Download Chapter 18 OXYGEN

Chapter 18
OXYGEN
18-1
Introduction
Oxygen compounds of all the elements except He, Ne, and possibly Ar are
known. Molecular oxygen (dioxygen, 02) reacts (at room temperature or on
heating) with all other elements except the halogens, a few noble metals, and
the noble gases.
The chemistry of oxygen involves the completion of the octet (neon config­
uration) by one of the following means:
1. Electron gain to form the oxide 0 2-.
2. Formation of two single covalent bonds, usually in bent AB 2E 2 systems,
such as water and ethers.
3. Formation of a double bond, as in ABE 2 systems, such as ketones or
CI 4 Re=0.
4. Formation of a single bond, as well as electron gain, as in ABE s systems,
such as OH- and RO-.
5. Formation of three covalent bonds, usually in pyramidal ABsE systems,
such as HsO+ and RsO+.
6. Formation in rare cases of four covalent bonds, as, for example, in
Be 4 0 (CH SC0 2)6'
The wide range of physical properties shown by the binary oxides of the el­
ements is due to the broad range of bond types from essentially ionic systems to
essentially covalent ones. Thus we distinguish the highly ionic oxides (such as
those of the alkali and alkaline earth metals) from the completely covalent, mo­
lecular oxides, such as CO 2, There are, however, intermediate cases such as the
oxides of boron, aluminum, or silicon.
Ionic Oxides
The formation of the oxide ion from molecular oxygen requires about 1000 kJ
mol-I:
! 02(g)
= O(g)
O(g) + 2e- = 0
2
-
!1H = 248 kJ mol- l
(18-1.1)
!1H = 752 kJ mol- 1
(18-1.2)
In forming an ionic metal oxide, energy must also be expended to vaporize and
to ionize the metal. Thus the stability of ionic metal oxides is a consequence only
of the high lattice energies that are obtained with the small and highly charged
oxide ion.
435
436
Chapter 18
I
Oxygen
Where the lattice energy is not sufficient to offset the energies for ionization,
and so on, oxides with substantial covalent character are formed. Examples of
oxides with some covalent character are BeO, Si0 2, and oxides of boron, such as
B2 0 g .
Covalent or Molecular Oxides
Covalent or molecular oxides are compounds, such as CO 2, S02, SOg, and N0 2,
in which covalent bonding is dominant. Such compounds are well described by
the ABxEy classification, as presented in Chapter 3, with some exceptions, as
noted in the following subsection. Use of the p orbitals in 1t bonding with other
atoms is an important aspect in the bonding of molecular oxides. This may be
pn-pn bonding as in the ketones (R 2C=0), or pn-d1t bonding as in phosphine
oxides (RgP=O) or linear M=O=M systems.
ABE g Systems
Terminal oxygen atoms that bear three lone pairs of electrons are found in
alkbxides (RO-), and hydroxide (OH-). Such oxygen atoms may be considered
to be spg hybridized.
AB 2 E 2 Systems
The compounds that fit into this class are usually angular due to the volume
requirements of two lone pairs of electrons. Examples include water, alcohols,
and ethers. The oxygen atoms are considered to be spS hybridized, but there are
wide variations from the tetrahedral bond angles due to electronic repulsions be­
tween the two lone pairs of electrons: H 20 (104.5°) and (CHg)20 (Ill 0). Where
the atoms bound to oxygen have d orbitals available, some pn-d1t character is
often present in the bond to oxygen, and the B-A-B angles may be even
larger, for example, the angle Si-O-Si in quartz is 142° and in HgSi-O-SiHg
it is greater than 150°.
A linear B-A-B situation at oxygen occurs in some AB 2E 2 systems con­
taining transition metals (e.g., [CIsRu-O-RuCIs ] 4-). The () bonds to Ru are
formed by sp hybrids on oxygen, thus leaving two pairs of 1t electrons on oxygen
in p orbitals that are oriented perpendicular to the Ru-O-Ru axis. These filled
p orbitals on oxygen interact with empty d orbitals on the Ru atoms, forming a
1t-bond system.
AB 3 E Systems
The third example containing spg hybridized oxygen atoms is that of the ox­
onium ions :OH; and :OR;. The formation of oxonium ions is analogous to for­
mation of ammonium ions (NH:). Oxygen is less basic than nitrogen, and the
oxonium ions are therefore less stable. Notice that ions of the type OH~+ are un­
likely (even though :OH; still has a lone electron pair), because of electrostatic
repulsion of the :OH; ion towards another proton. As for :NRg, the pyramidal
:OR; ions undergo rapid inversion.
