Chapter 22: Applying Chemical Principles to the Main Group Elements

Transcription

22
Applying Chemical
Principles to the
Main-Group Elements
Unit Outline
22.1
22.2
22.3
22.4
22.5
Structures of the Elements
Oxides and Halides of the Nonmetals
Compounds of Boron and Carbon
Silicon
Oxygen and Sulfur in the Atmosphere
In This Unit…
AlbertSmirnov/iStockphoto.com
The main-group elements, those in the s block and p block of the periodic table, vary greatly in their physical and chemical properties. They
are metals, nonmetals, and metalloids and solids, liquids, and gases. This
unit uses the chemical principles you have learned to describe some of
the chemical and physical properties of these elements. We will use, for
example, your knowledge of periodic trends, chemical structure, equilibria, and acid–base chemistry to understand the chemistry of some of the
main-group elements.
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22.1 Structures of the Elements
22.1a The Periodic Table
The periodic table is chemistry’s organizational framework. Periodic trends in atomic size,
orbital energy, and effective nuclear charge are important in understanding periodicity
and the properties of the elements, especially the nonmetals. Recall from Electron
Configurations and the Properties of Atoms (Unit 7) that atomic radii of elements generally decrease as you move from left to right across the periodic table and increase down a
group (Figure 22.1.1).
The decrease in size across a group is related to the increase in effective nuclear charge
and decrease in orbital energy (Figure 22.1.2). These differences result in significant differences in the properties of the elements and allow us to easily differentiate between metals and
1A
8A
He, 32
2A
3A
4A
5A
6A
7A
Li, 152
Be, 113
B, 83
C, 77
N, 71
O, 66
F, 71
Ne, 69
Na, 186
Mg, 160
Al, 143
Si, 117
P, 115
S, 104
Cl, 99
Ar, 97
K, 227
Ca, 197
Ga, 122 Ge, 123 As, 125 Se, 117 Br, 114 Kr, 110
Rb, 248
Sr, 215
In, 163
Sn, 141 Sb, 141 Te, 143
I, 133
Cs, 265
Ba, 217
Tl, 170
Pb, 154
At, 140 Rn, 145
Bi, 155 Po, 167
Xe, 130
© 2013 Cengage Learning
H, 37
Figure 22.1.1 Atomic radii (pm) of the main-group elements
Unit 22 Applying Chemical Principles to the Main-Group Elements
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Energy
3p
3s
2p
2s
1s
Na
4s
3p
3s
2p
2s
1s
Mg
4s
3p
3s
2p
2s
1s
Al
4s
3p
3s
2p
2s
1s
Si
4s
4s
4s
3p
4s
3p
3p
3p
3s
3s
3s
2p
2p
2p
2s
2s
2s
1s
1s
1s
3s
2p
2s
1s
P
S
Cl
Ar
4s
Figure 22.1.2 Orbital energies for elements in the third period
nonmetals. Main-group metals, located on the left and toward the bottom of the periodic table,
have larger radii, low effective nuclear charge, and low first ionization energies and generally
form cations when they react. Nonmetals, located on the right and toward the top of the periodic table, have smaller radii, large effective nuclear charge, and high first ionization energy
and generally form anions when they react. Those along the diagonal line from B to Te display
properties of both metals and nonmetals and are called metalloids (Interactive Figure 22.1.3).
22.1b Metals
Why are metals found to the left and toward the bottom of the periodic table? The answer lies
in the properties of metals and how those properties come about based on orbital energies
and the degree to which those orbitals are filled. Metals have important properties in common:
their structure involves atoms packed closely together without discrete covalent bonds, they
are good conductors of electricity, and they tend to lose electrons to form cations.
Bonding and Conduction in Metals
Elements on the left side of the periodic table have multiple vacancies in their valence orbitals, allowing metal atoms to form multiple bonding interactions with neighboring atoms. As
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Interactive Figure 22.1.3
Investigate metals and nonmetals in the periodic table.
8A
1A
2A
3A
4A
5A
6A
7A
He
2
Li
Be
B
C
N
O
F
Ne
3
Na
Mg
Al
Si
P
S
Cl
Ar
4
K
Ca
Ga
Ge
As
Se
Br
Kr
5
Rb
Sr
In
Sn
Sb
Te
I
Xe
6
Cs
Ba
Tl
Pb
Bi
Po
At
Rn
7
Fr
Ra
Period
H
Metal
Metalloid
Nonmetal
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a result, metals form closely packed structures, most often surrounded by 12 other like
atoms. Although these interactions have some covalent character, the atoms in a metallic
solid are not linked by traditional covalent bonds.
On the left side of the periodic table, atoms of a given element have few valence electrons but many orbital vacancies. As such, a single atom can share its electrons with many
other atoms because those other atoms have the orbital vacancies required to accept those
shared electrons. In the same way, atoms lower on the periodic table are more metallic. As
you move down a group in the periodic table, electrons in the highest-filled orbitals are
higher in energy and show much lower attraction to the nucleus. They can be shared more
readily, and the elements have greater metallic character.
The molecular orbital approach best describes these interactions (Figure 22.1.4). Recall
from The Solid State (Unit 12) that a set of valence orbitals from a set of bonded metal atoms
leads to a set of molecular orbitals spread over the metal structure. When a sample of metal
has many atoms, the structure has a large number of molecular orbitals over a range of energies. These closely spaced orbitals are called a “band” of orbitals, and in metals, the band of
orbitals is only partially filled with electrons. That is, lower-energy orbitals in the band are
occupied, whereas higher-energy orbitals are vacant.
The combination of closely spaced orbitals in a partially filled band leads to the highestfilled orbital being very closely spaced energetically to the lowest-energy vacant orbital.
Energy
The main-group metals, nonmetals, and metalloids
Half-filled
band of
N molecular
orbitals
3s
atomic
orbitals
(AO)
Na
Na2
Na3
Na8
NaN
Figure 22.1.4 Formation of bands of molecular
orbitals in a solid
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This in turn means that electrons need gain only a very small amount of energy to move
from one orbital to another. This small energy barrier to electron movement is the reason
metals are good conductors of electricity.
