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M O D U L E
1CHEMISTRY OF THE
NONMETALS
N.1 The Nonmetals
N.2 The Chemistry of Hydrogen
N.3 The Chemistry of Oxygen
Chemistry in the World Around Us: The Chemistry of the Atmosphere
N.4 The Chemistry of Sulfur
N.5 The Chemistry of Nitrogen
N.6 The Chemistry of Phosphorus
N.7 The Chemistry of the Halogens
N.8 The Chemistry of the Rare Gases
N.9 The Inorganic Chemistry of Carbon
More than 75% of the known elements have the characteristic properties of metals (seeFigure N.1). They have a metallic luster; they are malleable and ductile; and they conductheat and electricity. Eight other elements (B, Si, Ge, As, Sb, Te, Po, and At) are best de-scribed as semimetals or metalloids. They often look like metals, but they tend to be brit-tle, and they are more likely to be semiconductors than conductors of electricity.
Once the metals and semimetals are removed from the list of known elements, only 17are left to be classified as nonmetals. Six of these elements belong to the family of raregases in Group VIIIA, most of which are virtually inert to chemical reactions. Discussionsof the chemistry of the nonmetals therefore tend to focus on the following elements: H, C,N, O, F, P, S, Cl, Se, Br, I, and Xe.
One way of visualizing the difference between metals, semimetals or metalloids, andnonmetals is the plot of average valence electron energies (AVEE) in Figure N.2. Becausethe AVEE provides a measure of how tightly an atom holds onto its valence electrons, itcan be used to explore the dividing line between the metals and nonmetals. As noted inSection 3.24, elements with AVEE values below 1.06 MJ/mol are metals, whereas thosewith AVEE values above 1.26 MJ/mol are nonmetals. Elements with AVEE values in therange of 1.06–1.26 MJ/mol have properties between those of the metals and nonmetals andare therefore semimetals or metalloids.
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2 NONMETAL
N.1 THE NONMETALSThere is a clear pattern in the chemistry of the main-group metals discussed in Chapter 5:The main-group metals are oxidized in all of their chemical reactions. Aluminum, for ex-ample, is oxidized by bromine when these elements react to form aluminum bromide.
Al2Br62 Al � 3 Br2
Oxidation
Reduction
0 �3 �10
H
Li
Na
K
Rb
Cs
Fr
Be
Mg
Ca
Sr
Ba
Ra
Sc
Y
La
Ac Rf Db Sg Bh Hs Mt
Ti
Zr
Hf
V
Nb
Ta
Cr
Mo
W
Mn
Tc
Re
Fe
Ru
Os
Co
Rh
Ir
Ni
Pd
Pt
Cu
Ag
Au
Zn
Cd
Hg
B
Al
Ga
In
Tl
C
Si
Ge
Sn
Pb
N
P
As
Sb
Bi
O
S
Se
Te
Po
H
F
Cl
Br
I
At
He
Ne
Ar
Kr
Xe
Rn
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Metals
Nonmetals
Semimetals
FIGURE N.1 The elements can be divided into three categories: metals, semimetals,and nonmetals.
Ne
F
O
N
C
B
Ar
Kr
Xe
CI
S
P
Si
Br
Se
As
GeAl
Ga
I
Te
Sb
Sn
In
Be
Mg
Ca
Sr
Li
Na
K
Rb FIGURE N.2 Three-dimensional plot of the average valence electronenergies (AVEE) of the main-group elements versus position in the periodic table.
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NONMETAL 3
The chemistry of the nonmetals is more interesting because these elements can un-dergo both oxidation and reduction. Phosphorus, for example, is oxidized when it reactswith more electronegative elements, such as oxygen.
But it is reduced when it reacts with less electronegative elements, such as calcium.
The behavior of the nonmetals can be summarized as follows.
• Nonmetals tend to oxidize metals.
2 Mg(s) � O2(g) 88n 2 MgO(s)
• Nonmetals with relatively large electronegativities (such as oxygen and chlorine)oxidize substances with which they react.
2 H2S(g) � 3 O2(g) 88n 2 SO2(g) � 2 H2O(g)PH3(g) � 3 Cl2(g) 88n PCl3(l) � 3 HCl(g)
• Nonmetals with relatively small electronegativities (such as carbon and hydrogen)can reduce other substances.
Fe2O3(s) � 3 C(s) 88n 2 Fe(s) � 3 CO(g)CuO(s) � H2(g) 88n Cu(s) � H2O(g)
N.2 THE CHEMISTRY OF HYDROGENHydrogen is the most abundant element in the universe, accounting for 90% of the atomsand 75% of the mass of the universe. But hydrogen is much less abundant on earth. Evenwhen the enormous number of hydrogen atoms in the oceans is included, hydrogen makesup less than 1% of the mass of the planet.
The name hydrogen comes from the Greek stems hydro-, “water,” and gennan, “toform or generate.” Thus, hydrogen is literally the “water former.”
2 H2(g) � O2(g) 88n 2 H2O(g)
Although it is often stated that more compounds contain carbon than any other element,this isn’t necessarily true. Most carbon compounds also contain hydrogen, and hydrogenforms compounds with virtually all the other elements as well.
6 Ca � P4
Oxidation
Reduction
0 �2 �30
2 Ca3P2
P4O10P4 � 5 O2
Oxidation
Reduction
0 �5 �20
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4 NONMETAL
Compounds of hydrogen are frequently called hydrides, even though the name hydrideliterally describes compounds that contain an H� ion. There is a regular trend in the for-mula of the hydrides across a row of the periodic table, as shown in Figure N.3. This trendis so regular that the combining power, or valence, of an element was once defined as thenumber of hydrogen atoms bound to the element in its hydride.
H2
LiH
NaH
KH
RbH
CsH
BeH2
MgH2
CaH2
SrH2
BaH2
B2H6
AlH3
GaH3
CH4
SiH4
GeH4
SnH4
PbH4
NH3
PH3
AsH3
SbH3
BiH3
H2O
H2S
H2Se
H2Te
H2Po
H2
HF
HCl
HBr
HI
HAt FIGURE N.3 Hydrogen combines with every element in the peri-odic table except those in Group VIIIA. The formulas of the hydrides of the main-group elements are shown here.
Hydrogen has three oxidation states, corresponding to the H� ion, a neutral H atom,and the H� ion.
H� � 1s0
H � 1 s1
H� � 1s2
Because hydrogen forms compounds with oxidation numbers of both �1 and �1, manyperiodic tables include the element in both Group IA (with Li, Na, K, Rb, Cs, and Fr) andGroup VIIA (with F, Cl, Br, I, and At).
There are many reasons for including hydrogen among the elements in Group IA. Itforms compounds such as HCl and HNO3 that are analogs of alkali metal compounds suchas NaCl and KNO3. Under conditions of very high pressure, hydrogen has the propertiesof a metal. (It has been argued, for example, that any hydrogen present at the center ofthe planet Jupiter is present as a metallic solid.) Finally, hydrogen combines with a hand-ful of metals, such as scandium, titanium, chromium, nickel, and palladium, to form mate-rials that behave as if they were alloys of two metals.
There are equally valid arguments for placing hydrogen in Group VIIA. It forms com-pounds such as NaH and CaH2 that are analogs of halogen compounds such as NaF andCaCl2. It also combines with other nonmetals to form covalent compounds such as H2O,CH4, and NH3, the way a nonmetal should. Finally, the element is a gas at room temper-ature and atmospheric pressure, like other nonmetals such as O2 and N2.
It is difficult to decide where hydrogen belongs in the periodic table because of thephysical properties of the element. The first ionization energy of hydrogen (1312 kJ/mol),for example, is roughly halfway between the elements with the largest (He 2372 kJ/mol)and smallest (Cs 376 kJ/mol) first ionization energies. Hydrogen also has an electro-negativity (EN � 2.30) halfway between the extremes of neon, the most electronegativeelement (EN � 4.79), and cesium, the least electronegative (EN � 0.66) element. On thebasis of electronegativity, it is tempting to classify hydrogen as a semimetal, as shown inFigure N.4.
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NONMETAL 5
Hydrogen is oxidized by elements that are more electronegative to form compoundsin which it has an oxidation number of �1.
Hydrogen is reduced by elements that are less electronegative to form compounds in whichits oxidation number is �1.
At room temperature, hydrogen is a colorless, odorless gas with a density only one-fourteenth the density of air. Small quantities of H2 gas can be prepared in several ways.
• By reacting an active metal with water.
2 Na(s) � 2 H2O(l) 88n 2 Na�(aq) � 2 OH�(aq) � H2(g)
• By reacting a less active metal with a strong acid.
Zn(s) � 2 HCl(aq) 88n Zn2�(aq) � 2 Cl�(aq) � H2(g)
• By reacting an ionic metal hydride with water.
NaH(s) � H2O(l) 88n Na�(aq) � OH�(aq) � H2(g)
Reduction
Oxidation
2 NaH0 0 �1
2 Na � H2�1
Reduction
Oxidation
2 HCl0 0 �1
H2 � Cl2�1
Ne
F
O
N
C
B
Ar
Kr
Xe
CI
S
P
Si
Br
Se
As
GeAl
Ga
I
Te
Sb
Sn
In
Be
H
Mg
Ca
Sr
Li
Na
K
Rb
FIGURE N.4 Three-dimensional plot of the electronegativities of themain-group elements versus position in the periodic table.
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6 NONMETAL
• By decomposing water into its elements with an electric current.
electrolysis2 H2O(l) 888888n 2 H2(g) � O2(g)
Exercise N.1
Use oxidation numbers to determine what is oxidized and what is reduced in the follow-ing reactions, which are used to prepare H2 gas.(a) Mg(s) � 2 HCl(aq) n Mg2�(aq) � 2 Cl�(aq) � H2(g)(b) Ca(s) � 2 H2O(l) n Ca2�(aq) � 2 OH�(aq) � H2(g)
Solution
(a) Magnesium metal is oxidized in this reaction and the H� ions from hydrochloric acidare reduced.
(b) Calcium metal is oxidized in this reaction and the H� ions from water are reduced.
The covalent radius of a neutral hydrogen atom is smaller than any other element. Be-cause small atoms can come very close to each other, they tend to form strong covalentbonds. H2 therefore tends to be unreactive at room temperature. In the presence of a spark,however, a fraction of the H2 molecules dissociate to form hydrogen atoms that are highlyreactive.
sparkH2(g) 888n 2 H(g) �H° � 435.30 kJ/molrxn
The heat given off when these H atoms react with O2 is enough to catalyze the dissocia-tion of additional H2 molecules. Mixtures of H2 and O2 that are infinitely stable at roomtemperature therefore explode in the presence of a spark or flame.
N.3 THE CHEMISTRY OF OXYGENOxygen is the most abundant element on this planet. The earth’s crust is 46.6% oxygen bymass, the oceans are 86% oxygen by mass, and the atmosphere is 21% oxygen by volume.The name oxygen comes from the Greek stems oxys, “acid,” and gennan, “to form orgenerate.” Thus, oxygen literally means the “acid former.” The name was introduced by
Ca2� � 2 OH� � H2
Oxidation
Reduction
0 �2 0�1
Ca � 2 H2O
Mg2� � 2 Cl� � H2
Oxidation
Reduction
0 �2 0�1
Mg � 2 HCl
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NONMETAL 7
Lavoisier, who noticed that compounds rich in oxygen, such as SO2 and P4O10, dissolve inwater to give acids.
SO2(g) � H2O(aq) 88n H2SO3(aq)P4O10(s) � 6 H2O(aq) 88n 4 H3PO4(aq)
The electron configuration of an oxygen atom—[He] 2s2 2p4—suggests that O atomscan achieve an octet of valence electrons by sharing two pairs of electrons to form an OPOdouble bond, as shown in Figure N.5.
According to the Lewis structure, all of the electrons in the O2 molecule are paired. Thecompound should therefore be diamagnetic—it should be repelled by a magnetic field. Ex-perimentally, O2 is found to be paramagnetic—it is attracted to a magnetic field. This canbe explained using molecular orbital theory, which predicts that there are two unpairedelectrons in the �* antibonding molecular orbitals of the O2 molecule.
At temperatures below �183°C, O2 condenses to form a liquid with a characteristiclight blue color that results from the absorption of light with a wavelength of 630 nm. Thisabsorption isn’t seen in the gas phase and is relatively weak even in the liquid because itrequires that three bodies—two O2 molecules and a photon—collide simultaneously, whichis a very rare phenomenon, even in the liquid phase.
The Chemistry of Ozone
The O2 molecule isn’t the only elemental form of oxygen. In the presence of lightning oranother source of a spark, O2 molecules dissociate to form oxygen atoms.
sparkO2(g) 888n 2 O(g)
The O atoms can react with O2 molecules to form ozone, O3.
O2(g) � O(g) 88n O3(g)
Ozone is a resonance hybrid of two Lewis structures each of which contains one OPOdouble bond and one OOO single bond, as shown in Figure N.6. Because the valence elec-trons on the central atom are distributed toward the corners of a triangle, the O3 moleculeis angular or bent, with a bond angle of 116.5°.
FIGURE N.5 Lewis structure of the O2 molecule.
O
O O O O
O
FIGURE N.6 Lewis structures for ozone.
O O
When an element exists in more than one form—such as oxygen (O2) and ozone (O3)—the different forms of the element are called allotropes (from a Greek word meaning “inanother manner”). Because they have different structures, allotropes have different chem-ical and physical properties, as shown in Table N.1.
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Ozone is an unstable compound with a sharp, pungent odor that slowly decomposes tooxygen.
2 O3(g) 88n 3 O2(g)
At low concentrations, ozone can be relatively pleasant. (The characteristic clean odorassociated with summer thunderstorms is due to the formation of small amounts of O3.)Exposure to O3 at higher concentrations leads to coughing, rapid beating of the heart, chestpain, and general body pain. At concentrations above 1 ppm, ozone is toxic.
The most famous characteristic of ozone is its ability to absorb high energy radiationin the ultraviolet portion of the spectrum (� � 300 nm), thereby providing a filter that pro-tects us from exposure to high energy ultraviolet radiation emitted by the sun. We can un-derstand the importance of this filter if we think about what happens when radiation fromthe sun is absorbed by our skin.
Electromagnetic radiation in the infrared, visible, and low energy portions of the ul-traviolet spectrum carries enough energy to excite an electron into a higher energy or-bital. This electron eventually falls back into the orbital from which it was excited, andenergy is given off to the surrounding tissue in the form of heat. Anyone who has suf-fered from a sunburn can appreciate the painful consequences of excessive amounts ofthis radiation.
