb [NO3]

equilibrium constant Keq

Note: The reaction NO + O2 → NO3 is exothermic (gives off heat)

so Keq decreases as the temperature is raised.

eq

2

3

]O][NO[

]NO[K

k

k

b

f==

Actual Mechanism of the 2NO + O2 → 2NO2 Reactionk

Step II: NO3 + NO → 2NO2

but [NO3] = Keq[NO][O2]

so

and apparent rate constant k = kKeq

Significance What looks like a tri-molecular collision mechanism

(improbable) is actually a two-step mechanism involving bimolecular

collisions. Apparent rate constant k decreases with temperature because

of the decrease in equilibrium constant constant Keq with T.

]NO[]NO[d

]NO[d3

2 kt

=

]NO[NO][]NO[d

]NO[deq

2 Kkt

=

]NO[]NO[d

]NO[d 22 kt

=

1.9 SULFURIC ACID

(Water, methane, oxygen and nitrogen are used in larger

amounts, but are purified before use, not synthesized.)

• sulfuric acid is the go-to industrial acid.

• cheap (about $250 / tonne, 25 cents per kilogram)

• global production about 200 million tonnes per year

Main uses of sulfuric acid:

• 70 % for manufacturing fertilizers

• 9 % for mining (acid leaching of metal ores)

• 6 % for petroleum alkylation

• 5 % for the making inorganic chemicals (mostly sulfates)

sulfuric acid is the #1 industrial chemical

by weight produced by chemical synthesis

Industrial Production of Sulfuric Acid

Sulfuric acid is made from sulfur, air and water.

First, elemental sulfur is oxidized to form sulfur dioxide. Further

oxidation gives sulfur trioxide, which is hydrated to give sulfuric acid.

oxidize oxidize hydrate

S → SO2 → SO3 → H2SO4

Sulfur – the Starting Material for Sulfuric Acid

Frasch Mining Sulfur has been produced for many years by

pumping hot water, steam and compressed air into wells drilled into

geological sulfur deposits. Molten sulfur forced to the surface is

collected and dried. Drilling is expensive, and every tonne of Frasch

sulfur requires twenty to thirty tonnes of process water.

Sulfur – the Starting Material for Sulfuric Acid

Claus Process Sulfur

Sulfur is now produced mainly by the Claus process for the oxidation of H2S

available as a byproduct from desulfurization of natural gas.

Thermal Step: Hydrogen sulfide is burned in a non-catalytic combustion

chamber using atmospheric oxygen. The hot gas produced is cooled to 300 oC in a

waste-heat boiler to produce steam. 60 % to 70 % of the H2S is converted to sulfur:

2 H2S + O2 → S2 + 2 H2O

(also H2S + 3/2 O2 → SO2 + 2 H2O )

Catalytic Step: The cooled gas from the combustion chamber passes through

a reactor filled with Co-Mo catalysts on alumina (Al2O3). After condensing out

some sulfur at temperatures below 170 oC, the gas is heated to 230 oC and pumped

through a second and a third reactor. Overall 96 to 99 % yield of sulfur.

2 H2S + SO2 → 3 S + 2 H2O

The waste tail gas must be treated to remove unreacted H2S. • Why? • How?

Sulfur Stockpiled for Export in Vancouver Harbour

This sulfur is produced by the Claus process using H2S extracted from raw

natural gas. Smaller amounts of sulfur are produced using H2S from roasting

copper sulfide and other metal ores.

A miner in East Java, Indonesia, carries sulfur from a mine

3 km up the Ijen volcano, where “brimstone” (sulfur crystals)

condenses from sulfur vapor in volcanic gases.

The water in the crater of the Ijen volcano is the most

acidic lake water in the world, pH from about 0.1 to 0.5



S → SO2 → SO3 → H2SO4

Sulfur Dioxide Production S (l) + O2 (g) → SO2 (g)

SO2 is produced in “sulfur burners” (fire and brimstone!) by

oxidizing liquid sulfur in dry air. Liquid sulfur at 140 oC to 150 oC

(there is a viscosity minimum in this temperature range) is

“atomized” into small droplets using pressure spray nozzles or

mechanically driven spinning cups. Baffles and secondary air inlets

are used to promote good mixing. Catalysts are not needed.