Formation of Cations
Unlike metals, elemental forms of the nonmetals consist of either monoatomic species (the
noble gases) or atoms covalently bonded to one another. Some of the main-group elements
form structures containing multiply bonded atoms, whereas others have only single bonds
between atoms. Also, some elements form small, discrete molecules, and others form extended structures.
Two factors that control the molecular structure of the elements are the number of orbital
vacancies and the ability to form pi bonds with neighboring atoms. Interactive Table 22.1.1
and Figure 22.1.6 show the number of orbital vacancies, the number of bonds needed to form
to fill each atom’s octet, for the main-group nonmetals.
The number of possible bonds an atom of a given element can form plays an important role in determining the element’s structure. Starting with the Group 8A noble gases,
we see that these elements do not form molecules in their elemental state; they exist as
monatomic species because they have no valence orbital vacancies. The Group 7A elements, the halogens, have a single orbital vacancy and thus need only one bond to complete the octet. The elemental form of all halogens is diatomic molecules with a single
X—X bond.
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Explore the relationship between effective
nuclear charge and the formation of ions.
8
7
6
5
4
3
2
1
0
Na Mg
Al
Si
P
S
Cl
22.1c Nonmetals
Effective nuclear charge
Although metal atoms have multiple valence orbital vacancies, they tend to lose electrons
in their interactions with other atoms, forming cations. This is due to orbital energy
trends across each period in the periodic table. As you move across a period from left
to right, the nuclear charge of the elements increases. As electrons are added to the valence shell, they repel one another, increasing shielding. The effective nuclear charge, the
sum effects of nuclear positive charge attracting electrons and shielding electrons repelling valence electrons, increases as you move left to right because each added electron
does not effectively shield the full +1 charge of the added proton. As the effective nuclear
charge increases, the attraction of the outermost electrons to the nucleus increases and
the energy of the valence orbital decreases. Therefore, the highest-energy valence orbitals are found toward the left side of each period. Electrons found in these orbitals are
most easily lost, which is why metals tend to form cations (Interactive Figure 22.1.5).
Ar
Sodium has a low effective nuclear charge and therefore
high orbital energies. Argon has a high effective nuclear
charge and therefore low orbital energies.
Interactive Table 22.1.1
Orbital Vacancies for the Main-Group
Nonmetals
Group Number
4A
5A
6A
7A
8A
Number of Orbital
Vacancies
4
3
2
1
0
Number of Bonds
Needed to
Complete Octet
4
3
2
1
0
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2p
1s
2s
2p
1s
C
2s
2p
1s
N
2s
2p
1s
O
2s
2p
1s
F
2s
Ne
Figure 22.1.6 Electron configurations for the second-row main-group nonmetals
Element structure becomes more complex as we move to Groups 4–6A. For example,
oxygen has two allotropes, O2 and O3, each of which consists of small molecules containing
O=O double bonds. Sulfur, on the other hand, has multiple allotropes based on rings of
singly bonded S atoms and a polymeric form containing long chains of singly bonded S atoms.
Why does O form small molecules when S forms larger extended structures? The answer
lies in the size of the atoms and their ability to form pi bonds.
Bonds form when orbitals on adjacent atoms overlap. Sigma bonds can form between atoms
of any size because the bonds are the result of orbital overlap between the two atoms, along
the internuclear axis. Pi bonds, on the other hand, are the result of side-to-side overlap of
p or d orbitals. If the two atoms forming the pi bond are large, orbital overlap is not significant and pi bonds do not form.
Consider the Group 6A elements oxygen and sulfur. Recall that atomic radii increase
moving down a group. Sulfur atoms are larger than oxygen atoms, and as a result, pi bonds
form between oxygen atoms in elemental oxygen (O2 and O3) but not between sulfur atoms
in elemental sulfur (Figure 22.1.7).
In terms of thermodynamics, we can consider the effect of both enthalpy (bond energies) and entropy (disorder) on the structure of an element. Entropy will favor the formation of a large number of molecules; thus the formation of many smaller molecules of
an element is entropy favored. Enthalpy favors molecules with stronger bonds. In cases
where relatively strong pi bonds can form, which results in smaller molecules, entropy
“wins” and small, multiply bonded molecules are formed. When pi bonds are weak, enthalpy “wins” and extended structures containing only sigma bonds are formed. For oxygen and sulfur, the fact that oxygen forms O—O pi bonds results in small molecules.
Sulfur cannot easily form strong pi bonds with other sulfur atoms due to its larger atomic
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Charles D. Winters
Atomic Radius and Bonding in Nonmetals
Figure 22.1.7 Elemental forms of sulfur:
S8 and polymeric sulfur
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Charles D. Winters
Figure 22.1.8 Two allotropes of phosphorus
(white and red)
Section 22.1 Mastery
Charles D. Winters
radius, so in its elemental state, it exists as larger extended structures with only S—S
single bonds.
A similar situation exists in the elemental forms of the Group 5A elements. Nitrogen
is small, can form N—N pi bonds, and thus exists as small triply bonded N2 molecules.
Phosphorus has a larger atomic radius, cannot easily form strong P—P pi bonds, and
thus forms only P—P single bonds in its elemental structures. Phosphorus exists as multiple allotropes, the three most common of which are white (P4 molecules), red (polymeric chains of P4 monomers), and black (complex cross-linked polymeric structure)
(Figure 22.1.8).
Finally, the Group 4A elements continue the relationship between elemental form
and atomic radius. Carbon has a smaller atomic radius than silicon and can form C—C pi
bonds. Elemental forms of carbon include graphite (which contains both C—C sigma and
pi bonds), diamond (C—C sigma bonds only), and fullerenes such as buckyball (both
C—C sigma and pi bonds) (Figure 22.1.9). Elemental silicon is a network solid similar to
diamond that consists of a three-dimensional array of silicon atoms connected by Si—Si
sigma bonds.
