R. 1. Sanderson Arizona State University Tempe
Principles of Halogen Chemistry
The chemistry of the halogens covers a range far too wide to permit easy acquisition of a comprehensive yet well-organized picture. In an effort to pull together loose ends, to recognize unifying principles, and to help to compress the chief substance of this vast subject to something approaching the capacity of the average human brain, this paper has been written. It followsthe general pattern of preceding articles in this series on chemical bonding ( I ) , chemical reaction (I?), and hydrogen chemistry (3). The halogens, namely JEuorine, chlorine, bromine, and iodine,' are characterized by their similar electronic c a JiSurations which feature seven electrons in the outermost principal quantum level, and consequently the highest electronegatiuify in the respective period of each. The high electronegativity arises from the fact that the mutual shielding by electrons filling the same principal quantum level is relatively small compared to the increased nuclear charge. This is in agreement with the relatively high compactness of the electronic clouds of atoms of these elements. This outer shell similarity results in a reasonably close similaritu- o.f the halooens in their neoative oxidation state of -1, because each alom, having m e outermost half-filled orbital and a relatively high electronegativity, can form a single covalent b a d in which it usually becomes the negatively charged partner. The general similarity of all the alkali halides well illustrates this point. Significant dissimilarities among the halogens also occur, reflecting primarily the differences both in underlying and overlying alomic structures. The most important differences in underlying structure are in the penultimate shell, just below the outermost. In fluorine this contains only two electrons, in chlorine 8, in bromine and iodine, 18. In the overlying structures, stable d orbitals are unavailable in fluorine but available, and presumably increasingly so, in chlorine, hromine, and iodine. Since such orbitals appear to become stable mainly or only through partial removal of valence electrons and nothing can remove electrons from fluorine, the lack of d orbitals in fluorine has no practical importance except as an auxiliary explanation of the absence of positive oxidation states. But since both fluorine and oxygen are more electronegative than chlorine, bromine, and iodine, and chlorine more also than bromine and iodine, and bromine more than iodine (Table I), there are numerous ways in which outer d orbitals can be made more available in these other halogens through oxidation. In positive oxidation Little-known astatine is omitted from this paper because of its nuclear instability, scarcity, and present relative unimportance.
states, the halogens can be expected to diier according to whether the penultimate shell is 8 (chlorine) or 18 (bromine and iodine). The physical state of the halogens i s determined by their existence as diatomic molecules and by the slrength of van der Waals forces among these molecules. The diatomic molecules are the natural result of the limited bonding capacity of the halogen atoms, each of which can form hut one single covalent bond. The van der Waals attractions among these diatomic molecules, being in general proportional to the population of the electron clouds and also inversely proportional to the tightness with which these electrons are held, are weakest in fluorine and increase as the halogens become larger. Table 1 gives the polarizabilities of the halogens and some of the physical properties which reflect the strength of the intermolecular forces. It will he ohserved that these properties all vary in the predictable direction. Table 1 .
Some Properties of the Halogens
F, Atomic number 9 Electrons/molecule 18 Electronic configuration 2-7 Polaririrabilitv X 10".
First ionization of atom ca /mole) 402 Electronegativity 3.92 Melting point ( T ) -219.6 Heat of fusion (kcal/mole) 0.061 Boiline ~ o i n ('C) t -187.9
F"
Cix
Br*
12
17 34 2-8-7
300 3.28 -102.4 0.77 -34.0
The chemical reactivily of the halogens, to the extent thal it depends on dissociation of the diatomic molecules, i s generally high since the dismciation energies are relatively low, especially that offluorine. The halogens as a group are quite reactive under relatively mild conditions. Fluorine is especially reactive, which may in part he a result of its low dissociation energy (see Table 1). This in turn has been explained as resulting from interference of nonbondig electron pairs on the adjacent atoms when they are bonded together. Such an interpretation is consistent with the well-known tendency of carbon, nitrogen, and oxygen to unite by multiple bonds, showing that when a single bond is Volume 41, Number 7, July 1964
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as essentially donor-acceptor interaction, with all outermost electrons acting as valence electrons and all partially positive atoms furnishing all available vacant outer orbitals as acceptors (4). The coordination number of the halogen depends chiefly on its partial charge but also on its original electronegativity. The donating power is greater the higher the negative charge but lower, the more electronegative the donor atom was initially. For this reason, for a given amount of charge, halogen coordination number tends to increase in the order F-C1-Br-I. Properties such as melting point, heat of fusion, boiling point, and heat of vaporization reflect the magnitude of the forces of aggregation, and thus tend to be higher for more polar compounds in closely coordinated lattices, and also higher in the sequence, iodides, bromides, chlorides, and fluorides. Some illustrative examples are given in Table 3. The only known exceptions to this principle are the boron halides, which presumably have one vacant orbital if the three halogen atoms are attached to the boron by single covalent bonds, and yet are molecular, being monomeric in the gas phase. In these it is possible that the otherwise vacant orbital on boron becomes involved in partial double bond formation, using otherwise unshared electron pairs on the three halogen atoms. In part a t least, this hypothesis is supported by the shorter than expected bond length in BF3 and by the order of Lewis acid strength increasing from BFI to BBr8 (6),in keeping with the diminishing ability of the larger halogen to become multiple bonded to the boron. The aluminum halides appear to use the fourth aluminum orbital on the monomer to form dimeric molecules by halogen bridging (except for aluminum fluoride which condenses more highly), and all other halides condense even further if orbital vacancies would otherwise be left free. Halides in which the halogen has acquired relatively large negative charge tend to be high melting and low in volatility. This is very similar to the preceding principle because the kind of element from which halogen can acquire relatively high negative charge is that which is low in electronegativity and thus also a metal having more outermost vacancies than valence electrons, per atom. Binary halides that contain highly polar bonds therefore invariably would have also vacant low energy orbitals in the sin~plestpossible molecules. Condensation to make fuller use of these orbitals therefore occurs.
