Hydrogen Bonding and Proton Transfer
Melvin D. J o e s t e n
Vanderbilt University, Nashville, TN 37235
Historical Background
The first definitive paper on hydrogen bonding was pubthe
lished in 1920 bv Latimer and Rodehush (.1.) .Thevannlied
.
concept of hydrugen bonding to association of water nldecules
and acknowledged Hugginsd the same lahuratory inr the ide3
of hydrogen bonding. 1\11three were working in the laboratory
of C . N.I.c.w~swhu stated in his 19'23 clasilc hook on \.alenuv
hydrogen bonds include humo-intermolecular isclf-asswiattonj and hetero-intermolectl1:ir hsdruaen bonds. 'l'uhle ? rlves
a range of hydrogen hond strengths with an arbitrary c l k fication as weak, moderate, s t r o w , and verv strona. Most of
the emphasis in the literature has been on "moderate" or
"normal" hydrogen bonds. However, strong and very strong
hydrogen bonds are important in studies of the effect of hydrogen bonding on proton transfer.
.. . It seems to me that the most important addition to my theory of
Spectroscopic Studies of Hydrogen Bonding
A varietv of snectrosconic and diffraction methods have
been applied to the study of hydrogen bonding. Tables 3 and
4 list examnles with kev references.
Time scale differences (Table 5) must be kept in mind when
analyzing spectral data for hydrogen-bonded systems. For
example, NMR gives a time-average spectrum of free and
..
valence lies in the sueeestion
of what has became known as the hv....
d n w n hmd. The idea uas tint surgeit4 try L h M. L. Huzgin,, and
was alno nd\anrrd try Latimer and Hodt4)ush,whu ihuwru the grcnt
value of the idea in their paper.. .
A review by Huggins in 1971 describes his interests in hydroeen bonding over a 50-vear neriod (3).Thousands of naners
have been pubkshed on various aspect; of hydrogen boniing
since the 1920 naner hv Latimer and Rodebush. Fortunatelv.
a number of hboks (419) have been published which review
hydrogen bonding, and several sections of the most recent
series edited by Schuster, Zundel, and Sandorfy (9) will be
cited in this Daner.
T h e simultaneous development of hydrogen bonding concepts and Bronsted acid-base theory (10)led to early recognition of the influence of hydrogen bonding on proton transfer
reactions. The present paper will focus on spectroscopic and
diffraction studies which nrovide information about the effect
of hydrogen bonding od the extent of proton transfer in
A-H.. .B = -A,. .+H-B. This will include a consideration
of (1) the strength of the hydrogen hond and (2) the molecular
environment. A brief review of tvpes of hvdroaen bonds will
be followed by a discussion of the-applicati& orspectroscopic
and diffraction methods to studies of moderate and strong
hydrogen bonds, symmetric and asymmetric hydrogen bonds,
solvent effects, and hydrated proton species.
The hydrogen bond can be represented as A-H.. .B where
B is any a or a electron donor site (Lewis base) and A is an
atom more electroneeative than hvdroeen.
,
.. Hvdroeen
,
.. bonds
can he either intriimdtwllar ur intrrmulecular. 'Table 1 lists
rxi~mplesof differmt typt.s of hydrogt!n ln)nd.<.Intermolecular
-
A
Based on a talk given at the Symposium on Acid-Base Chemistry
in Celebration of the Centennial of Bronsted's Birth, 178th ACS National
Meeting, Washington, D.C., September, 1979.
362
Journal of Chemical Education
Table 1. Types of Hydrogen Bonds
I. lntermOlecular
A. Self-Association
cyclic (RCOOH)2,3
cyclic or linear (ROH)2e,4
8. Hetwo-asso~iation
CeHsOH. . .N(C2Hd3
(C2H&0...HCI
11. Intramolecular
2-chlorophenol
Pdiketones
pyruvic acid
A-H
Table 2. Strengths of Hydrogen Bonds
-AH kcal mole-'
+B=A-H.. . B = I A . . .H...B~=A-H+. . .B-
Cla~~ificatlon
Examples
.