The hot gases (~ 1000 oC) leaving the sulfur burner are cooled in a

waste-heat boiler which generates steam used for other processes.



S → SO2 → SO3 → H2SO4

Sulfur Trioxide Production SO2 (g) + 1/2O2 (g) → SO3 (g)

The Contact Process

Hot air + sulfur dioxide leaving the sulfur burner (~ 1000 oC) is

cooled to 440 oC before entering the first catalyst chamber where

about 65% of the SO2 is converted to SO3 on trays of vanadium oxide

(V2O5) catalyst pellets. The exothermic reaction heats the gas mixture

to about 600 oC. Heat exchangers are used to cool the gas to 440 oC

before it enters a second and a third catalyst chamber. After the third

chamber about 95% of the initial SO2 is converted to SO3, very close

to the thermodynamic limit. SO3 forms “on contact” with the catalyst.

Sulfur Trioxide Production SO2 (g) + 1/2O2 (g) → SO3 (g)

The Double Contact Process

The double contact process is used for even higher yields. SO3

produced after the second catalyst chamber is completely removed

from the gas mixture by absorption into liquid sulfuric acid. The gas

mixture leaving the second catalyst chamber is cooled to 200 oC and

fed into the bottom of a tower filled with ceramic rings (to provide a

large surface area for absorption) into which liquid sulfuric acid is fed

in at the top, countercurrent to the upward flow of gas.

The SO2-containing gas, stripped of SO3, is heated to 440 oC and fed

into a third V2O5 catalytic chamber in for an overall conversion of SO2

to SO3 of up to 99.7%. The gas mixture is cooled and passed through

a second and final absorption tower.

waterwater

SO2 to SO3

SO3 SO3

for hydration: SO3 + H2O → H2SO4:

Flow Sheet for Sulfuric Acid Production from Sulfur, Air and Water

• Why is it important to dry the process air? • What drying agent is used?

• Why does removal of SO3 before the 3rd catalytic reactor increase the yield of H2SO4?

• Why must essentially complete conversion of SO2 to H2SO4 be obtained?

• H2SO4 is cheap ($0.25/kg), but H2SO4 for the StFX Chem Dept costs $10/kg. • Why ?

10. PHOSPHORIC ACID

#9 in the “top 50” list.

Global production 50 million tonnes/year

About 85% is used for agricultural fertilizers. The main phosphorous-containing

fertilizers are prepared by reacting phosphoric acid and ammonia:

• AP (ammonium phosphate, (NH4)3PO4)

• DAP (diammonium hydrogen phosphate, (NH4)2HPO4)

• MAP (monoammonium hydrogen phosphate (NH4)H2PO4)

Also important:

• SP (superphosphate), a mixture of calcium dihydrogen phosphate Ca(H2PO4)2

and gypsum (CaSO42H2O) provides P and secondary Ca and S nutrients

Smaller but important uses of phosphoric acid:

• cleaning agents (such as Na3PO4)

• phosphatizing metals for corrosion protection

• food additives • nutritional supplement for livestock

Furnace Phosphoric Acid

Before the 1940s, most phosphoric acid was produced by the

furnace process: heating calcium phosphate ore to about 2000 oC

with quartz sand and carbon. The elemental phosphorus produced

is oxidized to phosphorous pentoxide (empirical formula P2O5 ,

molecular formula P4O10) which is hydrated to give phosphoric acid.

2Ca3(PO4)2 + SiO2 + 10C → P4 + 10CO + 6CaSiO3

P4 + 5O2 → P4O10

P4O10 + 6H2O → 4H3PO4

The phosphoric acid product is very pure, but fuel costs to reach

2000 oC makes the process expensive. The furnace process is still

used, mainly for high-purity food-grade phosphoric acid.

Wet-Process Phosphoric Acid

More than 95% of phosphoric acid is now made by the wet process.