formed between period 2 atoms, additional orbitals are brought close enough for interaction. Bond energies of halides, per halogen atom, consistently show that the bond strength decreases in the order, fluoride >>chloride > bromide > iodide. Some typical examples are given in Table 2. Table 2.
Bond Energies of Some Typical Binary Halides (Heats of atomizotion per equivalent)
Fluoride
Chloride
Bromide
Iodide
Bond energies of binary halides are in general directly proportional to the partial charge on halogen and the bond order, and inversely proportional to the bond length. An empirical equation representing this relationship is the following: 6x = A nH/R
+B
where ax is the partial charge on halogen, n is the bond order, H is the heat of atomization per equivalent, and A and B are constants characteristic of the halogen. This relationship holds for practically all solid binary halides and approximately for most liquid or gaseous molecular halides as well (4). Whenever the simplest possible binary halide molecule would possess unoccupied low ensrgy orbitals, in addition to the unshared electron pairs always present on the halogen atom, the molecules condense to a more highly associated state under ordinary conditions, and possess the physieal properties predictable for such a state. This condensation is to what is called the ionic state when the bond polarity is high and to what might be called coordination polymels when the polarity is less. These are probably different degrees of the same phenomenon, which is a general tendency for molecules to make maximum possible use of all available outer orbital vacancies (1). The coordimation number of the halogen atom in these condensed states is a t least partly dependent on the amount of partial negative charge it bears. Even in the most ionic crystals, bonding may usefully be regarded Table 3.
NaF NaCl NsBr Nd MgF9 MgCL MgBo CaF. CaCll CaRr, CaL ZnF* ZnCk ZnBn MnF, MnCL
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/
Properties of Some Halides with Relatively High Charge on Halogen
Charge on halogen
mp ('(2)
-0.75 -0.67 -0.62 -0.54 -0.34 -0.28 -0.25 -0.47 -0.40 -0.36 -0.30 -0.24 -0.18 -0.15 (-0.30) (-0.25)
995 808 755 661 1263 714 711 1418 782(P) 760 740 872 275 394 856 850
Journal of Chemical Educofion
AH"juaion (kcal/mole) 7.8 6.8
...
...
13.9 10.3
...
7.1 6.8
...
...
... 5.5 ... ... ...
AHoWon bp ("C)
(kcal/mole)
1704 1465 1390 1300 2227 1418
50 40.8
2500
...
... ...
810 714 1502 756 650
...
1190
... ...
65 32.7
...
...
... ... ... 15.5 ... ... ...
AH0! (kcal/equiv) -136.0 -98.2 -86.0 -68.8 -131.8 -76.7 -61.9 -145.2 -95.0 -80.7 -63.9 -48.2(?) -49.7 -39.1 -95 -57.7
Table 4. Charge on halogen SF* AsR IF, GeF, N h WFs
CC4 PClr TeClr ICI BiCls S.CL TaCk WCL BBn SiBrr SbBn TeBn A&
Te4 SiL
-0.05 -0.11 -0.06 -0.16 -0.07 -0.06 -0.10 -0.06 -0.13 -0.11 -0.09
...
...
-0.11 -0.11 -0.07 -0.07 0 -0.01 -0.07
~.
Prooerties of Some Halides with Relatively Low Charge - on Halogen mp
("c)
-507(P) -5.95 4.5 -15.0(.P) -208.5 2.3 -22.9 -92 224.1 27.3 232 -80 217 275 -46 5.2 96.6 25 146 150 120.5
1.20 2.5
...