(a)
(b)
(c)
Figure 22.1.9 Three allotropes of carbon: (a) graphite, (b) diamond,
and (c) buckyball
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22.2 Oxides and Halides of the Nonmetals
22.2a Nonmetal Oxides
The oxides of the nonmetals (Interactive Table 22.2.1) follow structural trends similar to
those found in the elements themselves. As we saw in the structures of the elements, atomic
radius and orbital vacancy both have a significant effect on the types of bonds formed by atoms of a given element. For example, the Group 7A elements have a single orbital vacancy
and generally form oxides with the formula OX2 (X = halogen). These are all small molecule
oxides, favored both in terms of bond energies (enthalpy) and multiple small molecules
(entropy).
The Group 6A elements sulfur and selenium both form small molecule oxides, EO2 and
EO3 (E = S, Se). In these molecules, the Group 6A element forms pi bonds to the smaller
oxygen atoms. In Group 5A, nitrogen forms both sigma and pi bonds to O and thus forms
many different small oxides. Phosphorus, however, has a larger radius and thus its compounds with oxygen tend to contain primarily P—O single bonds. Phosphorus forms a
Selected Oxides of the Nonmetals
Group 4A
Period 2
O
C
Group 5A
O
O
Group 6A
Group 7A
O
O
N
Group 8A
F
N
F
O
Si
O
O
O
O
O
P
O
n
Period 4
P
O
O
P
Cl
S
P
O
Cl
O
O
O
Se
O
Period 5
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O
Br
O
Br
O
Xe
O
O
Period 3
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O
As
O
O
P
O
O
P
O
O
P
O
As
O
P
O
O
As
O
O
O
As
O
As
O
O
O
O
As
O
O
Se
O
O
O
O
O
Se
O
Se
O
O
O
O
series of oxides that contain bridging oxygen atoms in P—O—P groups and terminal atoms
in P—O bonds that have some double-bond character. Similar oxide structures are found
for the larger nonmetals arsenic and selenium (Figure 22.2.1).
Finally, the Group 4A nonmetals carbon and silicon also follow this trend in oxide structure based on atomic radius. Carbon forms two molecular oxides, CO and CO2, both of
which contain C—O pi bonds and are entropy favored due to their small size. Silicon, however, forms an oxide containing only Si—O single bonds. Silicon dioxide, SiO2, has a formula
analogous to carbon dioxide, but it does not consist of individual SiO2 molecules. Silicon
dioxide (silica) is a covalently bonded network solid consisting of interlocking SiO4 tetrahedra. The larger silicon atoms cannot effectively form pi bonds with oxygen (Figure 22.2.2).
Figure 22.2.1 Structures of some nonmetal oxides containing
larger nonmetals
Figure 22.2.2 The structure of SiO2
Acid–Base Properties of Oxides
Metal oxides are generally basic oxides (they react with water to form bases or act as
bases in chemical reactions), and nonmetal oxides are generally acidic oxides (they react
with water to form acids or act as acids in chemical reactions). Again, we can explain these
experimental observations using orbital energies.
For example, consider what happens when magnesium oxide reacts with water.
MgO(s) + H2O(ℓ) S Mg(OH)2(s)  Mg2+(aq) + 2 OH−(aq)
Main-group metals have high-energy valence orbitals, so reaction with water leads to an
ionic hydroxide salt (a base).
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Nonmetal oxides react with water to form acids. For example, consider the reaction of
carbon dioxide with water to form carbonic acid.
CO2(g) + H2O(ℓ)  H2CO3(aq)
H2CO3(aq) + H2O(ℓ)  HCO3−(aq) + H3O+(aq)
There are two reasons this reaction results in the formation of carbonic acid. First, carbon
is a nonmetal with orbital vacancies that are relatively low in energy. Oxygen is very electronegative and electron withdrawing, which pulls electron density away from the carbon,
making it open to attack by the electronegative oxygen in the water molecule. Finally, in
carbonic acid, the electronegative oxygen atoms pull electron density away from hydrogen,
making it easy to remove (as the hydronium ion, H3O+).
22.2b Nonmetal Halides
All main-group elements form compounds with halogens, including the halogens themselves. (Compounds containing only halogens are called interhalogen compounds.) The
stoichiometry of nonmetal halide compounds is related to orbital vacancies and the maximum oxidation state of the nonmetal combined with the halogen.
As shown in Table 22.2.2, the second-row nonmetals in Groups 4–6A form halide compounds with stoichiometries based on orbital vacancies. These compounds contain only
sigma bonds to halogen atoms, and the structures of the compounds follow VSEPR theory.
Larger elements, those in periods 3 and higher, can have expanded valence structures. For
example, Table 22.2.3 shows some of the halides of the third-row nonmetals. The compounds formed are those expected based on orbital valencies, except that phosphorus,
sulfur, and chlorine can also have an expanded valence. For example, phosphorus forms
both PX3 and PX5 compounds with fluorine, chlorine, and bromine, and the PX3 compound
with iodine. In addition, compounds with the formula P2X4 are also known for X = F, Cl, Br,
and I. Sulfur combines with fluorine to make SF2, SF4, SF6, and S2F10. However, with the
exception of SCl2, it does not form analogous compounds with Cl, Br, or I (probably due to
the weaker S—X bonds).
Recall that the tendency to form expanded valence structures begins in the third period of the periodic table because these elements can make use of vacant d orbitals to form
bonds. The tendency increases as you move left to right in the periodic table because the
energy of orbitals decreases, making those vacant orbitals more attractive to bonding electrons. Thus, phosphorus forms many compounds with an expanded valence, but Si does not
(although it does form the ionic species SiF5− and SiF62−).