Radiation in the high energy portion of the ultraviolet spectrum carries enough energyto ionize atoms and molecules. Because living tissue is 70–90% water by weight, the mostcommon ionization reaction involves the loss of an electron from a neutral water moleculeto form an H2O� ion.
H2O 88n H2O� � e�
The H2O� ions formed in the reaction have an odd number of electrons and are extremelyreactive. They can cause permanent damage to the cell tissue and induce processes thateventually result in skin cancer. Relatively small amounts of this radiation can thereforehave drastic effects on living tissue. Therefore, the protection by the ozone (O3) layer whichprevents high energy radiation from reaching the earth is important to the health of livingorganisms.
In 1974 Molina and Rowland pointed out that chlorofluorocarbons, such as CFCl3 andCF2Cl2, which had been used as refrigerants and as propellants in aerosol cans, were be-ginning to accumulate in the atmosphere. In the stratosphere, at altitudes of 10 to 50 kmabove the earth’s surface, chlorofluorocarbons decompose to form Cl atoms and chlorineoxides such as ClO when they absorb sunlight. Cl atoms and ClO molecules also havean unpaired electron in the valence shell of the molecule. As a result, they are unusually
TABLE N.1 Properties of Allotropes of Oxygen
Property Oxygen (O2) Ozone (O3)
Melting point �218.75°C �192.5°CBoiling point �182.96°C �110.5°CDensity (at 20°C) 1.331 g/L 1.998 g/LOOO bond order 2 1.5OOO bond length 0.1207 nm 0.1278 nm
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NONMETAL 9
reactive. In the atmosphere, they react with ozone or with the oxygen atoms that are neededto form ozone.
Cl � O3 88n ClO � O2
ClO � O 88n Cl � O2
Molina and Rowland postulated that these substances would eventually deplete the ozoneshield in the stratosphere, with dangerous implications for biological systems that wouldbe exposed to increased levels of high energy ultraviolet radiation.
Oxygen as an Oxidizing Agent
Fluorine is the only element with which oxygen reacts that is more electronegative than oxy-gen. As a result, oxygen gains electrons in virtually all of its chemical reactions. Each O2
molecule must gain four electrons to satisfy the octets of the two oxygen atoms withoutsharing electrons.
Oxygen therefore oxidizes metals to form salts in which the oxygen atoms are formallypresent as O2� ions. Rust forms, for example, when iron reacts with oxygen in the pres-ence of water to give a salt that formally contains the Fe3� and O2� ions, with an averageof three water molecules coordinated to each Fe3� ion in the solid.
H2O4 Fe(s) � 3 O2(g) 88n 2 Fe2O3�3 H2O(s)
Oxygen also oxidizes nonmetals, such as carbon, to form covalent compounds in which theoxygen has an oxidation number of �2.
C(s) � O2(g) 88n CO2(g)
Oxygen is the perfect example of an oxidizing agent because it increases the oxidation stateof almost any substance with which it reacts. In the course of its reactions, oxygen is re-duced. The substances it reacts with are therefore reducing agents.
Exercise N.2
Identify the oxidizing agents and reducing agents in the following reactions.(a) Fe2O3(s) � 3 C(s) n 2 Fe(s) � 3 CO(g)(b) CH4(g) � 2 O2(g) n CO2(g) � 2 H2O(g)
Solution
(a) In this reaction, carbon reduces Fe2O3 to iron metal, which means carbon is the re-ducing agent. Fe2O3 oxidizes carbon to CO and is therefore the oxidizing agent.
2 Fe � 3 CO
Oxidation
Reduction0 �20�3
Fe2O3 � 3 C
¼
O³¼³PO � 4 e� 2 [¼O¼]2�þ
³
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10 NONMETAL
(b) Oxygen oxidizes CH4 in the reaction, so O2 is the oxidizing agent. The oxygen is re-duced by CH4, which means that CH4 is the reducing agent.
Each year between 75 and 80 quads, or quadrillion (1015) BTU (British thermal units),of energy is consumed in the United States.1 Less than 10% of this energy is provided bynuclear, solar, geothermal, or hydro power. The rest can be traced to a combustion reac-tion in which a fuel is oxidized by O2. The cars, trucks, and buses that fill our highways arepowered by gasoline engines that burn hydrocarbons such as octane, C8H18,
2 C8H18(l) � 25 O2(g) 88n 16 CO2(g) � 18 H2O(g)
or diesel engines that burn larger hydrocarbons such as cetane, C16H34.
2 C16H34(l) � 49 O2(g) 88n 32 CO2(g) � 34 H2O(g)
We heat our homes by burning the methane (CH4) in natural gas, the high molecular weighthydrocarbons in fuel oil, or the hydrocarbons in wood, or by using electricity generated ina power plant that burns either oil or coal.
The energy we use to fuel our bodies also comes from combustion reactions. Energyenters our bodies in the form of lipids, proteins, and carbohydrates. These “fuels” are con-verted into carbohydrates, such as glucose (C6H12O6), which react with oxygen to producethe energy we need to survive.
C6H12O6(aq) � 6 O2(g) 88n 6 CO2(g) � 6 H2O(l)
About 65% of the energy given off in the reaction is used to synthesize the ATP (adeno-sine triphosphate) that fuels biological processes. The remaining 35% is released as theheat that keeps our body temperatures higher than the temperature of the surroundings.
There is an ever-growing awareness that the earth contains a finite amount of fossil fuels, such as oil and coal. Nuclear, solar, and geothermal power will be increasingly im-portant sources of energy. But they won’t replace fossil fuels by themselves because theyare used to produce electrical energy, which is difficult to store. One possible solution tothis problem has been labeled the hydrogen economy.
The first step in the hydrogen economy is to use energy from nuclear, solar, or geo-thermal power to split water into its elements.
2 H2O(l) 88n 2 H2(g) � O2(g)
The oxygen is then released to the atmosphere, and the hydrogen is either burned as a fuel
2 H2(g) � O2(g) 88n 2 H2O(g) �H° � �483.64 kJ/molrxn
CO2 � 2 H2OCH4 � 2 O2
Oxidation
Reduction
0 �4�4 �2�2
1One BTU is the energy needed to raise the temperature of one pound of water by 1°F; 1 BTU � 1.055 kJ.
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NONMETAL 11
or used to reduce carbon monoxide to methanol or to gasoline, which can be stored andlater burned as a fuel.
CO(g) � 2 H2(g) 88n CH3OH(l)8 CO(g) � 17 H2(g) 88n C8H18(l) � 8 H2O(l)
Peroxides
It takes four electrons to reduce an O2 molecule to a pair of O2� ions. If the reaction stopsafter the O2 molecule has gained only two electrons, the O2
2� ion is produced.
The O22� ion has two more electrons than a neutral O2 molecule, which means that the
oxygen atoms must share only a single pair of bonding electrons to achieve an octet of va-lence electrons. The O2
2� ion is called the peroxide ion because compounds that containthe ion are unusually rich in oxygen. They are not just oxides—they are (hy-)peroxides.
The easiest way to prepare a peroxide is to react sodium or barium metal with oxygen.
2 Na(s) � O2(g) 88n Na2O2(s)Ba(s) � O2(g) 88n BaO2(s)
When the peroxides are allowed to react with a strong acid, hydrogen peroxide (H2O2) isproduced.
BaO2(s) � 2 H�(aq) 88n Ba2�(aq) � H2O2(aq)
The Lewis structure of hydrogen peroxide contains an OOO single bond.
The electron domain theory predicts that the geometry around each oxygen atom in H2O2
should be bent. But this theory can’t predict whether the four atoms should lie in the sameplane or whether the molecule should be visualized as lying in two intersecting planes. Theexperimentally determined structure of H2O2 is shown in Figure N.7.
O� 2 e� [ ]2�OOO O
H
94.8°111.5°
H
FIGURE N.7 Geometry of an H2O2 molecule.
The HOOOO bond angle in the molecule is only slightly larger than the angle between apair of adjacent 2p atomic orbitals on the oxygen atom, and the angle between the planesthat form the molecule is slightly larger than the tetrahedral angle. The geometry aroundeach oxygen atom is bent, or angular, with an OOOOH bond angle of 94.8°. In order tokeep the nonbonding electrons on the oxygen atoms as far apart as possible, the four atomsin the molecule lie in two planes that intersect at an angle of 111.5°.
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The oxidation number of the oxygen atoms in hydrogen peroxide is �1. H2O2 cantherefore act as an oxidizing agent and capture two more electrons to form a pair of hy-droxide ions, in which the oxygen has an oxidation number of �2.
H2O2 � 2 e� 88n 2 OH�
Or, it can act as a reducing agent and lose a pair of electrons to form an O2 molecule.
H2O2 88n O2 � 2 H� � 2 e�
CheckpointUse Lewis structures to explain what happens in the following reactions.
H2O2 � 2 e� 88n 2 OH�
H2O2 88n O2 � 2 H� � 2 e�
Reactions in which a compound simultaneously undergoes both oxidation and reduc-tion are called disproportionation reactions. The products of the disproportionation ofH2O2 are oxygen and water.
2 H2O2(aq) 88n O2(g) � 2 H2O(l)
Exercise N.3
Use the reactions that describe what happens when H2O2 loses a pair of electrons and whathappens when H2O2 picks up a pair of electrons to explain why the disproportionation ofhydrogen peroxide gives oxygen and water.
2 H2O2(aq) 88n O2(g) � 2 H2O(l)
Solution
Adding the half-reaction for the oxidation of H2O2 to the half-reaction for the reductionof the compound gives the following results.
H2O2 � 2 e� 88n 2 OH�
H2O2 88n O2 � 2 H� � 2 e�
2 H2O2 88n O2 � 2 H� � 2 OH�
The H� and OH� ions produced in the two halves of the reaction combine to form waterto give the following overall stoichiometry for the reaction.
2 H2O2 88n O2 � 2 H2O
The disproportionation of H2O2 is an exothermic reaction.
2 H2O2(aq) 88n O2(g) � 2 H2O(l) �H° � �189.3 kJ/molrxn
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NONMETAL 13
The reaction is relatively slow, however, in the absence of a catalyst. The principal uses ofH2O2 revolve around its oxidizing ability. It is used in dilute (3%) solutions as a disinfec-tant. In more concentrated solutions (30%), it is used as a bleaching agent for hair, fur,leather, or the wood pulp used to make paper. In very concentrated (40–70%) solutions,H2O2 has been used as rocket fuel because of the ease with which it decomposes to give O2.
Methods of Preparing O2
Small quantities of O2 gas can be prepared in a number of ways.
• By decomposing a dilute solution of hydrogen peroxide with dust or a metal sur-face as the catalyst.
2 H2O2(aq) 88n O2(g) � 2 H2O(l)
• By reacting hydrogen peroxide with a strong oxidizing agent, such as the perman-ganate ion, MnO4
�.
5 H2O2(aq) � 2 MnO4�(aq) � 6 H�(aq) 88n 2 Mn2�(aq) � 5 O2(g) � 8 H2O(l)
• By passing an electric current through water.
electrolysis2 H2O(l) 888888n 2 H2(g) � O2(g)
• By heating potassium chlorate in the presence of a catalyst until it decomposes.
MnO22 KClO3(s) 888n 2 KCl(s) � 3 O2(g)
CheckpointWhich of the following elements or compounds reacts with water to give a solutionthat could be used to produce O2?(a) Na (b) Na2O (c) Na2O2 (d) NaOH (e) NaCl
Chemistry in the World Around Us
The Chemistry of the Atmosphere
Although major changes occurred in the atmosphere during the early history of ourplanet, the chemistry of the earth’s atmosphere has been more or less constant duringthe time in which the human race evolved. This is no longer true. The amount ofmethane (CH4) in the atmosphere is increasing at a rate of more than 1% per year.The concentration of carbon dioxide has more than doubled since 1750 and seems to beincreasing at an exponential rate (see Figure N.8). In recent years, attention has beenfocused on one particular change in the atmosphere, the depletion of ozone above theAntarctic continent.
For over 25 years, a team of scientists collected data on variations in the amount ofO3 at different altitudes above the British Antarctic Survey station at Halley Bay. In1985, they reported that the O3 concentration declined after the return of solar radiation
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14 NONMETAL
each September. The effect was small when first noticed in the late 1970s, but it reachedsuch high levels by 1984 that the O3 concentration declined by 30% by the end of Octoberof that year.
The data from Halley Bay sampled only a small portion of the atmosphere over theAntarctic. Once the notion of ozone depletion was reported, however, scientists wereable to retrieve data that had been recorded by satellites over a period of years, whichconfirmed that the same effect occurred over virtually the entire Antarctic continent.Several features of the ozone depletion were particularly interesting.
• The drop in the ozone concentration occurred very rapidly, within a period of sixweeks each spring. (“Spring” occurs during September and October in the south-ern hemisphere.)
• Although the drop in O3 occurred above the Antarctic, it corresponds to a loss of3% of the total ozone concentration in the planet’s atmosphere.
• The decline in the O3 concentration became more serious each year. By 1989, theO3 concentration during the summer months dropped by 70%.
• The decline was temporary. During the winter months, the ozone level built backto normal levels.
Several possible explanations were available for the ozone holes. One of the mostpopular was the suggestion by Molina and Rowland that the ozone in the atmospherecould be destroyed by Cl and ClO radicals created when chlorofluorocarbons (CFC’s) inthe atmosphere decomposed. The question was: What evidence could be found to eithersupport or refute that hypothesis? The problem is complex, because more than 200chemical reactions have been included in the models used to explain the chemistry ofthe atmosphere. It is further complicated by the fact that ClO radicals exist in the atmo-sphere at concentrations of only about 1 part per trillion by volume (pptv).
In the 4 January 1991 issue of Science, James Anderson and co-workers reported datathat probed the link between the release of chlorofluorocarbons into the atmosphere andthe disappearance of ozone from the stratosphere above the Antarctic each spring. Thedata were obtained between 23 August and 22 September 1987, using special instrumentsmounted in high-altitude aircraft. Initial measurements after the plane took off from the tipof Chile (54° S latitude) suggested that the background concentration of ClO radicals wasat the threshold of detection: about 1 pptv. As the plane flew toward the south pole, theClO concentration increased slightly until about 65° S latitude, when it rapidly increased.
A plot of the concentration of ClO radicals versus latitude for the 16 September flightis shown in Figure N.9. Once the aircraft reached a latitude of 68° S the abundance of theradicals increased by three orders of magnitude, to a level of approximately 1200 pptv.These data by themselves are suggestive, but Figure N.9 contains more compelling
270
CO
2 c
once
ntra
tion
(pp
mv)
1700
280
290
300
310
320
330
340
360
1750 1800 1850 1900 1950 2000Year
FIGURE N.8 The CO2 concentration in the at-mosphere from the time of the industrial revo-lution to present in units of parts per million byvolume (ppmv).