...
...
2.5 0.60
...
...
1.83 2.6
This effect is usually indistinguishable from a so-called cation-anion attraction, or an attractive electrostatic interaction among polar bonds. The physical properties (Table 3) are those predictable for such highly condensed aggregates. Whenever the simplest possible binary halide molecule contains no available low energy orbital vacancy, no further condensation of the compound is possible except through relatively weak van der Waals inleractias. Consequently the physical properties of such compounds can be predicted to be consistent with the molecular state. All such compounds are molecular, and if solid, are relatively low melting, and volatile, involving only relatively small enthalpy changes in the usual changes of phase. Many are liquid or gaseous a t ordinary temperatures. Bonds in such compounds are of generally low polarity, which itself would limit the force of intermolecular attractions. But both low polarity and lack of unoccupied orbitals occur together, since elements whose atoms have as many or more outermost electrons as vacancies and therefore can leave no vacancies after covalent bonding are also elements which are relatively high in electronegativity and therefore cannot form very polar bonds with the halogens. Some typical molecular halides are listed with some of their properties in Table 4. Halides in which the halogen has not acquired much negative charge tend to be low melting and volatile. I n effect, this is nearly equivalent to the preceding principle. Compounds of halogens with elements that do not give up electrons easily have weak intermolecular attractions for two reasons. One is the low polarity of bonds between elements both of which are relatively high in electronegativity, affording little dipole-dipole interaction. The other is the absence of outer orbital vacancies on atoms of the more electronegative elementsouce they have formed their quota of conventional covalent bonds, thus eliminating the possibility of condensation through donor-acceptor interaction. Compounds of halogens with elements that do give up electrons easily may also have low negative charge per halogen atom if several halogen atoms are competing for the electrons of the same other atom. The electronegativity of the central atom is increased to the extent
... ...
... ... ... ... ... ... ... ...
bp ('C)
AHo,,
AffDr,
(keel/mole)
(kcel/equ~v)
-63.7(s) 58 5.5 -36.8(s) -129.1 17.1 76.7 76 388
5.46(s)
441 138 233 347 90.5 152.8 280
... ...
...
...
(400) 365 290
...
7.37 7.8(s) 2.93 6.2 7.17 7.28
... ...
13.1
... ... ...
... ... ... ... ...
-44 -72.8
... ...
-9.1
...
-8.3 -27.0 -19.3 4.2 -30.2 -7.2
...
... -17.6 -23.8 -20.7
...
-4.6 -25.5 -7.9
that electrons are withdrawn from it, so that when several halogen atoms are pulling on its electrons, the central atom relinquishes only a little to each. Not only is each bond therefore less polar, but also the central atom commonly is so well surrounded by the halogen atoms that the exterior of the molecule is essentially all halogen and there is no opportunity for strong intermolecular attractions. Some examples of halides of low charge on halogen and their physical properties are given in Table 4. The oxidizing (or halogenating) power of halides i s inversely proportional to the partial negative charge on halogen. An atom of halogen does not lose its potential electron-attracting power merely by forming a covalent bond. It must succeed, through forming this bond, in acquiring substantial partial negative charge to become reduced in electronegativity. Otherwise it remains, in the compound, a potential halogenating agent from which the halogen may break loose if more favorable control of electrons becomes available to it from some other atom. All compounds in which the partial negative charge on halogen is not very high are therefore potential halogenating agents. These include not only halides of elements initially high in electronegativity but also polyhalides of elements initially low in electronegativity, in which the competition among halogen atoms prevents acquisition of substantial negative charge by any one of them. The compounds of Table 4 are all examples of potential halogenating agents. The ease of reaction will of course vary widely, depending on the available mechanbms by which the reaction might occur. We would not think, for example, of CC14 or SF6 as active halogenating agents a t ordinary temperatures, not because they cannot give up halogen or because the halogen in them is already highly negative, but because the means of doing so is limited by their structures which protect the bond to halogen from outside influence under these conditions. Mixing such compounds with strong r e ducing agents in a reactive form is therefore certain to be potentially hazardous since once it starts, reaction may be expected to be highly exothermic and probably autocatalytic. On the other hand, halides having highly negative Volume 4 1, Number 7, July 1964
