Photoelectron
spectroscopy
IO-, OCIO-, and OIO-
of the halogen oxide anions FO-, CIO-, BrO-,
Mary K. Gilles, Mark L. Polak, and W. C. Lineberger
Joint Institute for Laboratory Astrophysics, University of Colorado and National Institute of Standards and
Technology, and Department of Chemistry and Biochemistry, University of Colorado, Boulder,
Colorado 80309-0440
(Received 29 January 1992; accepted 28 February 1992)
The 35 1 nm photoelectron spectra of FO-, ClO-, BrO-, IO-, OCIO-, and OIO- are
reported. The spectra of the halogen monoxides display transitions to both spin-orbit states of
the ‘Iii ground state neutrals. Anion vibrational frequencies are observed in the spectra and
bond lengths are obtained for the anions from Franck-Condon simulations. Spectra of the
halogen dioxides display two active vibrational modes-the symmetric stretch and the bend.
Anion symmetric stretching frequencies and normal coordinate displacements from the
corresponding neutral are reported. Adiabatic electron affinities found for the halogen oxides
are 2.272(6) eV (FO), 2.276(6) eV (ClO), 2.353(6) eV (BrO), 2.378(6) eV (IO), 2.140(8)
eV (OClO), and 2.577( 8) eV (010). The difference between the neutral and anion
dissociation energies [D, (X0) - De ( X0- ) ] is reported for each of the halogen monoxides.
Anion heats of formation (298 K) are also determined.
I. INTRODUCTION
Several of the halogen oxides play important roles in
atmospheric chemistry. Neutral Cl0 and BrO are formed in
the stratosphere from the reaction of Cl or Br with ozone,
resulting in the destruction of ozone.re3 Iodine undergoes a
similar reaction in the troposphere.4 Atomic fluorine reacts
rapidly with species containing hydrogen to form HF, which
acts as a sink for atomic fluorine in the atmosphere.5 At
higher altitudes, in the D region, negative ions such as Cl-,
ClO-, and OClO- are present.3 Despite the relative importance of the halogen oxides, basic physical properties, e.g.,
electron affinities, proton affinities, anion structures, and vibrational frequencies are poorly characterized.
The appearance potential of FO- was measured from
the electron bombardment6 of CF,OF and F,O.’ Haaland’
recently reported fourth-order unrestricted Msller-Plesset
single, double, and quadruple excitation [ UMP4( SDQ) ]
potentials for the lx+ and 311states of FO-. Peterson’ performed MP4SDQ calculations on FO-, predicting r, ( 1.5 13
A),@, (814cm-‘),
andw,x, (5.7cm-‘).
Although present in low abundance in the mesosphere,3
ClO- is highly reactive. Observations of ClO- in the gas
phase have been made by photodissociation, photodetachment, ‘O ion-molecule reactions, I’ charge transfer reactions,
and endoergic collisions.” Lee et al.” measured the ClOvisible photoabsorption spectrum and attributed it to two
components-a broad photodetachment feature beginning
-5700 A and a narrower photodissociation feature (to
Cl- + 0), peaking near 4300 A and about 400 A wide. Reaction rate constants have been measured in a flowing afterglow apparatus for ClO- with NO, NO,, SO?, CO,, and 0,
by Dotan et al.” Vogt et a1.12 measured the energy dependence of the ClO- product channel in the reaction of
Cl- + 0,. Baluev et all3 and Dibeler et alI4 investigated
8012
J. Chem. Phys. 96 (1 l), 1 June 1992
0021-9606/92/l
the appearance potentials of ClO- from electron impact on
ClO, and ClO,F. Each of these experiments on ClO- placed
upper or lower bounds on the electron affinity of ClO. A
solution phase vibrational frequency of 7 13 cm- ’ has been
reported.i5
Theoretical work’“” has focused on predictions of the
bonding character, the electron affinity, dipole moment, and
bond length of ClO-. In 1971, O’Hare and Wahl” calculated the dipole moment of the neutral and anion, and estimated the vertical electron affinity of Cl0 to be 2.2 f 0.5 eV.
Pershin and Boldyrev’* used second-order Moller-Plesset
perturbation theory to investigate the electronic structure,
bond length, and vibrational frequency of ClO-. Peterson
and Woods” have recently reported high level calculations
on ClO-, determining r,, w,, and w,x,, using the correlated
electron pair approximation (CEPA), singles and doubles
contiguration interaction (Cl-SD) with Moller-Plesset
fourth-order perturbation theory.
Few investigations of BrO- have been reported. The
BrO- ion has been observed2’ as a product of the reaction of
O- with CHBr,, CF,Br, and C,F,Br. The threshold of the
Br- + O,-+BrO- + 0 (4.58 f 0.2 eV) reaction was used
to estimate the electron affinity of BrO as ) 1.5 f 0.2 eV. l2 A
solution phase vibrational frequency of 620 cm- ’ has been
reported.21
IO- appears to have been first observed in the gas phase
by Henglein and Muccini22 as a product of O- reacting with
I,. Wren et a1.23 observed IR absorption frequencies of 430
and 560 cm- ’ in a solution containing IO-, along with several other absorbers. Although they assigned the lower frequency to IO-, the higher frequency is quite close to our
recently reported24 gas phase value of 58 l(25) cm-‘. Most
of the experimental work on IO- has focused on endoergic
ion-molecule reactions, which placed limits on the electron
affinity of IO, .20.22,25-27
these are discussed individually in the
section reporting electron affinities. Using photoelectron
18012-09$06.00
@ 1992 American Institute of Physics
Downloaded 04 Dec 2006 to 128.138.107.158. Redistribution subject to AIP license or copyright, see http://jcp.aip.org/jcp/copyright.jsp
Gilles, Polak, and Lineberger: Photoelectron spectroscopy of halogen oxide anions
spectroscopy, we recently determined the electron affinity of
IO to be 2.378( 6) eV.24
Both experimental and theoretical work on the halogen
dioxide anions is minimal. Appearance potentials for
OClO- from electron impact on chloryl fluoride (ClO,F) l3
and perchloryl fluoride (ClO,F) l4 have been reported. Each
of these displayed two low energy resonances. A recent study
of electron attachment to OClO found that dissociative attachment to form ClO- was the most abundant channel.28
BybergZ9 derived nuclear quadrupole coupling constants
from the electron-spin resonance (ESR) spectra of ClO; .