.$)
Weak <3
CeHsOH
Moderate 3-7
Strong 7-10
Very Strong >10
pF-CsH40H. S E t 2
CeHrOH- -NEB
CI- - -HCI
1.2
3.6
8.8
12.8
hydrogen-bonded HA while IR gives peaks for both free and
hydrogen-bonded HA. The properties of hydrogen-bonded
aggregates which are important for examination by spectroscopic and diffraction methods are (1) the H-A bond length
increases in the formation of A-H.. .B, (2) electron densities
of H in A-H.. .B decrease, (3) the polarity of HA and B are
increased in A-H.. .B, and (4) HA and B come much closer
in forming A-H . .B than the sum of van der Waals radii
allow.
Far-Infrared Spectroscopy
Direct spectroscopic d~servationof the hydrngen hond
H . .R in AH. . .Bis possible in the far-infrard rrgion 11WJ00
cm-'). Hydrogen honding modes whlch absorb in this reginn
include the hydrogen l r ~ n dstretching mwlt: u, and the hy.
drown
hond bendinr modes u 3 and u., tTal~le6).Advances in
"
instrumentation and experimental techniques have resulted
in an increased interest in the use of the far-infrared region
to study hydrogen bonding (21).
Infrared Spectroscopy
The major spectral changes that occur when weak and
normal hydrogen bonds form are (1) the A-H bending frequency increases, (2) the A-H stretching frequency decreases, and (3) the band width and band intensity of the
A-H stretching frequency increases. The most widely used
IR probe is (Z),the decrease in us when A-H.. .B forms.
Infrared spectra of strong O-H . .O hydrogen bonds are
of two types: type I consists of three main bands (A, B, C) in
the 1800-3000 cm-I region; type I1 has no bands above 1800
cm-I and a strong, broad band below 1500 cm-1. Type I is
believed to be characteristic of an asymmetric hydrogen bond,
and type I1 has been assigned to systems containing a symmetric hydrogen bond. Novak (33) uses IR spectra of acetic
acid, potassium hydrogen diacetate (type I), and sodium
hydrogen diacetate (type 11) to illustrate changes in the IR
spectrum of the O-H..O hydrogen bond as it becomes
progressively stronger.
Near-Infrared Spectroscopy
Near-infrared spectroscopy (0.7-3.0 p ) is used to examine
the influence of hydrogen bonding on the anharmonicity of
the A-H vibration. Overtones of O-H, N-H, and C-H
stretching vibrations occur in this region with little interferTable 3. Spectroscopic Studies of Hydrogen Bondlng
Reference
Method
Magnetic Resonance
13C, "0.
NMR Chemical Shifts 'H.
Nuclear Magnetic Relaxation
Nuclear QuadrupoieCoupling H , "N.
ESR, free radicals
ENDOR, solids
Wide-line NMR. solids
Far IR: W . .B stretch, bend
iR and RAMAN: A-H stretch, bend
UV-Visible: red a blue shills of
chromopbes
Ion C~clotronResonance: gas phase
Neutron Scanering: solids
lSN.
7, 8, 12-15
15. 16
17, 18
18-20
18
6. 18
21
7.8.22-24
"N, "F
"0.
8. 26, 27
28, 29
30
ence from other bands. However, the intensities of overtone
bands are quite low, with molar absorptivities of less than
2.
Electronic Absorption Spectroscopy
Ultraviolet and visible spectroscopy may be used to study
hydrogen bonding if the chromopboric portion of HA or B is
perturbed by the hydrogen bond. Generally, hydrogen
bonding causes a red (hathochromic) shift for chromophores
acting as proton donors and a hlue (hypsochromic) shift for
chromophores acting as proton acceptors. Kamlet and Taft
(27) have made extensive use of chromophore electronic shifts
in studies of solvated species and proton affinity of solvents.
Vinogradov and Linnell(7) have described the application
of electronic spectroscopic techniques to the study of the
equilibrium between hydrogen-bonded complexes and proton-transfer comolexes B.. .H-A = B-H+. . .A-. For example, the ahsorpiitm h d s inr undissncia~edand diswriakd
forms of P-nitro~henolat :317.9 and 4W.5 nm have been used
in spectrbphotometric titrations of p-nitrophenol with amines
to estimate the importance of solvent oolaritv. in nroton
.
transfer reactions. fn all of these studiesthe assumption is
made that the spectrum of the Droton transfer complex is the
same as that ofihep-nitrophehylate anion. ~ e n c e ultrilvio;
Iet-visiblt: spectral studies cannot distinauish htltween an ion
pair and a hydrogen-bonded ion pair.
