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Home , Allotropy, Dioxygen difluoride, Oxygen compounds, Oxygen difluoride, Perxenate, Triplet oxygen

Chapter 18

18-1 Introduction of all the elements except He, Ne, and possibly Ar are known. Molecular oxygen (dioxygen, 02) reacts (at room or on heating) with all other elements except the halogens, a few noble , and the noble . The of oxygen involves the completion of the octet ( config­ uration) by one of the following means: 1. gain to form the 0 2-. 2. Formation of two single covalent bonds, usually in bent AB2E2 systems, such as and . 3. Formation of a double bond, as in ABE2 systems, such as or CI 4Re=0. 4. Formation of a single bond, as well as electron gain, as in ABEs 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

Be40 (CHSC02)6' The wide range of physical properties shown by the binary 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 and alkaline earth metals) from the completely covalent, mo­ lecular oxides, such as CO2, There are, however, intermediate cases such as the oxides of , aluminum, or .

Ionic Oxides

The formation of the oxide from molecular oxygen requires about 1000 kJ mol-I:

! 02(g) = O(g) !1H = 248 kJ mol- l (18-1.1)

2 1 O(g) + 2e- = 0 - !1H = 752 kJ mol- (18-1.2)

In forming an ionic 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, Si02, and oxides of boron, such as B20 g .

Covalent or Molecular Oxides

Covalent or molecular oxides are compounds, such as CO2, S02, SOg, and N02, 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 is an important aspect in the bonding of molecular oxides. This may be pn-pn bonding as in the ketones (R2C=0), or pn-d1t bonding as in phosphine oxides (RgP=O) or linear M=O=M systems.

ABEg Systems Terminal oxygen atoms that bear three lone pairs of are found in alkbxides (RO-), and hydroxide (OH-). Such oxygen atoms may be considered to be spg hybridized.

AB2E2 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, , 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 (Ill0). 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 is 142° and in HgSi-O-SiHg it is greater than 150°. A linear B-A-B situation at oxygen occurs in some AB2E2 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 of1t electrons on oxygen in porbitals 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.

AB3 E Systems The third example containing spg hybridized oxygen atoms is that of the ox­ onium :OH; and :OR;. The formation of oxonium ions is analogous to for­ mation of ammonium ions (NH:). Oxygen is less basic than , 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.

ABE2 Systems Oxygen atoms of this type include those of ketones, , 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 . 18- 1 Introduction 437

The Sp2 hybridization of the carbon 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 are acidic. There are also a number of important amphoteric oxides.

Basic Oxides 2 Although X-ray studies show the existence of discrete oxide ions (0 -) [as well as peroxide (O~-) and (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.

2 0 - + H 20 -----c> 2 OH- (18-1.3) O~- + H 20 -----c> H02" + OH- (18-1.4)

2 02" + H 2 0 -----c> O 2 + H02" + OH- (18-1.5)

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 , 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 anhydrides. An example is given in Reaction 18-1.7, in which N 20 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 Sb20 5 ) , they will generally dissolve in bases (as in Reaction 18-1.8).

(18-1.8)

Acidic oxides will often combine directly, by fusion, with basic oxides to form salts, as in Reaction 18-1.9.

fu~n Na + SI·o 2 ) N a S·OI 3 (18-1.9) 2° 2 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

2 ZnO(s) + 2 H+(aq) ------+ Zn +(aq) + H 20 (18-1.10) ZnO + 2 OH- + H 20 ------+ Zn (OH)~- (18-1.11)

Other Oxides There are other oxides, some of which are relatively inert, which dissolve in neither acids nor bases (e.g., N20, CO, Pb02, and Mn02)' When Mn02 and Pb02 do react with acids (e.g., conc HCl) they do so by a rather than an acid-base reaction, as in Reaction 18-1.12.

