Transcript ch22_lecture_6e_final

Lecture PowerPoint
Chemistry
The Molecular Nature of
Matter and Change
Sixth Edition
Martin S. Silberberg
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Copyright  The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Chapter 22
The Elements in Nature and Industry
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The Elements in Nature and Industry
22.1 How the Elements Occur in Nature
22.2 The Cycling of Elements Through the Environment
22.3 Metallurgy: Extracting a Metal from Its Ore
22.4 Tapping the Crust: Isolation and Uses of Selected Elements
22.5 Chemical Manufacturing: Two Case Studies
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Figure 22.1
The layered internal structure of Earth.
The processes by which the Earth formed led to its differentiation
into different layers.
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Elemental Abundance
The abundance of an element is the amount of that
element in a particular region of the natural world.
Elements are not equally abundant in all regions –
abundances differ due to the differences in physical and
chemical behavior of the elements.
- The core of the Earth is rich in dense Group 8B(8) to 8B(10) metals.
- The Earth’s crust has the largest share of nonmetals, metalloids,
and light active metals.
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Table 22.1 Cosmic and Terrestrial Abundances (Mass %) of
Selected Elements.
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Compositional Phases of the Earth
As the Earth cooled to form its major layers, gravity and
convection caused materials of different densities to
separate, giving several phases.
Fe was the major component of the core or iron phase.
The outer silicate phase, containing oxygen combined
with Si, Al, Mg, and some Fe, separated into the
mantle and crust.
The sulfide phase, with intermediate density,
consisted mostly of iron sulfide mixed with parts of the
other phases.
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Figure 22.2
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Geochemical differentiation of the elements.
Distribution of the Elements
The distribution of the elements in the Earth’s layers was
controlled by their chemical affinity for one of the three
phases.
Elements with low or high electronegativity tended to
congregate in the silicate phase as ionic compounds.
These included active metals and nonmetals.
Metals with intermediate EN dissolved in the iron
phase.
Lower-melting transition metals, and many metals and
metalloids in Groups 11 to 16, became concentrated in
the sulfide phase.
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Impact of Life on Crustal Abundances
• Photosynthesis resulted in an increase in the O2
levels of the atmosphere.
– Oxidation became the major source of free energy in the
crust and biosphere.
• The [K+] of the oceans is much lower than the [Na+]
– since growing plants have absorbed dissolved K+.
• Subterranean deposits of organic carbon formed from
the decomposition of ancient plants under high
pressure.
• Fossilized skeletal remains of early marine organisms
have formed vast deposits of C, O and Ca.
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Figure 22.3
Ancient effect of an O2-rich atmosphere.
Banded-iron formations
containing Fe2O3
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Fossil of early multicellular
organism
Table 22.2 Abundances of Selected Elements in the Crust,
Its Regions, and the Human Body as
Representative of the Biosphere (Mass %)
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Sources of the Elements
O2, N2, and the noble gases (except He) are obtained
from the atmosphere, where they occur as the free
elements.
A few elements occur naturally in their uncombined
(native) state.
These include S, carbon in coal, and unreactive metals.
Most elements occur in ores, natural compounds or
mixtures from which an element must be extracted.
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Figure 22.4
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Sources of the elements.
Environmental Cycles
An environmental cycle is a natural process in which
elements are continuously cycled in various forms
between different regions of the Earth’s crust.
Elements are cycled through physical, biological, and
chemical pathways.
The most important of these cycles from the perspective
of living organisms are the carbon, nitrogen, and
phosphorus cycles.
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Sources of Carbon
• Carbon appears in all three portions of the Earth’s crust.
• Elemental C occurs as diamond and graphite in the
lithosphere.
• C is a component of carbonate minerals.
• C forms the basis of all organic substances.
• CO2(g) is found in the atmosphere and dissolved in
rivers, lakes and seas.
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The Carbon Cycle
Fixation is the process of converting a gaseous substance
into a condensed form. Through photosynthesis, plants
use sunlight to fix atmospheric CO2 into carbohydrates.