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Table 22.2.2 Halides of the Second-Row
Group 4–6A Nonmetals
4A
5A
6A
CX4 and other extended halogenated
organics (e.g., C3F8)
NF3, NCl3,
NBr3, NI3
OF2, OCl2,
OBr2
Table 22.2.3 Selected Halides
of the Third-Row Group 4–7A Nonmetals
4A
5A
6A
7A
SiF4
PF3, PF5
SF2
ClF, BrCl, ICl
SiCl4
PCl3, PCl5
SF4, SF6
ClF3, ClF5
SiBr4
PBr3, PBr5
S2F10
SiI4
PI3
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Melting Points of Selected Main-Group Halides
1A
2A
3A
4A
5A
6A
7A
Compound
NaF
MgF2
AlF3
SiF4
PF3
SF4
ClF
Melting point, °C
993
1263
1291
290
2151
2121
2155
If we expand our view out further to the left in the periodic table, we see an interesting
effect. Interactive Table 22.2.4 shows the melting points of the most common fluoride compound for the third-row main-group elements. Notice the large break between Group 3A
and Group 4A when we move from Al to Si. This sharp change in melting point is due to the
change from ionic compounds, which have strong ion–ion attractions and high melting
points, to molecular covalent compounds, which have only relatively weak intermolecular
forces. The compounds NaF, MgF2, and AlF3 are ionic, whereas those to the right in the
table are covalent (molecular).
22.3 Compounds of Boron and Carbon
22.3a Boron Compounds
Boron has the electron configuration [He]2s22p1; it has five valence orbital vacancies.
However, this does not translate into compounds with five bonds to boron. Instead, boron
tends to form compounds with the general formula BY3, where the boron atom is bonded to
three other atoms or groups. For example, in BH3, the central B atom has six valence electrons and therefore two vacancies. Because the boron is electron deficient in this compound, BH3 and compounds like it can act as Lewis acids, accepting a pair of electrons from
a Lewis base such as NH3 (Figure 22.3.1).
Other compounds of boron show similar reactivity. For example, consider boric acid,
B(OH)3. The chemical formula of boric acid is similar to the ionic hydroxide salts of the
Group 1A and 2A metals such as NaOH and Mg(OH)2. These ionic compounds are basic;
they dissociate in water to produce metal ions and hydroxide ions.
NaOH(s) S Na+(aq) + OH−(aq)
Mg(OH)2(s)  Mg2+(aq) + 2 OH−(aq)
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H
H
B
H
1
H
Lewis acid
H
N
H
H
Lewis base
H
H
B
H
N
H
H
Lewis acid-base
adduct
Figure 22.3.1 Reaction between BH3 and NH3
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H O
H
OH
B (aq)
HO
OH
OH
H2O
HO
B
OH
OH
2
(aq) 1 H3
O1(aq)
Explore bonding in sodium perborate.
Ka 5 7.3 3 10210
Structure of the perborate ion, B2(O2)2(OH)42−
(boron = purple, oxygen = red)
Explore stability and metastability.
Sodium Perborate
7 x 105
Solid III
6 x 105
5 x 105
Pressure (atm)
One of the most commercially important boron compounds is sodium perborate,
NaBO3 # 4H2O, which is the active ingredient in color-safe bleach. The structure of the
perborate ion (Interactive Figure 22.3.3) shows that it is actually a dimer with peroxide
groups linking the two boron atoms.
The active ingredient in bleach, sodium hypochlorite (NaClO), is added to laundry to
kill bacteria and destroy dyes, chemical compounds that are used to color fabrics. Colorsafe bleach contains the perborate ion because at high temperatures the ion decomposes
to produce, among other products, hydrogen peroxide. Hydrogen peroxide can remove
stains but, at the concentrations found in a typical wash containing color-safe bleach, it will
not destroy dye molecules. Color-safe bleaches also contain brightening agents that adhere
to fabrics and make them look brighter.
Diamond
4 x 105
Liquid
3 x 105
2 x 105
Diamond and
metastable
graphite
1 x 105
Graphite and Gr
metastable diamond
22.3b Elemental Carbon
Interactive Figure 22.3.4 shows a phase diagram for carbon under a wide range of temperatures and pressures. At room temperature and 1 bar, graphite is more stable than diamond.
C(s, diamond) S C(s, graphite)
∆Go = −2.9 kJ/mol
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0
1000
2000
3000
4000
Temperature (K)
5000
Boric acid, as its name implies, is an acid, not a base. It acts as a Lewis acid in water,
forming borate and hydronium ions (Figure 22.3.2). Why does B(OH)3 act differently than
metal hydroxides? The answer is once again orbital energies. Recall that as you move from
left to right across the periodic table, orbital energies decrease due to increasing effective
nuclear charge. Group 1A elements have valence s orbitals of very high energy that can
easily lose an electron to form a cation. Elements further to the right have lower orbital
energies and decreased ability to form a cation. Boron has relatively low valence orbital
energies and, instead of losing electrons, acts as an electron-pair acceptor.
Figure 22.3.2 Boric acid is a Lewis acid in water.
The phase diagram of carbon
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Under typical conditions, however, we know diamond does not convert into graphite. If it
did, expensive jewels would be converted into pencil lead. The process does not occur because diamond is kinetically stable. That is, even though carbon is more stable as graphite,
the activation energy for the conversion is so large that the reaction just does not occur.
Species such as diamond are called metastable, meaning they are thermodynamically unstable but highly kinetically stable.
Synthetic diamonds can be made from graphite, and a large industry exists that is
dedicated to this process. Two common methods are used to convert the more stable form
of carbon, graphite, into diamond. One method puts carbon under conditions where
diamond is more stable than graphite (high pressure, high temperature). The problem with
this idea is that even under fairly extreme conditions where diamond is more stable than
graphite, the graphite is metastable and reacts slowly, if at all. The key is to use a catalyst,
commonly a metal such as chromium, iron, or nickel. The graphite dissolves into the molten
metal, and experimental conditions are applied that favor the formation of diamond.
C(s, graphite)  C(s, dissolved in molten metal)  C(s, diamond)
A second method for synthesizing diamonds from graphite (or any carbon-containing
material) is chemical vapor deposition, where a carbon-containing material is heated to
about 1000 °C at relatively low pressures (∼50 mbar). Excess hydrogen atoms are used
to dilute the carbon and prevent the formation of graphite, and the gas-phase material
condenses into a thin film of diamond on a surface.
The chemistry of carbon compounds (specifically carbon–hydrogen compounds) was
discussed in Organic Chemistry (Unit 21). Carbon also has interesting inorganic chemistry, particularly the chemistry of carbon–oxygen compounds. Carbonate chemistry, and
particularly the chemistry involving carbonic acid in aqueous solution, is very important
in biology, environmental chemistry, and the chemical industry. Here, we explore the
carbonate equilibria that exist in caves such as the one shown in Interactive Figure 22.3.5.