ABE 2 Systems
Oxygen atoms of this type include those of ketones, aldehydes, and other or­
ganic carbonyls. The oxygen atoms are Sp2 hybridized and have a roughly trigo­
nal arrangement around the oxygen of the lone pairs E and the carbonyl carbon.
18- 1
437
Introduction
The Sp2 hybridization of the carbon atom leaves one p orbital available for for­
mation of a TC bond perpendicular to the trigonal plane.
Acid-Base Properties of Oxides
Generally, the oxides of the metals are basic, whereas those of the nonmetals are
acidic. There are also a number of important amphoteric oxides.
Basic Oxides
Although X-ray studies show the existence of discrete oxide ions (0 2-) [as
well as peroxide (O~-) and superoxide (02") to be discussed later], these ions
cannot exist in aqueous solution owing to the hydrolysis reactions shown in
Reactions 18-1.3 through 18-1.5.
+ H 20
-----c>
2 OH-
(18-1.3)
O~- + H 2 0
-----c>
H02" + OH-
(18-1.4)
2 02" + H 2 0
-----c>
O 2 + H02" + OH-
(18-1.5)
0
2
-
Consequently, only those ionic oxides that are insoluble in water are inert to it.
Ionic oxides function as basic anhydrides. When insoluble in water, they usually
dissolve in dilute acids, as in Reaction 18-1.6.
(18-1.6)
However, some ionic oxides (e.g., MgO) become very slow to dissolve in acids
after high-temperature ignition.
Acidic Oxides
The covalent oxides of the nonmetals are usually acidic, dissolving in water
to produce solutions of acids. They are termed acid anhydrides. An example is
given in Reaction 18-1.7, in which N 2 0 5 is seen to be the acid anhydride of nitric
acid.
(18-1.7)
Even when these oxides are insoluble in water (e.g., as in the case of Sb 2 0
will generally dissolve in bases (as in Reaction 18-1.8).
5) ,
they
(18-1.8)
Acidic oxides will often combine directly, by fusion, with basic oxides to form
salts, as in Reaction 18-1.9.
° + S·o
Na 2
I
2
fu~n
)
N a S·O
I 3
2
(18-1.9)
Amphoteric Oxides
These oxides behave acidicly towards strong bases and as bases towards
strong acids. The example ofZnO is illustrated in Reactions 18-1.10 and 18-1.11.
438
Chapter 18
I
Oxygen
ZnO(s) + 2 H+(aq) -------+ Zn 2+(aq) + H 20
ZnO + 2 OH- + H 20 -------+ Zn (OH)~-
(18-1.10)
(18-1.11)
Other Oxides
There are other oxides, some of which are relatively inert, which dissolve in
neither acids nor bases (e.g., N 20, CO, Pb0 2, and Mn02)' When Mn02 and
Pb0 2 do react with acids (e.g., conc HCl) they do so by a redox rather than an
acid-base reaction, as in Reaction 18-1.12.
(18-1.12)
18·2
Occurrence, Properties, and Allotropy
Oxygen has three isotopes, 160 (99.759%), 17 0 (0.0374%), and 180 (0.2039%).
Fractional distillation of water allows concentrates containing up to 97 atom %
180 or up to 4 atom % 170 to be prepared. Oxygen-18 is used as a tracer in study­
ing reaction mechanisms of oxygen compounds. Although 170 has a nuclear
spin (~), its low abundance means that even when enriched samples are used
spectrum accumulation and/or the Fourier transform method are required. An
example of 170 resonance studies is the distinction between H 20 in a complex,
for example, [Co(NH 3)sH 20]3+, and solvent water.
Oxygen has two allotropes; dioxygen (0 2) and trioxygen or ozone (0 3 ),
Dioxygen is paramagnetic in all phases and has the rather high dissociation en­
ergy of 496 kJ mol- 1. Simple valence bond theory predicts the electronic struc­
ture :0=0: which, though accounting for the strong bond, fails to account for
the paramagnetism. However, simple MO theory (Section 3-5) readily accounts
for the triplet ground state having a double bond. There are several low-lying sin­
glet states that are important in photochemical oxidations. Like NO, which has
one unpaired electron in an antibonding (11:*) MO, oxygen molecules associate
only weakly, and true electron pairing to form a symmetrical 0 4 species does not
occur even in the solid. Both liquid and solid O 2 are pale blue.
Ozone
The action of a silent electric discharge on O 2 produces 0 3 in concentrations up
to 10%. Ozone gas is perceptibly blue and is diamagnetic. Pure ozone obtained
by fractional liquefaction of 02- 0 3 mixtures gives a deep blue, explosive liquid.