Two sets of vibrational frequencies are quoted for OCIO(884,797, and 396cm-‘)‘6
(823.4, 789.7,418.4cm-‘).I*
In this paper, we report the electron affinities and vibrational frequencies of the halogen oxides FO-, ClO-, BrO-,
IO-, OCIO-, and OIO-. Geometries are also given for the
negative ions. This paper is divided into several sections. The
experimental section gives a brief overview of the experiment
and details on production of the individual halogen oxides.
Electron affinities are reported in the Discussion section and
results from the Franck-Condon analyses are provided. We
then discuss the thermochemistry and dissociation energies
of the halogen oxides.
II. EXPERIMENTAL METHODS
Only a brief discussion of the experimental apparatus is
given here since it has been described in detail in earlier publications.30*3’Negative ions produced in a flowing afterglow
source32 are gently extracted, accelerated, and focused into
an ion beam. Following mass selection in a Wien filter, the
ions are decelerated to increase the residence time for interaction with the 35 1.1 nm argon ion laser light. The present
UV laser system is described in detail in a recent paper.31
Photoelectrons emitted in a small solid angle perpendicular
to the ion and laser beams are focused into a hemispherical
analyzer. Calibration for the absolute electron kinetic energy3” is performed with O-, and a small correction (0.5%)
for the energy scale compression of the hemispherical energy
analyzer is obtained from the fme structure34 of the photoelectron spectrum of W -. Transitions from the anion initial
state to a number of neutral vibrational states are obtained
by measuring the kinetic energy of the electrons (eKE) from
the process
X0;
(v”) + hv+XO,(v’)
+ e-(eKE),
where u” and u’ denote the anion and neutral vibrational
states.
All of the ions in this experiment are produced in a flowing afterglow source.32 Several ion-molecule reactions using
O- were employed, as described below. For each of these
reactions a flow rate of approximately 7 std 1 min- ’ of buffer gas (He) was used with a total pressure of 0.5 Torr.
Although FO- has been reported as a product of electron bombardment of CF30F,6F,0,7 and C103F,35 the instability of these compounds prompted us to investigate other
sources for FO-. We were able to make an ion current of 25
pA of vibrationally hot ( =: 800 K) FO- by introducing O,,
trace NF,, and He buffer gas upstream of the microwave
cavity. Although the exact reaction responsible for produc-
8013
tion of FO- is unclear, the neutral spin-orbit splitting and
vibrational frequency36 evident in the photoelectron spectrum of the mass 35 anion easily verified the identity of FO- .
ClO-, BrO-, and IO- have all been reported as products of ion (O- )-molecule reactions. ’ls2’The requisite Owas made by flowing O2 ( 10 std cm3 min- ‘) and the helium
buffer gas through a microwave discharge cavity. The 0,
dissociatively attaches an electron to form O-, which reacts
with the reagent gas introduced downstream of the microwave cavity. Employing a flow rate of 6 std cm3 min-’ of
Ccl,, approximately 50 pA of 3*C10- was produced. Following the recipe of Streit, *’ CF,Br was introduced downstream of the microwave discharge and 45 pA each of
79Br0- and *‘BrO- were made. In a similar manner, using
CF,I, 110 pA of IO- was obtained along with substantial
currents (25 PA) of OIO-.
There was no simple reaction available to us for OClOproduction. Instead, a complex scheme was employed involving O-, SF,, and OClO. Maximum OClO- current was
obtained when O- was made as described earlier and SF,
was introduced just past the microwave discharge. Neutral
OClO, made from a heated (50-60 “CL!),moistened ( 1 ml
H,O) mixture of oxalic acid (7 g) and KClO, (2 g),37 was
introduced even further downstream. Although SF; does
not undergo charge transfer readily, charge transfer to OClO
has been suggested as a method of OClO- production.37
Several other negative ions were present in the mass spectrum, making a definitive determination of the reactions directly responsible for OClO- impossible. Nevertheless, ion
currents up to 150 pA of OClO- were produced.
III. RESULTS
The photoelectron spectra of FO-, ClO-, BIG-, and
IO- (taken at the “magic” angle) are shown in Fig. 1. Photoelectron intensity vs electron binding energy is plotted for
each of the halogen monoxides. The electron binding energy
(eBE) is obtained by subtracting the measured electron kinetic energy (eKE) from the photon energy (eBE = hv
- eKE) . The adiabatic electron affinity (EA) corresponds
to the transition from the lowest rovibrational state of the
anion to the lowest rovibrational state of the ground state
neutral. In Fig. 1, the data are plotted as points and FranckCondon simulations are plotted as solid lines.
Transitions from the ground state anion ( ‘X+ ) to each
of the 211i states are seen. The halogen monoxides possess
inverted spin-orbit states, so the 211’,2 states are higher in
energy than the 2113,2ground states. For FO, the spin-orbit
splitting is about a factor of 5 less than the vibrational frequency36 and a long progression of closely spaced doublets
appears in the photoelectron spectrum. The distance between the peaks forming the doublet is the spin-orbit splitting (193.8 cm- ’). 36In ClO, the spin-orbit splitting is about
one-third of the vibrational frequency3* and a distinct progression of doublets emerges. The spin-orbit splitting39 for
BrO is larger than the vibrational frequency4’ and the vibrational assignments are not immediately obvious because progressions from each of the spin-orbit states overlap. In the
photoelectron spectrum of IO-, two distinct electronic
states are detected because the spin-orbit
splitting
J. Chem. Phys., Vol. 96, No. 11,l June 1992
Downloaded 04 Dec 2006 to 128.138.107.158. Redistribution subject to AIP license or copyright, see http://jcp.aip.org/jcp/copyright.jsp
Gilles, Polak, and Lineberger: Photoelectron spectroscopy of halogen oxide anions
8014
l-ii-
(20)
FO-
(3,O)
n
0
z
t-GQ 6.0
n
i
BrO-
0 .o
3:l
3:0
2:9
2:4
2:5
217
216
2:6
ELECTRON
BINDING
ENERGY
r ao-
2:3
2:2
211
2
217
(eVJ
2.6
ELECTRON
2.5
BINDING
5.0,‘.““““““““““..““.‘,
2.4
ENERGY
2.3
2.2
feVJ
Ill 1,
(LO)
r7
1
ii
0 .o
2.6
2.7
2.6
ELECTRON
2.5
BINDING
ENERGY
ELECTRON
leVl
BINDING
ENERGY
leVJ
FIG. 1. The 35 1 nm photoelectron spectra of the halogen monoxides X0- (X = F, Cl, Br, I). Transitions from the ground state anion X0- ( Ix+) to each of
the X0(%,)
states are seen. Each vibrational assignment is indicated by a horizontal line, where the transition to the 2IlJ,2 state is the peak with the lower
electron binding energy. These lines are designated (v’,v”) corresponding to the neutral and anion vibrational states, respectively.
