Ionization and Dissociation Energies of the Hydrides and Fluorides of the First Row Elements in Relation to their Electronic Structures BY W. C . PRICE, T. R. PASSMORE AND D. M. ROESSLER The Wheatstone Laboratory, King’s College, London, W.C.2 Received 28th January, 1963 Curves of the bond dissociation energies of the hydrides and fluorides of the first row elements are given together with the curves for the corresponding processes in the isoelectronic ions. These show the dependence of the energies on the closed shell configurations appropriate to the shape of the radical or ion. When the closed shells are exceeded, stabilization of triatomic species is achievea by bending while planar tetratomic ones become pyramidal. The curves are also used to obtain the proton affinities of H20 and NH3.The availability of reliable data on the bond dissociation energies of some radicals and molecules together with spectroscopic or photoionization values for their ionization potentials enables the bond dissociation energies of the ionic species to be derived. With these data it is possible to plot isoelectronic curves such as those shown in fig. 1 and 2 for the dissociation energies of H,X-H and H,X+-H. FIG. 1.-Plots of dissociation energies of H,X-H and H,X+-H for the first row elements. The value for the dissociation energy of the isoelectronic ion is plotted along the same ordinate as that of the corresponding neutral molecule or radical; thus D(Be-H) and D(B+--H) are plotted at the same value of x, there being unit dis- placement along the x axis for X = Li, Be, B, etc.The curve connecting the dis- sociation energies of the ions should resemble in its general shape that connecting the dissociation energies of the corresponding isoelectronic neutral species. To get from the latter to the former all that has been done is to increase the charge on X by unity. It can be seen for the hydrides that the dissociation energies of the ions are uniformly greater than those of the isoelectronic species. This is mainly attributable to the polarization energy of the hydrogen atoms in the ionic species. G* 201202 IONIZATION A N D DISSOCIATION ENERGIES OF RADICALS While many of the points on these curves have been determined with certainty, others such as N-H and N+-H are less accurately known. In these cases the adjacent points CH, C+H, OH, O+H and FH, whose values are well established, enable the values for NH and N+H to be determined by virtue of the similarity which exists between the curves.The detailed discussion of the dissociation energy curves which follow refers with one or two exceptions to the ground states of the radicals or their ions. ktrahedrol FIG. 2.--Plots of dissociation energies of F,X-F and F,,X+-F for the first row elements. THE X--H AND X+-H CURVES Most of the values are well established. Those of B+-H, N+-H and Ff-H are interpolated. That of F+-H does not correspond to the lowest state of the ion which is formed from F+H+ not F++H. The value used corresponds to the latter pair of dissociation products, i.e., to the removal of a nF electron (the same may be true for N+-H).The ions are formed by the removal of sa electrons for Li, pa electrons for Be and B and non-bonding n electrons for C, N, 0 and F. Dissociation is into neutral H and the lowest state of X or X+. The general form of the XH curve can be understood as the passage from sa binding in LiH topa in BeH reduced by the promotion energy required by Be(1S). This is followed by a similar bond in BH, which is stronger because less promotion energy is involved. The ionization of BH may, in an oversimplified way, be described as that of the unbonded spa electron produced in the promotion which results in two oppositely directed spa orbitals. From CH to FH the ionization is of a non-bondingpn electron, the bond energies of these radicals increasing with increasing polarity. There appears to be no irregularity arising from the different amounts of energy required to decouple the spins in the ground states of C, N and 0.As can be shown for C this energy seems only to be about 0.2 eV with similar amounts for N and 0. [The accurate spectroscopic