The action of UV light on O 2 produces traces of 0 3 in the upper atmosphere.
The maximum concentration is at an altitude of about 25 km. It is of vital im­
portance in protecting the earth's surface from excessive exposure to UV light.
Ozone decomposes exothermically, as in Reaction 18-2.1:
I1H = -142 kJ mol- 1
(18-2.1)
but it decomposes only slowly at 250°C in the absence of catalysts and UV light.
The 0 3 molecule is symmetrical and bent; LO-O-O, 117°; 0-0,
1.28 A. Since the 0 - 0 bond distances are 1.49 Ain HOOH (single bond) and
1.21 A in O 2 (- double bond), it is apparent that the 0 - 0 bonds in 0 3 must
18-2
439
Occurence, Properties, and Allotropy
have considerable double-bond character. In terms of a resonance description,
this can be accounted for as in the resonance forms of Structures 18-1 and 18-11.
IS-I
IS-II
Chemical Properties of O 2 and 0 3
Ozone is a much more powerful oxidizing agent than O 2 and reacts with many
substances under conditions where O 2 will not. The reaction
(18-2.2)
is quantitative and can be used for analysis. Ozone is used for oxidations of or­
ganic compounds and in water purification. Oxidation mechanisms probably in­
volve free radical chain processes as well as intermediates with -OOH groups.
In acid solution, 0 3 is exceeded in oxidizing power only by F2, the perxenate ion
[H 2Xe0 6 ] 2-, atomic oxygen, OH radicals, and a few other such species.
The following poten tials indicate the oxidizing strengths of O 2 and 0 3 in or­
dinary aqueous solution.
O 2 + 4 H+(10-7 M) + 4e-
= 2 H 20
0 3 + 2 H+(10-7 M) + 2e- = O 2 + H 20
EO
= +0.815 V
(18-2.3)
EO
=
+1.65 V
(18-2.4)
The first step in the reduction of O 2 in aprotic solvents such as DMSO and pyri­
dine appears to be a one-electron step to give the superoxide anion:
(18-2.5)
whereas in aqueous solution a two-electron step occurs to give H0 2
(18-2.6)
It can also be seen from the potential given for Reaction 18-2.3 that neutral
water saturated with O 2 is a fairly good oxidizing agent. For example, although
Cr 2+ is just stable toward oxidation by pure water, in oxygen-saturated water it is
rapidly oxidized. Ferrous ion (Fe 2+) is oxidized (slowly in acid, but more rapidly
in base) to Fe 3+ in the presence of air, although in oxygen-free water it is quite
stable, as shown by the potential for Reaction 18-2.7.
(18-2.7)
Many oxidations by oxygen in acid solution are slow, but the rates of oxidation
may be vastly increased by catalytic amounts of transition metal ions, especially
Cu 2+, where a CUI_CUll redox cycle is involved.
The dioxygen molecule is readily soluble in organic solvents, and merely
pouring these liquids in air serves to saturate them with 02' This fact should be
kept in mind when determining the reactivity of air-sensitive materials in solu­
tion in organic solvents.
440
Chapter 18
/
Oxygen
Measurements of electronic spectra of alcohols, ethers, benzene, and even
saturated hydrocarbons show that there is reaction of the charge-transfer type
with the oxygen molecule. However, there is no true complex formation, since
the heats offormation are negligible and the spectral changes are due to contact
between the molecules at van der Waals distances. The classic example is that of
N,Ndimethylaniline, which becomes yellow in air or oxygen but colorless again
when the oxygen is removed. Such weak charge-transfer complexes make certain
electronic transitions in molecules more intense; they are also a plausible first
stage in photooxidations.
With certain transition metal complexes, O 2 adducts may be formed, some­
times reversibly (Section 18-7). Although the O 2 entity remains intact, the com­
plexes may be described as having coordinated O 2 or O~- ions, bound to the
metal in a three-membered ring or as a bridging group. Coordinated O 2 is more
reactive than free 02' and substances not directly oxidized under mild condi­
tions can be attacked in the presence of metal complexes.
The Excited State Chemistry of Oxygen
As discussed in Chapter 3, the oxygen molecule contains two unpaired electrons
in n* molecular orbitals. This electron configuration gives rise to three electronic
states, as shown in Table 18-1. The triplet state eL/) is the ground state, but two
excited states are also available at higher energies. These excited singlet states
(especially lLi g ) have sufficiently long lifetimes to allow them to be useful for re­
actions with a variety of substrates, where they cause specific oxidations, a very
typical example being 1,4 addition to a 1,3-diene, as in Reaction 18-2.8.