[ 209 l(40) cm- ‘1 24 is significantly larger than the neutral
vibrational frequency (68 1 cm- ’) .41 In Fig. 1, each vibrational assignment is indicated by a horizontal line. These
lines are designated (v’,u” ), corresponding to the neutral
and anion vibrational states, respectively For each horizontal line, the transition to the *IIX,* state is the peak at lower
electron binding energy, indicated by an arrow.
Vibrational quanta up to n = 6 are seen in the photoelectron spectrum of FO- and only up to n = 2 for IO-. In
fact, the number of vibrational quanta observed in the neutral species decreases systematically in the order FO> ClO- > BrO- > IO-. Because the anion is formed by the
addition of an electron into an antibonding orbital, the negative ions all possess longer bond lengths than their corresponding neutrals. Since the neutral bond length is significantly longer for IO (1.8677 A)4’ than for FO ( 1.354A),36
the addition of an electron into the antibonding orbital results in a smaller increase in bond length for IO- than for
FO-. The d@rence between the anion and neutral bond
lengths decreases with the larger halogens, giving a narrower
Franck-Condon envelope.
The halogen monoxide anion ground states ( ‘Z+ > are
expected to be strongly bound.8~‘0*‘9Lee et al.” discussed the
possibility that one or more of the anion states (32-, 311,3A,
‘II, and ‘A), arising from the Cl(*P) + O-( ‘D) asymptote
of ClO-, could be slightly bound. We searched at lower electron binding energies for transitions from excited anion electronic states of the halogen monoxides, but did not detect
photoelectron signal. However, our inability to discern excited anionic states does not preclude their existence. Poor
Franck-Condon overlap between excited states of the anion
and the neutral, or collisional cooling of anion excited states
in the source, could limit our ability to observe these transitions. All of the halogen monoxide photoelectron spectra
displayed in this paper are of the 2113,2+‘X+ and
*I-L/*+ ‘Z+ transitions.
The photoelectron spectra of OClO- and OIO- taken
at the “magic” angle, where the photodetachment signal is
proportional to the total photodetachment cross section, appear in Fig. 2. Each of the origins, the transition from u = 0
of the anion to u = 0 of the neutral, is indicated by Oz. Both
the neutral ground state (‘B,) and the anion ground state
J. Chem. Phys., Vol. 96, No. 11,l June 1992
Downloaded 04 Dec 2006 to 128.138.107.158. Redistribution subject to AIP license or copyright, see http://jcp.aip.org/jcp/copyright.jsp
Gilles, Polak, and Lineberger: Photoelectron spectroscopy of halogen oxide anions
00
29
2
28
2.7
2.6
ELECTRON
2 5
2.4
2.3
BINDING
ENERGY
2.2
(eV1
2:7
BINDING
2 :5
IeVJ
2.1
2 .o
30
3
s
520
I
?
0”
z 10
B
00
3 ‘0
219
2il
ELECTRON
2:6
ENERGY
2.4
FIG. 2. The photoelectron spectra of OClO- and OIO- taken at the magic
angle showing the OXO(‘B,) -0X0(‘A ,) transitions. The vibrational
origins, the transition from v = 0 of the anion to v = 0 of the neutral, are
indicated by @. Each contains progressions in the symmetric stretch and in
the bending mode.
( ‘A, ) possessC 2U,symmetry. The symmetry allowed vibrational modes, those with totally symmetric transition moments, are the symmetric stretch, the bend, and even quanta
of the asymmetric stretch. The photoelectron spectrum of
OClO- extends over nearly 1 eV, while that of OIO- covers
only about half that energy range. W e expect that similar to
the halogen monoxides, the longer halogen-xygen bond
length of 010 is perturbed less by the addition of an electron
into an antibonding orbital than OClO. In the spectrum of
OClO-, two vibrational progressions are observed-the
symmetric stretch and the bend. A similar trend is observed
in the photoelectron spectrum of OIO-.
IV. DISCUSSION
A. Electron affinities
Due to the differences in vibrational frequencies between the neutrals and their respective anions, origin assignments for the halogen oxides are unambiguous. Rotational
corrections, accounting for the difference between the center
of a vibrational peak and the position of the rotationless origin, were calculated for most of the halogen oxides by simulating the rotational spectrum. Parameters in this simulation
include rotational constants, rotational temperatures, and
8015
state symmetries. The rotational origins were found to lie 34 meV lower in binding energy than the vibrational peak
center for the halogen monoxides and 5 meV lower for
OClO. Because anion and neutral geometries were not available from the literature, a rotational contour was not generated for the photoelectron spectrum of OIO-. Instead, the
photoelectron spectrum was fit using a single Gaussian with
a full width at half-maximum (FWHM) of 15 meV for each
vibrational peak.
Electron affinities for the halogen oxides are as follows:
2.272(6) eV for FO; 2.276(6) eV for ClO; 2.353(6) eV for
BrO; 2.378(6) eV for IO; 2.140(8) eV for OClO; and
2.577( 8) eV for 010. Larger uncertainties are included for
OClO and 010 due to vibrational congestion from sequence
bands in the OClO- spectrum and the lack of a rotational
correction for the OIO- spectrum.
The electron affinities of the halogen monoxides increase slightly with the larger halogens. This stabilization is
probably due to the longer bond lengths in the larger halogens, delocalizing the charge density. Although calculations19 have shown that the photodetached electron is largely localized on the oxygen, this electron is evidently quite
stabilized by the presence of the halogen since the electron
affinities of the halogen monoxides are all significantly larger than the electron affinity of oxygen ( 1.46 11 eV) .33
Similar to the halogen monoxides, the halogen dioxide
electron affinities show the ordering EA(OI0) > EA
(OClO) . Although the photoelectron spectrum of OBrO- is
not reported here, it is expected that EA(OCl0) (EA
(OBrO) <EA( 010). This would bracket the electron affinity of OBrO between 2.140 and 2.577 eV.