data on the ionization energies of C(3P), CH,1 CH22 and CH3 can be explained assuming a polarization stabilization in the ions of 0.4 eV per hydrogen atom, a spin correlation energy in C(3P) of 0-225 eV and a spin correlation energy in CH3(3E;) of 0.16 eV according to the following equations : I(C3P) -Z(CH) = 1 1.265 - 10.64 = 0-625 = 0.40 + 0.225 eV, I(CH)-I(CH2) = 10.64-1040 = 0-24 = 0-40-0.16 eV, I(CH2)-I(C€€3) = 1040-9.84 = 0-56 = 0.40+0.16 eV.]W. C .PRICE, T. R . PASSMORE A N D D. M. ROESSLER 203 The low value for Li+H represents a common phenomenon found when the ionization can be regarded as that of an electron from a single X-H or X-F bond. When X has a low ionization potential, e.g., X = Li (5.39 eV), HBe (8.4 eV), H2B (8.2 eV), H3C (9.84 eV) the low values found for the dissociation energy of the ion are 0.8, 0.9, 0.9 and 1.26 eV respectively. This can be explained in terms of the poor exchange to be expected, the electron spending most of its time on the hydrogen. It also explains the diminishing proportion of parent RHf ions in fragmentation in mass spectrometry with increase in the size of the alkyl radical R, this being associ- ated with the diminishing ionization potential (i.p.) of R, resulting in a low bond energy R+ .H. The values are even lower for the fluoride ions as expected from the increase in the i.p. of F relative to H which leads to a further reluctance to exchange its electron with an attached radical of much lower ionization energy. The polarity of the bonds is increased in going from XH to X+H with correspond- ing increases in their strengths. For the fluorides it is diminished or reversed with in most cases reduction in bond dissociation energy. The bond energy of HeH is taken as zero. That of Ne+H could be evaluated by extrapolation to be about 9 eV but this would be an excited state which might rapidly dissociate into Ne+H+. There are, however, many cases of radicals containing heavier ionized inert gases where this dissociation does not occur and which are strongly bound.For example, the energies of A+-H and Mr+-H are greater than 4-5 eV since they are formed by the interaction of the inert gas ion with H2.3 These are to be compared with 4.43 eV for HCl and 3.75 eV for HBr their isoelectronic analogues. THE HX-H A N D HX+-H CURVES The values of HC-H, HCf-H, HN-H, HN+-H, HO-H and HO+-H which establish the major parts of the curve are fairly well known. They are given in table 1. HLi-H is assumed to be zero and HBe+-H determined from the curve mentioned in the discussion on Li+H. Its low value is to be correlated with the high value of D(Be+-H). The values of HBe-H and HB-H presented some difficulty. They are based on interpolations from the accepted values of H,C-H, H,N-H and their ions through the isoelectronic curves.That of HB-H depends, for example, on taking the value of the total bond energy of BH3 as 11.6 eV,4-6 subtracting B-H to give the sum of HB-H plus H2B-H and dividing this between them as indicated by the XH2 and XH3 curves. The high value of HB-H (4.7 eV) arises from the large overlap of its colinear sp bonding. The actual magnitude is important in that when compared with HC-H (5.45eV) it shows that the major factor in the binding of CH2 is also sp bonding and the additional stabilization arising from the spin correlation in the 3& ground state of CH2 is only a few tenths of a volt. The low value of the i.p. of BH2 (8.2 eV) which is confirmed by a value of 8.1 eV from an analysis of appearance potential data,31 is to be associated with the fact that a non-bonding p n electron is removed (8.3 eV in B, 2P+) and that the dissociation energy of the ion HB+-H acquires 0.4 eV polarization stability relative to the neutral radical.In this latter respect it differs from CH;! where the gain of 0.4 eV polar- ization energy is offset by loss of spin correlation energy (0-16 eV). Note that the upward kinks in the Be, B, C and B+, C+, N+ curves of the triatomic radicals are due to the effect of the spin correlation energy. The downward ones for the tetratomic species arise from departures from planarity. Maximum bond energy is achieved for the linear