~ + °2(Singlet) ~
n
(18-2.8)
0-0
Singlet oxygen molecules may be generated either by photochemical or
chemical means. The photochemical route typically employs a sensitizer, which
first absorbs energy from the light source and then transfers an appropriate
amount of that energy to triplet oxygen to give an oxygen molecule in an excited
(singlet) state. The sensitizer molecule or ion must be in an excited triplet state
for this energy transfer to be spin allowed.
The chemical generation of singlet oxygen may be accomplished as in
Reactions 18-2.9 and 18-2.10:
Table 18-1 The Three Electronic States Arising
from the (n:*) 2 Electron Configuration of
Molecular Oxygen
State
1:[;
1
6g
3:[;
n:
n:*b
Energy
l'
i1
l'
1
155 kJ (-13,000 em-I)
92 kJ (-8,000 cm- l )
o (Ground state)
i
18-3
441
Hydrogen Peroxide
H 20 2 + Cl 2 ----> 2 CI- + 2 H+ + O 2
H 20 2 + ClO- ----> Cl- + H 20 + O 2
(18-2.9)
(18-2.10)
which are accompanied by a red chemiluminescent glow.
18·3
Hydrogen Peroxide
Pure hydrogen peroxide (H 20 2) is a colorless liquid (bp 152.1 cC, fp - 0.41 cc).
It resembles water in many of its physical properties and is even more highly as­
sociated via hydrogen bonding and 40% denser than is H 20. It has a high di­
electric constant, but its utility as an ionizing solvent is limited by its strong oxi­
dizing nature and its ready decomposition in the presence of even traces of many
heavy-metal ions according to the reaction:
!',.H = -99 kJ mol- 1
(18-3.1)
In dilute aqueous solution it is more acidic than water.
(18-3.2)
The molecule H 20 2 has a skew, chain structure (Fig. 18-1).
There are two methods for large-scale production of H 20 2. One is by autox­
idation of an anthraquinol, such as 2-ethylanthraquinol.
OH
°
C 2H s
C 2H s + H 20
2
°2
H 2 /Pd
(18-3.3)
°
OH
The resulting quinone is reduced with H 2 gas. The H 20 2 is obtained as a 20%
aqueous solution. Only 02' H 2 and H 20 are required as raw materials.
An older and more expensive method is electrolytic oxidation of sulfuric
acid or ammonium sulfate-sulfuric acid solutions to give peroxodisulfuric acid,
which is then hydrolyzed to yield H 20 2:
(18-3.4)
H 2S20
S
+ H 20
---->
H 2SO S + H 2S0 4
H 2SO S + H 20 ----> H 20 2 + H 2S0 4
(Rapid)
(Slow)
(18-3.5)
(18-3.6)
Fractional distillation can then give 90-98% H 20 2.
The redox chemistry of H 20 2 in aqueous solution is summarized by the po­
tentials.
(18-3.7)
442
Chapter 18
I
Oxygen
Figure 18-1 The structure of
hydrogen peroxide.
O 2 + 2 H+ + 2 e- = H 20 2
H0 2 + H 20 + 2 e-
= 3 OH-
EO = 0.68 V
(18-3.8)
= 0.87 V
(18-3.9)
EO
These show that H 20 2 is a strong oxidizing agent in either acid or basic solution.
It behaves as a reducing agent only toward very strong oxidizing agents such as
MnO';.
Dilute or 30% H 20 2 solutions are widely used as oxidants. In acid solution,
oxidations with H 20 2 are slow, whereas in basic solution, they are usually fast.
Decomposition to H 20 and 02' which may be considered a self-oxidation, or
disproportionation, occurs most rapidly in basic solution; hence an excess of
H 20 2 may best be destroyed by heating in basic solution.
Many reactions involving H 20 2 (and also 02) in solutions involve free radi­
cals. Metal-ion catalyzed decomposition of H 20 2 and other reactions form radi­
cals of which H0 2 and OH are most important. The hydroperoxo radical (H0 2)
has been detected in aqueous solutions where H 20 2 interacts with Ti 3 +, Fe 2+, or
Ce 4 + ions.
18-4
Peroxides and Superoxides
These substances are derived formally from O~- (peroxides) and 0; (superox­
ides).
Ionic Peroxides
Ionic peroxides are formed by alkali metals, Ca, Sr, and Ba. Sodium peroxide is
made commercially by air oxidation of sodium. Sodium peroxide is a yellow pow­
der that is very hygroscopic, though thermally stable to 500°C. It contains, ac­
cording to electron spin resonance (ESR) studies, about 10% of the superoxide.