Animals eat the plants, incorporating the carbon into
their systems.
All living organisms release CO2 by respiration and by
decomposition.
The portions of the carbon cycle are linked by the
atmosphere – one link between oceans and air, and the
other between land and air.
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Figure 22.5
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The carbon cycle.
The Nitrogen Cycle
• The nitrogen cycle involves a direct interaction between
land and sea.
• Atmospheric N2 must be fixed to enter the land and sea.
– Atmospheric fixation occurs when lightning provides the energy
for the reaction between N2(g) and O2(g).
– Industrial fixation results from the production of NH3, which is
used to make fertilizers.
– Biological fixation involves the conversion of atmospheric N2 to
NO3- by blue-green algae and nitrogen-fixing bacteria.
• Denitrifying bacteria reduce NO3- back to N2 in various
stages, allowing it to cycle back to the atmosphere.
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Figure 22.6
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The nitrogen cycle.
The Phosphorus Cycle
• P cycles between land and sea in the form of
phosphate (PO43-) and the various hydrogen phosphate
ions.
• There are three interlocking subcycles:
– The inorganic cycle involves slow weathering of phosphatecontaining rocks, which causes PO43- to leach into the rivers
and seas.
– The land-based biological cycle involves incorporation of
PO43- into organisms and its release through excretion and
decomposition.
– The water-based biological cycle consists of the cycling of
phosphate oxoanions by aquatic organisms.
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Figure 22.7
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The phosphorus cycle.
Figure 22.8
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Industrial uses of phosphorus.
Metallurgy
Metallurgy is the technology of metals, and is concerned
with their extraction and utilization.
• Pyrometallurgy uses heat to obtain the metal.
• Electrometallurgy employs an electrochemical step.
• Hydrometallurgy relies on the metal’s aqueous solution
chemistry
Extraction of a pure element from its ore involves a
series of steps, beginning with mining the ore.
The mineral contains the element while the gangue is the
portion of the ore with no commercial value.
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Table 22.3 Common Mineral Sources of Some Elements
Element Mineral, Formula
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Al
Gibbsite (in bauxite), Al(OH)3
Ba
Barite, BaSO4
Be
Beryl, Be3Al2Si6O18
Ca
Limestone, CaCO3
Fe
Hematite, Fe2O3
Hg
Cinnabar, HgS
Na
Halite, NaCl
Pb
Galena PbS
Sn
Cassiterite, SnO2
Zn
Sphalerite, ZnS
Figure 22.9
Steps in metallurgy.
Pretreating
Mining
Magnetic attraction
Cyclone separation
Flotation
Leaching
Refining
Alloying
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Electrorefining
Distillation
Zone refining
Converting
(mineral to compound)
Pyrometallurgy
(roasting, etc.)
Hydrometallurgy
Converting
(compound to metal)
Chemical redox
(smelting, etc.)
Electrochemical redox
Pretreating the Ore
• Magnetic attraction can be used to separate magnetic
minerals from the gangue.
• Density separation exploits large differences in densities.
– Panning for gold relied on the density differences between gold
and sand.
• Flotation uses an oil-detergent mixture to form a slurry,
which is mixed rapidly with air to give an oily, mineralrich froth that floats.
• Leaching involves the formation of a complex ion that is
water soluble.
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Figure 22.10
The cyclone separator.
Panning for gold also relies
on density differences.
The cyclone separator blows high-pressure air through the
pulverized mixture to separate the particles.
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Figure 22.11
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The flotation process.
Conversion of Minerals by Roasting
Minerals are often converted to their oxides because
oxides can be reduced easily.
CaCO3(s)
Δ
CaO(s) + CO2(g)
Metal sulfides are converted to oxides by roasting in air:
2ZnS(s) + 3O2(g)
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Δ
2ZnO(s) + 2SO2(g)
Conversion of Minerals by Chemical Redox
Reduction with carbon is common. Heating an oxide with a
reducing agent to obtain the metal is called smelting.
ZnO(s) + C(s) → Zn(g) + CO(g)
Oxides of less active metals are reduced with hydrogen
instead of carbon.