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NONMETAL 15
evidence. The plot at the top of Figure N.9 shows that the concentration of O3 droppedby a factor of about 2.5 at virtually the same time that the ClO concentration increased.
Anderson and co-workers concluded as follows: “When taken independently, eachelement in the case contains a segment of the puzzle that in itself is not conclusive. Whentaken together, however, they provide convincing evidence that the dramatic reductionin O3 over the Antarctic continent would not have occurred had CFC’s not been synthe-sized and then added to the atmosphere.”2
N.4 THE CHEMISTRY OF SULFURBecause sulfur is directly below oxygen in the periodic table, these elements have similarelectron configurations.
O [He] 2s2 2p4
S [Ne] 3s2 3p4
As a result, sulfur forms many compounds that are analogs of oxygen compounds, as shownin Table N.2. The last two examples in Table N.2 show how the prefix thio- can be used todescribe compounds in which sulfur replaces an oxygen atom.
2J. G. Anderson, D. W. Toohey, and W. H. Brune, Science, 251, 45 (1991).
1200
600
062 64 66 68 70 72
0
1000
2000
3000
O3 c
once
ntra
tion
in p
pbv
ClO
• co
ncen
trat
ion
in p
ptv
Latitude (degrees south)
16 SEPT
ClOO3
FIGURE N.9 Variations in the ClO and O3 concentrations in the atmosphere during the September 16, 1987, flight toward the Antarcticcontinent.
TABLE N.2 Oxygen Compounds and Their Sulfur Analogs
Oxygen Compound Sulfur Compound
Na2O (sodium oxide) Na2S (sodium sulfide)H2O (water) H2S (hydrogen sulfide)O3 (ozone) SO2 (sulfur dioxide)CO2 (carbon dioxide) CS2 (carbon disulfide)OCN� (cyanate) SCN� (thiocyanate)OC(NH2)2 (urea) SC(NH2)2 (thiourea)
There are four principal differences between the chemistry of sulfur and oxygen.
• OPO double bonds are much stronger than SPS double bonds.• SOS single bonds are almost twice as strong as OOO single bonds.• Sulfur is much less electronegative than oxygen.• Sulfur can expand its valence shell to hold more than eight electrons; oxygen cannot.
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16 NONMETAL
These seemingly minor differences have important consequences for the chemistry of theseelements.
The Effect of Differences in the Strength of XOOX and XPPX Bonds
The radius of a sulfur atom is about 60% larger than that of an oxygen atom.
� � 1.6
As a result, it is harder for sulfur atoms to come close enough together to form doublebonds. SPS double bonds are therefore much weaker than OPO double bonds.
Double bonds between sulfur and oxygen or carbon atoms can be found in compoundssuch as SO2 and CS2. But these double bonds are much weaker than the equivalent dou-ble bonds to oxygen atoms in O3 or CO2. The CPS double bonds in CS2, for example, areabout 65% as strong as the CPO double bonds in CO2.
Elemental oxygen consists of O2 molecules in which each atom completes its octet ofvalence electrons by sharing two pairs of electrons with a single neighboring atom. Becausesulfur doesn’t form strong SPS double bonds, elemental sulfur consists of cyclic S8 mole-cules in which each atom completes its octet by forming single bonds to two different neigh-boring atoms, as shown in Figure N.10.
0.104 nm��0.066 nm
Covalent radius of sulfur���Covalent radius of oxygen
S
SS S
S
S
SS
S8 FIGURE N.10 Structure of a cyclic S8 molecule.
S8 molecules can pack to form more than one crystal. The most stable form of sulfurconsists of orthorhombic crystals of S8 molecules, which are often found near volcanos. Ifthe orthorhombic crystals are heated until they melt and the molten sulfur is then cooled,an allotrope of sulfur consisting of monoclinic crystals of S8 molecules is formed. The mon-oclinic crystals slowly transform themselves into the more stable orthorhombic structureover a period of time.
The tendency of an element to form bonds to itself is called catenation (from the Latinword catena, “chain”). Because sulfur forms unusually strong SOS single bonds, it is bet-ter at catenation than any element except carbon. As a result, the orthorhombic and mono-clinic forms of sulfur are not the only allotropes of the element. Allotropes of sulfur alsoexist that differ in the size of the molecules that form the crystal. Cyclic molecules thatcontain 6, 7, 8, 10, and 12 sulfur atoms are known.
Sulfur melts at 119.25°C to form a yellow liquid that is less viscous than water. If theliquid is heated to 159°C, it turns into a dark red liquid that can’t be poured from its con-tainer. The viscosity of the dark red liquid is 2000 times greater than that of molten sulfurbecause the cyclic S8 molecules open up and link together to form long chains of as manyas 100,000 sulfur atoms.
When sulfur reacts with an active metal, it can form the sulfide ion, S2�.
16 K(s) � S8(s) 88n 8 K2S(s)
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NONMETAL 17
This is not the only product that can be obtained, however. A variety of polysulfide ionswith a charge of �2 can be produced that differ in the number of sulfur atoms in the chain.
2 K(s) � S8(s) 88n K2S2 � [K�]2[SOS]2�
K2S3 � [K�]2[SOSOS]2�
K2S4 � [K�]2[SOSOSOS]2�
K2S5 � [K�]2[SOSOSOSOS]2�
K2S6 � [K�]2[SOSOSOSOSOS]2�
K2S8 � [K�]2[SOSOSOSOSOSOSOS]2�
Exercise N.4
Use the tendency of sulfur to form polysulfide ions to explain why iron has an oxidationnumber of �2 in iron pyrite, FeS2, one of the most abundant sulfur ores.
Solution
If the oxidation number of iron in FeS2 is �2, the sulfur must be present as the S22� ion.
FeS2 � [Fe2�][S22�]
The disulfide ion, S22�, is the sulfur analog of the peroxide ion, O2
2�, and has an analo-gous Lewis structure.
The Effect of Differences in the Electronegativities of Sulfur and Oxygen
Because sulfur is much less electronegative than oxygen, it is more likely to form com-pounds in which it has a positive oxidation number (see Table N.3).
TABLE N.3 Common Oxidation Numbers for Sulfur
Oxidation Number Examples
�2 Na2S, H2S�1 Na2S2, H2S2
0 S8
�1 S2Cl2�2 S2O3
2�
�2�12
� S4O62�
�3 S2O42�
�4 SF4, SO2, H2SO3, SO32�
�5 S2O62�
�6 SF6, SO3, H2SO4, SO42�
In theory, sulfur can react with oxygen to form either SO2 or SO3, whose Lewis struc-tures are given in Figure N.11.
SO2 is therefore a resonance hybrid of two Lewis structures analogous to the structureof ozone in Figure N.6. SO3 can be thought to result from the donation of a pair of non-bonding electrons on the sulfur atom in SO2 to an empty orbital on a neutral oxygen atomto form a covalent bond.
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18 NONMETAL
In practice, combustion of sulfur compounds gives SO2, regardless of whether sulfuror a compound of sulfur is burned.
S8(s) � 8 O2(g) 88n 8 SO2(g)CS2(l) � 3 O2(g) 88n CO2(g) � 2 SO2(g)
3 FeS2(s) � 8 O2(g) 88n Fe3O4(s) � 6 SO2(g)
Although the SO2 formed in the reactions should react with O2 to form SO3, the rate ofthat reaction is very slow. The rate of the conversion of SO2 into SO3 can be greatly in-creased by adding an appropriate catalyst.
V2O5/K2O2 SO2(g) � O2(g) 888888n 2 SO3(g)
Enormous quantities of SO2 are produced by industry each year and then convertedto SO3, which can be used to produce sulfuric acid, H2SO4. In theory, sulfuric acid can bemade by dissolving SO3 gas in water.
SO3(g) � H2O(l) 88n H2SO4(aq)
In practice, this isn’t convenient. Instead, SO3 is absorbed in 98% H2SO4, where it reactswith the water to form additional H2SO4 molecules. Water is then added, as needed, tokeep the concentration of the solution between 96% and 98% H2SO4 by weight.
Sulfuric acid is by far the most important industrial chemical. It has even been sug-gested that there is a correlation between the amount of sulfuric acid a country consumesand its standard of living.
Sulfuric acid dissociates in water to give the HSO4� ion, which is known as the hy-
drogen sulfate, or bisulfate, ion.
H2SO4(aq) 88n H�(aq) � HSO4�(aq)
Roughly 10% of the hydrogen sulfate ions dissociate further to give the SO42�, or sulfate, ion.
HSO4�(aq) 88n H�(aq) � SO4
2�(aq)
Sulfur dioxide dissolves in water to form sulfurous acid.
SO2(g) � H2O(l) 88n H2SO3(aq)
Sulfurous acid doesn’t dissociate in water to as great an extent as sulfuric acid.Sulfuric acid and sulfurous acid are examples of a class of compounds known as oxy-
acids because they are literally acids that contain oxygen. Because they are negative ions
S
O O O O
S
O O
O
SS
O O
SO2
SO3
O
S
O O
O
FIGURE N.11 Lewis structures of SO2
and SO3.
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NONMETAL 19
(or anions) that contain oxygen, the SO32� and SO4
2� ions are known as oxyanions. TheLewis structures of some of the oxides of sulfur that form oxyacids and oxyanions are givenin Figure N.12. One of the oxyanions deserves special mention. This ion, which is knownas the thiosulfate ion, is formed by the reaction between sulfur and the sulfite (SO3
2�) ion.
8 SO32�(aq) � S8(s) 88n 8 S2O3
2�(aq)
CheckpointUse the following Lewis structures to explain why the S2O3
2� ion is literally a thio-sulfate.
O
S
O
O HOH
Sulfuric acid, H2SO4
S
S
O
O HOH
Thiosulfuric acid, H2S2O3
O
S O HOH
Sulfurous acid, H2SO3
O
S
O
O OOH
O
S
O
O H
Peroxydisulfuric acid, H2S2O8
O
S
O
O OO
O
S
O
O
Peroxydisulfate, S2O82–
2–
O
S
O
S SO
O
S
O
O
Tetrathionate, S4O62–
2–
2–O
S OO
Sulfite, SO32–
S
S
O
OO
Thiosulfate, S2O32–
2–
O
S
O
OO
Sulfate, SO42–
2–
Oxyacids Oxyanions
FIGURE N.12 Oxyacids of sulfurand their oxyanions.
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20 NONMETAL
The Effect of Differences in the Abilities of Sulfur and Oxygen to Expand TheirValence Shell
Oxygen reacts with fluorine to form OF2.
O2(g) � 2 F2(g) 88n 2 OF2(g)
The reaction stops at this point because oxygen can hold only eight electrons in its valenceshell.
Sulfur, however, reacts with fluorine to form SF4 and SF6 because sulfur can expand its va-lence shell to hold 10 or even 12 electrons.
S8(s) � 16 F2(g) 88n 8 SF4(g)S8(s) � 24 F2(g) 88n 8 SF6(g)
There are 10 valence electrons on the sulfur atom in SF4, so the structure of the moleculeis based on a distorted trigonal bipyramid, as shown in Figure N.13. The 12 valence elec-trons on the central atom in SF6 are distributed toward the corners of an octahedron.
FF
F
F
F
S S
F
F
F
F F
186.9°
101.6°
SF4 SF6
N.5 THE CHEMISTRY OF NITROGENThe chemistry of nitrogen is dominated by the ease with which nitrogen atoms form dou-ble and triple bonds. A neutral nitrogen atom contains five valence electrons: 2s2 2p3. Anitrogen atom can therefore achieve an octet of valence electrons by sharing three pairsof electrons with another nitrogen atom.
Because the covalent radius of a nitrogen atom is relatively small, nitrogen atoms comeclose enough together to form very strong multiple bonds. The NqN triple bond (�H°ac ��945.41 kJ/molrxn) is almost twice as strong as the OPO double bond (�H°ac � �498.34kJ/molrxn).
The strength of the NqN triple bond makes the N2 molecule so inert that lithium isone of the few elements with which it reacts at room temperature.
6 Li(s) � N2(g) 88n 2 Li3N(s)
In spite of the fact that the N2 molecule is unreactive, compounds containing nitrogen ex-ist for virtually every element in the periodic table except those in Group VIIIA (He, Ne,
FIGURE N.13 Structures of SF4 and SF6.
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NONMETAL 21
Ar, and so on). This can be explained in two ways. First, N2 becomes significantly morereactive as the temperature increases. At high temperatures, nitrogen reacts with hydro-gen to form ammonia and with oxygen to form nitrogen oxide.
N2(g) � 3 H2(g) 88n 2 NH3(g)N2(g) � O2(g) 88n 2 NO(g)
Second, a number of catalysts found in nature can overcome the inertness of N2 at lowtemperatures.
The Synthesis of Ammonia
It is difficult to imagine a living system that doesn’t contain nitrogen, which is an essentialcomponent of the proteins, nucleic acids, vitamins, and hormones that make life possible.Animals pick up the nitrogen they need from the plants or other animals in their diet. Al-though there is an abundance of N2 in the atmosphere, plants cannot use nitrogen in itselemental form. Plants have to pick up their nitrogen from the soil or absorb it as N2 fromthe atmosphere. The concentration of nitrogen in the soil is fairly small, so the process bywhich plants reduce N2 to NH3 (nitrogen fixation) is extremely important.
Although 200 million tons of NH3 are produced by nitrogen fixation each year, plants,by themselves, cannot reduce N2 to NH3. The reaction is carried out by blue-green algaeand bacteria that are associated with certain plants. The best understood example of ni-trogen fixation involves the Rhizobium bacteria found in the root nodules of legumes suchas clover, peas, and beans. The bacteria contain a nitrogenase enzyme that is capable ofthe remarkable feat of reducing N2 from the atmosphere to NH3 at room temperature.
Ammonia is made on an industrial scale by a process first developed between 1909 and1913 by Fritz Haber. In the Haber process, a mixture of N2 and H2 gas at 200 to 300 atmand 400°C to 600°C is passed over a catalyst of finely divided iron metal.
FeN2(g) � 3 H2(g) 88n 2 NH3(g)
Almost 20 million tons of NH3 are produced in the United States each year by this process.About 80% of it, worth more than $2 billion, is used to make fertilizers for plants that can’tfix nitrogen from the atmosphere. On the basis of weight, ammonia is the second most im-portant industrial chemical in the United States. (Only sulfuric acid is produced in largerquantities.)