The ionic peroxides give H 20 2 on reaction with H 20 or dilute acids. All of
the ionic peroxides are powerful oxidizing agents, converting organic materials
to carbonate even at moderate temperatures. Sodium peroxide will also vigor­
ously oxidize some metals (e.g., Fe, which violently gives FeO~~). The peroxides
of the alkali metals also react with CO 2 according to Reaction 18-4.1 to give car­
bonates:
(18-4.1 )
Other electropositive metals such as Mg and the lanthanides also yield per­
18-5
443
Other Peroxo Compounds
oxides; these are intermediate in character between the ionic ones and the es­
sentially covalent peroxides of metals such as Zn, Cd, and Hg.
Many ionic peroxides form well-crystallized hydrates such as Na202'8H20
and M II 0 2·8H 20. These contain discrete O~- ions to which water molecules are
hydrogen bonded, giving chains of the type shown in Structure 18-111.
--- -O~- - - -(H 2 0)s- - - -O~- - - -(H 2 0)s- --­
IS-III
The formation of such stable hydrates accounts for the extreme hygroscopic na­
ture of the crystalline peroxides.
Ionic Superoxides
Ionic superoxides, M0 2, are formed by the interaction of O 2 with K, Rb, or Cs as
yellow-to-orange crystalline solids. Na0 2 can be obtained by reaction of Na 20 2
with O 2 at 300 atm and 500°C. Li0 2 cannot be isolated. Alkaline earth, Zn, and
Cd superoxides occur only in small concentrations as solid solutions in the per­
oxides. The O 2 ion has one unpaired electron. Superoxides are very powerful
oxidizing agents. They react vigorously with water.
(18-4.2)
(Slow)
(18-4.3)
The reaction with CO 2, which involves peroxocarbonate intermediates, is
used for removal of CO 2 and regeneration of O 2 in closed systems (e.g., sub­
marines). The overall reaction is
(18-4.4)
18·5
Other Peroxo Compounds
There are many organic peroxides and hydroperoxides. Peroxo carboxylic acids, for
example peracetic acid, CH 3 C(0)00H, can be obtained by the action ofH 20 2
on acid anhydrides. The peroxo acids are useful oxidants and sources offree rad­
icals, for example by treatment with Fe 2+(aq). Benzoyl peroxide and cumyl hy­
droperoxide are moderately stable and widely used where free radical initiation
is required, as in polymerization reactions.
Organic peroxo compounds are also obtained by autoxidation of ethers,
alkenes, and the like, on exposure to air. Autoxidation is a free radical chain re­
action initiated by radicals generated by interaction of oxygen and traces of met­
als such as Cu, Co, or Fe. The attack on specific reactive C-H bonds by a radi­
cal (XO) , first gives RO and then hydroperoxides that can react further.
RH+X"~K+HX
K+02~RO;
RO;+ RH
~
ROOH + K
(18-5.1 )
(18-5.2)
(18-5.3)
444
Chapter 18
I
Oxygen
Explosions can occur on distillation of oxidized solvents. These solvents should
be washed with acidified FeS0 4 solution or, for ethers and hydrocarbons, passed
through a column of activated alumina. Peroxides are absent when the Fe 2+ +
SCN- reagent does not give a red color indicative of the Fe(SCN)2+ ion.
There are also many inorganic peroxo compounds where - 0 - is
replaced by - 0 - 0 - groups, such as peroxodisulfuric acid,
(HO)2S(0)00S(0) (OHh mentioned previously. Potassium and ammonium
peroxodisulfates (Section 19-5) are commonly used as strong oxidizing agents in
acid solution, for example to convert C into CO 2, Mn 2+ into MnO.!, or Ce 3 + into
Ce 4 +. The last two reactions are slow and normally incomplete in the absence of
silver ion as a catalyst.
It is important to make the distinction between true peroxo compounds,
which contain - 0 - 0 - groups, and compounds that contain H 20 2 of crys­
tallization, such as 2Na 2C0 3 '3H 20 2 or Na4P207'nH202'
18-6
The Dioxygenyl Cation
The interaction ofPtF6 with O 2 gives an orange solid (02PtF6) isomorphous with
KPtF6, which contains the paramagnetic 0; ion. This reaction was of importance
in that it lead N. Bartlett to treat PtF6 with xenon (Section 21-2). A number of
other salts of the 0; ion are known.
It is instructive to compare the various O;± species, since they provide an in­
teresting illustration of the effect of varying the number of antibonding elec­
trons on the length and stretching frequency of a bond, as shown by the data in
Table 18-2.