XH2 structures for CH2 and N+H2. It appears that in addition to the closed shell (ug)2(au)2 system the structure can accept one p electron in each plane on the central atom.These two electrons in mutually perpendicular orbitals are virtually non-bonding because of their204 IONIZATION AND DISSOCIATION ENERGIES OF RADICALS TABLE TA TABLE OF DISSOCIATION AND IONIZATION ENERGIES OF THE HYDRDES OF THE FIRSTROW ELEMENTS EXPRESSED AS D(HnX-H) : D(HnX+- H) 1WnX-H) : I(HnX) LiH BeH BeH2 2.43 (0.8) 2-3 3.2 (4.4) (0.9) (7-0) 5.39 8-4 9.32 (11.9) 8.4 BH BH2 BH3 3.39 (3.0) (4.7) (5.2) (3.2) (0.9) (8.7) 8.296 (8.2) 8-7 (10.5) (8.2) CH CH2 CH3 CH4 3-47 4.10 5-45 5-69 3.90 4.46 4-40 1.26 10.64 11-265 10.396 10.64 9.84 10-396 12.98 9.84 NH NH2 NH3 NH4 3.7 (4.3) 4.2 (6.8) 4.4 5.6 (0.) (5.5) (13*9)* 14.545 11-3 (13.9) 10.15 11.3 (4.7) 10.15 OH OH2 OH3 FH 440 (5.0) 5-20 (5-6) (09 (6.4) 5.87 (6.5) 13-0 13.615 12-61 13.0 (6.2) 12-61 (1 6*8)* 17.42 Brackets indicate values interpolated from isoelectronic similarity curves of radicaIs and ions.* Starred values do not correspond to lowest i.p.s. The interpolated values differ in many cases from those given previously (Price, Harris and Passmore, J . Quant. Spectrosc. Radiat. Transfer, 1962, 2, 327, due to the inadequacy of the data then available for plotting the isoelectronic curves. central position in the molecule though they do occupy the lowest nu bonding orbitals. (Only terminal p n electrons entering this orbital can be bonding.) The system does, however, derive some stabilization from the unpairing of their spins. This is lost when they are paired in the singlet system. For this system the most stable configur- ation is the bent one. In this respect the XH2 radicals differ from the XF2 radicals which cannot accept anypn: electrons on the central atom in the linear configuration because of the large repulsions which would arise with thepn electrons on the terminal F atoms.These repulsions are comparable with those in the diatomic species NO or CF and are of several eV in magnitude. This is equivalent to saying that in XF2 the pF electrons push the pX electrons out of the lowest bonding n orbital to the first antibonding one. By symmetry they cannot enter the second bonding n orbital. A radical change in structure occurs in going from CH2 and N+H2 to NH2 and O+H2. It is no longer energetically profitable to promote to a linear (sp) structure which could not be assisted by spin correlation due to the necessary pairing of the additional electron.The repulsion between the rc2 and the bond orbitals is probably also important. The change from linear to bent structure shows up strongly as a sudden drop in HN-H relative to HC-H (and HOf-H relative to HN+--H). The fall is about 1.5 eV. The changes that occur as bent NH2 is ionized to the linear triplet N+H2 are in striking contrast to what happens in ionizing CH2 or H20 where little change in shape or bond type accompanies the ionization. This is re- flected in the changes in i.p. for XH and XH2 set out below : I(CH) - I(CH2) = 10.64 - 10.40 = 0.24 eV, I(OH)-I(OHz) = 13-0- 12-61 = 0.39 eV, I(NH) - Z(NH2) = 13-9 - 11 -3 = 2.6 eV.W. C. PRICE, T . R. PASSMORE AND D. M. ROESSLER 205 The first five ionizations involve simply the removal of a non-bonding electron but with NH2 a reorganization of the remaining electrons accompanies this removal.The increase of HO-H relative to HN-H is presumably due to the greater electronegativity of 0 relative to N as in the diatomic case and also because no spin decoupling is required in its formation from OH(2II). In spite of its high predicted total bond energy HF+-H is not stable because of dissociation into HF+H+. It is expected, however, that €€2C1+ should be stable. THE H2X-H AND H2X+-H CURVES The known points on these curves are HzC-H, H2Cf-H, HzNf-H and H2Ot-H ; the last is derived from the proton affinity of H20 7 given as 7.35 Ifr 0.1 eV in a private communication, a value which leads to D(HzO+-H) = 6.4+0-1 eV. The dissociation energy of H2B+. H can be expected to be low because it dissociates into the strongly bound " closed shell " linear B'H2, the stability of which is