Two-thirds of the ammonia used for fertilizers is converted into solids such as ammo-nium nitrate, NH4NO3; ammonium phosphate, (NH4)3PO4; ammonium sulfate, (NH4)2SO4;and urea, H2NCONH2. The other third is applied directly to the soil as anhydrous (liter-ally, “without water”) ammonia. Ammonia is a gas at room temperature. It can be han-dled as a liquid when dissolved in water to form an aqueous solution. Alternatively, it canbe cooled to temperatures below �33°C, in which case the gas condenses to form the an-hydrous liquid, NH3(l).
The Synthesis of Nitric Acid
The NH3 produced by the Haber process that isn’t used as fertilizer is burned in oxygento generate nitrogen oxide.
4 NH3(g) � 5 O2(g) 88n 4 NO(g) � 6 H2O(g)
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Nitrogen oxide—or nitric oxide, as it was once known—is a colorless gas that reacts rapidlywith oxygen to produce nitrogen dioxide, a dark brown gas.
2 NO(g) � O2(g) 88n 2 NO2(g)
Nitrogen dioxide dissolves in water to give nitric acid and NO, which can be captured andrecycled.
3 NO2(g) � H2O(l) 88n 2 HNO3(aq) � NO(g)
Thus, by a three-step process developed by Friedrich Ostwald in 1908, ammonia can beconverted into nitric acid.
4 NH3(g) � 5 O2(g) 88n 4 NO(g) � 6 H2O(g)2 NO(g) � O2(g) 88n 2 NO2(g)
3 NO2(g) � H2O(l) 88n 2 HNO3(aq) � NO(g)
The Haber process for the synthesis of ammonia combined with the Ostwald processfor the conversion of ammonia into nitric acid revolutionized the explosives industry. Ni-trates have been important explosives ever since Friar Roger Bacon mixed sulfur, saltpeter,and powdered carbon to make gunpowder in 1245.
16 KNO3(s) � S8(s) � 24 C(s) 88n 8 K2S(s) � 24 CO2(g) � 8 N2(g)�H° � �4575 kJ/molrxn
Before the Ostwald process the only source of nitrates for use in explosives was naturallyoccurring minerals such as saltpeter, which is a mixture of NaNO3 and KNO3. Once a de-pendable supply of nitric acid became available from the Ostwald process, a number of ni-trates could be made for use as explosives. Combining NH3 from the Haber process with HNO3
from the Ostwald process, for example, gives ammonium nitrate, which is both an excellentfertilizer and an inexpensive, dependable explosive commonly used in blasting powder.
2 NH4NO3(s) 88n 2 N2(g) � O2(g) � 4 H2O(g)
The destructive power of ammonium nitrate is apparent in photographs of the Alfred P.Murrah Federal Building in Oklahoma City, which was destroyed with a bomb made fromammonium nitrate on April 19, 1995.
Intermediate Oxidation Numbers
Nitric acid (HNO3) and ammonia (NH3) represent the maximum (�5) and minimum (�3)oxidation numbers for nitrogen. Nitrogen also forms compounds with every oxidation num-ber between these extremes (see Table N.4).
Negative Oxidation Numbers of Nitrogen besides �3
At about the time that Haber developed the process for making ammonia and Ostwaldworked out the process for converting ammonia into nitric acid, Raschig developed a processthat used the hypochlorite ion (OCl�) to oxidize ammonia to produce hydrazine, N2H4.
2 NH3(aq) � OCl�(aq) 88n N2H4(aq) � Cl�(aq) � H2O(l)
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NONMETAL 23
This reaction can be understood by noting that the OCl� ion is a two-electron oxidizingagent. Let’s therefore imagine a hypothetical mechanism for the reaction in which the firststep is the loss of a pair of nonbonding electrons from an ammonia molecule to form anNH3
2� ion, as shown in the first step in Figure N.14. A pair of nonbonding electrons froma second NH3 molecule could then be donated into the empty valence shell orbital on theNH3
2� ion to form an NON bond. In step 3, the product of the reaction in step 2 couldthen lose a pair of H� ions to form a hydrazine molecule.
TABLE N.4 Common Oxidation Numbers for Nitrogen
Oxidation Number Examples
�3 NH3, NH4�, NH2
�, Mg3N2
�2 N2H4
�1 NH2OH��
13
� NaN3, HN3
�0 N2
�1 N2O�2 NO, N2O2
�3 HNO2, NO2�, N2O3, NO�
�4 NO2, N2O4
�5 HNO3, NO3�, N2O5
Hydrazine is a colorless liquid with a faint odor of ammonia that can be collected whenthe solution is heated until N2H4 distills out of the reaction flask. Many of the physicalproperties of hydrazine are similar to those of water.
H2O N2H4
Density 1.000 g/cm3 1.008 g/cm3
Melting Point 0.00°C 1.54°CBoiling Point 100°C 113.8°C
Step 1
Step 3
Step 2
H
H
H
N –2 e–H
H
H
N
2+
2+H
H
H
H
HH
N N
2+
+
+
H
H
H
N
H
H
H
N
NNH
H H
H2+H
H
H
N
H
H
H
N 2 H�
FIGURE N.14 Hydrazine is prepared by reacting NH3 with a two-electronoxidizing agent.
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24 NONMETAL
There is a significant difference between the chemical properties of the compounds, how-ever. Hydrazine burns when ignited in air to give nitrogen gas, water vapor, and largeamounts of energy.
N2H4(l) � O2(g) 88n N2(g) � 2 H2O(g) �H° � �534.3 kJ/molrxn
The principal use of hydrazine is as a rocket fuel. It is second only to liquid hydrogenin terms of the kilograms of thrust produced per kilogram of fuel burned. Hydrazine hasseveral advantages over liquid H2, however. It can be stored at room temperature, whereasliquid hydrogen must be stored at temperatures below �253°C. Hydrazine is also moredense than liquid H2 and therefore requires less storage space.
Pure hydrazine is seldom used as a rocket fuel because it freezes at the temperaturesencountered in the upper atmosphere. Hydrazine is mixed with N,N-dimethylhydrazine,(CH3)2NNH2, to form a solution that remains a liquid at low temperatures. Mixtures ofhydrazine and N,N-dimethylhydrazine were used to fuel the Titan II rockets that carriedthe Project Gemini spacecraft, and the reaction between hydrazine derivatives and N2O4
is still used to fuel the small rocket engines that enable the space shuttle to maneuverin space.
The product of the combustion of hydrazine is unusual. When carbon compounds burn,the carbon is oxidized to CO or CO2. When sulfur compounds burn, SO2 is produced.When hydrazine is burned, the product of the reaction is N2 because of the unusually strongNqN triple bond in the N2 molecule.
N2H4(l) � O2(g) 88n N2(g) � 2 H2O(g)
Positive Oxidation Numbers for Nitrogen: The Nitrogen Halides
Fluorine, oxygen, and chlorine are the only elements more electronegative than nitrogenthat form compounds with nitrogen. As a result, positive oxidation numbers of nitrogenare found in compounds that contain one or more of these elements.
In theory, N2 could react with F2 to form a compound with the formula NF3. In prac-tice, N2 is too inert to undergo the reaction at room temperature. NF3 is made by reactingammonia with F2 in the presence of a copper metal catalyst. The HF produced in the re-action combines with ammonia to form ammonium fluoride. The overall stoichiometry forthe reaction is therefore written as follows.
Cu4 NH3(g) � 3 F2(g) 88n NF3(g) � 3 NH4F(s)
The Lewis structure of NF3 is analogous to the Lewis structure of NH3, and the moleculeshave similar shapes.
Ammonia reacts with chlorine to form NCl3, which seems at first glance to be closelyrelated to NF3. But there is a significant difference between the compounds. NF3 is es-sentially inert at room temperature, whereas NCl3 is a shock-sensitive, highly explosive liq-uid that decomposes to form N2 and Cl2.
2 NCl3(l) 88n N2(g) � 3 Cl2(g)
Ammonia reacts with iodine to form a solid that is a complex between NI3 and NH3. Thismaterial is the subject of a popular, but dangerous, demonstration in which freshly pre-pared samples of NI3 in ammonia are poured onto filter paper, which is allowed to dry ona ring stand. After the ammonia evaporates, the NH3/NI3 crystals are touched with a feather
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NONMETAL 25
attached to a meter stick, resulting in detonation of the shock-sensitive solid, which de-composes to form a mixture of N2 and I2.
2 NI3(s) 88n N2(g) � 3 I2(g)
Positive Oxidation Numbers for Nitrogen: The Nitrogen Oxides
Lewis structures for seven oxides of nitrogen with oxidation numbers ranging from �1 to�5 are given in Figure N.15. These compounds all have two things in common: they con-tain NPO double bonds, and they are less stable than their elements in the gas phase.
FIGURE N.15 The oxides of nitrogen.
N N N NO O
Dinitrogen oxide, N2O(nitrous oxide)
N O
Nitrogen oxide, NO(nitric oxide)
O N
N O
Dinitrogen dioxide, N2O2
Dinitrogen trioxide, N2O3
N N
O
O
NO
O
Nitrogen dioxide, NO2
NNOO
OO
Dinitrogen tetroxide, N2O4
NN OO
O
O
O
Dinitrogen pentoxide, N2O5
O
Dinitrogen oxide, N2O, which is also known as nitrous oxide, can be prepared by care-fully decomposing ammonium nitrate.
170–200°CNH4NO3(s) 888888n N2O(g) � 2 H2O(g)
Nitrous oxide is a sweet-smelling, colorless gas best known to nonchemists as “laughinggas.” As early as 1800, Humphry Davy noted that N2O, inhaled in relatively small amounts,produced a state of apparent intoxication often accompanied by either convulsive laugh-ter or crying. When taken in larger doses, nitrous oxide provides fast and efficient relieffrom pain. N2O was therefore used as the first anesthetic. Because large doses are needed
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26 NONMETAL
to produce anesthesia, and continued exposure to the gas can be fatal, N2O is used todayonly for relatively short operations.
Nitrous oxide has several other interesting properties. First, it is highly soluble in cream;for that reason, it is used as the propellant in whipped cream dispensers. Second, althoughN2O does not burn by itself, it is better than air at supporting the combustion of other ob-jects. This can be explained by noting that N2O can decompose to form an atmospherethat is one-third O2 by volume, whereas normal air is only 21% oxygen by volume.
2 N2O(g) 88n 2 N2(g) � O2(g)
For many years, the endings -ous and -ic were used to distinguish between the lowestand highest of a pair of oxidation numbers. N2O is nitrous oxide because the oxidationnumber of the nitrogen is �1. NO is nitric oxide because the oxidation number of thenitrogen is �2.
Enormous quantities of nitrogen oxide, or nitric oxide, are generated each year by thereaction between the N2 and O2 in the atmosphere, catalyzed by a stroke of lightningpassing through the atmosphere or by the hot walls of an internal combustion engine.
N2(g) � O2(g) 88n 2 NO(g)
One of the reasons for lowering the compression ratio of automobile engines in recentyears is to decrease the temperature of the combustion reaction, thereby decreasing theamount of NO emitted into the atmosphere.
NO can be prepared in the laboratory by reacting copper metal with dilute nitric acid.
3 Cu(s) � 8 HNO3(aq) 88n 3 Cu(NO3)2(aq) � 2 NO(g) � 4 H2O(l)
The NO molecule contains an odd number of valence electrons. As a result, it is impossi-ble to write a Lewis structure for the molecule in which all of the electrons are paired (seeFigure N.15). When NO gas is cooled, pairs of NO molecules combine in a reversible re-action to form a dimer (from Greek, meaning “two parts”), with the formula N2O2, inwhich all of the valence electrons are paired, as shown in Figure N.15.
NO reacts rapidly with O2 to form nitrogen dioxide (once known as nitrogen perox-ide), which is a dark brown gas at room temperature.
2 NO(g) � O2(g) 88n 2 NO2(g)
NO2 can be prepared in the laboratory by heating certain metal nitrates until they de-compose.
2 Pb(NO3)2(s) 88n 2 PbO(s) � 4 NO2(g) � O2(g)
It can also be made by reacting copper metal with concentrated nitric acid.
Cu(s) � 4 HNO3(aq) 88n Cu(NO3)2(aq) � 2 NO2(g) � 2 H2O(l)
NO2 also has an odd number of electrons and therefore contains at least one unpaired elec-tron in its Lewis structures. NO2 dimerizes at low temperatures to form N2O4 molecules,in which all the electrons are paired, as shown in Figure N.15.
Mixtures of NO and NO2 combine when cooled to form dinitrogen trioxide, N2O3,which is a blue liquid. The formation of a blue liquid when either NO or NO2 is cooled
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NONMETAL 27
therefore implies the presence of at least a small portion of the other oxide because N2O2
and N2O4 are both colorless.By carefully removing water from concentrated nitric acid at low temperatures with a
dehydrating agent we can form dinitrogen pentoxide.
4 HNO3(aq) � P4O10(s) 88n 2 N2O5(s) � 4 HPO3(s)
N2O5 is a colorless solid that decomposes in light or on warming to room temperature. Asmight be expected, N2O5 dissolves in water to form nitric acid.
N2O5(s) � H2O(l) 88n 2 HNO3(aq)
N.6 THE CHEMISTRY OF PHOSPHORUSPhosphorus was the first element whose discovery can be traced to a single individual. In1669, while searching for a way to convert silver into gold, Hennig Brand obtained a white,waxy solid that glowed in the dark and burst spontaneously into flame when exposed toair. Brand made the substance by evaporating the water from urine and allowing the blackresidue to putrefy for several months. He then mixed the residue with sand, heated themixture in the presence of a minimum of air, and collected under water the volatile prod-ucts that distilled out of the reaction flask.
Phosphorus forms a number of compounds that are direct analogs of nitrogen-containing compounds. However, the fact that elemental nitrogen is virtually inert at roomtemperature, whereas elemental phosphorus can burst spontaneously into flame when ex-posed to air, shows that there are differences between the elements as well. Phosphorusoften forms compounds with the same oxidation numbers as the analogous nitrogen com-pounds, but with different formulas, as shown in Table N.5.
TABLE N.5 Nitrogen and Phosphorus Compounds with the Same Oxidation Numbers but Different Formulas
Nitrogen Compound Phosphorus Compound Oxidation Number
N2 P4 �0HNO2 (nitrous acid) H3PO3 (phosphorous acid) �3N2O3 P4O6 �3HNO3 (nitric acid) H3PO4 (phosphoric acid) �5NaNO3 (sodium nitrate) Na3PO4 (sodium phosphate) �5N2O5 P4O10 �5
The same factors that explain the differences between sulfur and oxygen can be usedto explain the differences between phosphorus and nitrogen.