Phosphate rock (mostly calcium and phosphate with fluoride,

hydroxide, silicate and other impurities) is ground up, mixed with

aqueous sulfuric acid, and heated to about 80 oC. The reactions are

complex. Simplified versions using calcium phosphate or apatite are

Ca3(PO4)2 + 3H2SO4 + 6H2O → 2H3PO4 + 3(CaSO4•2H2O)(s)

Ca5F(PO4)3 + 5H2SO4 + 10 H2O → 3H3PO4 + HF + 5(CaSO4•2H2O)(s)apatite gypsum

Precipitated gypsum can be filtered off. HF (a useful byproduct) is

removed as vapor using air streams. The aqueous phosphoric acid can

be concentrated by water evaporation. The product is impure, but

suitable for making fertilizers. Heavy metals (such as arsenic, copper,

uranium) can be precipitated as sulfides.

1.11 LIMESTONE CHEMICALS

CaCO3 CaO Ca(OH)2

Limestone A naturally-occurring sedimentary rock, mostly

calcium carbonate (CaCO3) from the skeletal remains of coral, shellfish

and other marine organisms. An excellent building material.

Limestone and Niagara Falls? What’s the Connection?

Limestone (CaCO3)

Chemical applications (directly or indirectly, CaCO3 is used in more

industries than any other naturally-occurring compound):

• used to make calcium oxide and sodium carbonate to be discussed

• flux in steel making. SiO2, a common impurity in iron ore, does

not melt in blast furnaces unless CaCO3 is added to form liquid

lava-like “slag” that floats to the top of the furnace for removal.

• fertilizer and soil conditioner (especially for acidic soils)

• desulfurization of flue gases for pollution control:

CaCO3(aq. slurry) + SO2 + 1/2H2O → CaSO3.1/2H2O(s) + CO2(g)

• CaCO3 powder is used as a filler for paper, rubber, plastics, paints

• glass making • water treatment • nutritional supplement

Calcium Oxide (CaO) (also called lime and quicklime)

Produced by calcining limestone at temperatures above 900 oC.

CaCO3(s) → CaO(s) + CO2 (gas)

Historically, CaO was made by burning piles of wood, coal or peat

with pieces of limestone placed on top (“lime burning”).

Modern technology uses crushed limestone heated in rotary kilns

fired by natural gas, fuel oil, or coal.

CaO production is energy intensive. About 4 MJ of heat (equivalent

to burning 1/3 tonne of coal) is required to make one tonne of CaO.

In some facilities, CO2 released by calcining limestone is recovered

from flue gas and used to make Na2CO3.

heat

Coal/Natural Gas-Fired Rotary Kiln for CaO Production from Limestone

Note the counter-current flow of limestone feed and hot combustion gases from burning fuel.

Calcium Hydroxide (Ca(OH)2) (also called slaked lime)

CaO (lime) is unstable in air. It reacts with moisture to form calcium

hydroxide (slaked lime because its thirst for water has been quenched).

CaO + H2O → Ca(OH)2

Because the hydration of CaO is strongly exothermic, and dangerous if

uncontrolled, CaO is usually converted to Ca(OH)2 for shipping.

Just as H2SO4 is the cheap go-to industrial acid, Ca(OH)2 is the go-to

alkali for many applications (cheaper than NaOH). A saturated aqueous

solution of Ca(OH)2 is called lime water. A suspension of (white)

Ca(OH)2 particles in water is called lime milk.

Lime also absorbs carbon dioxide from air, reverting back to limestone.

CaO + CO2 → CaCO3

Applications of Calcium Oxide and Calcium Hydroxide

The applications of lime and slaked lime in the chemical industry are

too numerous to list. Main uses:

• steel production, flux for the removal of sulfur and phosphorous

• pollution control, CO2 and SO2 lime scrubbers for stack gases

• water treatment (Ca(OH)2 + Ca2+ + 2HCO3- → 2CaCO3(ppt) + 2H2O)

• chemical manufacture (such as Na2CO3 (soda ash) to be discussed)

• construction materials (concrete and cement)

• fertilizers (such as calcium nitrate) and soil conditioners

• paint (whitewash)

• paper production, for the regeneration of NaOH (caustic soda)

used to digest and to pulp wood fiber

Na2CO3 + CaO + H2O → CaCO3 + 2NaOH

Cement and Concrete They are different !

Concrete buildings, roads, bridges, dams and other structures are the

most visible applications of limestone chemicals. It’s difficult to over-

estimate the importance of cement and concrete for modern civilization.