18·7
Dioxygen as a Ligand
Although the most common mode of reaction of molecular oxygen with transi­
tion metal complexes is oxidation (i.e., extraction of electrons from the metal or
from its ligands), under appropriate circumstances the dioxygen molecule may,
instead, become a ligand. Such reactions are termed oxygenations, because the
dioxygen ligand retains its identity, whereas oxidation reactions are those in
which the O 2 molecule loses its identity through reduction.
Oxygenation reactions are often reversible. That is, upon increasing tem­
perature and/ or reducing the partial pressure of 02' the dioxygen ligand is lost
by dissociation or by transfer to another acceptor (which may become oxidized).
The process of reversible oxygenation plays an essential role in life processes. In
humans or other higher animals, oxygen molecules are "carried" from the lungs
Table 18-2
Bond Values for Oxygen Species
Species
0-0 distance
A
Number of
n* Electrons
0;
O2
O·2
1.12
1.21
1.33
1.49
1
2
3
4
O~·
vo_o(cm- I )
1860
1556
1145
-770
445
18-8 Oxygen Compounds as ligands
/0
°MI
0-0
\
\
O-M
/
(b)
(a)
M-O
M
~O~
M
I M
'------0/
(e)
(d)
Figure 18-2 The five structural types of
dioxygen ligands.
to the various tissues by hemoglobin and myoglobin molecules, in which 1:1
02-Fe complexes are formed. In lower animals, there are molecules such as
hemerythrins and hemocyanins, that serve similar functions. More detail con­
cerning these biological complexes will be given in Chapter 31.
Broadly speaking, there are two types of 1:1 02-M complexes, the "end-on"
and the "slide-on" types, as shown in Fig. 18-2, types (a) and (b). In addition,
there are many 1:2 02-M complexes, as shown in Fig. 18-2, types (c) and (d).
The hemoglobin and myoglobin complexes are of type (a), and there are a num­
ber of synthetic examples in which O 2 fiIls one position in an octahedral com­
plex. Most of these can be considered to contain a coordinated superoxide ion
0;, and thus have an unpaired electron formaIly present on the coordinated
dioxygen unit. Many of these complexes form reversibly.
The "side-on" complexes, type (b) in Fig. 18-2, are also numerous. Many are
formed reversibly, as with Vaska's compound in Reaction 18-7.1.
(18-7.1)
These compounds are generally best regarded as peroxide complexes, that is,
compounds containing the O~- ligand. The complexes in Fig. 18-2, types (c) and
(d), are also best regarded as peroxide complexes.
18·8 Oxygen Compounds as Ligands
Water: Aqua Ligands
Hydration of transition metals has already been discussed in Chapter 6, as have
the rates and thermodynamics of water ligand exchange in solution. In some
cases, such as the alkali metal cations, the water ligands are weakly bound (and
446
Chapter 18
/
Oxygen
H,
/0",,­
M
(a) terminal
M
(b) fL2-Bridging
H,
H,
/o~
M
°
~o~
M
M
H
M
~o.-/'
~o.-/'
~
"
H
H
(c) Bis(fL2)-bridging
(d) Tris(fL2)-bridging
H
I
/o~
/ ~ I.---M
M~l/
~M
(e) (fL3)-Bridging
Figure 18-3 The common structural
types of hydroxo ligands.
rapidly substituted), whereas in cases such
they are firmly bound and exchange with
(Chapter 6).
Ligand water molecules can be acidic,
high charge, giving hydroxo complexes, as
as [Cr(H 20)6]3+ and [Rh(H 20)6]3+,
solvent water molecules only slowly
especially when bound to cations of
in Reactions 18-8.1 and 18-8.2.
[Pt(NH 3)4(H 20)2]4+ -----? [Pt(NH 3)4(H 20) (OH)] 3+ + H+
(18-8.1)
[Co(NH 3h(H 20)]3+ -----? [Co(NH 3)5(OH)]2+ + H+
(18-8.2)
Hydroxide: Hydroxo Ligands
Many important hydroxo complexes are known, the hydroxo ligand serving in
some cases as a simple terminal ligand, and in other cases as a bridging ligand,
examples of which are shown in Fig. 18-3. Double (f.l2) bridges are most com­
mon. For complexes containing only terminal hydroxo ligands, there has been
particular interest in the structural changes that are apparent when comparing
the octahedral aqua ions (e.g., [M(H 20)6]3+, where M = Co Ul or AJIII) with the
corresponding hydroxo complexes [M(OH)4r, which are tetrahedral.