• NqN triple bonds are much stronger than PqP triple bonds.• POP single bonds are stronger than NON single bonds.• Phosphorus is much less electronegative than nitrogen.• Phosphorus can expand its valence shell to hold more than eight electrons, but ni-
trogen cannot.
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28 NONMETAL
The Effect of Differences in the XOOX and XqqX Bond Strengths
The ratio of the radii of phosphorus and nitrogen atoms is the same as the ratio of the radiiof sulfur and oxygen atoms, within experimental error.
� � 1.6
As a result, PqP triple bonds are much weaker than NqN triple bonds, for the same rea-son that SPS double bonds are weaker than OPO double bonds, namely, phosphorusatoms are too big to come close enough together to form strong multiple bonds.
Each atom in an N2 molecule completes its octet of valence electrons by sharing threepairs of electrons with a single neighboring atom. Because phosphorus doesn’t form strongmultiple bonds with itself, elemental phosphorus consists of tetrahedral P4 molecules inwhich each atom forms single bonds with three neighboring atoms, as shown in FigureN.16.
0.110 nm��0.070 nm
Covalent radius of phosphorus����
Covalent radius of nitrogen
P4, white phosphorus
P
P
P
P
FIGURE N.16 Tetrahedral P4 molecule.
Phosphorus is a white solid with a waxy appearance, which melts at 44.1°C and boilsat 287°C. It is made by reducing calcium phosphate with carbon in the presence of silica(sand) at very high temperatures.
2 Ca3(PO4)2(s) � 6 SiO2(s) � 10 C(s) 88n 6 CaSiO3(s) � P4(s) � 10 CO(g)
White phosphorus is stored under water because the element spontaneously bursts intoflame in the presence of oxygen at temperatures only slightly above room temperature.Although phosphorus is insoluble in water, it is very soluble in carbon disulfide. Solutionsof P4 in CS2 are reasonably stable; as soon as the CS2 evaporates, however, the phospho-rus bursts into flame.
The POPOP bond angle in a tetrahedral P4 molecule is only 60°. This very small an-gle produces a considerable amount of strain in the P4 molecule, which can be relieved bybreaking one of the POP bonds. Phosphorus therefore forms other allotropes by openingup the P4 tetrahedron. When white phosphorus is heated to 300°C, one bond inside eachP4 tetrahedron is broken, and the P4 molecules link together to form a polymer (from theGreek words pol, “many,” and meros, “parts”) with the structure shown in Figure N.17.This allotrope of phosphorus is dark red, and its presence in small traces often gives whitephosphorus a light yellow color. Red phosphorus is more dense (2.16 g/cm3) than whitephosphorus (1.82 g/cm3) and is much less reactive at normal temperatures.
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NONMETAL 29
Red phosphorus
P P
P
P
P P
P
P
P P
P
P
The Effect of Differences in the Strengths of PPPX and NPPX Double Bonds
The size of a phosphorus atom also interferes with its ability to form double bonds to otherelements such as oxygen, nitrogen, and sulfur. As a result, phosphorus tends to form com-pounds that contain two POO single bonds, where nitrogen would form an NPO doublebond. Nitrogen forms the nitrate (NO3
�) ion, for example, in which it has an oxidationnumber of �5. When phosphorus forms an ion with the same oxidation number, it is thephosphate (PO4
3�) ion, as shown in Figure N.18.
O
N
O O
P
O
O
OO
– 3–
Similarly, nitrogen forms nitric acid, HNO3, which contains an NPO double bond, whereasphosphorus forms phosphoric acid, H3PO4, which contains POO single bonds, as shownin Figure N.19.
P
O
O
O HHH NO OO
O H
The Effect of Differences in the Electronegativities of Phosphorus and Nitrogen
Because phosphorus is less electronegative than nitrogen, it is more likely to exhibit pos-itive oxidation numbers. The most important oxidation numbers for phosphorus are �3,�3, and �5 (see Table N.6).
FIGURE N.17 Portion of the polymeric chain in red phosphorus.
FIGURE N.18 Lewis structures for the NO3� and PO4
3� ions.
FIGURE N.19 Lewis structures for nitric acid (HNO3) andphosphoric acid (H3PO4).
TABLE N.6 Common Oxidation Numbers of Phosphorus
Oxidation Number Examples
�3 Ca3P2
�3 PF3, P4O6, H3PO3
�5 PF5, P4O10, H3PO4, PO43�
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30 NONMETAL
Because it is more electronegative than most metals, phosphorus reacts with metals atelevated temperatures to form phosphides, in which it has an oxidation number of �3.
6 Ca(s) � P4(s) 88n 2 Ca3P2(s)
The metal phosphides react with water to produce a poisonous, highly reactive, colorlessgas known as phosphine (PH3), which has one of the foulest odors the authors have en-countered.
Ca3P2(s) � 6 H2O(l) 88n 2 PH3(g) � 3 Ca2�(aq) � 6 OH�(aq)
Samples of PH3, the phosphorus analog of ammonia, are often contaminated by traces ofP2H4, the phosphorus analog of hydrazine. As if the toxicity and odor of PH3 were notenough, mixtures of PH3 and P2H4 burst spontaneously into flame in the presence ofoxygen.
Compounds such as Ca3P2 and PH3, in which phosphorus has a negative oxidation num-ber, are far outnumbered by compounds in which the oxidation number of phosphorus ispositive. Phosphorus burns in O2 to produce P4O10 in a reaction that gives off extraordi-nary amounts of energy in the form of heat and light.
P4(s) � 5 O2(g) 88n P4O10(s) �H° � �2984 kJ/molrxn
When phosphorus burns in the presence of a limited amount of O2, P4O6 is produced.
P4(s) � 3 O2(g) 88n P4O6(s) �H° � �1640 kJ/molrxn
P4O6 consists of a tetrahedron in which an oxygen atom has been inserted into each POPbond in the P4 molecule (see Figure N.20). P4O10 has an analogous structure, with an ad-ditional oxygen atom bound to each of the four phosphorus atoms.
O
PP
OO
O
P
O P
O
P4O6
O
OO
PP
OO
O
P
O P
O
O
OP4O10
P4O6 and P4O10 react with water to form phosphorous acid, H3PO3, and phosphoricacid, H3PO4, respectively.
P4O6(s) � 6 H2O(l) 88n 4 H3PO3(aq)P4O10(s) � 6 H2O(l) 88n 4 H3PO4(aq)
(P4O10 has such a high affinity for water that it is commonly used as a dehydrating agent.)Phosphorous acid, H3PO3, and phosphoric acid, H3PO4, are examples of a large class ofoxyacids of phosphorus. Lewis structures for some of these oxyacids and their relatedoxyanions are given in Figure N.21.
FIGURE N.20 Structures of P4O6 and P4O10.
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NONMETAL 31
Hypophosphite, H2PO2–
O
H
OH P
Phosphite, HPO32–
O
O
OH P
2–
Phosphate, PO43–
O
O
OO P
3–
–
O O
O O
OO OP PHH
O O
O O
OO OH P P
H H
O
O
O HP
HTriphosphoric acid, H5P3O10 Triphosphate, P3O10
5–
Diphosphate, P2O74–
(pyrophosphate)
5–
4–
O O
O O
OO OP P
O
O
OP
O O
O O
OO OP P
H HDiphosphoric acid, H4P2O7
(pyrophosphoric acid)
O
H
OH P H
Hypophosphorous acid, H3PO2
O
O
OH P H
HPhosphorous acid, H3PO3
O
O
OO PH H
HPhosphoric acid, H3PO4
Oxyacid Oxyanion
The Effect of Differences in the Abilities of Phosphorus and Nitrogen to ExpandTheir Valence Shell
The reaction between ammonia and fluorine stops at NF3 because nitrogen uses the 2s, 2px,2py, and 2pz orbitals to hold valence electrons. Nitrogen atoms can therefore hold a maxi-mum of eight valence electrons. Phosphorus, however, can expand the valence shell to hold10 or more electrons. Thus, phosphorus can react with fluorine to form both PF3 and PF5.Phosphorus can even form the PF6
� ion, in which there are 12 valence electrons on thecentral atom, as shown in Figure N.22.
FIGURE N.21 Oxyacids of phosphorus and their oxyanions.
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32 NONMETAL
F
FF
F
P
F
F F
F
F
P
F
F
PF6–
PF5
N.7 THE CHEMISTRY OF THE HALOGENSThere are six elements in Group VIIA, the next-to-last column of the periodic table. Asexpected, these elements have certain properties in common. They all form diatomic mol-ecules (H2, F2, Cl2, Br2, I2, and At2), for example, and they all form negatively chargedions (H�, F�, Cl�, Br�, I�, and At�).
When the chemistry of these elements is discussed, hydrogen is separated from the oth-ers and astatine is ignored because it is radioactive. (The most stable isotopes of astatinehave half-lives of less than a minute. As a result, the largest samples of astatine compoundsstudied to date have weighed less than 50 ng.) Discussions of the chemistry of the elementsin Group VIIA therefore focus on four elements: fluorine, chlorine, bromine, and iodine.These elements are called the halogens (from the Greek words hals, “salt,” and gennan,“to form or generate”) because they are literally the salt formers.
None of the halogens can be found in nature in their elemental form. They are invari-ably found as salts of the halide ions (F�, Cl�, Br�, and I�). Fluoride ions are found inminerals such as fluorite (CaF2) and cryolite (Na3AlF6). Chloride ions are found in rocksalt (NaCl), in the oceans, which are roughly 2% Cl� ion by weight, and in lakes that havea high salt content, such as the Great Salt Lake in Utah, which is 9% Cl� ion by weight.Both bromide and iodide ions are found at low concentrations in the oceans, as well as inbrine wells in Louisiana, California, and Michigan.
The Halogens in Their Elemental Form
Fluorine (F2), a highly toxic, colorless gas, is the most reactive element known—so reac-tive that asbestos, water, and silicon burst into flame in its presence. It is so reactive it evenforms compounds with Kr, Xe, and Rn, elements that were once thought to be inert. Flu-orine is such a powerful oxidizing agent that it can coax elements into unusually high ox-idation numbers, as in AgF2, PtF6, and IF7.
Fluorine is so reactive that it is difficult to find a container in which it can be stored.F2 attacks both glass and quartz, for example, and it causes most metals to burst into flame.Fluorine is handled in equipment built out of certain alloys of copper and nickel. It stillreacts with the alloys, but it forms a layer of a fluoride compound on the surface that pro-tects the metal from further reaction.
Fluorine is used in the manufacture of Teflon—or poly(tetrafluoroethylene), (C2F4)n—which is used for everything from linings for pots and pans to gaskets that are inert tochemical reactions.
Chlorine (Cl2) is a highly toxic gas with a pale yellow-green color. Chlorine is a verystrong oxidizing agent, which is used commercially as a bleaching agent and as a disinfec-tant. It is strong enough to oxidize the dyes that give wood pulp its yellow or brown color,for example, thereby bleaching out this color, and strong enough to destroy bacteria andthereby act as a germicide. Large quantities of chlorine are used each year to make solventssuch as carbon tetrachloride (CCl4), chloroform (CHCl3), dichloroethylene (C2H2Cl2), andtrichloroethylene (C2HCl3).
FIGURE N.22 Structures of PF5 and the PF6� ion.
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NONMETAL 33
Bromine (Br2) is a reddish-orange liquid with an unpleasant, choking odor. The nameof the element, in fact, comes from the Greek stem bromos, “stench.” Bromine is used toprepare flame retardants, fire-extinguishing agents, sedatives, antiknock agents for gaso-line, and insecticides.
Iodine is an intensely colored solid with an almost metallic luster. The solid is relativelyvolatile, and it sublimes when heated to form a violet-colored gas. Iodine has been used formany years as a disinfectant in “tincture of iodine.” Iodine compounds are used as catalysts,drugs, and dyes. Silver iodide (AgI) plays an important role in the photographic processand in attempts to make rain by seeding clouds. Iodide is also added to salt to protect againstgoiter, an iodine deficiency disease characterized by a swelling of the thyroid gland.
Some of the chemical and physical properties of the halogens are summarized inTable N.7.
TABLE N.7 Properties of F2, Cl2, Br2, and I2
First a
Melting Boiling Ionization Electrona Ionica
Point Point Energy Affinity Radius Density(°C) (°C) Color (kJ/mol) (kJ/mol) (nm) (g/cm3)
F2 �219.6 �188.1 Colorless 1681.0 328.0 0.136 1.513Cl2 �101 �34.0 Pale green 1251.1 348.8 0.181 1.655Br2 �7.2 59.5 Dark red-brown 1139.9 324.6 0.196 3.187I2 112.9 185.2 Dark violet, 1008.4 295.3 0.216 3.960
almost black
There is a regular increase in many of the properties of the halogens as we proceed downthe column from fluorine to iodine, including the melting point, boiling point, intensity ofthe color of the halogen, the radius of the corresponding halide ion, and the density of theelement. On the other hand, there is a regular decrease in the first ionization energy as wego down the column. As a result, there is a regular decrease in the oxidizing strength ofthe halogens from fluorine to iodine.
F2 � Cl2 � Br2 � I2oxidizing strength
This trend is mirrored by an increase in the reducing strength of the corresponding halides.
I� � Br� � Cl� � F�
reducing strength
Exercise N.5
Use the fact that Cl2 is a stronger oxidizing agent than Br2 to devise a way to prepare elemental bromine from an aqueous solution of the Br� ion.
Solution
When the chemistry of the main-group metals was introduced in Chapter 5, we found thatmetals can be prepared by reacting one of their salts with a metal that is a stronger reducingagent. Titanium metal, for example, can be prepared by reacting TiCl4 with magnesium metal.
TiCl4(l) � 2 Mg(s) 88n Ti(s) � 2 MgCl2(s)
aFor the atomic species.
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34 NONMETAL
Extending this argument to oxidizing agents suggests that we can produce Br2 by reactinga solution that contains the Br� ion with something that is an even stronger oxidizing agentthan Br2, such as Cl2 dissolved in water.
2 Br�(aq) � Cl2(aq) 88n Br2(aq) � 2 Cl�(aq)
Methods of Preparing the Halogens from Their Halides
The halogens can be made by reacting a solution of the corresponding halide ion with asubstance that is a stronger oxidizing agent than the halogen being prepared. Iodine, forexample, can be made by reacting the iodide ion with either bromine or chlorine.
2 I�(aq) � Br2(aq) 88n I2(aq) � 2 Br�(aq)
Bromine can be prepared by reacting bromide ions with a solution of Cl2 dissolved in water.