Cement is prepared by mixing crushed limestone (CaCO3), silicate

minerals (such as clay), and other ingredients, then heating the

mixture in a kiln to form clinkers, which are cooled and ground

to form cement powder. Before grinding, gypsum (CaSO4•2H2O) is

added to delay the hardening of the cement when water is added.

Concrete is prepared by mixing cement with sand, gravel and water,

and allowing the mixture to harden. The hydrated cement solidifies

to bind together the sand and gravel (called aggregate), forming a

strong artificial stone, often reinforced with steel rods (rebar).

Portland Cement By far the most common and important cement.

It was first manufactured in England in 1845 by heating mixtures of

limestone and clay. The concrete produced resembled building stone

quarried from the nearby Isle of Portland. The name stuck. The

chemical composition of portland cement is roughly CaO-SiO2.

A more precise composition is 64% CaO, 21% SiO2, 6% Al2O3,

3% Fe2O3, 2% MgO, 2% alkali oxides, and 2% SO3.

Specialty Cements

• hydraulic cement solidifies even under water

• alumina cement retains strength at high temperatures

• white cement low iron content, decorative

• magnesia cement conducts electricity, no static buildup

• asbestos cement very durable because of asbestos fibers,

good for roofing tiles

A typical cement kiln rotates about once per minute. It takes about

two hours for solids to pass through a kiln.

• Why does the kiln rotate? • Why are chunks of cement leaving

the kiln called clinkers? • Why are the clinkers ground to powder

before shipping. • Can you suggest why cement kilns are used to

dispose of used tires, plastic, steel slag and other waste products?

Rotating Cement Kiln

cement

Concrete Cement, sand, gravel and

water are mixed to form a slurry

called concrete, that is poured into

frames (often filled with reinforcing

steel rods) and allowed to solidify.

How does concrete solidify? When cement powder is mixed with

water, part of the CaO, Al2O3 and gypsum react to form very small

crystals of the mineral ettringite on the surface of the cement particles.

3CaO + Al2O3 + 3CaSO4 + 32H2O → 3CaO.Al2O3.3CaSO4.32H2O

These crystals are too small to bridge the gap between the cement

particles, so the cement slurry remains fluid. Without added gypsum,

tricalcium aluminate (3CaO + Al2O3 ) immediately hydrates, which

fills the gap between the cement particles, with rapid solidification.

concrete

1.12 SODIUM CARBONATE (Na2CO3)

CaCO3 + 2NaCl → Na2CO3 + CaCl2 Solvay Process

# 11 on the top-50 list. Global production about 50 million tonnes/year.

Sodium carbonate (also called soda ash) is comparable in importance

to sodium hydroxide (caustic soda).

Main Uses (too numerous to list in detail !):

50 % used to manufacture glass (mostly bottle glass)

25 % used for chemical manufacture (sodium silicate, NaHCO3 ...)

10 % for soaps and detergents

10 % miscellaneous (water treatment, pulp and paper, ...)

5 % for desulfurization of flue gases (for pollution control)

Solvay Process Outside North America, most sodium carbonate

is produced by an ingenious process invented by Ernest Solvay in the

1860s using simple raw materials (salt and limestone) and ammonia

as a cyclic reagent. The overall reaction looks simple:

CaCO3 + 2NaCl → Na2CO3 + CaCl2

The actual chemical processes are more complicated:

• CaCO3 is heated in a lime kiln to form CaO and CO2

• the CO2 is reacted with aqueous NH3 to form aqueous (NH4)HCO3

• NaCl is added to the aqueous (NH4)HCO3 solution to precipitate NaHCO3

and form aqueous NH4Cl

• NaHCO3 is collected and heated to form NaCO3 and CO2 (which is recycled)

• CaO is added to the aqueous NH4Cl solution to regenerate NH3 for recycle

and to form CaCl2 (the other main Solvay product)

Solvay Process Reactions

CaCO3(s) → CaO(s) + CO2(g)

2NH3(g) + H2O(l) → NH4OH(aq)

2NH4OH(aq) + 2CO2(g) → 2NH4HCO3(aq)

2NaCl(s) + 2NH4HCO3(aq) → 2NaHCO3(s) + 2NH4Cl(aq)

2NaHCO3(s) → Na2CO3(s) + CO2(g) + H2O(g)

CaO(s) + H2O(l) → Ca(OH)2(aq)

Ca(OH)2(aq) + 2NH4Cl(aq) → CaCl2(aq) + 2NH3(g) + 2H2O(l)

________________________________________________________

CaCO3 + 2NaCl → Na2CO3 + CaCl2 (overall)

Note the regeneration and

recycle of NH3 and CO2.