Caves are formed when large bodies of limestone rock (CaCO3) are dissolved. Calcium
carbonate is an insoluble inorganic compound with a small Ksp.
CaCO3(s)  Ca2+(aq) + CO32−(aq)
Explore cave chemistry.
salajean/Shutterstock.com
22.3c Cave Chemistry
Ksp = 3.8 × 10−9
solubility 5 "3.8 3 1029 5 6.2 3 1025 mol/L 5 0.0062 g/L
However, the solubility of calcium carbonate in natural conditions is about 80 times greater
than its solubility in pure water. We can explain this by exploring the effect of carbon dioxide on carbonate compound solubility.
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Caves are large underground chambers
that usually contain limestone formations.
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Carbon dioxide reacts with water to form carbonic acid.
CO2(g) + H2O(ℓ)  H2CO3(aq)
When calcium carbonate is exposed to water with a high carbonic acid concentration, it
dissolves.
H2CO3(aq) + CaCO3(s)  Ca2+(aq) + 2 HCO3−(aq)
The equilibrium constant for this reaction can be determined by adding reactions for which
we know the equilibrium constants:
CO32−(aq) + H2O(ℓ)  HCO3−(aq) + OH−(aq)
Kb(CO32−) = 2.1 × 10−4
H2CO3(aq) + H2O(ℓ)  HCO3−(aq) + H3O+(aq)
Ka(H2CO3) = 4.2 × 10−7
H3O+(aq) + OH−(aq)  2 H2O(ℓ)
1/Kw = 1.0 × 1014
H2CO3(aq) + CaCO3(s)  Ca2+(aq) + 2 HCO3−(aq)
K = Ksp × Ka × Kb × 1/Kw = 3.4 × 10−5
Although the equilibrium constant for the net reaction is not large (not greater than 1),
it is about 104 times greater than Ksp. The equilibrium is pushed further to the right by large
concentrations of carbonic acid (and other organic acids) in soil.
Cave formations such as stalactites and stalagmites result from the reverse of the dissolution of calcium carbonate in the presence of carbonic acid. Consider a drop of water,
saturated with CaCO3, hanging inside a cave. The precipitation of CaCO3 is not due to water
evaporation; the air inside most caves is saturated with water vapor, and evaporation happens little, if at all. Instead, deposition of CaCO3 is again dependent on carbonate equilibria
involving CO2.
When a cave is forming, the partial pressure of CO2 in the cave air is very high, much
higher than in the outside air, where CO2 pressure is nearly zero. The CaCO3 is much
more soluble in that acidic environment. When the cave is open to the outside air, the CO2
pressure in the cave air drops to near zero. When this happens, the amount of H2CO3 dissolved in the cave water also decreases, as does the solubility of CaCO3. Cave formations
occur due to degassing of CO2 out of droplets saturated with CaCO3.
22.3d Carbon Dioxide and Global Warming
The greenhouse effect is caused by gases in Earth’s atmosphere that trap energy that would
otherwise escape Earth. All objects give off blackbody radiation, the wavelength of which
depends on the temperature of the body. The sun gives off UV and visible radiation; Earth,
being cooler, gives off infrared (IR) radiation (Figure 22.3.6).
18995_ch22_rev01.indd 748
90
80
70
60
50
40
Ksp = 3.8 × 10−9
Radiative flux (in W/m2/mm)
CaCO3(s)  Ca2+(aq) + CO32−(aq)
30
20
10
0
Sun
(scaled by
a factor of 10-6)
0.2
Earth
0.4 0.7 1
2 3 45 7 9
20 30 50 80
Wavelength (l) in mm
Figure 22.3.6 Blackbody radiation curves
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Our planet is warmed by the radiation from the sun. Our atmosphere contains greenhouse gases, gases that can absorb IR radiation given off by Earth and send it back to
Earth. This greenhouse effect warms the surface of our planet. Without greenhouse
gases, the average temperature of Earth would be much lower than the actual average
temperature.
We can quantify the impact of the greenhouse effect on the temperature of a planet by
calculating the temperature of a planet such as Earth in the absence of greenhouse gases
(Equation 22.1).
TE 5
4
Å
Es 11 2 a2
4s
(22.1)
In Equation 22.1, TE is the equilibrium average surface temperature, Es is the solar constant (the amount of solar light that reaches the planet), α is the fraction of sunlight
reaching the planet that is reflected before reaching the planet surface, and σ is a constant (5.67 × 10−8 W/m2K4).
For Earth, Es = 1380 W/m2 and α = 0.28. Solving for TE,
TE 5
11380 W/m22 11 2 0.282
5 257 K 1216° C2
Å 4 15.67 3 1028 W/m2K42
4
Earth’s average temperature is about 15 ºC, approximately 31 degrees warmer than what
the average temperature would be without the greenhouse effect.
We can also calculate the fraction of radiation that leaves a planet but is trapped by
greenhouse gases and reflected back to the planet surface. This quantity, called the IR
transmission factor ( f ), is a measure of the fraction of IR light that escapes the planet’s
atmosphere (Equation 22.2).
Es
11 2 a2
4
(22.2)
f5
s 1TE2 4
To calculate this quantity for our planet, we use the actual measured temperature for TE
(288 K).
1380 W/m2
11 2 0.282
4
5 0.64
f5
15.67 3 1028 W/m2K42 1288 K2 4
This means that about 36% of the IR radiation leaving Earth’s surface is trapped by greenhouse gases and radiated back to Earth.
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Explore molecular vibrations.
O
O
O
C
C
C
O
O
O
Symmetric stretching
Bending
Asymmetric stretching
(one bond stretches while
the other compresses)
Atmospheric compounds that are “greenhouse gases” absorb IR radiation when they
can excite a vibration in the molecule that leads to a change in the molecule’s dipole moment, changing its polarity. The three most abundant gases in the atmosphere, N2, O2, and
Ar, are nonpolar and do not absorb IR light; they are not greenhouse gases. The two most
important atmospheric gases that act as greenhouse gases are water, H2O, and carbon
dioxide, CO2. Although CO2 is linear and nonpolar, it has vibrational modes that cause it to
become polar. The three vibrational modes of CO2 are shown in Interactive Figure 22.3.7.