Oxide: Oxo Ligands
Oxo compounds can be of several structural types, as shown in Fig. 18-4. The
multiply bonded oxo group (M=O) is found not only in oxo compounds and
oxo anions of the nontransition elements (e.g., SO~-, Chapter 5, and CI 3P=0,
Chapter 17), but also in transition metal compounds, such as vanadyl (V=O),
uranyl (U=O), permanganate (MnO;), and osmium tetroxide (OS04)' In cases
involving metals, the bond distance to oxygen (1.59-1.66 A) corresponds to a
double bond, which is best formulated as arising from Oprr ~ Mdn donation.
18-9
447
Oxygen Fluorides
o
II
o
II
M-O-M
M
II
M
o
Bent
(a) Terminal
~O~
M
~O/
M
Linear
(b) Jl-2-Bridging
M---
___ --0,,-­
_M
~O-------
Symmetrical
Unsymmetrical
(c) Bis(Jl-2)-bridging
M
M
/
I
M--O
\ M
Pyramidal
M'"
~M
(e) Jl-4-Bridging
/O~
I
o
M--O--M
M=O
O=M-O-M=O
I
M
Planar
(d) Jl-g-Bridging
O=M
\\\\\\\I~
I
o
(j) Mixed-bonding modes
Figure 18-4
The common structural types of oxo ligands.
Thus the metal oxo complexes are most stable when the metal is in a high oxi­
dation state. In contrast, for the oxides of the nonmetals (e.g., CO and S02), low
oxidation states of the nonmetal are preferred.
The M=O bond is commonly affected by the nature of the group trans to
oxygen. Donors that increase electron density on the metal tend to reduce the
metal's acceptor ability, thus lowering the M=O 1t-bond character. Conse­
quently, the MO stretching frequency in such complexes is found to be lower
than when the oxo ligand is trans to a weak donor ligand.
8·9
Oxygen Fluorides
Most oxygen compounds are properly called oxides and, therefore, are discussed
under the chemistry of the other elements. However, since fluorine is more elec­
tronegative than oxygen, it is logical to treat oxygen fluorides in this chapter.
While these compounds are sometimes called fluorine oxides, it is best to call
them oxygen fluorides. These compounds have been intensively studied as
rocket fuel oxidizers.
448
Chapter 18
I
Oxygen
Oxygen Difluoride (OF 2)
This compound can be prepared by passing fluorine rapidly through a 2%
NaOH solution, by electrolysis of aqueous HF-KF solutions, or by reaction of
fluorine with moist KF. It is a pale yellow, poisonous gas (bp 145°C), which is rel­
atively unreactive as far as this class of compounds is concerned. It can be mixed
without reaction with H 2, CH 4 , or CO, although an electrical spark in such mix­
tures will cause a violent explosion. When mixed with C1 2, Br2, or 12, OF 2 will ex­
plode at room temperature. It reacts only slowly with water, as in Reaction
18-9.1,but explodes with steam. Oxygen difluoride will liberate other halogens
from their acids or salts, as in Reaction 18-9.2.
OF 2 + H 2 0
OF 2 + 4 HX(aq)
~
~
O 2 + 2 HF
2 X 2 + 2 HF + H 20
(18-9.1 )
(18-9.2)
Oxygen difluoride will oxidize most metals and nonmetals, and even reacts with
Xe in an electric discharge to give xenon fluorides and xenon oxide fluoride
(Chapter 21).
Dioxygen Difluoride: (02F2)
This compound is a yellow-orange solid (mp 109.7 K) that is made by high-volt­
age electric discharge on mixtures of O 2 and F2 at low temperature and pressure.
It decomposes into the elements in the gas at -50 °C, and is a potent fluorinat­
ing and oxidizing agent. Many substances explode on exposure to 02F2, even at
low pressures.
The structure of 02F2 is bent, one fluorine atom being about 87° out of the
plane of the other three atoms (Structure 18-IV). The 0 - 0 bond is quite short
(1.217 A) compared to the value for H 2 0 2 (1.48 A).
IS-IV
STUDY GUIDE
Study Questions
A. Review
1. Give the electron configuration of the oxygen atom.
2. Give two examples of oxonium ions. What is their structure?
3. Describe the carbon-oxygen bond in acetone.
449
Study Guide
4. Describe the interaction with water of acidic, basic, and neutral oxides. Give two ex­
amples of each case.
5. Explain why the oxygen molecule is paramagnetic.
6. Write out the electron configurations of the two excited state singlets found in Table
18-I.
7. Describe the preparation in the laboratory of ozone.
8. How is H 20 2 made?
9. Write balanced equations for the following reactions: (a) H 20 2 and KMn0 4 in acidic
solution; (b) Fe(OH)2 and O 2 in basic solution; (c) sodium peroxide and CO 2 ; and
(d) potassium superoxide and water.