Ingenious!



Solvay Process Like the Haber process for ammonia production

from air and hydrocarbons, the Solvay process for making sodium

carbonate from widely available salt and limestone was one of the

first triumphs of industrial chemistry. Hundreds of Solvay plants

operate worldwide, but not in North America.

The first (1884) and last North

American Solvay plant (Solvay,

NY - named after a chemical

process!), shut down in 1986.

• Why? Economics. A cheaper

process for sodium carbonate

production was developed.

North American Sodium Carbonate Production from Trona

Huge deposits of the mineral trona Na2CO3.NaHCO3.2H2O

(trisodium hydrogendicarbonate dihydrate, more commonly called

sodium sesquicarbonate) were discovered in southwestern Wyoming

in the 1940s. Using this naturally-occurring raw material, high quality

sodium carbonate can be produced with very little processing:

• dissolution in water

• filtration

• crystallization

• heating to drive off H2O and CO2

Sodium carbonate production costs using trona are significantly lower

than the Solvay process. By 1986, all North American sodium

carbonate production switched to trona. Worldwide, however, the

Solvay process is still very important.

1.13 GLASS

The terminology “glass” refers to a rigid supercooled liquid (no

ordered crystalline structure) having no definite melting point and a

very high viscosity (more than 1015 times larger than the viscosity of

water at room temperature) which prevents recrystallization.

Most industrial glass is a mixture of inorganic oxides (such as SiO2,

Na2O, CaO, Al2O3, B2O3, PbO ...). These glasses have many uses

because of their low cost, transparency, resistance to chemical attack,

and ability to act as electrical insulators.

Glassmaking in North America is a 10 billion dollar / year industry:

• 40 % for glass bottles and other containers

• 25 % for blown glass (such as light bulbs)

• 20 % for flat glass (such as windows).

Industrial Glassmaking

Ingredients such as silica sand, sodium carbonate, calcium oxide,

boric acid, aluminum silicates, and cullet (crushed recycled glass)

are melted in large tank furnaces constructed from refractory blocks

and heated by burning natural gas or another fuel.

Regenerative furnaces are used to save fuel costs. Hot gas from

the burner passes over the molten glass and then flows through a

three-dimensional checker-work of refractory bricks, which are

heated to 600 to 1500 oC. Simultaneously, air for the burner is

preheated by passing through a previously heated regenerative

Checker-work on the other side of the furnace. The flow of air

for combustion is switched from one regenerative chamber (which

has cooled) to the other (now hot) every 20 to 30 minutes.


In addition to regenerative heating to save fuel costs, the crown-shaped

roof over the tank of molten glass provides significant radiative heating.


Bottles and other glass containers are made by using jets of air

to force slugs of molten glass into hollow steel molds. The glass

cools and hardens, and the objects are removed from the molds.

Industrial glass is manufactured in fully automated plants

with high throughputs. Manual forming of glass products, still

used on a smaller scale for high-value complicated shapes and

esthetic pieces, requires considerable skill and craftsmanship.


Flat glass for windows was produced for many years by squeezing

hot, soft glass between steel rollers, or by flattening curved glass

from the walls of blown glass cylinders. Today, industrial flat glass is

is high-quality float glass produced by cooling sheets of molten glass

floating on the surface tanks filled with liquid tin.


Flat glass is an important construction material for modern buildings.

Main Categories of Industrial Glass

1. Soda-Lime Glass

About 95% of all glass manufactured is soda-lime glass consisting

of 70 to 74% SiO2, 13 to 18% Na2O, 8 to 13% CaO, 1 to 2% Al2O3.

Soda-lime glass is used for windows, bottles, drinking glasses, etc.

It is relatively cheap, and can be melted and processed at lower

temperatures (~1200 oC) than other glasses. Higher Al2O3 content

increases the chemical resistance of soda-lime glass by reducing the

leaching of sodium ions, but the melting point is raised.