Earlier experimental and theoretical work has provided
upper and lower limits on the electron affinities of the halogen monoxides. The appearance potential of FO- from
CF,OF placed a lower limit of 1.4 * 0.5 eV on EA( F0).6
Another estimate of EA( FO) 22.24 eV comes from the dissociation3’ of ClO,F+FO+ ClO,. Alekseev er al7 obtained a value of 2.05 + 0.08 eV by measuring the appearance potential of FO- from F,O. Early Hartree-Fock
calculations predicted the vertical electron affinity of FO to
be 1.4 eV.42 More recent higher level calculationsa predict a
lower bound of 2.08 f 0.2 eV for the EA(FO), consistent
with our result of 2.272(6) eV.
Previous estimates of the EA( ClO) come from both experiment and theory. Vogt et al.‘* ascertained that the
EA( ClO) ) 1.6 f 0.2 eV from endoergic ion-molecule collisions of Cl- + 0, forming ClO-. Dotan ef al.” measured
the charge transfer rates of ClO- with NO, and ozone, giving 2.2 eV as an upper limit on EA(Cl0). Lee et al.” measured the photodestruction spectrum of ClO- and concluded that the EA(C10) is below 2.3 eV. Another estimate of
the electron affinity [EA(ClO) >2.35 f 0.12 eV] arises
from the appearance potential13 of ClO- from OClO, together with a calculated dissociation energy for OClO. An
early calculation17 predicted the vertical electron affinity of
Cl0 to be 2.2 -& 0.5 eV. Peterson and Woods” recently obtained 2.16 eV using MP4SDTQ dissociation energies with
CEPA-1 vibrational constants in a thermodynamic cycle.
The value found in the present experiment [ EA(C1O)
J. Chem. Phys., Vol. 96, No. 11,l June 1992
Downloaded 04 Dec 2006 to 128.138.107.158. Redistribution subject to AIP license or copyright, see http://jcp.aip.org/jcp/copyright.jsp
Gilles, Polak, and Lineberger: Photoelectron spectroscopy of halogen oxide anions
8016
= 2.276( 6) ] is certainly the most accurate measurement
and is in reasonable agreement with previous results.
Vogt et al.‘* established that the EA(Br0) > 1.5 f: 0.2
eV from the endoergic reaction of Br- + 0, + BrO-, in accord with the value reported here 2.353(6) eV. No other
values for the EA( BrO) have been reported.
Several endoergic ion-molecule reactions determined
limits on the EA (IO). Refaey and Franklin*’ examined the
endoergic reactions of I- with O,, CO, and CO, and established that the EA ( IO ) z 2.6 eV. Their earlier work26 on Iwith SO, gave a value of 2.3 eV. Vogt and Mischke27 studied
ion-molecule collision processes of I- on CO and found
EA(I0) 22.5 eV. Later Vogt et al** examined the reaction
of I- with 0, and lowered their limit to 22.1 f 0.3 eV. Our
recently reported24 value of 2.378 (6) eV is consistent with
the lower limits previously established.
Previously reported electron affinities for OClO range
from 1.3 (Ref. 35) to 22.45 eV.43 From the appearance potential of ClO; from ClO,F, Alekseev et ai.35 determined
that the electron affinity of C10,>2.45 f 0.17 eV, quite
close to the vertical electron affinity. Babcock et aZ.37examined the charge transfer reactions of OClO- with several
neutral species and bracketed the EA( OClO) between those
of SF, and NO, [ 1.5 + 0.2 eV<EA(OClO)<2.273<0.005
eV] .44*45There do not appear to be any reported values for
the EA(OI0).
B. Franck-Condon
analysis
Franck-Condon analyses were completed for each of
the halogen oxides to determine the change in geometry between the ground state neutral and the anion. Neutral frequencies, anharmonicities, and bond lengths were fixed at
known spectroscopic values for the Franck-Condon simulation. In addition, for the halogen monoxides, the spin-orbit
splittings were restricted to spectroscopic values. Variable
parameters were anion vibrational frequencies, anion vibrational temperatures, and normal coordinate displacements.
Hot bands, which arise from the transitions of vibrationally
excited anions to neutral vibrational states (e.g., transitions
from u = 1 of the anion to u = 0 of the neutral) provided
anion vibrational frequencies. The intensity of the hot bands
gave a measure of the anion vibrational temperatures. These
were all ~450 K, with the exception of FO-, at ~800 K.
Anion anharmonicities were not determined in this experiment. Although they were constrained to the values given in
Table I they do not have a significant effect on the reported
fits. For the halogen monoxides, the anion and neutral states
were treated as Morse oscillators and Franck-Condon factors were calculated by numerically integrating the products
of the Laguerre wave functions. To simulate the FranckCondon envelope for the halogen monoxides, the equilibrium bond length of the anion was varied until the best fit was
obtained. The anion bond length was chosen to be longer
than the neutral bond length because the photodetached
electron originates in an antibonding P orbital.17 Results of
the Frank-Condon simulations and the neutral values used
for the simulations are given in Table I. Both spin-orbit
states exhibit similar intensity profiles in FO- and ClO(Fig. 1 ), showing that the neutral bond lengths are nearly
the same. For BrO- and IO-, the vibrational progression of
the *II,,,+ ‘X+ transition is longer than that of the
‘X+ transition. This indicates that the *II,,* state
*I-L,*+
has a bond length closer to that of the anion than that of the
% 3,2 state.
Vibrational frequencies of the halogen monoxide anions
observed were 73%-85% of their respective neutral values,
in accord with the idea that the photodetached electron
comes from an antibonding orbital. The change in bond
length between fluorine monoxide anion and neutral flu-
TABLE I. Spectroscopic constants for the halogen monoxides.
State
FO
FOCl0
c10BrO
BlO
BrOIO
IO
IO-
r, (A,
?
‘P+
-1
*x+
?n*,*
zh
‘Et
*n1/2
%*
‘Hf
1.354 12”
1.516(6)b
1.569 6d
1.673(8)b
1.733(9)b
1.720 72h
1.814(9)”
1.887( 10)’
1.867 7’
1.929( 10)’
0, (cm-‘)
1052.99”
769(25)b
853.72’
665(25)”
713(25)b
727.05h
575(25)b
658(25)’
681.6k
581(25)’
w,x, (cm-‘)
Spin orbit (cm- ’)
9.90’
7.95’
5.58’
3.36’
4.74g
4.74h
4.74’
4.378
4.37k
4.378
- 193.8W
- 320.31’
- 967.98’
- 2091(40)’
“Reference 36.
bValues obtained from this experiment.