(18-1.12)

18·2 Occurrence, Properties, and Oxygen has three isotopes, 160 (99.759%), 170 (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 (~), 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(NH3)sH20]3+, and water. Oxygen has two allotropes; dioxygen (02) and trioxygen or (03), 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 . 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 . Both 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 is perceptibly blue and is diamagnetic. Pure ozone obtained

by fractional liquefaction of02-0 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 Occurence, Properties, and Allotropy 439 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 O2 and 0 3 Ozone is a much more powerful 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 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 [H2Xe06 ] 2-, atomic oxygen, OH radicals, and a few other such species. The following potentials 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 EO = +0.815 V (18-2.3) 0 3 + 2 H+(10-7 M) + 2e- = O 2 + H 20 EO = +1.65 V (18-2.4)

The first step in the reduction of O 2 in aprotic 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 H02

(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 Cr2+ is just stable toward oxidation by pure water, in oxygen-saturated water it is rapidly oxidized. Ferrous ion (Fe2+) is oxidized (slowly in acid, but more rapidly in base) to Fe3+ 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 ions, especially Cu2+, where a CUI_CUll redox cycle is involved. The dioxygen molecule is readily soluble in organic solvents, and merely pouring these 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 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 O2 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 gives rise to three electronic states, as shown in Table 18-1. The eL/) is the ground state, but two excited states are also available at higher energies. These excited singlet states

(especially lLig) 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 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 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 n: n:*b Energy 1:[; l' 1 155 kJ (-13,000 em-I) 1 l 6g i1 92 kJ (-8,000 cm- ) 3:[; l' i o (Ground state) 18-3 Peroxide 441

H 20 2 + Cl2----> 2 CI- + 2 H+ + O 2 (18-2.9) H 20 2 + ClO- ----> Cl- + H 20 + O 2 (18-2.10)

which are accompanied by a red chemiluminescent glow.

18·3

Pure hydrogen peroxide (H20 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 ofeven 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

C2Hs ° C2Hs + H20 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 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 2SOS + H 2S04 (Rapid) (18-3.5)

H 2SOS + H 20 ----> H 20 2 + H 2S04 (Slow) (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 EO = 0.68 V (18-3.8) H02+ H 20 + 2 e- = 3 OH- EO = 0.87 V (18-3.9)

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 2may 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 H02 and OH are most important. The hydroperoxo radical (H02) 3 2 has been detected in aqueous solutions where H 20 2 interacts with Ti +, Fe +, or Ce4+ ions.

18-4 Peroxides and 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. peroxide is made commercially by air oxidation of sodium. 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 . 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 CO2 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 Other Peroxo Compounds 443

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 II and M 0 2·8H20. These contain discrete O~- ions to which water molecules are hydrogen bonded, giving chains of the type shown in Structure 18-111.

-O~- -O~------(H20)s------(H2 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, M02, are formed by the interaction of O 2with K, Rb, or Cs as yellow-to-orange crystalline . Na02 can be obtained by reaction of Na20 2 with O 2 at 300 atm and 500°C. Li02 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 CO2, which involves peroxocarbonate intermediates, is used for removal of CO2 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 , CH3C(0)00H, can be obtained by the action ofH20 2 on acid anhydrides. The peroxo acids are useful oxidants and sources offree rad­ icals, for example by treatment with Fe2+(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 (18-5.1 )

K+02~RO; (18-5.2)

RO;+ RH ~ ROOH + K (18-5.3) 444 Chapter 18 I Oxygen

Explosions can occur on distillation of oxidized solvents. These solvents should

be washed with acidified FeS04 solution or, for ethers and hydrocarbons, passed through a column of activated alumina. Peroxides are absent when the Fe2+ + 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. and ammonium peroxodisulfates (Section 19-5) are commonly used as strong oxidizing agents in 2 3 acid solution, for example to convert C into CO2, Mn + into MnO.!, or Ce + into Ce4+. The last two reactions are slow and normally incomplete in the absence of 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 2Na2C03 '3H20 2 or Na4P207'nH202'