2. Borosilicate Glass (tradename Pyrex® (translation?))

Borosilicate glass contains 10 to 20% B2O3, 80 to 87% SiO2, and

less than 10% Na2O. It is tougher (more resistant to mechanical

shock and brittle fracture) and more chemically resistant than

cheaper soda-lime glass. Most laboratory glassware is borosilicate

glass. It has a low thermal expansivity and is used for optical lenses.


3. Alkali Silicate Glass (Silica Gel)

Alkali silicates are the only industrially important two-component

glasses. Manufactured by melting quartz sand (crystalline SiO2) and

soda ash (Na2CO3). Carbon dioxide is expelled, producing glass

with compositions from Na2O.SiO2 to Na2O.4SiO2. Water glass

(more accurately termed “water soluble glass”) is an aqueous solution

of Na2O•SiO2 used for fireproofing and an adhesive for cardboard.

4. Fused Silica

Fused silica, also called vitreous silica, is manufactured by melting

quartz sand (SiO2) or by high-temperature pyrolysis of SiCl4. It has

a higher softening point than other glasses and can therefore be

used at higher temperatures, but is more difficult to melt and

process. Unlike other glasses, fused silica is transmits ultraviolet

radiation and is used as windows and lenses uv spectrometers.


5. Lead Glass

Lead glass is made by replacing some or all of the CaO with PbO.

These glasses are important for the construction of lenses in optical

instruments because of their high refractive index (ability to “bend”

light rays). “Cut glass” or “crystal” (actually lead glass) owes its

brilliance to the high dispersion (change in refractive index with

wavelength). Lead glass is also used for nuclear radiation shielding.

6. Glass Fibers

Mats or bales of fine glass fibers (typical diameter 0.001 cm) are

widely used to provide fireproof thermal insulation. These fibers

have a very large surface area and are vulnerable to attack by

atmospheric moisture, so their Na2O content is kept low (< 0.5%).

Glass fibers are also used for filters and to reinforce plastics (usually

epoxies and polyesters).


7. Laminated Safety Glass

Windshields for cars and trucks are manufactured by laminating two

sheets of glass (about 3 mm thick) with a sheet of flexible plastic

glued in between. Lamination makes windshields strong and tough,

part of the vehicle structure. The plastic layer holds glass fragments

together in a crash, a very important safety feature.

8. Tempered Glass

This kind of glass is heat treated by warming the glass to just below

its softening temperature, then rapidly cooling it in air or oil. The

outer skin of the glass cools and hardens rapidly. The interior cools

more slowly and gradually contracts, pulling inward on the skin and

compressing it for greater strength. When tempered glass breaks,

the built-in stresses cause the glass to “crumble” into small pieces

which are much less dangerous than large sharp fragments.


9. Photochromic Glass

Photochromic glass darkens when exposed to light, and reversibly

bleaches at low light intensity levels. This behavior is used in

sunglasses and windows to control the amount of transmitted light.

Photochromic behavior is caused by manufacturing glass with

submicroscopic particles of AgCl (about 5 nm diameter. Light

provides energy to release electrons from Cl- ions, which are picked

up by Ag+ ions to forming particles of silver metal which darkens

the glass. The glass provides a rigid inert matrix which does not allow

the silver particles to diffuse together to irreversibly form large

stable silver particles, as in photographic film.

n AgCl (low light) = Agn0 + n Cl0 (bright light)

Organic dyes are used for photochromic plastic lenses.


10. Colored Glass (also called stained glass)

Hundreds of different colors can be produced:

a) dissolving chemicals in the glass that absorb certain wavelengths

(often oxides of Ti, V, Cr, Mn, Fe, Co or Ni)

b) precipitating colloidal particles in the glass (e.g., colloidal gold

produces a ruby-red glass)

c) adding microscopic particles of colored pigment material

(e.g., red SeO2 particles produce red glass for traffic lights)

Red gold, copper, selenium

Yellow cadmium, calcium

Green iron, chromium, tin, arsenic

Blue cobalt

Purple manganese, nickel

Stained Glass

Chartres Cathedral,

northern France

manufactured in the 1230s

1.14 IRON and STEEL

Golden Gate Bridge, San Francisco. Constructed 1933-37. 75,000 tons of steel.