‘Reference 9.
dReference 67.
‘Reference 38.
‘Reference 19.
Unstrained to this value for the Franck-Condon optimization.
“Reference 40.
‘Reference 39.
‘From this experiment and Ref. 24.
‘Reference 41.
J. Chem. Phys., Vol. 96, No. 11,i June 1992
Downloaded 04 Dec 2006 to 128.138.107.158. Redistribution subject to AIP license or copyright, see http://jcp.aip.org/jcp/copyright.jsp
Gilles, Polak, and Lineberger: Photoelectron spectroscopy of halogen oxide anions
orine monoxide Ar, = r,(FO-)
- r,(FO) of 0.162 A is
nearly three times larger than Ar, for iodine monoxide
(0.061 A). This shows clearly that the longer bond length
(due to the larger halogen) is less affected by the addition of
an electron into an antibonding orbital.
The anion molecular constants determined from this experiment are in good agreement with recent quantum chemical calculations. H&and* predicted a bond length 1.53 A,
vibrational frequency w, (767 cm-‘), and anharmonicity
w,x, (7.95 cm- ’) for FO- based ab initio calculations. Because the observed experimental frequency of FO[ 769 ( 25 ) cm- ’] was closest to the value calculated by Haaland,* the anharmonicity was fixed to this value for the
Franck-Condon analysis. The anion bond length calculated
by Haaland* is slightly larger than the bond length from our
Franck-Condon simulation [ 1.5 16( 6) A]. The calculated
FO- bond length ( 1.513 A) of Peterson’ agrees well with
the results of our Franck-Condon simulation.
For ClO-, the ‘X+ (ground state anion) bond length of
1.673( 8) A concurs with the MP4SDQ value ( 1.6780 A) of
Peterson and Woods. I9 Their CI-SD( s) vibrational frequency of w, = 668 cm-’ (w,x, = 3.36 cm-‘) also agrees with
our observed value of 665(25) cm-’ for ClO-. This vibrational frequency is also consistent with the second-order
Moller-Plesset value (685.4 cm-‘).”
For BrO-, there are no gas phase experimental or theoretical results with which to compare the anion bond length
[ 1.814(9) A] and vibrational frequency [575(25) cm-‘].
A solution phase2’*46frequency of 620 cm- ’has been reported for BrO-. Wren ef a1.23 observed (solution phase)
stretching frequencies of 430( 2) and 560( 2) cm-’ and after
some discussion assigned these to IO- and I,OH-, respectively. Only the later value (560 cm-‘) is consistent with
our observed24 vibrational frequency for IO- of 581(25)
cm-‘, and we conclude that their assignment was incorrect.
The photoelectron spectra of the halogen dioxides exhibit two active vibrational modes. The more intense progression, the symmetric stretch, was modeled as a Morse
oscillator. The bending mode was modeled as an independent harmonic oscillator. Franck-Condon factors were calculated by numerically integrating the Laguerre wave functions (Morse oscillator) and using the recursion formula
method of Hutchisson for the harmonic oscillators.47 The
Franck-Condon simulation was perforemd by varying the
normal coordinate displacements and optimizing the neutral
and anion vibrational frequencies. The anion symmetric
stretching frequency was obtained from the difference between the positions of the origin and the 1: hot band. Anion
and neutral spectroscopic parameters used to simulate the
spectra of the halogen dioxides are given in Table II.
The method used to make ClO, produces only the OClO
isomer. Conclusive proof that the spectrum is that of OClOand not ClOO- comes by comparing the well-known frequencies of OClO with those observed in the photoelectron
spectrum. The ground state *B’ frequencies are known to be
942 cm- ’ (symmetric stretch), 445 cm- ’ (bend), and 1110
cm- ’ (asymmetric stretch). 48*49The neutral frequencies observed in the photoelectron spectrum were 945(25) and
445 (25 ) cm- ‘. The anion bending frequency was not ob-
TABLE II. Vibrational frequencies (cm-‘)
OIO-, and 010.
Species
OClO
oao010
010-
8017
observed for OClO-, OClO,
Symmetric stretch
Bend
945(25)”
774(25)
765(25)
675(25)
445(25)”
192(35)
“More accurate frequencies are reported (942 and 445 cm- ’) in Refs. 48
and 49.
served in the photoelectron spectrum and was set at the calculated value5’ 378 cm-‘. The Franck-Condon simulation
does not account well for the intensity of the bending mode.
While a single quantum of the bending mode is barely visible,
at higher vibrational quanta, the bend intensity increases to
nearly l/4 of the stretch intensity. Normal coordinate displacements found in the Franck-Condon analysis are 0.478
amu”* A for the symmetric stretch and 0.139 amu”* A for
the bending mode. These were used with the neutral geometry r(Cl-0) = 1.4698 A, L(OCl0) = 117.4”, and force field
of Miyazaki et al.” to determine the anion geometry. Because only a single hot band is observed, only the magnitude
and not the direction of the normal coordinate displacement
can be determined directly from the simulation. With two
normal coordinate displacements, there are four possible
sign permutations. Two of these geometries resulted in longer Cl-O bond lengths: (a) r(Cl-0) = 1.563(2) A, LOClO
= 112(2)O and (b) r(C1-0) = 1.566(2) A, LOClO
= 117(2)“. These error bars reflect maximum assumed errors of 20% in the normal coordinate displacements found
from the Franck-Condon analysis and do not consider the
error in the force field or neutral geometry used. Preliminary
results of configuration interaction (CI) calculations by Peterson and Werner” predict a Cl-O bond length of 1.5728 A
and a bond angle of 113.9” for the anion. This calculation
suggests that the first geometry (a) is the preferred anion
geometry.
Since the mechanism of OIO- production was uncertain, it was necessary to ascertain that OIO-, and not IOO-,
was the isomer which we observed. Reported frequencies for
the O-O stretch in FOO and Cl00 are 1486.93 and 1443
cm-‘, respectively. 52 Extrapolating from these frequencies,
a frequency of z 1400 cm- ’ would be expected for the O-O
stretch in 100, which does not agree with any of the observed frequencies. Wren et a1.23assigned a frequency of 685
cm-’ to the symmetric stretch of OIO- in the liquid phase.
This assignment is compatible with our observed frequency
of 675(25) cm-’ for the symmetric stretch of the anion.