2 Br�(aq) � Cl2(aq) 88n Br2(aq) � 2 Cl�(aq)
To prepare Cl2, we need an unusually strong oxidizing agent, such as manganese dioxide(MnO2) or the permanganate ion (MnO4
�).
2 Cl�(aq) � MnO2(aq) � 4 H�(aq) 88n Cl2(aq) � Mn2�(aq) � 2 H2O(l)
The synthesis of fluorine escaped the efforts of chemists for almost 100 years. Part ofthe problem involved finding an oxidizing agent strong enough to oxidize the F� ion to F2.The task of preparing fluorine was made even more difficult by the extraordinary toxicityof both F2 and the hydrogen fluoride (HF) used to make it.
The best way of producing strong reducing agents is to pass an electric current througha salt of the metal. Sodium, for example, can be prepared by the electrolysis of moltensodium chloride.
electrolysis2 NaCl(l) 888888n 2 Na(s) � Cl2(g)
The same process can be used to generate strong oxidizing agents, such as F2.Attempts to prepare fluorine by electrolysis, however, were initially unsuccessful.
Humphry Davy, who prepared potassium, sodium, barium, strontium, calcium, and mag-nesium by electrolysis, repeatedly tried to prepare F2 by the electrolysis of fluorite (CaF2),and succeeded only in ruining his health. Joseph Louis Gay-Lussac and Louis JacquesThenard, who prepared elemental boron for the first time, also tried to prepare fluorineand suffered from very painful exposures to hydrogen fluoride. George and Thomas Knoxwere badly poisoned during their attempts to make fluorine, and both Paulin Louyet andJerome Nickles died from fluorine poisoning.
Finally, in 1886 Henri Moissan successfully isolated F2 gas from the electrolysis of amixed salt of KF and HF and noted that crystals of silicon burst into flame when mixedwith the gas. Electrolysis of KHF2 is still used to prepare fluorine today.
electrolysis2 KHF2(s) 888888n H2(g) � F2(g) � 2 KF(s)
A schematic drawing of the cell in which KHF2 is electrolyzed to produce F2 gas at the an-ode and H2 gas at the cathode is shown in Figure N.23.
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NONMETAL 35
Common Oxidation Numbers for the Halogens
Fluorine is the most electronegative element in the periodic table. As a result, it has anoxidation number of �1 in all its compounds. Because chlorine, bromine, and iodine areless electronegative, it is possible to prepare compounds in which those elements have ox-idation numbers of �1, �3, �5, and �7, as shown in Table N.8.
Anodeconnection
F2 outlet
HF inlet H2 outlet
Cellcover
Carbonanode
Coolingjacket
Steelcathode
Gas separation skirt
Molten solutionof KF dissolvedin HF
General Trends in Halogen Chemistry
There are several patterns in the chemistry of the halogens.
• The chemistry of fluorine is simplified by the fact that it is the most electronegative element in the periodic table that forms compounds and by the fact that it cannotexpand its valence shell to hold more than eight valence electrons.
• Chlorine, bromine, and iodine can expand their valence shells to hold as many as14 valence electrons.
• The chemistry of the halogens is dominated by oxidation–reduction reactions.
The Hydrogen Halides (HX)
The hydrogen halides are compounds that contain hydrogen attached to one of the halo-gens (HF, HCl, HBr, and HI). The compounds are all colorless gases that are soluble in
TABLE N.8 Common Oxidation Numbers for the Halogens
Oxidation Number Examples
�1 CaF2, HCl, NaBr, AgI�0 F2, Cl2, Br2, I2
�1 HClO, ClF�3 HClO2, ClF3
�5 HClO3, BrF5, BrF6�, IF5
�7 HClO4, BrF6�, IF7
FIGURE N.23 Schematic drawing of the cell in which KHF2 is elec-trolyzed.
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36 NONMETAL
water. Up to 512 mL of HCl gas can dissolve in 1 mL of water at 0°C and 1 atm, for ex-ample. Each of the hydrogen halides ionizes to at least some extent when it dissolves inwater.
H2OHCl(g) 88n H�(aq) � Cl�(aq)
Exercise N.6
Explain why it is easy to believe that HCl is an ionic compound. Describe the best evi-dence that it isn’t an ionic compound.
Solution
It is easy to believe HCl is an ionic compound because it forms ions when it dissolves inwater.
H2OHCl(g) 88n H�(aq) � Cl�(aq)
At first glance, the reaction seems to be similar to the reaction that occurs when ionic com-pounds dissolve in water.
H2ONaCl(s) 88n Na�(aq) � Cl�(aq)
The best evidence that HCl isn’t an ionic compound is the fact that it is a gas at room tem-perature, and ionic compounds are invariably solids at room temperature.
Some chemists try to distinguish between the behavior of HCl and NaCl when they dis-solve in water as follows. They argue that HCl ionizes when it dissolves in water becauseions are created by the reaction. NaCl, on the other hand, dissociates in water because NaClalready consists of Na� and Cl� ions in the solid.
Several of the hydrogen halides can be prepared directly from the elements. Mixturesof H2 and Cl2, for example, react with explosive violence in the presence of light to formHCl.
H2(g) � Cl2(g) 88n 2 HCl(g)
Because chemists are usually more interested in aqueous solutions of hydrogen halidesthan they are in the pure gases, the compounds are usually synthesized in water. Aqueoussolutions of the hydrogen halides are often called mineral acids because they are literallyacids prepared from minerals. Hydrochloric acid is prepared by reacting table salt with sul-furic acid, for example, and hydrofluoric acid is prepared from fluorite and sulfuric acid.
2 NaCl(s) � H2SO4(aq) 88n 2 HCl(aq) � Na2SO4(aq)CaF2(s) � H2SO4(aq) 88n 2 HF(aq) � CaSO4(aq)
The acids are purified by taking advantage of the ease with which HF and HCl gas boilout of the solutions. The gas given off when one of the solutions is heated is collected andthen redissolved in water to give relatively pure samples of the mineral acid.
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NONMETAL 37
The Interhalogen Compounds
Interhalogen compounds are formed by reactions between different halogens. All possibleinterhalogen compounds of the type XY are known. Bromine reacts with chlorine, for ex-ample, to give BrCl, which is a gas at room temperature.
Br2(l) � Cl2(g) 88n 2 BrCl(g)
Interhalogen compounds with the general formulas XY3, XY5, and even XY7 are formedwhen pairs of halogens react. Chlorine reacts with fluorine, for example, to form chlorinetrifluoride.
Cl2(g) � 3 F2(g) 88n 2 ClF3(g)
The compounds are easiest to form when Y is fluorine. Iodine is the only halogen thatforms an XY7 interhalogen compound, and it does so only with fluorine.
ClF3 and BrF5 are extremely reactive compounds. ClF3 is so reactive that wood, as-bestos, and even water spontaneously burn in its presence. The compounds are excellentfluorinating agents, which tend to react with each other to form positive ions such as ClF2
�
and BrF4� and negative ions such as ClF4
� and BrF6�.
2 BrF5(l) 88n [BrF4�][BrF6
�](s)
Neutral Oxides of the Halogens
Under certain conditions, it is possible to isolate neutral oxides of the halogens, such asCl2O, Cl2O3, ClO2, Cl2O4, Cl2O6, and Cl2O7. Cl2O7, for example, can be obtained by de-hydrating perchloric acid, HClO4. The oxides are notoriously unstable compounds that ex-plode when subjected to either thermal or physical shock. Some are so unstable they det-onate when warmed to temperatures above �40°C.
Oxyacids of the Halogens and Their Salts
Chlorine reacts with the OH� ion to form chloride ions and hypochlorite (OCl�) ions.
Cl2(aq) � 2 OH�(aq) 88n Cl�(aq) � OCl�(aq) � H2O(l)
This is a disproportionation reaction in which one-half of the chlorine atoms are oxidizedto hypochlorite ions and the other half are reduced to chloride ions.
When the solution is hot, this reaction gives a mixture of the chloride and chlorate (ClO3�)
ions.
3 Cl2(aq) � 6 OH�(aq) 88n 5 Cl�(aq) � ClO3�(aq) � 3 H2O(l)
Under carefully controlled conditions, it is possible to convert a mixture of the chlorateand hypochlorite ions into a solution that contains the chlorite (ClO2
�) ion.
ClO3�(aq) � ClO�(aq) 88n 2 ClO2
�(aq)
Cl� � OCl� � H2O
Reduction
Oxidation
0 �1
Cl2 � 2 OH�
�1
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38 NONMETAL
The last member of this class of compounds, the perchlorate ion (ClO4�), is made by elec-
trolyzing solutions of the chlorate ion.The names of the oxyanions of the halogens use the endings -ite and -ate to indicate
low and high oxidation numbers and the prefixes hypo- and per- to indicate the very low-est and very highest oxidation numbers, as shown in Table N.9. Each of the ions can beconverted into an oxyacid, which is named by replacing the -ite ending with -ous and the-ate ending with -ic.
TABLE N.9 Oxyanions and Oxyacids of Chlorine
Oxidation StateOxyanion Name Oxyacid Name of Chlorine
ClO� Hypochlorite HClO Hypochlorous acid �1ClO2
� Chlorite HOClO Chlorous acid �3ClO3
� Chlorate HOClO2 Chloric acid �5ClO4
� Perchlorate HOClO3 Perchloric acid �7
The hypochlorite (OCl�) ion is the active ingredient in liquid bleaches, such as Clorox.Calcium salts of the ion can be found in dry bleaches, such as Clorox 2. Ca(OCl)2 is alsothe active ingredient in most commercial products used to “chlorinate” swimming pools.
N.8 THE CHEMISTRY OF THE RARE GASESIn 1892 Lord Rayleigh noted that nitrogen isolated from air was more dense than nitro-gen prepared by decomposing ammonia. William Ramsay attacked this problem by puri-fying a sample of nitrogen gas to remove any moisture, carbon dioxide, and organic cont-aminants. He then passed the purified gas over hot magnesium metal, which reacts withnitrogen to form the nitride.
3 Mg(s) � N2(s) 88n Mg3N2(s)
When he was finished, Ramsay was left with a small residue of gas that occupied roughly1/80th of the original volume. He excited the gas in an electric discharge tube and foundthat the resulting emission spectrum contained lines that differed from those of all knowngases. After repeated discussions of the results of these experiments, Rayleigh and Ramsayjointly announced the discovery of a new element, which they named argon from the Greekword meaning “lazy one” because the gas refused to react with any element or compoundthey tested.
Argon didn’t fit into any of the known families of elements in the periodic table, butits atomic weight suggested that it might belong to a new group that could be insertedbetween chlorine and potassium. Shortly after reporting the discovery of argon in 1894,Ramsay found another unreactive gas when he heated a mineral of uranium. The lines inthe spectrum of the gas also occurred in the spectrum of the sun, which led Ramsay toname the element helium (from the Greek word helios, “sun”). Experiments with liquidair led Ramsay to a third gas, which he named krypton (“the hidden one”). Experimentswith liquid argon led him to a fourth gas, neon (“the new one”), and finally a fifth gas,xenon (“the stranger”).
These elements were discovered between 1894 and 1898. Because Moissan had onlyrecently isolated fluorine for the first time and fluorine was the most active of the known
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NONMETAL 39
elements, Ramsay sent a sample of argon to Moissan to see whether it would react withfluorine. It did not. The failure of Moissan’s attempts to react argon with fluorine, coupledwith repeated failures by other chemists to get the more abundant of the gases to undergochemical reaction, eventually led to their being labeled inert gases. The development of theelectronic theory of atoms did little to dispel this notion because it was obvious that thesegases had very stable electron configurations. As a result, the elements were labeled “in-ert gases” in almost every textbook and periodic table until about 30 years ago.
In 1962 Neil Bartlett found that PtF6 was a strong enough oxidizing agent to removean electron from an O2 molecule.
PtF6(g) � O2(g) 88n [O2�][PtF6
�](s)
Bartlett realized that the first ionization energy of Xe (1170 kJ/mol) was slightly smallerthan the first ionization energy of the O2 molecule (1177 kJ/mol). He therefore predictedthat PtF6 might also react with Xe. When he ran the reaction, he isolated the first com-pound of a Group VIIIA element.
Xe(g) � PtF6(g) 88n [Xe�][PtF6�](s)
A few months later, workers at the Argonne National Laboratory near Chicago foundthat Xe reacts with F2 to form XeF4. Since that time, more than 200 compounds of Kr, Xe,and Rn have been isolated. No compounds of the more abundant elements in the group(He, Ne, and Ar) have yet been isolated. However, the fact that elements in the family canundergo chemical reactions has led to the use of the term rare gases rather than inert gasesto describe these elements.
Compounds of xenon are by far the most numerous of the rare gas compounds. Withthe exception of XePtF6, rare gas compounds have oxidation numbers of �2, �4, �6, and�8, as shown by the examples cited in Table N.10.
There is some controversy over whether the rare gases should be viewed as having theoutermost shell of electrons filled (in which case they should be labeled Group VIIIA) orempty (in which case they should be labeled Group 0). The authors believe these elementsshould be labeled Group VIIIA because they behave as if they contribute eight valenceelectrons when they form compounds.
TABLE N.10 Compounds of Xenon and Their Oxidation Numbers
Compound Oxidation Number Compound Oxidation Number
XeF� �2 XeO3 �6XeF2 �2 XeOF4 �6Xe2F3
� �2 XeO2F2 �6XeF3
� �4 XeO3F� �6XeF4 �4 XeO4 �8XeOF2 �4 XeO6
4� �8XeF5
� �6 XeO3F2 �8XeF6 �6 XeO2F4 �8Xe2F11
� �6 XeOF5� �8
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The synthesis of most xenon compounds starts with the reaction between Xe and F2 athigh temperatures (250–400°C) to form a mixture of XeF2, XeF4, and XeF6.
3 Xe(g) � 6 F2(g) 88n XeF2(s) � XeF4(s) � XeF6(s)
The positively charged XeFn� ions are then made by reacting XeF2, XeF4, or XeF6 with
either AsF5, SbF5, or BiF5.
XeF2(s) � SbF5(l) 88n [XeF�][SbF6�](s)
2 XeF2(s) � AsF5(g) 88n [Xe2F3�][AsF6
�](s)XeF4(s) � BiF5(s) 88n [XeF3
�][BiF6�](s)
2 XeF6(s) � AsF5(g) 88n [Xe2F11�][AsF6
�](s)
Oxides of xenon, such as XeOF2, XeOF4, XeO2F2, XeO3F2, XeO2F4, XeO3, and XeO4, areprepared by reacting XeF4 or XeF6 with water. The XeO6
4� ion, for example, is producedwhen XeF6 dissolves in strong base.