1.14 IRON and STEEL

Impossible to overestimate the industrial importance of iron.

No other element defines our age more than iron, the main

ingredient of steel, which is used to make:

• reinforced concrete

• structural beams for buildings and bridges

• engines, pumps, compressors

• cars, trucks, ships, trains, other vehicles

• railroads

• pipe, tanks, valves

• reactors for chemical production

• machine parts, cutting tools, fasteners

• tin cans (actually tin-coated steel)

• countless other products

About 1.7 billion tonnes of steel are produced per year. Smaller

amounts of iron are used to make other iron alloys and catalysts.

Steel Production

Historical

For centuries, steel knives, swords,

armor, horseshoes and other items

were produced one-at-a-time by

skilled blacksmiths using techniques

developed by trial and error.

• iron ore containing silicates and other impurities is mixed with charcoal

and heated in a furnace, producing a porous spongy material containing

particles of steel and slag (silicates and other minerals)

• hammering (working) the sponge released grains of (wrought) steel

• labor-intensive, unsuitable for large work pieces

Good steel was found to be made by repeatedly, heating, and cooling the

workpieces. Not initially understood, the charcoal served as the carbon source

needed to make steel. Heating and cooling modified the steel crystal structure

for improved hardness and toughness.

Steel Production

Bessemer Process - the breakthrough

In the 1850s, Henry Bessemer developed the Bessemer converter

for the large-scale production of high-quality low-cost steel.

The Industrial Age (The Steel Age) had begun.

• iron ore smelted in blast furnaces, converting iron oxides to impure iron

• 10 to 30 tonnes batches of molten iron loaded into Bessemer converters

• air is blasted through the melt to burn off carbon as CO and CO2

• dolomite (CaMg(CO3)2) and manganese added to bind S, P, Si impurities

• slag containing the impurities floats to the top

• denser liquid steel tapped off from below

Bessemer

converter

(retired)

Iron Ores

• major naturally-occurring iron ore deposits contain

hematite (Fe2O3) magnetite (Fe3O4)

goethite (Fe2O3) limonite (FeO(OH).nH2O))

• the iron minerals are mixed with sand, clay and other impurities

• high-grade ores (> 60% Fe) can be fed directly into blast furnaces

• lower grade ores are concentrated using magnetic separation or

froth flotation (more important)

Froth Flotation

Ores are crushed and ground to a fine powder. Powdered ore is added to tanks of

water through which air is bubbled. With the help of surface-active (surfactant)

additives, such as alkyl quaternary ammonium salts, iron ore grains preferentially

stick to the bubbles and “float” to the top of the tanks and overflow into collection

tanks or are skimmed off. Silicates and other impurities collect as sludge in the

bottom of the tanks. Ore concentration by flotation is often performed at mine sites

to reduce shipping costs.

Coke

Blast furnaces use heat and reducing agents such as CO to convert

iron oxides to steel: iron containing several per cent carbon and smaller

amounts of other elements. A highly simplified overall reaction is

Fe2O3 + 3CO → 2Fe + 3CO2

The reducing agents are derived from coal, oil or natural gas.

Coke, the most commonly used reducing agent, is a hard carbonaceous

residue prepared by heating coal in coke ovens. Important: the coking

process drives off most of the sulfur and other impurities.

Petroleum coke (called petcoke), a byproduct from petroleum refining,

is the solid carbon-rich residue left over after the distillation and

processing of lower molecular weight hydrocarbon fractions (Part 3).

red-hot coke produced

the pyrolysis of coal

in a coking oven

petcoke (also used a fuel)

residue from petroleum refining

Blast Furnace – Converts Iron Ore to Pig Iron (Impure Iron)

Blast Furnace Design

• steel shell about 30 m tall, diameter about 10 m, open at the top

• lined with refractory bricks (mostly Al2O3) to withstand high temperatures and

harsh chemical reactions

Blast Furnace Operation

• continuous countercurrent flow (not batch)

• coke and iron ore pellets added at the top of a vertical furnace descend slowly

in an upward current of hot gases produced by blasting high-pressure air

or oxygen into the base of the furnace

• hydrocarbons (oil or natural gas) can be added to the blast to produce H2 to decrease

the amount of coke required to reduce the iron oxides

• descending iron oxides release oxygen into the rising gas stream, finally

producing molten pig iron at the bottom of the furnace, which is poured off

• limestone (mostly CaCO3) flux is also added at the top of the furnace to absorb

sulfur, phosphorus, silicon and other impurities in slag, which is less dense

and floats on top of the molten pig iron (impure iron) and can be tapped off

Blast Furnace Chemistry

Complicated! Fe2O3, Fe3O4 and FeO are reduced to Fe.