shown by its high abundance in mass spectrometry apart from the dissociation energy derived from the previous set of curves and the low i.p.of BH2. The value taken for H2B+ . H is 1.6 eV. This derives from a Rydberg series in BH, going to 9.8 eV found by Herzberg (private communication) which accords with an a.p. of 11.9 eV for BH+ and a value of 1.8 eV for BzH642BH3. In addition it requires I(BH2) = 8-2 and D(H2B-H) = 3.2 eV which comes from the known values of B-H 8 and the total bond energy of BH3 by dividing the difference between HB-H and H2B-H in the only way that allows the curves of the energies of the neutral and isoelectronic radicals to retain their feature of similarity.The bonding in BH3 is not as high as in BH2 presumably because it corresponds to sp2 as against sp bonding. H2C-H (3-9 eV) is weaker than expected (mean bond energy CH3 = 4.27 eV) because of the extra binding due to spin correlation in CH2, i.e., one of its dissoci- ation products has extra stability. Here is an independent indication that spin correlation amounts to only a few tenths of a volt. If it were larger, then H2C-H should be much weaker than it is found to be since all stabilization from this cause is lost in going from CH:! to CH3. Although it might be expected from the change in shape in going from BH3 and CH3 to NH3 that there is appreciable change in bond character, this is accompanied by little change in mean bond energies.These are 3.87, 4.27 and 4.1 eV, respectively, for BH3, CH3 and NH3. This is not sur- prising since it only requires 0.25 eV, the barrier height in NH3, to convert this molecule to the planar structure similar to BH3 and CH3. The fact that H2N-H is greater than H2C-H although their mean bond energies are in the reverse order is mainly a result of HN-H (bent) being lower than HC-H (linear). The value of H20+-H could be fixed by the known points involving Cf, C, N and N+ and it is satisfying that a value thus predicted should agree so well with results from ion impact 7 and considerations of heats of formation of crystalline perchlorates and lattice energies.99 10 It enables an effective ionization energy of 64eV to be associated with H3O (see table 1) which explains how this group can act as a metal in combination with electronegative atoms.The soundness of this extrapolation from the neutral NH3 to the isoelectronic positive ion O+H3 can also be supported from considerations of bond lengths and force constants. The positive charge contracts the bonds and tightens the binding of the system. The similar extrapolation of the properties of NH$ from those of CH4 will be demonstrated in the next section. THE H3X-H AND H3X+-H CURVES The known values are H3C-H and H3C+-H. Assuming H3B-H to be zero and drawing parallel curves gives 5.5 eV for H3N+-H, a value which agrees with206 IONIZATION A N D DISSOCIATION ENERGIES OF RADICALS the heat of formation of NHi(g) of 150 kcal calculated by a cyclic process from the heat of formation of crystalline ammonium chloride and the electron affinity of chlorine 8 .9 and with the value obtained in ion impact studies.7 Assuming (H3N-H) = 0 and using the i.p. of NH3 = 10.15 eV a value of 4.7 eV is obtained for the ionization energy of NH4 which explains why it acts as an alkali metal. By a similar method the ionization energy of PH4 can be found to be 6.4 eV, the higher value correlating with its lower electropositive character. The bond distances of NHt, CH4 and BH, are 1-03, 1.093 and 1.26 A and their v3 frequencies 3146,11 3020 and 2300 cm-1 respectively. These values fall in line with the progressive decrease in dissociation energies. BOND DISSOCIATION ENERGIES OF THE FLUORIDES OF THE FIRST ROW ELEMENTS A N D THEIR IONS The isoelectronic curves of the fluorides XFn do not show such close quantitative similarity as those of the hydrides because the bonds are strongly polar and in- creasing the charge on X greatly diminishes and possibly reverses the bond polarity.As a result, the bond dissociation energies of the ions are mostly less than those of the corresponding isoelectronic neutral species. The positive charge is no longer located mainly on the