2 XeF6(s) � 16 OH�(aq) 88n XeO64�(aq) � Xe(g) � O2(g) � 12 F�(aq) � 8 H2O(l)
Some xenon compounds are relatively stable. XeF2, XeF4, and XeF6, for example, arestable solids that can be purified by sublimation in a vacuum at 25°C. XeOF4 and Na4XeO6
are also reasonably stable. Others, such as XeO3, XeO4, XeOF2, XeO2F2, XeO3F2, andXeO2F4, are unstable compounds that can decompose violently.
The principal use of rare gas compounds at present is as the light-emitting componentin lasers. Mixtures of 10% Xe, 89% Ar, and 1% F2, for example, can be “pumped,” or ex-cited, with high energy electrons to form excited XeF molecules, which emit a photon witha wavelength of 354 nm.
N.9 THE INORGANIC CHEMISTRY OF CARBONFor more than 200 years, chemists have divided compounds into two categories. Those thatwere isolated from plants or animals were called organic, while those extracted from oresand minerals were inorganic. Organic chemistry is often defined as the chemistry of car-bon. But that definition would include calcium carbonate (CaCO3) and graphite, whichmore closely resemble inorganic compounds. We will therefore define organic chemistryas the study of compounds such as formic acid (HCO2H), methane (CH4), and vitamin C(C6H8O6) that contain both carbon and hydrogen. This section focuses on inorganic car-bon compounds.
Elemental Forms of Carbon: Graphite, Diamond, Coke, and Carbon Black
Carbon occurs as a variety of allotropes. There are two crystalline forms—diamond andgraphite—and a number of amorphous (noncrystalline) forms, such as charcoal, coke, andcarbon black.
References to the characteristic hardness of diamond (from the Greek adamas, “in-vincible”) date back at least 2600 years. It was not until 1797, however, that SmithsonTennant was able to show that diamonds consist solely of carbon. The properties of diamondare remarkable. It is among the least volatile substances known (MP � 3550°C, BP �4827°C), it is the hardest substance known, and it expands less on heating than any othermaterial.
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The properties of diamond are a logical consequence of its structure. Carbon, with fourvalence electrons, forms covalent bonds to four neighboring carbon atoms arranged towardthe corners of a tetrahedron, as shown in Figure N.24. Each of the sp3-hybridized atoms isthen bound to four other carbon atoms, which form bonds to four other carbon atoms, andso on. As a result, a perfect diamond can be thought of as a single giant molecule. Thestrength of the individual COC bonds and their arrangement in space give rise to the un-usual properties of diamond.
FIGURE N.24 The simplest repeating unit in diamond.
In some ways, the properties of graphite are like those of diamond. Both compoundsboil at 4827°C, for example. But graphite is also very different from diamond. Diamond(3.514 g/cm3) is significantly more dense than graphite (2.26 g/cm3). Whereas diamond isthe hardest substance known, graphite is one of the softest. Diamond is an excellent insu-lator, with little or no tendency to carry an electric current. Graphite is such a good con-ductor of electricity that graphite electrodes are used in electrical cells.
The physical properties of graphite can be understood from the structure of the solidshown in Figure N.25. Graphite consists of extended planes of sp2-hybridized carbon atomsin which each carbon is tightly bound to three other carbon atoms. (It takes 477 kJ to breaka mole of the bonds within the planes.) The strong bonds between carbon atoms withineach plane explain the exceptionally high melting point and boiling point of graphite. Thebonds between planes of carbon atoms, however, are relatively weak. (The force of attrac-tion between planes is only 17 kJ/mol.) Because the bonds between planes are weak, it iseasy to deform the solid by allowing one plane of atoms to move relative to another. As aresult, graphite is soft enough to be used in pencils and as a lubricant in motor oil.
The characteristic properties of graphite and diamond might lead you to expect thatdiamond would be more stable than graphite. This isn’t what is observed experimentally.Graphite at 25°C and 1 atm pressure is slightly more stable than diamond. At very high
FIGURE N.25 Portion of the structure of extendedplanes of carbon atoms found in graphite.
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42 NONMETAL
temperatures and pressures, however, diamond becomes more stable than graphite. In 1955General Electric developed a process to make industrial-grade diamonds by treatinggraphite with a metal catalyst at temperatures of 2000 to 3000 K and pressures above 125,000atm. Roughly 40% of industrial-quality diamonds are now synthetic. Although gem-quality diamonds can be synthesized, the costs involved are prohibitive.
Both diamond and graphite occur as regularly packed crystals. Other forms of carbonare amorphous—they lack a regular structure. Charcoal, carbon black, and coke are allamorphous forms of carbon. Charcoal results from heating wood in the absence of oxygen.To make carbon black, natural gas or other carbon compounds are burned in a limitedamount of air to give a thick, black smoke that contains extremely small particles of car-bon, which can be collected when the gas is cooled and passed through an electrostatic pre-cipitator. Coke is a more regularly structured material, closer in structure to graphite thaneither charcoal or carbon black, which is made from coal.
Carbides: Covalent, Ionic, and Interstitial
Carbon reacts with less electronegative elements at high temperatures to form compoundsknown as carbides. When carbon reacts with an element of similar size and electronega-tivity, a covalent carbide is produced. Silicon carbide, for example, is made by treating sil-icon dioxide from quartz with an excess of carbon in an electric furnace at 2300 K.
SiO2(s) � 3 C(s) 88n SiC(s) � 2 CO(g)
Covalent carbides have properties similar to those of diamond. Both SiC and diamond areinert to chemical reactions, except at very high temperatures; both have very high melt-ing points; and both are among the hardest substances known. SiC was first synthesized byEdward Acheson in 1891. Shortly thereafter, Acheson founded the Carborundum Com-pany to market the material. Then, as now, materials in this class are most commonly used asabrasives.
Compounds that contain carbon and one of the more active metals are called ioniccarbides.
CaO(s) � 3 C(s) 88n CaC2(s) � CO(g)
It is useful to think about ionic carbides as if they contained negatively charged carbonions: [Ca2�][C2
2�] or [Al3�]4[C4�]3. This model is useful because it explains why these car-bides burst into flame when added to water. Ionic carbides such as Al4C3 that formally con-tain the C4� ion react with water to form methane, which is ignited by the heat given offin the reaction.
C4� � 4 H2O 88n CH4 � 4 OH�
The ionic carbides such as CaC2 that formally contain the C22� ion react with water to
form acetylene, which is ignited by the heat of reaction.
C22� � 2 H2O 88n C2H2 � 2 OH�
At one time, miners’ lamps were fueled by the combustion of acetylene prepared from thereaction of calcium carbide with water.
The difference between covalent carbides and ionic carbides can be understood byadding compounds such as SiC, Al4C3, and CaC2 to the bond type triangle introduced in
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Chapter 5. When those compounds are added to Figure 5.10, for example, we find that SiCfalls well into the region expected for covalent compounds. CaC2, on the other hand, isclearly an ionic compound. Al4C3 falls on the borderline between ionic and covalent, whichis consistent with the fact that the compound is hard—as one would expect for a covalentcarbide—and yet reacts with water to form methane—as might be expected for an ioniccarbide.
Interstitial carbides, such as tungsten carbide (WC), form when carbon combineswith a metal that has an intermediate electronegativity and a relatively large atomic ra-dius. In these compounds, the carbon atoms pack in the holes (interstices) betweenplanes of metal atoms. The interstitial carbides, which include TiC, ZrC, and MoC, re-tain the properties of metals. They act as alloys, rather than as either salts or covalentcompounds.
The Oxides of Carbon
Although the different forms of carbon are essentially inert at room temperature, theycombine with oxygen at high temperatures to produce a mixture of carbon monoxide andcarbon dioxide.
2 C(s) � O2(g) 88n 2 CO(g) �H° � �221.05 kJ/molrxn
C(s) � O2(g) 88n CO2(g) �H° � �393.51 kJ/molrxn
CO can also be obtained by reacting red-hot carbon with steam.
C(s) � H2O(g) 88n CO(g) � H2(g)
Because the mixture of gases is formed by the reaction of charcoal or coke with water itis often referred to as water gas. It is also known as town gas because it was once made bytowns and cities for use as a fuel. Water gas, or town gas, was a common fuel for both homeand industrial use before natural gas became readily available. The H2 burns to form wa-ter, and the CO is oxidized to CO2. Eventually, as our supply of natural gas is depleted, itwill become economical to replace natural gas with other fuels, such as water gas, that canbe produced from our abundant supply of coal.
Both CO and CO2 are colorless gases. CO boils at �191.5°C, and CO2 sublimes (passesdirectly from the solid to the gaseous state) at �78.5°C. Although CO has no odor or taste,CO2 has a faint, pungent odor and a distinctly acidic taste. Both are dangerous substancesbut at very different levels of exposure. Air contaminated with as little as 0.002 gram ofCO per liter can be fatal because CO binds tightly to the hemoglobin that carries oxygenthrough the blood. CO2 is not lethal until the concentration in the air approaches 15%. Atthat point, it has replaced so much oxygen that a person who attempts to breathe the at-mosphere suffocates. The danger of CO2 poisoning is magnified by the fact that CO2 isroughly 1.5 times more dense than the air in our atmosphere. Thus, CO2 can accumulateat the bottom of tanks or wells.
CO2 in the Atmosphere
Carbon dioxide influences the temperature of the atmosphere by a phenomenon knownas the greenhouse effect. The glass walls and ceilings of a greenhouse absorb some of thelower energy, longer wavelength radiation from sunlight thereby inevitably raising the tem-perature inside the building. CO2 in the atmosphere does exactly the same thing, it absorbs
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low energy, long wavelength radiation from the Sun that would otherwise be reflected backfrom the surface of the planet. Thus, CO2 in the atmosphere traps heat. Although thereare other factors at work, it is worth noting that Venus, whose atmosphere contains a greatdeal of CO2, has a surface temperature of roughly 400°C, whereas Mars, with little or noatmosphere, has a surface temperature of �50°C.
There are many sources of CO2 in the atmosphere. Over geologic time scales, the largestsource has been volcanos. Within the twentieth century, the combustion of petroleum, coal,and natural gas has made a significant contribution to atmospheric levels of CO2 (see Fig-ure N.8). Between 1958 and 1978, the average level of CO2 in the atmosphere increasedby 6%, from 315.8 to 334.6 ppm.
At one time, the amount of CO2 released to the atmosphere wasn’t a matter for con-cern because natural processes that removed CO2 from the atmosphere could compensatefor the CO2 that entered the atmosphere. The vast majority of the CO2 liberated by vol-canic action, for example, was captured by calcium oxide or magnesium oxide to form cal-cium carbonate or magnesium carbonate.
CaO(s) � CO2(g) 88n CaCO3(s)MgO(s) � CO2(g) 88n MgCO3(s)
CaCO3 is found as limestone or marble, or mixed with MgCO3 as dolomite. The amountof CO2 in deposits of carbonate minerals is at least several thousand times larger than theamount in the atmosphere.
CO2 also dissolves, to some extent, in water.
H2OCO2(g) 88n CO2(aq)
It then reacts with water to form carbonic acid, H2CO3.
CO2(aq) � H2O(l) 88n H2CO3(aq)
As a result of the reactions, the sea contains about 60 times more CO2 than the atmo-sphere.
Can the sea absorb more CO2 from the atmosphere, or is it near its level of saturation?Is the rate at which the sea absorbs CO2 greater than the rate at which we are adding it tothe atmosphere? The observed increase in the concentration of CO2 in recent years sug-gests pessimistic answers to those two questions. A gradual warming of the earth’s atmo-sphere could result from continued increases in CO2 levels, with adverse effects on the cli-mate and therefore the agriculture of at least the northern hemisphere.
The Chemistry of Carbonates: CO32� and HCO3
�
Eggshells are almost pure calcium carbonate. CaCO3 can also be found in the shells ofmany marine organisms and in both limestone and marble. The fact that none of thosesubstances dissolves in water suggests that CaCO3 is normally insoluble in water. Calciumcarbonate will dissolve in water saturated with CO2, however, because carbonated water(or carbonic acid) reacts with calcium carbonate to form calcium bicarbonate, which is sol-uble in water.
CaCO3(s) � H2CO3(aq) 88n Ca2�(aq) � 2 HCO3�(aq)
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When water rich in carbon dioxide flows through limestone formations, part of thelimestone dissolves. If the CO2 escapes from the water, or if some of the water evapo-rates, solid CaCO3 is redeposited. When this happens as water runs across the roof of acavern, stalactites, which hang from the roof of the cave, are formed. If the water dropsbefore the carbonate reprecipitates, stalagmites, which grow from the floor of the cave,are formed.
The chemistry of carbon dioxide dissolved in water is the basis of the soft drink in-dustry. The first artificially carbonated beverages were introduced in Europe at the end ofthe nineteenth century. Carbonated soft drinks today consist of carbonated water, a sweet-ening agent (sugar, saccharin, or aspartame), an acid to impart a sour or tart taste, flavor-ing agents, coloring agents, and preservatives. As much as 3.5 liters of gaseous CO2 dis-solve in a liter of soft drink to provide the characteristic “bite” associated with carbonatedbeverages.
Carbonate chemistry plays an important role in other parts of the food industry as well.Baking soda, or bicarbonate of soda, is sodium bicarbonate, NaHCO3, a weak base, whichis used to neutralize the acidity of other ingredients in a recipe. Baking powder is a mix-ture of baking soda and a weak acid, such as tartaric acid or calcium hydrogen phosphate(CaHPO4). When mixed with water, the acid reacts with the HCO3
� ion to form CO2 gas,which causes the dough or batter to rise.
HCO3�(aq) � H�(aq) 88n H2CO3(aq) 88n H2O(l) � CO2(g)
Before commercial baking powders were available, cooks obtained the same effect by mix-ing roughly a teaspoon of baking soda with a cup of sour milk or buttermilk. The acidsthat give sour milk and buttermilk their characteristic taste also react with the bicarbon-ate ion to give CO2.
AllotropeAnhydrousCarbideDiamagneticDimerDisproportionation
reaction
Haber processHalideHalogenHydrideOstwald processOxidizing agent
OxyacidOxyanionParamagneticPeroxideRare gasesReducing agent
PROBLEMS
Metals, Nonmetals, and Semimetals
1. List the elements that are nonmetals. Describe where these elements are found in theperiodic table.
2. Explain why semimetals (such as B, Si, Ge, As, Sb, Te, Po, and At) exist and describesome of their physical properties.