2C(coke) + O2 → 2CO

2CH4 + O2 → 2CO + 4H2

3Fe2O3 + CO → 2Fe3O4 + CO2

3Fe2O3 + H2 → 2Fe3O4 + H2O

Fe3O4 + CO → 3FeO + CO2

Fe3O4 + H2 → 3FeO + H2O

FeO + CO → Fe + CO2

FeO + H2 → Fe + CO2

___________________________________________

Overall: Fe2O3 + 3CO → 2Fe + 3CO2

Blast Furnace Chemistry

Temperatures in a blast furnace can be as high as 1600 oC. The

residence time for rising gas is only a few seconds, but residence

times for the descending solids are typically hours. Blast furnaces

can be 10 stories tall. Once started, a blast furnace can’t be turned

off and must be operated 24/7/365. • Why? • Why called “blast” furnace?

Typical blast-furnace material balance to make 1000 kg pig iron:

Inputs: Outputs:

1600 kg iron ore 1000 kg pig iron (impure iron)

1300 kg air 2300 kg gases

450 kg coke 300 kg slag

150 kg limestone 10 kg dust

60 kg oxygen

50 kg hydrocarbons


Adding iron ore, coke and limestone to the top of the furnace.


Pouring liquid pig iron.

• Why is it called pig iron?

Basic Oxygen Steelmaking (BOS)

• used to convert pig iron to steel

• significant improvement in steel production introduced in the 1950s

• pure O2 (instead of air) is blasted into BOS furnaces

• BOS furnaces therefore run hotter (1700 oC) than older furnaces (• Why?)

• therefore faster too ( 1 hr per batch compared to 10 hr for older furnaces)

• CaO, MgO, CaMg(CO3)2 added to absorb impurities in slag

• no N2 impurity to make the steel brittle

• easier to use scrap steel (almost pure steel)

• used for 60 % of global steel production

• How is the pure oxygen produced?

• Why called “basic” oxygen process?

Steelmaking is the subject of continuous research and development:

Taking advantage of steel recycling, electric arc furnaces are

becoming increasingly important for steel production. Because

scrap steel (almost pure steel) is used instead of impure iron ore,

less chemical processing is required to make high-quality steel.

• electrical heating of the batch

• 500 volt and 50,000 amperes

• graphite electrodes

• temperatures up to1800 oC

• therefore fast ( 1 hr per batch)

• pig iron also used as feed

_____________________________________________________________

pig iron after BOS or EAF processing

C 4 to 5 % 0.3 to 1.0 %

Si 0.2 to 1.0 % 0.001 to 0.004 %

P 0.1 to 0.2 % 0.01 to 0.03 %

S 0.02 to 0.06 0.01 to 0.03 %

Electric Arc Furnace (EAF) Steelmaking

Hundreds of Different Kinds of Steel are Manufactured using

Different Chemical Compositions and Different Heat Treatments

(Quenching, Annealing, Tempering)

• low-carbon steels (< 0.25 % C)Sometimes called mild steel. Relatively ductile and easily formed

into structural beams, cold-rolled sheets, etc. Many applications.

• medium-carbon steels (0.25 % to 0.70 % C)Stronger but less ductile (more brittle) than mild steel. Usually

heat-treated to produce items such as machine parts.

• high-carbon steels (> 0.70 % C)Very hard, but lack toughness and can be brittle. Used for cutting

tools and bearings for which wear resistance is very important.

• stainless steelsContain several percent chromium, nickel, vanadium and other

elements to improve corrosion resistance by forming a protective

layer of chromium oxide and other compounds. More expensive and

more difficult to work than carbon steels.

Documents

Top 50 Industrial Chemicals (by weight) · Chem 471 Part 1: Primary Industrial Inorganic Chemicals 1.1 water, purification and treatment 1.2 methane (organic, but obtained from natural