WO3(s) + H2(g) → W(s) + 3H2O(g)
A more active metal may be used as the reducing agent.
Cr2O3(s) + 2Al(s) → 2Cr(l) + Al2O3(s) DH° << 0
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Figure 22.12
The thermite reaction.
Al powder is used to reduce a metal oxide in a spectacular
exothermic reaction, which produces the molten metal.
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Electrochemical Redox in Mineral Conversion
The mineral is converted to the element in an electrolytic
cell. Often the pure mineral is used to prevent unwanted
side reactions.
The cation is reduced to the metal at the cathode, and the
anion is oxidized to the nonmetal at the anode:
BeCl2(l) → Be(s) + Cl2(g)
Specially designed cells separate the products to
prevent recombination.
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Figure 22.13
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The redox step in converting a mineral to the element.
Refining and Alloying the Element
• Refining (purifying) the element is carried out by one of
the following three methods:
– Electrorefining employs an electrolytic cell, in which the impure
metal acts as the anode and the pure metal as the cathode.
– Distillation is used to refine metals with relatively low boiling
points.
– Zone refining exploits freezing point depression.
• Many metals are in alloys rather than their pure forms.
Alloying involves combining the metals in the correct
proportions to produce a material of the desired physical
properties.
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Figure 22.14
Zone refining of silicon.
• The impure metal or metalloid is passed slowly through a heating
coil in an inert atmosphere.
• A small zone of the impure solid melts, and as the next zone melts,
the dissolved impurities from the first zone lowers the freezing point.
• The purer solid of the first zone refreezes.
• Gradually impurities from each zone move into the adjacent zone,
producing pure solid at the end of the rod.
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Table 22.4 Some Important Alloys and their Composition
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Figure 22.15
vanadium carbide
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Three binary alloys.
Cu3Au
b-brass
Producing Sodium and Potassium
Sodium ore is halite, which is obtained either by the
evaporation of brines or by mining salt deposits.
Na is extracted and purified in an electrolytic apparatus
called the Downs cell.
Sylvite (mostly KCl) is the major ore of potassium.
Chemical reduction of K+ ions by Na at high
temperature produces the metal:
Na(l) + K+(l)
Na+(l) + K(g)
The reaction is carried out at 850°C so that gaseous K is produced.
As the K(g) is removed the equilibrium shifts to produce more.
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Figure 22.16
The Downs cell for production of sodium.
CaCl2 is mixed with the NaCl to lower the melting point.
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Evaporative salt beds near
San Francisco Bay.
Firefighter with KO2
breathing mask.
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Metallurgy of Iron
• Iron ores contain iron oxides, carbonates and sulfides.
• Iron is recovered from its ores by reduction with C in a
blast furnace.
– This process produces impure pig iron.
• Pig iron is converted to steel by means of the basicoxygen process.
– High pressure O2 is blown over and through the molten iron,
oxidizing the impurities, which are converted to a molten slag
and removed.
• Steel is produced by alloying iron with other metals to
prevent corrosion and increase both strength and
flexibility.
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Table 22.5 Important Minerals of Iron
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Mineral Type
Mineral, Formula
Oxide
Hematite, Fe2O3
Magnetite, Fe3O4
Ilmenite, FeTiO3
Carbonate
Siderite, FeCO3
Sulfide
Pyrite, FeS2
Pyrrhotite, FeS
Figure 22.17
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The major reactions in a blast furnace.
Figure 22.18
The basic-oxygen process for making steel.
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
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Isolation of Copper
• The most common copper ore is chalcopyrite, a mixed
ore of CuFeS2, which contains 0.5% Cu by mass.
• Flotation of the ore removes some of the iron, enriching
the mass % of Cu.
• Roasting oxidizes the FeS to FeO but leaves the CuS
unreacted.
– 2FeCuS2(s) + 3O2(g) → 2CuS(s) + 2FeO(s) + 2SO2(g)
• Heating in sand converts FeO to FeSiO3, a molten slag.
– FeO(s) + SiO2(g) → FeSiO3(l)
– CuS → Cu2S at this high temperature.