X atom but spreads out to the fluorine atoms. Because of the nature of the closed shells available to R: electron systems, the shapes of the radicals differ in certain cases from those of the corresponding hydride radicals (cf. CF2 and CF3 as compared with CH2 and CH3). Also, whereas it can be assumed that He-H, HLi-H, H2Be-H and H3B-H have zero dissociation energy, this is not so for the fluorine analogues because of the high electron affinity of F.How- ever, it does appear that these attachment energies are small. In the later discus- sions it is assumed that in the structures of the radicals and their ions only the ground states of X, X+ and F are involved. TABLE 2.-TABLE OF DISSOCIATION AND IONIZATION ENERGIES OF THE FLUORIDES OF THE FIRST ROW ELEMENTS EXPRESSED AS D(F,X-F) : D(FnX+-F) I(F,X-F) : I(F,X) LiF BeF BeF2 6-0 0.3 5.4 (5.6) (7.5) (0.5) 1 1 - 1 5.39 (9-1) 9.32 (16.1) (9.1) BF BF2 BF3 8.5 (5-2) (4-2) (6.3) (7.0) 0.8 (11-6) 8-296 (9.5) (11.6) (15-7) (9-5) CF CF2 CF3 CF4 4-9 7.3 5.2 3.0 (4.8) (5-8) 5.3 0 8-91 11.265 11.1 8.91 10.1 11.1 15.4 10.1 NF NF2 NF3 NF4 (3-15) (4.8) 3.0 (3.8) 2-5 (1.4) 0 (4.4) OF OF2 OF3 F2 (12-9) 14.545 (121) (12-9) 13-2 (12.1) (8-8) 13.2 (2.25) (3.7) (1.7) (0-3) (0.7) (0.0) 1-63 (3.3) (12-2) 13.615 13-6 (12.2) (14.3) 13.6 (15-7) 17-42 Brackets indicate values interpolated from isoelectronic similarity curves of radicals and ions.W.C . PRICE, T. R . PASSMORE AND D. M. ROESSLER 207 THE X-F AND X+-F CURVES The points well-established on these curves are those for LiF,12 Li+F,13 BF, CF, C+F 14 and F2. With a lower accuracy we have BeF, Be+F, OF and O+F, some of the values of which will be derived. A value of 3-3 eV for D(F,C) given by Iczkows and Margrave 15 and based on an ionization potential of 15-7 eV for F2 from a tentative Rydberg series was at first thought to be too high to be acceptable.However, it fits in with the flatter nature of the ion curve as compared with that of the neutral molecules. We take 3.3 and 15.7 eV respectively for D* and I. The first pair of points on these curves have low values. The bond energy of Li+ . F is low as is apparent from its appearance potential and in agreement with its expected small exchange. That of HeF is not known but in view of the electron affinity of fluorine and the high ionization potential of helium an attachment value of a few tenths of a volt is expected, slightly more than Li+. (Note I(He) = 24.6 eV. I(Li+) = 75 eV.) The next pair Li-F (6.0 eV) and Be+-F (5.6 eV) about equal in magnitude though probably different in character. Whereas in LiF the bond is mainly ionic, in BefF it is mainly covalent since the ionization energy of Be+ is 18.21 eV as compared with 17-42 eV for F.The near equality of these differen- bonds is not unexpected since the bond energy in HF (5.87 eV) which is probably nearly 50 % ionic is also almost identical with that in LiF. As in the hydrides, the dissociation energies of BeF and BfF are lower than the preceding pair due to the promotion energy necessary to convert the 1S to the valence state. These molecules can also be regarded as having “ closed shell minus one” configurations and the values are consistent with this approach also since they are slightly higher than those of the corresponding “ closed shell plus one ” structures. This agrees with the expectation that the ratio of the antibonding to the bonding power of an electron in these orbitals is as (1 + S)/(l- S ) , where S is the overlap integral. The bond energies reach their peak in BF and C+F which have closed shell structures of bonding electrons.They are isoelectronic with N2, COi, NO+ and have comparable binding energies. The next pair, CF and N+F, have structures involving a closed shell plus one antibonding n electron. The antibonding power of this electron is large whereas the effect of additional antibonding electrons be- comes progressively smaller. A simplified theory based on the assumption that in CF, for example, “ antibonding ” involves resonance to structures C=F+, predicts relative total antibonding powers of approximately 1 : 1.5 : 1.75 and 1-88 for 1 , 2, 3 and 4 antibonding electrons respectively (when two antibonding electrons are present as in NF, two p electrons cannot simultaneously leave the fluorine, thus restricting the operative time for repulsion).This is close to the ratio found for a number of sets of such molecules for which dissociation energies are known. The radicals NF and OF+ are isoelectronic with 0 2 , the 3Zg ground state of which is stabilized by interaction with triplet states of the same symmetry arising from 3P+1D atomic terms. The stabilization in 0 2 is about 1 eV as can be determined by its departure from a plot of the dissociation energies of BeO, BO, CO, NO, (00), OF as ordinate with unitary abscissa displacement for successive molecules. A curve through B2, CZ, N2, (02), and F2 also gives this value for the extra stabilization.However, an examination of the term values of the low states of N,O and F indicates that it is unlikely that stabilization of this magnitude is present in N F and OF+. The former is estimated as 3-15 eV both from the BF, CF, (NF), OF, FF curve and also from a BN, CN, NN, NO, (NF) curve. The dissociation energy of OF+ is Its large value is partly ionic in origin, cf Ne’F. -208 IONIZATION AND DISSOCIATION ENERGIES OF RADICALS taken as 3.7 eV from the CfF, N+F, (O+F), F+F curve and also a C+O, N+O, OfO, (O+F) curve. Both these values fit in well with the energy schemes and ionization potential data for NF3,2,1 and OF1,2 radicals.16917 The value 2.3 eV has been taken for OF partly because it fits in with the division expected for the total bond energy of OF2 (3.9 ev) between the fist and second bond dissociations when compared with the other halogen oxides.In these, steric factors make the first dissociation less than the second by amounts which are greater in C120 and Br2O than F20. The value taken for OF fits the present curve. It can hardly exceed the well-established value for F2 by more than 0.6 eV in view of the progressively diminishing effect of added antibonding electrons. THE FX-F AND FX+-F CURVES These curves show the dependence of binding on closed shell structures. They build up to a high maximum in BeF2 which is a linear closed shell structure analogous to C02. The additional p electron in BF2 causes the radical to bend as in the iso- electronic radical N02, thep orbital lying along the bisector in the BF2 plane.This structure can take another p electron which is paired with its predecessor giving a fairly stable CF2(1A) which is a subsidiary closed shell structure. The source of this stabilization is the high inductive effect of the F atoms on paC electrons. Their orbitals are oriented in a suitable way to be taken in by the fluorine atom to complete its closed shell and so to satisfy its electron affinity. This is not the case for sub- sequently added p n electrons which have to stand perpendicular to the plane and become involved in strong resonance repulsion with the pn electrons on the fluorine atoms. This leads to rapid reductions of bond energy for FN-F and FO-F relative to FC-F similar in magnitude to the repulsions found in the analogous diatomic cases.The first neutral radical and isoelectronic ion pair on these curves is FLi-F and FBe+-F. Both are assumed linear 2E. Although no work on FBef-F has yet been reported the ions of MgF2 and MgC12 are stable enough to be observed.18 The appearance potentials of these ions combined with bond dissociation data in- dicate values of up to 1 eV for XMg+ . X. A value of 0.5 eV is taken for FBef . X and 1.4 eV for FLi-F. This may seem arbitrary but they are both small and yet they cannot be zero as they correspond to structures which are only one electron less than linear triatomic closed shells. For comparison with FBef . F, we have Lif . F = 0.3 eV, F2B+ . F = 0.5 eV 19 and F3Cf. F = 0.0 eV. Their exact magnitudes do not greatly affect the discussions which follow.The next pair FBe-F and FB+-F are the linear closed shell species. The energy of the first of these is found from the total energy of BeF2 and the spectroscopic value for BeF. There is no experimental value for the former but a value can be obtained by extrapolation from experimental values for BeC12, BeBr2, Be12 