The photoelectron spectrum of OIO- shown in Fig. 2
exhibits an intense progression in the symmetric stretch and
a weaker progression in the bending mode. The frequency of
the bending mode is significantly lower than that for OClO,
and is seen only as a shoulder on the symmetric stretch peak,
making it difficult to determine the exact peak location.
Neutral vibrational frequencies for 010 were not found in
the literature, but could be determined from the photoelec-
J. Chem. Phys.,
Vol. 96,
No. license
11,i June
Downloaded 04 Dec 2006 to 128.138.107.158. Redistribution
subject
to AIP
or 1992
copyright, see http://jcp.aip.org/jcp/copyright.jsp
Gilles, Polak, and Lineberger: Photoelectron spectroscopy of halogen oxide anions
8018
tron spectrum to be 192 (35) cm-’ for the bend and 765 (25)
cm-’ for the symmetric stretch. The symmetric stretching
frequency for OIO- was observed at 675( 25) cm- ‘. Normal coordinate displacements obtained from the FranckCondon simulation were 0.176 amu”’ A for the bend and
0.295 amu1’2 A for the symmetric stretch. Without a neutral
geometry and force constants, we were unable to determine
the anion geometry.
TABLE
mol) .
C. Thermochemistry
“A,H(A),,,
Dissociation energies are not directly measured in photoelectron spectroscopy. Instead, the following thermodynamic relation was used to derive the halogen monoxide anion bond dissociation energies given in Table III:
D,(X--0)
= D,(XO)
- EA(X)
+ EA(X0).
This experiment combined with the well-known atomic halogen electron affinities very accurately measures the difference between the bond energy of the anion and neutral species. Electron affinities for the halogens were taken from
Hotop and Lineberger. 53 Since several different dissociation
energies have been reported for F0,54,55 C10,5”58 BIO,~~*~~
the anion binding energies given in Table III
and I0,56*60*61
rely heavily on the choice of dissociation energy for the neutral species. If the neutral dissociation energies are improved, it is trivial to make the corresponding corrections for
the anion dissociation energies.
As seen in Table III, all of the atomic monoxides have
smaller dissociation energies than their corresponding neutrals, as would be anticipated from adding an electron into
an antibonding orbital. The weakening of the anion bond
energies relative to the neutral bond energies decreases from
IO- to FO-. The bond energy of FO- is half that of FO,
while for IO-, the bond energy is only 36% less than the
neutral IO.
Anion heats of formation given in Table IV are calculated by means of the following relationship:
A&4
-) = A+(A)
- EA(A).
Anion heats of formation were calculated at 298 K. We employ the “ion convention,” which is equivalent to taking the
integrated heat capacity of the electron to be zero.62 In addition, we assume the integrated heat capacities of the ions and
neutrals are the same. The error from this assumption is
normally less than 0.5 kcal mol-’ because of the similarity
IV. Halogen oxide anion and neutral heats of formation (kcal/
Species
A/H(A - 1zps
Wf(A
)m
FO
Cl0
BrO
IO
OClO
-
26.0 +
24.19 f
30.05 f
41.84 f
25.00 f
3.5”
0.5b
0.7”
4.6=
lSb
26.3 f 3.5
28.29 -& 0.5
24.20 & 0.7
13.00*4.6
24.35 f 1.5
taken from Ref. 68; the error bars given are the same as those on
&(X0)
from Table III.
bA,fW )m taken from Ref. 69.
between the vibrational frequencies and structures of the ion
and neutral species. This assumption is acceptable here since
the error bars on the heats of formation for the neutrals exceed this value.
D. Angular distributions
of photoelectrons
As a consequence of the linear polarization of the laser
light, the differential cross section in the electric dipole approximation is given by63
g=
(o,,/‘4rr)[l
+ (B/2)(3cos20-
l)],
where a, is the total photodetachment cross section, 8 is the
angle between the electric vector of the laser light and the
photoelectron collection direction, and /3 is the asymmetry
parameter ( - 1 <PC + 2). Spectra taken at the “magic”
angle (3 cos28 - 1 = 0) have a photodetachment signal
proportional to the total photodetachment cross section.
Details on the measurements of photoelectron angular distributions are given elsewhere.@
A general description of angular distributors of photoelectrons is given by Cooper and Zare63*65and in a recent
paper by Hanstorp et a1.66Basically, the angular distributions of photoelectrons are functions of both the orbital from
which the electron is detached and the kinetic energy of the
photodetached electron. To determine whether asymmetry
parameters for photodetachment to different spin-orbit
states would be the same, we measured the asymmetry parameters for transitions to the 2113,2and 211,,2 states of the
halogen monoxides. These are given in Table V. The differ-
TABLE V. Asymmetry parameters (/3) measured in this experiment.
TABLE III. Electron afhnities (EA) and dissociation energies of the halogen monoxides (eV). Electron athnities and the dissociation energy differences [&(X0)
- &(X0-)]
are from this work.
Species
FO
Cl0
BrO
IO
EA
2.272(6)
2.276(6)
2.353(6)
2.378(6)
*Reference 54.
‘Reference 57.
“Reference 56.
dDerived using D,( X0).
D,(XO,
- D,(xo-,
1.127(7)
1.341(7)
1.012(7)
0.681(7)
D&X0)
2.25( 15)”
2.7504(4)”
2.39(3)’
1.8(2)”
Do(XO-)d
1.12( 15)
1.409(S)
1.38(3)
1.1(2)
Species
FO *bz
FO 2~m
Cl0 *IIs,*
Cl0 Q,,*
BIQ %,z
BOO *~I,z
10 2~s/2
10 zh2
OClO
010
Peak position
eKE (eV)
0.997
0.973
1.253
1.212
1.174
1.055
1.152
0.893
1.137
0.857
B
-
0.77(
0.75(
0.72(
0.70(
0.87(
0.80(
0.83(
0.80(
0.49(
0.71(
15)
15)
10)
10)
10)
10)
10)
10)
10)
10)
J. Chem. Phys., Vol. 96, No. 11,l June 1992
Downloaded 04 Dec 2006 to 128.138.107.158. Redistribution subject to AIP license or copyright, see http://jcp.aip.org/jcp/copyright.jsp
Gilles, Polak, and Lineberger: Photoelectron spectroscopy of halogen oxide anions
ences we observe in fi for photodetachment to the two spinorbit states are small enough that they could arise from the
dissimilar kinetic energies of the photodetached electrons,
making it difficult to determine if there are spin-orbit effects
in the asymmetry parameters. In addition, the asymmetry
parameters measured for several vibrational levels to each of
the ‘lIJ,z and 2111,2neutral states of Cl0 showed similar
results-asymmetry
parameters for transitions to each of
the spin states are the same within our error bars; those to
different vibrational levels were also nearly the same within
our experimental error. The asymmetry parameters measured for the halogen monoxides all approach the minimum
value possible ( - 1) for photodetachment from a pure p
orbital. This observation is consistent with the calculations
of Peterson and Woods’9 and O’Hare and Wah117*43who
have predicted that this electron is quite localized on the
oxygen.