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Isolation of Copper
• The final smelting step converts some Cu2S to Cu2O by
roasting in air. Cu2S and Cu2O react:
– 2Cu2S(s) + 3O2(g) →2Cu2O(s) + 2SO2(g)
– Cu2S(s) + 2Cu2O(s) → 6Cu(l) + SO2(g)
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Mining chalcopyrite.
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Figure 22.19
The electrorefining of copper.
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Copper for wiring needs to be 99.99% pure. This is achieved by
electrorefining. The impure Cu is used as a series of anodes, and the
cathodes are made from purified Cu. An acidic solution of CuSO4
serves as the electrolyte.
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Isolation of Aluminum
• Al is the most abundant metal in the Earth’s crust.
– The major ore of Al is bauxite, which contains Al2O3.
• The Bayer process is the pretreatment of bauxite by
boiling with NaOH.
– Al2O3 and SiO2 are amphoteric and dissolve in the base, while
Fe2O3 and TiO2 do not.
– The silicate is precipitated by heating, and the filtrate acidified to
precipitate Al3+ as Al(OH)3.
– Drying at high temperature converts this to Al2O3.
• Al2O3 is converted to the free metal by the Hall-Heroult
electrolytic process.
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Mining bauxite.
Figure 22.20
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The electrolytic cell in the manufacture of aluminum.
Figure 22.21
The many familiar and essential uses of aluminum.
Although Al is a very active metal, it does not corrode readily because
when Al2O3 forms on the metal surface it adheres to the metal and
prevents more O2 from penetrating. Al can be anodized to form a
thicker protective oxide layer.
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Isolation of Magnesium
• Magnesium is readily obtained from sea water, although
it is also abundant on land.
• The Dow process is used to isolate Mg.
– Ca(OH)2 is generated from crushed seashells, then added to the
seawater.
– Ca(OH)2(aq) + Mg2+(aq) → Mg(OH)2(aq) + Ca2+(aq)
– Mg(OH)2(s) + 2HCl(aq) → MgCl2(aq) + 2H2O(l)
– Heating evaporates the water to give hydrated MgCl2·nH2O(s).
– The solid is heated and melted, and electrolysis is used to
convert the MgCl2 to molten Mg and Cl2(g).
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Figure 22.22 The Dow process for isolation of elemental Mg
from seawater.
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Sources of Hydrogen
Very pure H2 is prepared through electrolysis of water
with Pt (or Ni) electrodes:
2H2O(l) + 2e- → H2(g) + 2OH-(aq)
E = -0.42 V [cathode; reduction]
H2O(l) → ½O2(g) + 2H+(aq) + 2e- E = 0.82 V [anode; oxidation]
H2O(l) → H2(g) + ½O2(g)
E = -1.24 V
H2 can also be produced by thermal methods. The most
common of these use H2O and a simple alkane like CH4.
H2O(g) + CH4(g) → CO(g) + 3H2(g)
DH° = 206 kJ
H2O(g) + CO(g)
DH° = -41 kJ
CO2(g) + H2(g)
This second reaction, called the water-gas shift reaction, is carried
out using an iron or cobalt oxide catalyst and high temperatures.
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Uses of Hydrogen
• Ammonia is synthesized using N2 and H2.
• H2 is used in the hydrogenation of C=C double bonds in
liquid oils to form C-C bonds in solid fats and margarine.
– Partially hydrogenated vegetable fats and oils are present in a
wide variety of foods.
• H2 is essential in the manufacture of “bulk” chemicals,
those produced in large amounts.
– Methanol, used as a gasoline additive, is produced by the
reaction of CO(g) and H2(g) over a copper-zinc catalyst.
CO(g) + 2H2(g)
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Cu-ZnO catalyst
CH3OH(l)
Hydrogen Isotopes
• Hydrogen has three isotopes;
– 1H (protium), 2H or 2D (deuterium), and 3H or 3T (tritium).
• The relative difference in the mass of these isotopes is
very large, leading to significant differences in bond
energies and reactivities.