NMR Spectroscopy
Proton NMR has been important to studies of hydrogen
bondine since the earlv. davs of NMR (341.
, , Generallv.
, ". t,he
~~-~
proton resonance signal of HA is shifted to lower field in
A-H. ..Band only one peak intermediate between those for
pure HA and A-H.. .B is observed. During the past fifteen
vears. 'H. 2H. I3C.. '70. , 14N. 15N. and 19F chemical shifts in
H-A. . .B syskms have been examined for self-associated alcohols and carboxylic acids as well as a variety of hetero-associated adducts. Most of these studies (12) have used NMR
data to (1) determine equilibrium constants, or (2) check the
validity of correlations of NMR chemical shifts with pKa, IR
frequency shifts, AH, and suhstituent constants.
-
Correlations
The sensitivity of the A-H stretching 1rt.quenry to hgdrogen bond formation led to early investigations of whrther
1 ,
shifts wrn:lat.ed with hydrogen bond
the 1 ~ ~ frequency
strengths. In 1937 Hadger and Hauer I351 propoced that a
linear rehtilmshi~exis16betweetl theenthalova>fthe hvdrogen bond, AH,,&d AVAH.Since that time, L u m b e r bf reTable 5. Tlme Scales lor SpectroscopicTechnlques (32)
UV-Visible
IR-Raman
Microwave
sec
lo-'' sec
ESR
NMR
10-4 to
t0-l I0
Table 6. A-H.
Region
3500-2500 em-'
Table 4. X-Rav and Neutron Dillractlon Methods ( 6 . 311
1.
Location of hydrogen atoms in A-H.
.
A,. .H. .B
2. Linearily of hydrogen bonds
. .B.
3. Isotope effects on hydrogen bond l e n g h
4. Hydrogen bands in hydrates
5. H1H+3Lt
to lo-''
':.
A-H..
t
.B
SBC
see
see
. .B Vibrations
Description
Y,
A-H
Stretch
In-plane bend
1700-1000 cm-'
A-t?
YO
A-H
900-300 cm-'
i
A-H. . .B
Ut
A-H
A-H.'.-B-
Y,
Out-of-plane
bend'
H.. -0 Stretch
un,v7
H-. .B Bend
250-100
cm-'
Below 200 cm-'
A-H..
..B
T
.B
1
i.Indicates vibrallonal movement pependicular to the A-H
Volume 59
Number 5
.'6 plane.
May 1982
363
search groups have ohtained AH and AVAHdata to test the
Badger-Bauer proposal (36-38). Data for the reference acids
phenol (39), substituted phenols (39),2,2,2-trifluoroethanol
(40), 1,1,1,3,3,3-hexafluoro-2-propanol
(41). t-butanol (42),
oerfluoro-t-hutanol(43).and ."
ovrrole (44) with hases of different structural typesgave a series of essentially parallel
AH-Av lines. Equations for the least-squares lines (based on
7 to 16 data points) were considered reliable for predicting
hydrogen b n d enereies to within 0.2 kcallmole for moderate
hidrogen honds of these acids with nitrogen and oxygen donor
bases.
Arnett and coworkers (38) have questioned the reliability
of AH-Av equations for making quantitative predictions.
l'hev used the AH-Au eouatiou orooosed for substituted
phenols (39) to calculate & values f i r adducts of p-fluorophenol with 65 bases. The average agreement between experimental and calculated AH values was f0.6 kcallmole.
Their AH-Av olot for D-fluoroohenoladducts eives a series
of lines, some barallel.and some of different slope, for each
structural tvoe of base. On the basis of this work Arnett (38)
proposed that reliable quantitative (f0.2 kcal/mole) predictions of AH are possible only for bases of a given structural
type interacting with the same reference acid, hut use of the
same AH-Av equation for all base types is not reliable. An
analvsis of literature data hv this author 136) indicates that
the A K A v equations given in references (4-9) are useful for
predicting AH values within f0.5 kcallmole for moderate
hydrogen bonds formed by a reference acid with nitrogen and
oxveen donor bases.