18-6 The Cation

The interaction ofPtF6with O 2gives an orange solid (02PtF6) isomorphous with KPtF6, which contains the paramagnetic 0; ion. This reaction was of importance in that it N. Bartlett to treat PtF6 with (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 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

0-0 distance Number of I Species A n* Electrons vo_o(cm- )

0; 1.12 1 1860

O 2 1.21 2 1556 O·2 1.33 3 1145 O~· 1.49 4 -770 18-8 Oxygen Compounds as ligands 445

/0 0-0 °I \/ M M (a) (b)

~O~ M-O \ M I M O-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 cations, the water ligands are weakly bound (and 446 Chapter 18 / Oxygen H, /0",,­ M M (a) terminal (b) fL2-Bridging H, H, /o~ ~o~° M M M H M ~o.-/' ~o.-/' ~ H "H (c) Bis(fL2)-bridging (d) Tris(fL2)-bridging

H I /o~ / ~ I.---M M~l/ ~M Figure 18-3 The common structural (e) (fL3)-Bridging types of hydroxo ligands.

rapidly substituted), whereas in cases such as [Cr(H20)6]3+ and [Rh(H20)6]3+, they are firmly bound and exchange with solvent water molecules only slowly (Chapter 6). Ligand water molecules can be acidic, especially when bound to cations of high charge, giving hydroxo complexes, as in Reactions 18-8.1 and 18-8.2.

[Pt(NH3)4(H20)2]4+ -----? [Pt(NH3)4(H20) (OH)] 3+ + H+ (18-8.1) [Co(NH3h(H20)]3+ -----? [Co(NH3)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 Ul the octahedral aqua ions (e.g., [M(H20)6]3+, where M = Co 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 CI3P=0, Chapter 17), but also in transition metal compounds, such as vanadyl (V=O), uranyl (U=O), (MnO;), and (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 Oxygen 447 o o II II M M-O-M M II o Bent Linear (a) Terminal (b) Jl-2-Bridging

~O~ ___ --0,,-­ M M M--- _M ~O/ ~O------Symmetrical Unsymmetrical (c) Bis(Jl-2)-bridging

M M / I M--O \\\\\\\I~ \ M'" I ~M M Pyramidal Planar M (d) Jl-g-Bridging (e) Jl-4-Bridging o /O~ I O=M M=O M--O--M I O=M-O-M=O 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 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 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 (OF2) This compound can be prepared by passing fluorine rapidly through a 2% NaOH solution, by 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, CH4 , or CO, although an electrical spark in such mix­ tures will cause a violent explosion. When mixed with C1 2, Br2, or 12, OF2 will ex­ plode at room temperature. It reacts only slowly with water, as in Reaction 18-9.1,but explodes with steam. will liberate other halogens from their acids or salts, as in Reaction 18-9.2.

~ OF2 + H 20 O 2 + 2 HF (18-9.1 ) OF2 + 4 HX(aq) ~ 2 X2 + 2 HF + H 20 (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 (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 ofO 2and F2at 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 . 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 20 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 . Study Guide 449

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 KMn04 in acidic solution; (b) Fe(OH)2 and O 2 in basic solution; (c) sodium peroxide and CO2 ; and (d) 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 (~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) O2 and 0 3 (b) 0i and O~- (c) CH30H and H 20 (d) CO2 and S03

(e) H 20 2 and OH- (f) (CH3)20 and CH3C02 H (g) CH3C(0)00H (h) peroxodisulfuric acid 3. Draw the orbitals as they interact to form the 1t-bond systems in (a) ketones (b) carbonate ion (c) [ClsRu-O-RuCls]4- (d) ozone

(e) triphenylphosphine oxide (f) H 3Si-0-SiH3

(g) OSCl2 4. Calculate the standard redox potential for the air oxidation of Fe2+ in aqueous solu­ tion.

C. Questions from the Literature of 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' 0i, or ot ligands? (b) Explain how magnetic data support or conflict with your answer to (a). (c) What should be the approximate 0-0 distances in the -02 compounds of Crumbliss? 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 20 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 20 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.