– A bond to H is much weaker and easier to break than a bond to
deuterium.
• Any reaction that includes a bond to H or D in the ratedetermining step will occur much faster with H than with
D. This is called the kinetic isotope effect.
– This effect is exploited in the isolation of deuterium, and is also
used to study reaction mechanisms.
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Table 22.6 Some Molecular and Physical Properties of
Diatomic Protium, Deuterium, and Tritium.
Property
H2
D2
T2
Molar mass (g/mol)
2.016
4.028
6.032
Bond length (pm)
74.14
74.14
74.14
Melting point (K)
13.96
18.73
20.62
Boiling point (K)
20.39
23.67
25.04
DH°fus (kJ/mol)
0.117
0.197
0.250
DH°vap (kJ/mol)
0.904
1.226
1.393
Bond energy (kJ/mol 432
at 298 K)
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443
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Producing Sulfuric Acid
Sulfur can be produced by the Claus process from H2S
in “sour” natural gas:
2H2S(g) + 2O2(g) → ⅛S8(g) + SO2(g) + 2H2O(g) at low temperature
2H2S(g) + SO2(g) → ⅛S8(g) + 2H2O(g)
using an Fe2O3 catalyst
The Frasch process is used when natural deposits of S
are found:
Superheated water is pumped into the deposit to melt the sulfur,
which is forced to surface by compressed air.
Sulfur is burned in air to form SO2:
⅛S(s) + O2(g) → SO2(g) DH° = -297 kJ
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Figure 22.23
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The Frasch process for mining elemental sulfur.
Producing Sulfuric Acid
The contact process oxidizes SO2 to SO3:
SO2(g) + ½O2(g)
SO3(g)
DH° = -99 kJ
This reaction is exothermic and very slow at room temperature. A V2O5
catalyst is used to optimize rate, and the yield of SO3 is increased by
increasing the pressure and removing the SO3 as it forms.
SO3 cannot be added directly to water to form H2SO4
because it will polymerize. A small amount of H2SO4 is
used instead:
SO3(g) → H2S2O7(l)
H2S2O7(l) + H2O(l) → 2H2SO4(l)
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Figure 22.24 The many indispensable applications of sulfuric acid.
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The Chlor-Alkali Process
The chlor-alkali process yields chlorine, which ranks
among the top 10 chemicals in the United States.
This process involves the electrolysis of concentrated
aqueous NaCl.
Diaphragm-Cell Method:
2Cl-(aq) → Cl2(g) + 2eE° = 1.36 V [anode; oxidation]
2H2O(l) + 2e- → 2OH-(aq) + H2(g) + 2e- E° = -1.0 V [cathode; reduction]
2Cl-(aq) + 2H2O(l) → 2OH-(aq) + H2(g) + Cl2(g)
Ecell = -2.4 V
To maximize the yield of Cl2, a voltage almost twice this value and a
current in excess of 3x104A are used.
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Figure 22.25
A diaphragm cell for chlor-alkali process.
Overall reaction:
2Na+(aq) + 2Cl-(aq) + 2H2O(l) → 2Na+(aq) + 2OH-(aq) + H2(g) + Cl2(g)
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Mercury-Cell Method:
High purity NaOH can be obtained with this method, which uses a
mercury cathode. This creates such a large overvoltage for the
reduction of H2O that Na+ is reduced instead:
2Cl-(aq) → Cl2(g) + 2e2Na+(aq) + 2e- → 2Na(Hg)
[anode; oxidation]
[cathode; reduction]
The sodium amalgam is pumped out of the system and treated with
H2O, which is reduced by the Na:
2Na(Hg) + 2H2O(l) → 2Na+(aq) + 2OH-(aq) + H2(g)
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Membrane-Cell Method:
This method replaces the diaphragm with a polymeric membrane to
separate the cell compartments.
The membrane allows only cations to move through it, and only from
the anode to the cathode compartment
As Cl- ions are oxidized at the anode, Na+ cations in the anode
compartment move through the membrane to the cathode
compartment and form an NaOH solution.
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