10. What is the difference between oxygenation and oxidation?
B. Additional Exercises
1. Prepare MO energy-level diagrams for all of the ions O~+ that are chemically impor­
tant, and determine the bond order and the expected magnetic moment (~eff in
Bohr magnetons, as discussed in Chapter 2).
2. Classify the oxygen atoms in the following systems according to the ABxEy scheme of
Chapter 3, and, where appropriate, discuss the geometry about oxygen in terms of
the VSEPR theory:
(a) O 2 and 0 3
(b) 0i and O~(c) CH 30H and H 20
(d) CO 2 and S03
(e) H 20 2 and OH(f) (CH3)20 and CH 3C0 2 H
(h) peroxodisulfuric acid
(g) CH 3C(0)00H
3. Draw the orbitals as they interact to form the 1t-bond systems in
(b) carbonate ion
(a) ketones
(c) [Cl sRu-O-RuCl s ]4(d) ozone
(e) triphenylphosphine oxide
(f) H 3 Si-0-SiH 3
(g) OSCl 2
4. Calculate the standard redox potential for the air oxidation of Fe 2+ in aqueous solu­
tion.
C. Questions from the Literature of Inorganic Chemistry
1. Compare the structures and properties of two very different "reversible oxygen com­
plexes" as reported by S. ]. La Placa and]. A. Ibers, j. Am. Chem. Soc., 1965, 87,
2581-2586, and as reported by A. L. Crumbliss and F. Basolo,j. Am. Chem. Soc., 1970,
92,55-60. See also L. Vaska, Science, 1963, 140,809.
(a) Should the oxygen ligands in these complexes be considered to be 02'
ligands?
ot
0i, or
(b) Explain how magnetic data support or conflict with your answer to (a).
(c) What should be the approximate 0-0 distances in the cobalt-0 2 compounds
of Crum bliss?
2. Consider the work by M. M. Morrison,]. L. Roberts,jr., and D. T. Sawyer, lnorg. Chem.,
1979,18,1971-1973.
(a) What reaction takes place between OH- and H 2 0
2
in pyridine solution?
(b) What is formed upon electrochemical reduction of H 2 0 2 in pyridine solution?
(c) Mter electrochemical reduction of H 2 0 2 in pyridine solution, what reaction
takes place between HOi and H 20 2 ?
(d) How are the reactions for (c) and (a) related?
450
Chapter 18
/
Oxygen
(e) What role does solvent play in these reactions? What is different about these
redox reactions in water and in pyridine?
SUPPLEMENTARY READING
Bailey, P. S., Ozonation in Organic Chemistry, Academic, ew York, Vol. 1, 1978, Vol. 2,
1982.
Dotto, L. and Schiff, H., The Ozone War, Doubleday, New York, 1978.
Golodets, G. I., Heterogeneous Catalytic Reactions Involving Oxygen, Elsevier, Amsterdam,
1983.
Greenwood, G. and Hill, H. O. A., "Oxygen and Life," Chem. Br:, 1982, 194.
Hayaishi, 0., Molecular Oxygen in Biology, North-Holland, Amsterdam, 1974.
Hoare, P.]., The Electrochemistry of Oxygen, Wiley, New York, 1968.
Horvath, M., Bilitzky, L., and Huttner,]., Ozone, Elsevier, Amsterdam, 1985.
Martell, A. E. and Sawyer, D. T., Eds., Oxygen Complexes and Oxygen Activation, Plenum,
New York, 1988.
Murphy,]. S. and Orr,]. R., Ozone Chemistry and Technology, Franklin Institute Press,
Philadelphia, 1975.
Oberley, L. W., Ed., Superoxide Dismutase, Vol. 3. CRC Press, Boca Raton, FL, 1985.
Patai, S., Ed., The Chemistry of the Hydroxyl Group, Wiley-Interscience, New York, 1971.
Schaap, A. P., Ed., Singlet Molecular Oxygen, Wiley, New York, 1976.
Severn, D., Organic Peroxides, Vols. I-III, Wiley-Interscience, New York, 1972.
Spiro, T. G., Metal Ion Activation of Oxygen, Wiley, ew York, 1983.
Toft-Sorensen, 0., Ed., Nonstoichiometric Oxides, Academic, ew York, 1981.
Valentine,]. S., "The Dioxygen Ligand in Mononuclear Group VIII Transition Metal
Complexes," Chem. Rev., 1973, 73,235.
Vaska, L., "Dioxygen Metal Complexes," Acc. C1um. Res.,1976, 9, 175.