Because two of these modes (bending and asymmetric stretching) result in a change in
dipole moment, CO2 absorbs IR radiation and is a greenhouse gas.
Vibrational modes for CO2
22.4 Silicon
22.4a Silicon Semiconductors
Si(s) + 2 Cl2(g) S SiCl4(ℓ)
The reason for forming SiCl4 is that it has a relatively low boiling point and can be purified
by distillation to a purity of greater than 99%.
Next, the SiCl4 is reduced with zinc.
SiCl4(ℓ) + 2 Zn(s) S Si(s) + 2 ZnCl2(s)
Finally, the Si is cast into a cylindrical shape and highly purified by zone refining (Interactive
Figure 22.4.1). The zone refining process is based on the same principle that causes the solvent
in a solution to solidify in its pure form as temperature is lowered. When a solution freezes, the
solvent forms a solid that does not contain any particles of any solute. Because of this, the solute
(in this case the impurities in the Si) stay in solution. Zone refining works by using radio waves to
heat a thin layer of Si in the cylinder until it melts. The region of melted Si is slowly moved down
the cylinder. As this happens, each area of the Si cylinder melts and then, as the liquid layer
passes, solidifies. The key is that as the liquid layer passes, it takes the impurities away with it. The
impurities dissolve in the solution when the solid melts but stay in solution when it refreezes.
18995_ch22_rev01.indd 750
Explore the zone refining process.
Impure solid
Impurities are
concentrated
here
Molten zone
Heater
Purified
solid
Silicon is the basis of the huge semiconductor industry. Interestingly, in a typical room,
the Si in a silicon computer chip is the purest substance present, but it functions as a semiconductor precisely because it is impure. The key to the function of a silicon semiconductor
is that the impurities are very specific and have very low concentrations.
Semiconductor chips are formed by first purifying Si to a very high purity of 99.9999%.
The first step is to obtain SiCl4 from relatively pure silicon.
Diagram showing the zone refining process for silicon
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04/06/14 8:51 AM
The pure Si surface is then exposed to either a Group 3A element to form a p-type
semiconductor or a Group 5A element to form an n-type semiconductor, as described in
The Solid State (Unit 12).
22.4b Silicates
Most of the minerals in Earth’s crust are silicates, silicon compounds that contain extensive
networks of silicon–oxygen bonds. These compounds form the basis of all rocks and rock
chemistry. Unlike carbon, whose chemistry is dominated by compounds containing chains
and rings of carbon atoms, the chemistry of silicon is dominated by Si—O bonds. The relatively weak Si—Si bond energy (Si—Si, 222 kJ/mol; Si—O, 452 kJ/mol) means that the same
types of structures found in carbon compounds are just not found in stable silicon
compounds.
The structures of silicates follow the following rules:
1. Every Si is bonded to four O atoms.
2. Oxygen atoms can either be bonded to a single Si atom (terminal O) or bridge between
two Si atoms (bridging O).
For example, consider Mg2SiO4. The silicon oxygen unit is the silicate ion, SiO44−. The
charge on the ion is the sum of the silicon and oxygen oxidation states [Si + 4 × (O) =
(+4) + 4 × (−2) = −4]; the structure of the ion is tetrahedral, and all oxygen atoms are
terminal (Figure 22.4.2).
Now consider the Si2O76− ion, which consists of two [SiO4] tetrahedra sharing a single
(bridging) oxygen. The charge on the ion is the sum of the silicon and oxygen oxidation states
[2 × (Si) + 7 × (O) = 2 × (+4) + 7 × (−2) = −6]; the geometry is tetrahedral around each
silicon atom, and all but one oxygen atom are terminal (Figure 22.4.3).
As the O:Si ratio decreases, more O atoms bridge between Si atoms in order to maintain four oxygen atoms around each silicon atom. Interactive Table 22.4.1 shows the types
of formulas and structures of the various silicates as a function of O:Si ratio.
As the ratio decreases from 4 O:1 Si to 2 O:1 Si, the silicate structures change from
monomeric [SiO4] units to dimers, to chains and rings, to double chains, to two-dimensional
sheets, and finally to three-dimensional networks.
18995_ch22_rev01.indd 751
Figure 22.4.2 The SiO44− ion
© 2013 Cengage
Learning
4. The charge on a silicon–oxygen unit is determined using oxidation numbers of
+4 for Si and −2 for O.
© 2013 Cengage
Learning
3. The ratio of O atoms to Si atoms in a structure controls how branched and interlinked
the structure is.
Figure 22.4.3 The Si2O76− ion
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04/06/14 8:51 AM
Silicate Structures
Structure
Corners Shared
at Each Si Atom
Repeating Unit
Si:O Ratio
Example
Tetrahedra
0
SiO44−
1:4
Olivines
Pairs of tetrahedra
1
Si2O76−
1:3.5
Thortveitite
Closed rings
2
SiO32−
1:3
Beryl
Infinite single chains
2
SiO32−
1:3
Pyroxenes
Infinite double chains
2.5
Si4O116−
1:2.75
Amphiboles
Infinite sheets
3
Si2O52−
1:2.5
Talc
Infinite network
4
SiO2
1:2
Quartz
22.4c Silicones
Most Americans first encounter silicones when they play with Silly Putty as a child. Silly
Putty, a compound made by mistake during research into synthetic rubber during World
War II, has as its main ingredient a silicone polymer. Modern uses of silicones include
18995_ch22_rev01.indd 752
Charles D. Winters
Because of the wide variation in structure and arrangement of the [SiO4] units, silicates
have a wide range of physical and chemical properties. For example, the single- and doublechain structures result in fibrous minerals such as asbestos (Figure 22.4.4).
Because all silicates (with the exception of SiO2) are anions, they form ionic compounds with various metal ions. Some of these, those with relatively high O:Si ratios, are
traditional ionic compounds with relatively small silicate anions whose negative charge is
balanced by the positive charges of the cations. Those with lower O:Si ratios have extended,
negatively charged Si—O chains or networks, along with accompanying cations. These
compounds tend to have layered structures such as those found in clays such as kaolinite
and montmorillonite (Figure 22.4.5).