using as a comparison the series MgF2, MgC12, MgBr2 and MgI2.18 This leads to a value of 12.9 eV from which we obtain 7-5 eV for FBe-F by subtracting 5-4 eV for Be-F. This high value is to be expected in view of the linear closed shell structure of BeF2. For the dissociation energy of FB+-F it is necessary to choose a value of 6-3 eV to fit in with several experimental facts. The derivation consists in fitting in values of the dissociation and ionization energies of the BF3 section of table 2.In this section all the values for BF are established and in addition the i.p. of BF3 is known (15-7 eV) and the dissociation energy of F2B+. F (0.5 eV) 19 can be derived from its a.p. The total bond energy of BF3 using the latest value for the heat of sub- limation of boron 20 is 19.7 eV from which subtraction of 8-5 eV for B-F givesW. C. PRICE, T. R . PASSMORE AND D. M. ROESSLER 209 11.2 eV for the sum of FB-F and F2B-F. The latter of these is expected to be much greater than the former because of its closed shell configuration, Also the former is weak because it dissociates into the strongly bound closed shell BF. It is also expected to be bent (cf. N02). The dissociation energy of FBf-F is expected to be high because it is a linear closed shell species, a fact supported by the high abundance of BF; in mass spectrometry.19 Fairly well-established values for the dissociation energies of the adjacent radicals CF2, C+F2 and N+F2 help to fix F2B-F at 7.0 eV and FB-F at 4.2 eV.These lead to a value of FB+-F of 6.3 eV, an ex- pected high value, and an i.p. of 9.5 eV for BF2, a low value which can be compared with the i.p. of 8-2 eV derived for BH2, also a low value (compare the i.p.s CH2, 10.40 eV and CF2, 11.1 eV, also OH2, 12.61 eV and OF2, 13.6 eV). It is clear that the variation in ionization and dissociation energies of the boron fluorides exemplify four types of closed shell, namely, BF, B+F2, BF3 and BF:. The values of the heat of formation 21922 of CF2 lead to values of FC-F which vary from 4-85 to 5.65 eV.We take a value of 5.2 eV as most probable and fitting our curve best. Combined with an i.p. of CF2 of 11-1 eV (originally given by Mar- grave 23) but supported by spectroscopic and photoionization work on C2F4 2425 and appearance potential data23 also comparison with the i.p.s of BH2, BF2 and OH2, OF2 given above), this leads to a value of 3.0 eV for FC+-F. For the nitrogen fluorides the values are established by taking the diatomic values from the curve, accepting a value of 2.5 eV for F2N-F 2 6 2 7 and thus obtaining FN-F by sub- traction from the total bond energy. The ionization energies of NF2 (12-1 eV) 28 and NF3 16 (13.2) permit the calculation of FN+-F and F2N+-F. The former of these together with FC-F and FC+-F enables FB-F to be located at 4.2eV by assuming similar behaviour of the isoelectronic curves.This then gives the line between FBe-F and FB-F and a roughly parallel line through FC+--F locates the value of FB+-F at 6.3 eV. THE F2X-F AND FzX+-F CURVES The curves of F2X-F and F2X+-F rise to maxima for the planar closed shell n:nz structures BF3 and C+F3. (n1 has no transverse nodal plane, n2 has a central nodal plane perpendicular to the molecular plane.) Similar closed shell structures are responsible, for example, for the stability of the carbonate and the guanadinium ions and govern chemical activity in a wide range of substances of similar structure. We would even include benzene with these structures, particularly since none of these structures can have a pn electron located on the central carbon atom.This can only happen in the hydrides XH, where a non-bondingp electron is the sole occupant of the n1 orbital. Because of the central nodal plane in 712, a centralpn electron can- not contribute to a n2 molecular orbital. Thus, when there are terminalpn electrons also contributing to the n1 molecular orbital, the central pn electron is pushed largely into the antibonding n3 molecular orbital the nodal surface of which passes through all three bonds. (The bonding of the n1 orbital formed from a centrally located pn atomic orbital in XH, can only be between the terminal H