3. Which member of each of the following pairs of elements is more nonmetallic?(a) As or Bi (b) As or Se (c) As or S (d) As or Ge (e) As or P
KEY TERMS
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46 NONMETAL
The Chemistry of the Nonmetals
4. Which of the following elements can exist as a triatomic molecule?(a) hydrogen (b) helium (c) sulfur (d) oxygen (e) chlorine
5. Which of the following elements should form compounds with the formulas Na2X, H2X,XO2, and XF6?(a) B (b) C (c) N (d) O (e) S
6. Which of the following elements should form compounds with the formulas XH3, XF3,and Na3XO4?(a) Al (b) Ge (c) As (d) S (e) Cl
7. Which of the following can’t be found in nature? Explain why.(a) MgCl2 (b) CaCO3 (c) F2 (d) Na3AlF6 (e) NaCl
The Role of Nonmetal Elements in Chemical Reactions
8. Explain why more electronegative elements tend to oxidize less electronegative elements.9. Which of the following ions or molecules can be oxidized?
(a) H2SO3 (b) P4 (c) Cl� (d) SiO2 (e) PO43� (f) Mg2�
10. Which of the following ions or molecules can be reduced?(a) H2O (b) H2SO3 (c) HCl (d) CO2 (e) Mg2� (f) Na
Deciding What Is Oxidized and What Is Reduced
11. For each of the following reactions, identify what is oxidized and what is reduced.(a) Fe2O3(s) � 3 CO(g) n 2 Fe(s) � 3 CO2(g)(b) H2(g) � CO2(g) n H2O(g) � CO(g)(c) CH4(g) � 2 O2(g) n CO2(g) � 2 H2O(g)(d) 2 H2S(g) � 3 O2(g) n 2 SO2(g) � 2 H2O(g)
12. For each of the following reactions, identify what is oxidized and what is reduced.(a) PH3(g) � 3 Cl2(g) n PCl3(g) � 3 HCl(g)(b) 2 NO(g) � F2(g) n 2 NOF(g)(c) 2 Na(s) � 2 NH3(l) n 2 NaNH2(s) � H2(g)(d) 3 NO2(g) � H2O(l) n 2 HNO3(aq) � NO(g)
13. Hydrazine is made by a reaction known as the Raschig process.
2 NH3(aq) � NaOCl(aq) 88n N2H4(aq) � NaCl(aq) � H2O(l)
Decide whether this is an oxidation–reduction reaction. If it is, identify the compoundoxidized and the compound reduced.
14. The thiosulfate ion, S2O32�, is prepared by boiling solutions of sulfur dissolved in
sodium sulfite.
8 SO32�(aq) � S8(s) 88n 8 S2O3
2�(aq)
Is this an oxidation–reduction reaction? If it is, identify the compound oxidized andthe compound reduced.
15. Chlorine dioxide, ClO2, is used commercially as a bleach or a disinfectant because ofits excellent oxidizing ability. ClO2 is prepared by decomposing chlorous acid as follows.
8 HOClO(aq) 88n 6 ClO2(g) � Cl2(g) � 4 H2O(l)
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NONMETAL 47
Is this an oxidation–reduction reaction? If it is, identify the oxidizing agent and the re-ducing agent.
16. Nitric acid sometimes acts as an acid (as a source of the H� ion) and sometimes as anoxidizing agent. For each of the following reactions, decide whether HNO3 acts as anacid or as an oxidizing agent.(a) Na2CO3(s) � 2 HNO3(aq) n 2 NaNO3(aq) � CO2(g) � H2O(l)(b) 3 P4(s) � 20 HNO3(aq) � 8 H2O(l) n 12 H3PO4(aq) � 20 NO(g)(c) Al2O3(s) � 6 HNO3(aq) n 2 Al(NO3)3(aq) � 3 H2O(l)(d) 3 Cu(s) � 8 HNO3(aq) n 3 Cu(NO3)2(aq) � 2 NO(g) � 4 H2O(l)
17. Which of the following would you expect to be the best oxidizing agent?(a) Na (b) H2 (c) N2 (d) P4 (e) O2
18. Which of the following would you expect to be the best reducing agent?(a) Na� (b) F� (c) Na (d) Br2 (e) Fe3�
19. For each of the following pairs of elements, determine which is the better reducing agent?(a) P4 or As (b) As or S8 (c) P4 or S8 (d) S8 or Cl2 (e) C or O2
Predicting the Products of Chemical Reactions
20. Predict the products of the following reactions.(a) Mg(s) � N2(g) n(b) Li(s) � O2(g) n(c) Br2(l) � I�(aq) n
21. Predict the products of the following reactions.(a) SO2(g) � H2O(l) n(b) Cl2(g) � OH�(aq) n(c) CO2(g) � H2O(l) n
22. Predict the products of the following reactions.(a) HCl(g) � H2O(l) n(b) P4O10(s) � H2O(l) n(c) NO2(g) � H2O(l) n
23. Predict the products of the following reactions.(a) S8(s) � O2(g) n(b) Al(s) � I2(s) n(c) P4(s) � F2(g) n
The Chemistry of Hydrogen
24. Describe three ways of preparing small quantities of H2 in the lab.25. Explain why it is not a good idea to prepare H2 by reacting sodium metal with a strong
acid.26. Give an example of a compound in which hydrogen has an oxidation number of �1;
of 0; of �1.27. Which of the following reactions produce a compound in which hydrogen has an oxi-
dation number of �1?(a) Li � H2 n(b) O2 � H2 n(c) S8 � H2 n(d) Cl2 � H2 n(e) Ca � H2 n
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48 NONMETAL
28. Use tables of first ionization energies and electronegativities to explain why it is so dif-ficult to decide whether hydrogen belongs in Group IA or Group VIIA of the peri-odic table.
29. Which of the following substances can be used as evidence for placing hydrogen inGroup IA? Which can be used as evidence for including hydrogen in Group VIIA?(a) CaH2 (b) AlH3 (c) H2S (d) H3PO4 (e) H2
30. The earth’s atmosphere once contained significant amounts of H2. Explain why onlytraces of H2 are left in the earth’s atmosphere, whereas the atmospheres of otherplanets—such as Jupiter, Saturn, and Neptune—contain large quantities of H2.
31. Use Lewis structures to explain what happens in the four reactions described in Sec-tion N.2 that can be used to prepare small quantities of H2 gas.
The Chemistry of Oxygen and Sulfur
32. Describe three ways of preparing small quantities of O2 in the lab.33. Describe the relationships among oxygen (O2), the peroxide ion (O2
2�), and the oxideion (O2�). Explain why the number of electrons shared by a pair of oxygen atomsdecreases as the oxidation number of the oxygen becomes more negative.
34. Which of the following elements or compounds could eventually produce O2 when itreacts with water?(a) Ba (b) BaO (c) BaO2 (d) Ba(OH)2 (e) BaNO3
35. Explain why the only compounds in which oxygen has a positive oxidation number arecompounds, such as OF2, that contain fluorine.
36. Explain why hydrogen peroxide can be either an oxidizing agent or a reducing agent.Describe at least one reaction in which H2O2 oxidizes another substance and one re-action in which it reduces another substance.
37. Write the Lewis structures for ozone, O3, and sulfur dioxide, SO2. Discuss the rela-tionship between the compounds.
38. Explain why elemental oxygen exists as O2 molecules, whereas elemental sulfur formsS8 molecules.
39. Explain why sulfur forms compounds such as SF4 and SF6, when oxygen can only formOF2.
40. Describe the relationship between the thiosulfate and sulfate ions and between thethiocyanate and cyanate ions. Use that relationship to predict the formula of the trithio-carbonate ion.
41. Write the Lewis structures of the following products of the reaction between sodiumand sulfur.(a) Na2S (b) Na2S2 (c) Na2S3 (d) Na2S8
42. Explain why sulfur readily forms compounds in the �2, �4, and �6 oxidation states,but only a handful of compounds exist in which oxygen is in a positive oxidation state.
43. Explain why sulfur-containing compounds such as FeS2, CS2, and H2S form SO2 in-stead of SO3 when they burn.
44. Use Lewis structures to explain why solutions of the SO32� ion react with sulfur to
form thiosulfate, S2O32�.
45. Use Lewis structures to explain why a two-electron reduction of the S2O62� ion gives
SO32�.
S2O62� � 2 e� 88n 2 SO3
2�
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NONMETAL 49
46. Use Lewis structures to explain why a two-electron oxidation of the S2O32� ion gives
the S4O62� ion.
47. Which of the following does not have a reasonable oxidation number for sulfur?(a) Na2S (b) H2S (c) SO3
2� (d) SO4 (e) SF4
48. Explain why SO2 plays an important role in the phenomenon known as acid rain.49. Explain why problems with acid rain would be much more severe if sulfur compounds
burned to form SO3 instead of SO2.
The Chemistry of Nitrogen and Phosphorus
50. Nitrogen has a reasonable oxidation number in all of the following compounds, andyet one of them is still impossible. Which one is impossible?(a) NF3 (b) NF5 (c) NO3
� (d) NO2� (e) NO
51. Earth’s atmosphere contains roughly 4 1016 tons of nitrogen, and yet the biggest prob-lem facing agriculture in the world today is a lack of “nitrogen.” Explain why.
52. Explain why elemental nitrogen is almost inert, but nitrogen compounds such asNH4NO3, NaN3, nitroglycerin, and trinitrotoluene (TNT) form some of the most dan-gerous explosives.
53. Which of the following oxides of nitrogen are paramagnetic?(a) N2O (b) NO (c) NO2 (d) N2O3 (e) N2O4 (f) N2O5
54. Use Lewis structures to explain what happens in the following reaction.
2 NO � O2 88n 2 NO2
55. Use Lewis structures to explain why NO reacts with NO2 to form N2O3 when a mix-ture of the compounds is cooled.
56. Use the fact that nitrous oxide decomposes to form nitrogen and oxygen to explainwhy a glowing splint bursts into flame when immersed in a container filled with N2O.
2 N2O(g) 88n 2 N2(g) � O2(g)
57. Describe ways of preparing small quantities of each of the following compounds in thelaboratory.(a) N2O (b) NO (c) NO2 (d) N2O4
58. Describe how to tell the difference between a flask filled with NO gas and a flask filledwith NO2.
59. Lightning catalyzes the reaction between nitrogen and oxygen in the atmosphere toform nitrogen oxide, NO.
N2(g) � O2(g) 88n 2 NO(g)
Explain how lightning acts as one source of acid rain.60. Which of the following elements or compounds is not involved at some stage in the
preparation of nitric acid?(a) O2 (b) N2 (c) NO (d) NO2 (e) H2
61. Explain why phosphorus forms both PCl3 and PCl5 but nitrogen forms only NCl3.62. Explain why nitrogen is essentially inert at room temperature, but white phosphorus
bursts spontaneously into flame when it comes into contact with air.63. Explain why red phosphorus is much less reactive than white phosphorus.
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50 NONMETAL
64. Explain why nitrogen forms extraordinarily stable N2 molecules at room temperature,but phosphorus forms P2 molecules only at very high temperatures.
65. Explain why nitric acid has the formula HNO3 and phosphoric acid has the formula H3PO4.66. Write the Lewis structures for phosphoric acid, H3PO4, and phosphorous acid, H3PO3.
Explain why phosphoric acid can lose three H� ions to form a phosphate ion, PO43�,
whereas phosphorous acid can lose only two H� ions to form the HPO32� ion.
67. Explain why only two of the four hydrogen atoms in H4P2O5 are lost when the oxy-acid forms an oxyanion.
68. Describe the role of carbon in the preparation of elemental phosphorus from calciumphosphate.
69. Predict the product of the reaction of phosphorus with excess oxygen and then predictwhat will happen when the product of the reaction is dissolved in water.
70. Explain why the most common oxidation states of antimony are �3 and �5.71. Which of the following compounds should not exist?
(a) Na3P (b) (NH4)3PO4 (c) PO2 (d) PH3 (e) POCl3
The Chemistry of the Halogens
72. Which of the halogens is the most active, or reactive? Explain why.73. Describe the difference between halogens and halides. Give examples of each.74. Fe3� ions can oxidize Br� ions to Br2, but they can’t oxidize Cl� ions to Cl2. Use this
information to determine where the Fe3� ion belongs in the following sequence of de-creasing oxidizing strength: F2 � Cl2 � Br2 � I2.
75. HBr can be prepared by reacting PBr3 with water.
PBr3(l) � 3 H2O(l) 88n 3 HBr(aq) � H3PO3(aq)
Use this information to explain what happens in the following reaction.
P4(s) � 6 Br2(s) � 12 H2O(l) 88n 12 HBr(aq) � 4 H3PO3(aq)
76. Explain why chlorine reacts with fluorine to form ClF3 but not FCl3.77. Chlorine reacts with base to form the hypochlorite ion.
Cl2(aq) � 2 OH�(aq) 88n Cl�(aq) � OCl�(aq) � H2O(l)
Use this information to explain why people who make the mistake of mixing Cloroxwith hydrochloric acid often suffer damage to their lungs from breathing chlorine gas.
The Inorganic Chemistry of Carbon
78. Use the structure of graphite to explain why the bonds between carbon atoms are sostrong that it is difficult to boil off individual carbon atoms, yet the material is so softit can be used as a lubricant.
79. Explain why silicon forms a covalent carbide but calcium forms an ionic carbide.80. Write balanced equations for the combustion of both CO and H2 that explain why a
mixture of the gases can be used as a fuel.81. Use Lewis structures to explain the following reaction.
CO2(g) � H2O(l) 88n H2CO3(aq)
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Integrated Problems
82. A recent catalog listed the following prices: $21.40 for 450 grams of sodium, $18.00 for1 kilogram of zinc, and $52.80 for 250 grams of sodium hydride. Which reagent wouldbe the least expensive source of H2 gas?
83. A solution of hydrogen peroxide in water that is 30% hydrogen peroxide by weightsells for $15.95 per 500 g, and potassium chlorate sells for $12.75 per 500 grams. Is itless expensive to generate oxygen by decomposing H2O2 or KClO3?
84. Write a sequence of reactions for the conversion of elemental nitrogen into nitric acid.Calculate the weight of nitric acid that can be produced from a ton of nitrogen gas.
85. At 1700°C, P4 molecules decompose partially to form P2.
P4(g) 88n 2 P2(g)
If the average molecular weight of phosphorus at that temperature is 91 g/mol, whatfraction of the P4 molecules decompose?
86. Uranium reacts with fluorine to form UF6, which boils at 51°C. The relative rate ofdiffusion of 235UF6 and 238UF6 in the gas phase was used in the Manhattan project toseparate the more abundant 238U isotope from 235U. Predict which substance diffusesmore rapidly, and calculate the ratio of the rate of diffusion of the compounds.
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