Silica, the neutral silicon–oxygen compound with the empirical formula SiO2, is a network solid consisting of a three-dimensional network of interlocking [SiO4] tetrahedra
(Figure 22.4.6).
Figure 22.4.4 A sample of the mineral asbestos
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1 (Na) + x (H2O)
9 (O)
6 (Si)
5 (Al) + 1 (Mg)
© Cengage Learning
3 (OH) + 6 (O)
9.6 – 21.4 Å
3 (OH) + 6 (O)
O
Si
Na
Al
Layer of water
Figure 22.4.6 A small portion of the silica
(SiO2) structure
Mg
H
6 (Si)
9 (O)
Figure 22.4.5 A small portion of the montmorillonite structure
lubricants, adhesives, insulation, and medicine (implants); their wide use is related to the
desirable properties of silicones such as chemical inertness, lack of reactivity toward UV
light, high gas permeability, and low toxicity.
Silicones combine the high stability of silicates with the hydrophobic nature of hydrocarbons. Silicones are small molecules or polymers with a silicon–oxygen backbone and
carbon-containing side groups. The simplest silicone contains only two silicon atoms.
CH3
H3C
Si
CH3
O
Si
CH3
CH3
CH3
Notice that each silicon atom is bonded to both carbon and oxygen. Larger silicones exist,
as do silicone polymers (polysilicones).
R
Si
R
O
n
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Silicones are synthesized from compounds containing relatively reactive Si—Cl bonds.
2 (CH3)3SiCl + H2O S (CH3)3Si—O—Si(CH3)3 + 2 HCl
Reagent
Example
Result
Chain capping
(CH3)3SiCl
(CH3)3Si—O—
Chain extending
(CH3)2SiCl2
CH3
Si
O
Si
O
CH3
As shown in Interactive Table 22.4.2, three types of reagents are used when forming
silicones.
Unbranched silicones are formed by the reaction between a capping reagent such as
(CH3)3SiCl and a chain-extending reagent such as (CH3)2SiCl2. As the reaction progresses,
each Cl is replaced by a bridging oxygen atom. Therefore, each extending unit lengthens
the growing chain by a (CH3)2Si—O— unit. Each time a capping unit reacts, that end of the
chain loses any active Si sites and growth in that direction is halted. The length of the chain
can be controlled by the ratio of capping to extending reagents used. When [extending] >>
[capping], the structures will be relatively long because many extending reactions will occur
before the less concentrated capping unit halts growth. When [extending] is not much
greater than [capping], the chain lengths will be shorter.
Likewise, the use of branching reagents leads to branches and cross-links that form in
the growing structure. The greater the amount of branching reagent used, the more crosslinked the resulting structure. In this way, a wide variety of structures with predictable
properties can be formed. Liquid silicone oils can be produced by using no branching reagents, and their boiling points can be controlled by controlling the chain lengths by altering the [extending]:[capping] ratio. Solid silicones with varying toughness and hardness can
be formed by controlling the amount of branching reagent used. This way, very stable solids
ranging from very soft to very hard and tough can be formed.
Reagents Used to Synthesize Silicones
Chain branching
(CH3)SiCl3
CH3
Si
O
Si
O
O
22.5 Oxygen and Sulfur in the Atmosphere
22.5a Atmospheric Ozone
Ozone, O3, is a highly reactive form of oxygen. In the upper atmosphere (stratosphere),
ozone is beneficial to life on Earth because it blocks short-wavelength, highly energetic UV
radiation. Ozone in the lower atmosphere is toxic and is the principal pollutant that defines
harmful air quality alerts in cities.
Ozone is less stable than dioxygen, O2. Ozone is formed in the upper atmosphere when
highly energetic photons from the sun cause an O2 molecule to split. The free oxygen atoms
(O radicals) formed can react with O2 molecules to form ozone.
O2(g) + photon S 2 O(g)
2 O(g) + 2 O2(g) S 2 O3(g)
Net reaction:
3 O2(g) + photon S 2 O3(g)
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Atmospheric ozone is measured in units called the Dobson, with the symbol DU. One
DU is equivalent to a 10-mm-thick layer of pure ozone at 1 atm pressure and a temperature
of 0 °C. In the 1980s, scientists discovered that the amount of ozone in the atmosphere had
decreased significantly, especially over Antarctica around October each year (Interactive
Figure 22.5.1). This area with significantly decreased ozone is called the ozone hole (even
though it is not really a hole, as the ozone concentration is not zero within the “ozone
hole”). The decrease in stratospheric ozone is primarily due to the effect of free radicals,
compounds that react with stratospheric ozone and convert it to oxygen, specifically the
chlorine free radical.
Chlorine free radicals enter the atmosphere in the form of chlorofluorocarbons, CFCs.
CFCs, compounds containing carbon, fluorine, and chlorine, are a class of compounds developed in the 1890s and put into use as refrigerants, industrial cleaners, and firefighting
materials in the 20th century. They are extraordinarily nontoxic and have just the right
boiling points and enthalpy of vaporization values to make them useful in refrigeration and
as propellants in spray cans.
Long after CFCs were put into commercial use, it was found that they are so stable that
they move into the upper atmosphere without being destroyed. In the upper atmosphere,
a CFC such as Freon-11 (CF3Cl) absorbs high-energy photons from the sun, breaking the
carbon–chlorine bond.
Investigate the mechanism of
ozone depletion.
The Cl radicals are highly reactive, and they react with ozone in a three-step mechanism where Cl is a catalyst.
Cl + O3 S ClO + O2
Courtesy of NASA
CCl3F(g) + photon S CCl2F(g) + Cl(g)
Total Ozone (Dobson Units)
110 220 330 440 550
Ozone concentrations over Antarctica on September 24,
2006. The size of the hole was the largest ever observed.
O3 S O2 + O (in presence of light)
ClO + O S Cl + O2
Net reaction
2 O3 S 2 O2
A single chlorine free radical can destroy up to 105 ozone molecules before it is converted
to a more inert form.