The asymmetry parameters measured for the halogen
dioxides, given in Table V, are also negative, but smaller in
magnitude than those of the halogen oxides. The photodetached electron comes from an out-of-plane molecular orbital where the halogen p orbital has a phase opposite of the
oxygen p orbitals. Only a few reports of measurements and
theoretical work have been made on asymmetry parameters
of photodetached electrons from anions; we speculate that
the smaller negative asymmetry parameter for the halogen
dioxides arises from the delocalization of the extra electron.
V. CONCLUSION
The electron affinities of the halogen monoxides are reported to be 2.272(6) eV (FO, 2.276(6) eV (CIO),
2.353(6) eV (BrO), and 2.378(6) eV (IO). The halogen
dioxides electron affinities are 2.140(8) eV for OClO and
2.577( 8) eV for 010. Using the halogen monoxide electron
affinities with thermodynamic relationships, we have determined the difference in the dissociation energies of the anions and their corresponding neutrals. Frequencies for the
halogen monoxide anions are 769(25) cm-’ for FO-,
665(25) cm-’ for CIO-, 575(25) cm-’ BrO-, and
58 l(25) cm-’ IO-. The Franck-Condon analyses yielded
anion bond lengths of 1.516(6) A (FO-),
1.673(8) A
(ClO-),
1.814(9) A (BrO-), and 1.929(10) 8, (IO-).
The symmetric stretching frequency of OCIO- is 774( 25)
cm-‘. The recommended anion geometry for OCIO- is
r( Cl-O) = 1.563 8, and L( OClO) = 112”. Neutral 010 frequencies of 765(25) cm-’ (symmetric stretch) and
192 (35) cm- ’ (bend) are observed. A symmetric stretching
frequency of 675 (25) cm- ’ was determined for OIO-. The
electron angular distributions for photodetachment to each
of the spin-orbit states of the halogen monoxides were found
to be quite similar.
ACKNOWLEDGMENTS
Dr. J. Ho and Robert Gunion have provided experimental assistance and expertise. We wish to acknowledge the
generosity of Dr. K. Peterson for providing us with the results of calculations on FO-, ClO,, and CIO; and numerous discussions. Dr. E. Riihl also provided scientific advice
8019
and discussions. This work was supported by National
Science Foundation Grant Nos. CHE88- 19444 and PHY9012244.
‘M. J. Molina and F. S. Rowland, Nature 249, 810 (1977).
*F. S. Rowland and M. J. Molina, Rev. Geophys. Space. Phys. 13, 1
(1975).
3R. P. Turco, J. Geophys. Res. 82,3585 (1977).
4W. L. Chameides and D. D. Davis, J. Geophys. Res. 85,7383 ( 1980).
‘J. A. Kerr, in Frontiers of Free Radical Chemistry, edited by W. A. Pryor
(Academic, New York, 1980), pp. 171-193.
6J. C. J. Thynne and K. A. G. MacNeil, Int. J. Mass Spectrom. Ion Phys. 5,
95 (1970).
‘V. I. Alekseev, V. M. Volkov, L. I. Fedorova, and A. V. Baluev, Izv. Akad.
Nauk SSSR Ser. Khim 6, 1302 ( 1984) [English translation: Aka. Nauk
SSSR Bull. Chem. Sci. 33, 1195 ( 1984) 1.
sP. Haaland, Chem. Phys. Lett. 176,287 (1991).
9K. A. Peterson, Ph.D. thesis, University of Wisconsin-Madison, 1990.
“L. C. Lee, G. P. Smith, J. T. Moseley, P. C. Cosby, and J. A. Guest, J.
Chem. Phys. 70,3237 ( 1979).
“I. Dotan, D. L. Albritton, F. C. Fehsenfeld, G. E. Streit, and E. E. Ferguson, J. Chem. Phys. 68,5414 (1978).
‘*D. Vogt, W. Dreves, and J. Mischke, Int. J. Mass Spectrosc. Ion Phys. 24,
285 (1977).
13A. V. Baluev, Z. K. Nikitina, L. I. Fedorova, and V. Y. Rosolovskii, Izv.
Akad. Nauk SSSR Ser. Khim. 9, 1963 ( 1980) [English translation: Aka.
Nauk SSSR Bull. Chem. Sci. 29, 1385 ( 1980) 1.
I%‘. H. Dibeler, R. M. Reese, and D. E. Mann, J. Chem. Phys. 27, 176
(1957).
“G. Kujumzelis, Phys. Z. 39, 665 ( 1938).
16E.L. Wagner, J. Chem. Phys. 37,751 ( 1962).
“P. A. G. O’Hare and A. C. Wahl, J. Chem. Phys. 54,377O ( 1971).
“V. L. Pershin and A. I. Boldyrev, Izv. Akad. Nauk SSSR 4,796 (1987)
[English translation: Aka. Nauk SSSR Bull. Chem. Sci. 36,722 ( 1987) 1.
19K. A. Peterson and R. C. Woods, J. Chem. Phys. 92,7412 ( 1990).
“G. E. Streit, J. Phys. Chem. 86, 2321 (1982).
*‘B. Sombret and F. Wallart. C. R. Acad. Sci. Paris Ser. B 277,663 ( 1973).
**A. Henglein and G. A. Muccini, J. Chem. Phys. 31, 1426 (1959).
23J.C. Wren, J. Paquette, S. Sunder, and B. L. Ford, Can. J. Chem. 64,2284
(1986).
24M. K. Gilles, M. L. Polak, and W. C. Lineberger, J. Chem. Phys. 95,4723
(1991).
*‘K. M. A. Rafaev and J. L. Franklin. Int. J. Mass Spectrom. Ion Phys. 20,
19 (1976).
26K. M. A. Rafaey and J. L. Franklin, J. Chem. Phys. 65, 1994 ( 1976).