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halogen cannot he very active oxidizing agents because their halogen has lost most of its electronegativity by combining. Sodium chloride, for example, is in no ordinary sense an oxidizing agent. But as the negative charge becomes less, a halide gains in oxidizing power, and becomes capable of oxidizing strong, and then weaker, reducing agents. The reducing power of halides i s directly proportional to the partial negative charge on halogen, andfor the sum. partial charge increases in the order F-Cl-BPI. The halide negative ions have the highest reducing power possible for halogen, hut since the initial electronegativity of all halogens is relatively high, they cannot he expected to give up electrons very easily once they have acquired them. Iodide ion is the best reducing agent of the group in keeping with its lowest initial electronegativity. Fluoride ion cannot he reduced except by a positively charged electrode because no chemical agent can remove the extra electron in competition with fluorine. The electrode oxidation potentials for the reactions 2 X - = Xz 2 e are: fluorine -2.87, chlorine -1.36, bromine -1.07, and iodine -0.54 v. The greater bond energy of the smaller and more electronegative oxygen combined with other elements makes their bromides and iodides especially susceptible to atmospheric oxidation. As might he expected, the halogens may also he displaced from their halides by more electronegative halogens. Although halogen atoms themselves cannot ordinarily function as electron acceptors, their compounds can ij' the simplest molecules would contain outer shell low eneryy vacant orbitals. This property is, however, primarily that of the element other than halogen, although the property is enhanced by electron withdrawal from the central atom by the halogen atoms. It results in the formation of many complexes, including complex halides where the donor is more negative halogen from another compound, such as the halide ion forming BFd- or PF6-. Some of the best known Lewis acids are boron and aluminum halides; any other metallic halide might so function, to the extent that the metal atom can act as acceptor. As mentioned earlier, such reaction among like molecules can lead to association such as dimerization of AICll or polymerization of ZnCl*. Between unlike molecules the reaction can produce a large number of molecular addition compounds suoh as BFa .O(CHs)t. The complexing ability of halides as donors i s proportional to the partial negative charge on halogen. The charge on fluorine (-0.25) in HF, for example, is insufficient to permit the fluorine to form stable complexes as donor. Acids such as HBF4 or HPFB are only hypothetical and exist only as the ionized form in solution, where the negative fluoride ion is the donor. Similarly, complex chlorides are not formed by HCl in which the charge on chlorine is -0.16 although more negative chlorine, especially chloride ion, readily forms them. As already mentioned, such complexing results in condensation of metal halides to ionic or polymeric crystallme forms, in many of which halogen bridges hold the metal atoms together. Only hydrogenJluorirEe of all the hydrogen halides is associated at ordinary temperatures by protonic bridging, because Jluorine alone of the halogens offers its outer
+
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Journal of Chemical Education
electron pairs in su&ient spatial concentration to attract a partially positive hydrogen or a proton. The requirements for protonic hridging are attached hydrogen bearing partial positive charge, and a small negatively charged atom having unshared outer electron pairs available. All the hydrogen halides satisfy the first requirement (although hydrogen iodide would satisfy it least). The second requirement is not met by larger atoms because, apparently, an electron pair spreads out to the greatest possible extent and thus becomes less available to a protonic hydrogen of another molecule. I n fluorine the smaller area available per electron pair restricts the negative charge to a region where more effectiveinteraction with a positive hydrogen can occur. The availability of an electron pair on fluoride ion is greater than on the fluorine in HF, so if protonic hridging can occur in H F it certainly ought to occur between H F and a fluoride ion. If it should, however, there would he, in effect, two fluorine atoms competing for the same hydride ion or two fluoride ions competing for the same proton (6). Neither of these hypothetical possibilities could result in an unsymmetrical bonding arrangement, because to the hydrogen the two fluorines must be exactly equivalent. A symmetrical bifluoride ion therefore results: (F-H-F)-. Hydrogen here is clearly divalent, which apparent violation of the Pauli exclusion principle may be rationalized by pointing out that the average occupancy of the hydrogen orbital does not exceed two electrons, but is only 1.04. This interpretation obviously distinguishes the bifluoride ion from examples of ordinary protonic bridging, a distinction borne out by the much higher bond energy (25-30 kcal instead of 5-10) in the bifluoride ion and the even spacing of the hydrogen halfway between the fluorines. Evidence of similar hihalide ions of the other halogens has recently been reported (7). The hydrogen halides increase in acid strength froin hydrogen Jluoride to hydrogen iodide. The same factors that contribute to protonic bridging appear to iufluence the action of a halide ion as electron donor to a bare proton. If the electron pair on the halide ion is concentrated within a small enough area, as in fluoride, the proton can he effectively coordinated to it. Otherwise, the coordination is weaker. The acid strength of a hydrogen compound in a solvent is a function of the competition between solvent molecules and hydrogen compound anion for the proton. The greater the relative donor ability of the solvent, the greater the acidity. I n a solvent such as water, only fluoride of the halide ions is an effective competitor a t ordinary dilutions, with the result that hydrogen fluoride is a weak acid but the other hydrogen halides are strong acids. I n water these others become leveled to the strengthof hydronium ion, hut in more acidic solvents the strength (as proton donor) is shown to increase in the order HF