In Antarctic winter, there is no light from the sun hitting the pole, and because of
weather patterns, there is a vortex around the South Pole that prevents the atmosphere
over the pole from mixing with the air at lower latitudes. During winter, the chlorine radicals are captured in “reservoir compounds” such as HCl and chlorine nitrate ClONO2. These
compounds hold the chlorine until the sun returns, which it does each year in the southern
hemisphere spring, in October. At that point a large amount of chlorine radicals are produced and a marked decrease in ozone occurs.
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Ozone depletion has decreased in recent years, as evidenced by a decrease in the size
of the “ozone hole” over Antarctica. This decrease is due primarily to the international agreement on phasing out substances that deplete the ozone layer (the Montreal Protocol, first
signed by participating nations in 1987) and research into the creation of CFC substitutes.
Explore the acid-base properties
of oxides.
22.5b Sulfur and Acid Rain
Acid rain is an environmental problem that was identified as early as the 17th century and was
addressed by government regulations in the 1970s. As we saw earlier in this unit oxides of
nonmetals are acidic. In particular, all coal contains sulfur, sometimes in significant proportions,
in the form of metal sulfide minerals. When coal is burned, largely in coal-fired power plants,
the sulfur is also burned.
S(s) + O2(g) S SO2(g)
If the SO2 is ejected into the atmosphere, it slowly, over a span of days, is further
oxidized.
SO2(g) + ½ O2(g) S SO3(g)
Sulfur trioxide then reacts with water in clouds to form sulfuric acid.
The sulfuric acid dissolves in cloud droplets and when those droplets fall to Earth as rain,
that rain is acidic. As shown in Interactive Figure 22.5.2, acid rain can damage statues,
buildings, and plants, and it is also very damaging to lakes, streams, and the species that
live in these bodies of water.
The solutions to this problem are conceptually simple: use a fuel other than coal (natural gas contains very little sulfur), remove the sulfur from the coal prior to burning, or
capture the SO2 formed after burning but before it escapes the smokestack. This latter
process is called scrubbing.
Scrubbing can be categorized as wet or dry. Both methods often involve the use of
limestone (CaCO3), slaked lime [Ca(OH)2], soda lime [a mixture of Ca(OH)2, KOH, and
NaOH], or soda ash (Na2CO3). In all cases, the scrubber reacts with the SO2 gas to form a
sulfate or sulfite salt that is removed as a solid. For example,
© Vanessa Miles/Alamy
SO3(g) + H2O(g) S H2SO4(aq)
Acid rain can damage limestone statues.
CaCO3(s) + SO2(g) S CaSO3(s) + CO2(g)
Ca(OH)2(s) + SO2(g) S CaSO3(s) + H2O(ℓ)
The calcium sulfite is further oxidized to calcium sulfate, which can be sold as gypsum.
CaSO3(s) + H2O(ℓ) + ½ O2(g) S CaSO4 # H2O(s)
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Unit Recap
Key Concepts
22.1 Structures of the Elements
●
●
●
The physical and chemical properties of metals are related to the fact that they have
multiple orbital vacancies in valence orbitals, low effective nuclear charge, and highenergy valence orbitals (22.1b).
The physical and chemical properties of nonmetals are related to the fact that they
form covalent bonds with other nonmetals, have high effective nuclear charge, and
have low-energy valence orbitals (22.1c).
The structure of elemental nonmetals is related to atomic radius and the ability of atoms of an element to form pi bonds (22.1c).
●
●
●
The type of oxide formed by a nonmetal is related to orbital vacancies and atomic radius (ability to form pi bonds) (22.2a).
Metal oxides are basic because the metals have high-energy valence orbitals; nonmetal
oxides are acidic because of the low-energy vacant orbitals on the nonmetal (22.2a).
The type of halide formed by a nonmetal is related to orbital vacancies, atomic radius
(ability to form pi bonds), and ability to form compounds with expanded octets (22.2b).
22.3 Compounds of Boron and Carbon
●
●
●
●
Boron forms a wide variety of compounds with other nonmetals, many of which are
electron deficient (22.3a).
The most stable elemental form of carbon is graphite (at 1 bar and 25 °C), but synthetic
diamond can be synthesized by a number of techniques (22.3b).
The creation of caves and cave formations is controlled by the effect carbon dioxide
has on the solubility of calcium carbonate in water (22.3c).
Carbon dioxide is an important greenhouse gas, a gas that reflects infrared radiation
back to Earth’s surface (22.3d).
22.4 Silicon
●
Silicon is purified using the zone refining process to a very high purity for use in the
semiconductor industry (22.4a).
18995_ch22_rev01.indd 757
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●
●
●
●
Silicates form the basis of all rocks and rock chemistry and are found in most of the
minerals in Earth’s crust (22.4b).
Silicates contain structures based on SiO4 tetrahedra connected in chains, rings,
sheets, and three-dimensional networks (22.4b).
Silicones are silicon–oxygen polymers with hydrophobic hydrocarbon side groups
(22.4c).
The structure of a silicone can be controlled by using capping, extending, and branching reagents in the synthesis process (22.4c).
22.5 Oxygen and Sulfur in the Atmosphere
●
●
●
●
Ozone in the stratosphere is beneficial to life on Earth because it shields us from harmful UV radiation (22.5a).
Stratospheric ozone concentrations have decreased due to free radicals such as those
formed by CFCs when they enter the upper atmosphere (22.5a).
Burning coal ejects large amounts of sulfur into the atmosphere, which reacts with
water vapor to form sulfuric acid and falls to Earth in the form of acid rain (22.5b).
Scrubbing is used to remove sulfur oxides from plant emissions before they can enter
the atmosphere (22.5b).
Key Equations
TE 5
Es 11 2 a2
Å
4s
4
(22.1)
Key Terms
Es
11 2 a2
4
f5
s 1TE2 4
(22.2)
basic oxide
acidic oxide
Unit 22
Review and Challenge Problems
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Chapter 22: Applying Chemical Principles to the Main Group Elements

Transcription

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