*‘D. Vogt and J. Mischke, Phys. Lett. A 60, 19 ( 1977).
“M. Meinke, C. Schmale, E. Riihl, and E. Illenberger (unpublished resuits).
29J. R. Byberg, Chem. Phys. Lett. 81, 156 (1981).
“‘D G Leopold, K. K. Murray, A. E. Stevens Miller, and W. C. Linebeige;, J. Chem. Phys. 83,4849 (1985).
3’K. M. Ervin, J. Ho, and W. C. Lineberger, J. Chem. Phys. 91, 5974
(1989).
3ZE.E. Ferguson, F. C. Fehsenfeld, and A. L. Schmeltkopf, Adv. At. Mol.
Phys. 5, i (1969).
“D. M. Neumark, K. R. Lykke, T. Andersen, and W. C. Lineberger, Phys.
Rev. A32, 1890 (1985).
MC. E. Moore, AtomicEnergy Levels (Natl. Bur. Stand., U.S. Government
Printing Office, Washington, D. C., 1958), Circ. 467.
“V. I. Alekseev, L. I. Fedorova, and A. V. Baluev, Izv. Aka. Nauk SSSR 5,
1084 ( 1983) [English translation: Aka. Nauk SSSR Bull. Chem. Sci. 32,
980 (1983)].
“P. D. Hammer, A. Sinha, J. B. Burkholder, and C. J. Howard, J. Mol.
Spectrosc. 129,99 ( 1988).
37L. M. Babcock, T. Pentecost, and W. H. Koppenol, J. Phys. Chem. 93,
8126 (1989).
38J.B. Burkholder, P. D. Hammer, C. J. Howard, A. G. Maki, G. Thompson, and C. Chackerian, Jr., J. Mol. Spectrosc. 124, 139 ( 1987).
39A. R. W. McKellar, J. Mol. Spectrosc. 86,43 ( 1981).
4oJ. E. Butler, K. Kawaguchi, and E. Hirota, J. Mol. Spectrosc. 104, 372
(1984).
41J.P. Bekooy, W. L. Meerts, and A. Dymanus, J. Mol. Spectrosc. 102,320
(1983).
J. Chem. Phys., Vol. 96, No. 11,l June 1992
Downloaded 04 Dec 2006 to 128.138.107.158. Redistribution subject to AIP license or copyright, see http://jcp.aip.org/jcp/copyright.jsp
8020
Gilles, Polak, and Lineberger: Photoelectron spectroscopy of halogen oxide anions
‘*P. A. G. O’Hare and A. C. Wahl, J. Chem. Phys. 53,2469 (1970).
43D. Wecker, A. A. Christodoulides, and R. N. Schindler, Int. J. Mass Spectrosc. Ion. Phys. 38,391 (1981).
“A. A. Viggiano, T. M. Miller, A. E. Stevens Miller, R. A. Morris, J. M.
Van Doren, and J. F. Paulson, Int. J. Mass Spectrom. Ion Proc. 109,327
(1991).
45K. M. Ervin, J. Ho, and W. C. Lineberger, J. Phys. Chem. 92, 5405
(1988).
46J.C. Evans and G. Y-S. Lo, Inorg. Chem. 6, 1483 (1967).
“E. Hutch&on, Phys. Rev. 36,410 (1930); 37,45 (1931).
48M. G. Krishna Pillai and R. F. Curl, Jr., J. Chem. Phys. 37,292l ( 1962).
@R. F. Curl, Jr., K. Abe, J. Bissinger, C. Bennett, and F. K. Tittel, J. Mol.
Spectrosc. 48,72 ( 1973).
“‘K. A. Peterson and H.-J. Werner (to be published).
“K Miyaxaki, M. Tanoura, K. Tanaka, and T. Tanaka, J. Mol. Spectrosc.
U-6,435 (1986).
s2M. A. Jacox, J. Phys. Chem. Ref. Data 17, 343 ( 1988).
5sH. Hotop and W. C. Lineberger, J. Phys. Chem. Ref. Data 14,73 1 ( 1985).
“M. A. A. Clyne and R. T. Watson, Chem. Phys. Lett. 12,344 (1971).
ssD. H. Levy, J. Chem. Phys. 56, 1415 (1972).
56R. A. Durie and D. A. Ramsay, Can. J. Phys. 36,35 (1958).
“5. A. Coxon and D. A. Ramsay, Can. J. Phys. 54, 1034 (1976).
58G. Porter, Discuss. Faraday Sot. 9,60 ( 1950).
s9E. H. Coleman and A. G. Gaydon, Discuss. Faraday Sot. 2,166 ( 1947).
60E. H. Coleman, A. G. Gaydon, and V. M. Vaidya, Nature 162, 108
(1948).
‘IV. M. Trivedi and V. B. Gohel, J. Phys. B 5, L38 ( 1972).
62S.G. Liis, J. E. Bartmess, J. F. Liebman, J. L. Holms, R. D. Levin, and W.
G. Mallard, J. Phys. Chem. Ref. Data 17, Suppl. 1 ( 1988).
63J Cooper and R. N. Zare, J. Chem. Phys. 48,942 (1968); (errata) ibid.
49,4252 (1968).
MM. K. Gilles, K. M. Ervin, J. Ho, and W. C. Lineberger, J. Phys. Chem.
96, 1130 (1992).
65J. Cooper and R. N. Zare, Lectures Theor. Phys. llc, 323 ( 1969).
&D. Hanstorp, C. Bengtsson, and D. J. Larson, Phys. Rev. A 40, 670
(1989).
“F. K. Kakar, E. A. Cohen, and M. J. Geller, J. Mol. Spectrosc. 70, 243
(1978).
68D. D. Wagman, W. H. Evans, V. B. Parker, R. H. Schumm, W. Halow, S.
M. Bailey, K. L. Chumey, and R. L. Nuttall, J. Phys. Chem. Ref. Data 11,
Suppl. 2 (1982).
-M. W. Chase, Jr., J. L. Cumutt, J. R. Downey, Jr., R. A. McDonald, N. N.
Syverud, and E. A. Valenxuela, J. Phys. Chem. Ref. Data 11,695 ( 1982).
J. Chem. Phys., Vol. 96, No. 11,i June 1992
Downloaded 04 Dec 2006 to 128.138.107.158. Redistribution subject to AIP license or copyright, see http://jcp.aip.org/jcp/copyright.jsp