atoms. It is doubtful whether this can be different from zero.) Apn electron introduced on to the central atom of XF3, as in CF3 becomes involved in strong resonance repulsion with the p n orbitals of the terminal atoms.The interaction is between the ynC electron and electrons in a n(F + F + F) orbital. (It is convenient to consider sub-molecular orbitals made up only of F orbitals. Only those of the right symmetry can interact with the central C orbital which is the two electron orbital given above.) The inter- action is thus similar to that in CF where the antibonding relative to BF is indicated by the difference in dissociation energy, viz., 8-5-5.0 = 3.5 eV. Without exception,210 IONIZATION AND DISSOCIATION ENERGIES OF RADICALS this repulsion is relieved by bending out of linearity in the triatomic case or out of the plane in the tetra-atomic case. In this way the resonance causing the repulsion is reduced by the consequent departure from parallelism of the axes of the p orbitals.The following comparison giving the differences in the bond dissociation energies for " closed shell " and " closed shell plus one " structures illustrates this : BF-CF = 8.5-5.0 = 3.5 eV, (FBe-F)-(FB-F) = 7.5-4-2 = 3.3 eV, (F2B-F)- (F2C-F) = 7.0-4.8 = 22eV. Several factors affect these differences but the antibonding power of the additional electron is undoubtedly the major one. The antibonding power of a second antibonding electron does not fall off relative to that of the first in CF3 and NF3 as in CF and NF, the antibonding being almost equal for the tetra-atomics instead of reduced by half as it is in going from CF to NF. This is thought to be connected with the possibility of resonance structures occurring in the tetra-atomic case in which several pF electrons can interact simultaneously with the p X electrons without invoking structures involving doubly or triply charged fluorine atoms as would be necessary in the diatomic cases. The only points on these curves which have not been considered are those of Be, C and Cf.The first of these FZBe-F would be expected to be low but not necessary zero since the structure corresponds to one electron less than the closed shell BF3. A value of about 2 eV obtained by drawing the BBe line parallel to the C+B+ line seems to be about right though it might be less than this. The value of FZC-F depends on the heat of formation of CF4,29 that of F3C-F 30 and the values of FC-F and C-F already discussed.The i.p. of CF; 30 together with the above data gives FzC+--F. THE F3X-F AND THE F3Xf-F CURVES The two points known on these curves are F3C-F 30 and F3C+-F, the latter being zero. A low value of a few tenths of a volt is assumed for F3B-F (it might be zero). Joining points B to C and drawing the C+F+ line parallel indicates a value of F3N+-F only slightly less than that of F3C-F. Little is known of the NfF4 ion, in contrast to the N+H4 ion. Its ionization potential can be derived as ca. 8 eV from NF3+F = 0, N+F3+F = 5 eV and i.p. NF3 = 13-2 eV. It is clear that its ionization potential is too high and its dimensions too large for it to give enough lattice energy to form crystalline salts. We should like to acknowledge support from the Institute of Petroleum, Imperial Chemical Industries, the Department of Scientific and Industrial Research and the U.S.Dept. of the Army through its European Research Office under Contract DA-9 1-59 1 EUC-1683. Note added in proof: If a Rydberg series converging to 9.8 eV found for BH (Herzberg and Johns, private communication) corresponds to the lowest state of the ion then the following changes have to be made in table I : D(B+-H) = 1-9 eV, I(BH) = 9.8 eV, Z(BH2)=9.3 eV and I(BH3) = 11.6 eV. 1 Herzberg, J . Quunt. Spectrosc. Radiat. Transfer, 1962, 2, 319, and private communication. 2 Herzberg, Can. J. Physics, 1961, 39, 155. 3 Stevenson and Schlissler, J. Chem. Physics, 1955, 23, 1353. 4 Shepp and Bauer, J. Amer. Chem. SOC., 1954,76,265. 5 McCoy and Bauer, J. Amer. Chem. Soc., 1956,78,2061. 6 Gunn and Green, J. Chem. Physics, 1962, 36, 1 1 18. 7 Tal'rose and Frankevitch, Dokf. Akad. Nuuk. 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