%tensive calorimetric studies by Arnett and coworkers (38)
have also shown that different base tvoes do not corresoond
in the same way to hydrogen bonding g d proton transfer: The
protonation enthalpy, AH;, was determined by measuring the
heat of transfer of a base from carbon tetrachloride to high
dilution in fluorosulfuric acid at 2 5 T . Plots of AHf for hydrogen-bonded p-fluorophenol adducts versus AHi give
separate parallel lines for bases with different functional
groups. However, AHi does give a good linear correlation with
pK. values of the bases. Arnett and coworkers conclude (38)
that no single relationship exists to correlate protonation and
hydrogen bonding.
Excellent linear correlations for adducts of reference acids
with a given structural type of base are also ohtained for AYAH
versus a*.
versus AGO. h a w versus DK.AYAH
versus a. AYAU
'H chemical shift versus ;K. and&', and loi&,,, versus
'9F chemical shift. A review of these correlations is aiven in
reference (8).
In recent years Kamlet, Taft, and coworkers have published
a series of papers (27, 45-47) on the use of chromophore
electronic shifts in hydrogen bonded adducts to derive a - and
&scales of hydrogen-bond donor acidities and hydrogen-bond
acceptor basicities, respectively. Chromophoric shifts of 4nitroaniline. the reference acid. are measured in a series of
hydrogen bonding solvents. A correction is made for nonhydrogen bonding solvent effects by subtracting chromophoric shifts for N,N-diethyl-4-nitroanilinein the same solvents. The resulting "enhanced solvatochromic shift (-AAu)"
is used to establish the @-scaleof solvent basicity (46).
The a-scale of solvent acidity was determined in a similar
fashion by using chromophoric shifts for reference hases in
and
a series of solvents (47).A linear correlation hetewen AVAH
pvalues was observed for bases with similar bonding sites (45).
This is in agreement with earlier reports of linear correlations
between -AAu and the log K,
and '9F NMR chemical shifts
for p-fluorophenol adducts of bases of a given structural type
(46).
Asymmetric and Symmetric Hydrogen Bonds (48)
Weak and moderate hydrogen bonds always are asymmetric
with A-H shorter than H.. .B. For strong or very strong hydrogen bonds with short H. . .B distances (2.S2.6 A), the environment around A and B determines whether A-H.. .B is
364
Journal of Chemical Education
a symmetric or asymmetric hydrogen hond. For example, the
anions in KHF? and NaHFz have a symmetric environment
and the [F...H.. .F]- hydrogen bond is symmetric with a F.. .F
distance of 2.26 A (49). However, an asymmetric hydrogen
hond is found in D-toluidiniumhvdrosen bifluoride (50) where
the two fluorine koms of HF2- have iifferent F. . .H distances
to neighboring -NH3+ groups. The two H-F distances of
1.025 and 1.235 A in HF2- are quite different even though the
F . .F distance of 2.26 A is the same as that found in KHF?.
The short, symmetric or asymmetric hydrogen bonds a;e
generally found in hydrogen bihalide ions, HX2-, or in acid
salts of monocarboxylic acids, [M+[H(RCOO-121. Olovssou
and Jonsson (31) have reviewed X-ray and neutron diffraction
studies of these short hydrogen bonds. Their survey of
structures with 0 . . .O hydrogen bond distances in the 2.402.50 A ranee (Table 7) shows that both svmmetric and
asymmetriciydrogen bonds are found in very ihort hydrogen
honds. They suggest that "centered" rather than "symmetric"
he used to describe the position of the hydrogen atom since
it may he equidistant in 0 . . .H. . .O without k i n g in a crystallographically symmetric position. For example, the hydrogen
atom in the 0 . . .H.. .O hond of ootassium hvdroaen chloromaleate is effectively centered (0;. .H distancks ofi.999,1.206
A) but does not have crvstallozra~hic
svmmetrv).
. .
('onversely, a hydrogcn atom at a symmrtry renter may he
truly centered or may he statically or dynamically disordrred
with the hydrogen atom distrihuted between alterl~ativepw
sitions close to the center ofthe 0 . O Imnd. For example, the
twu halves uf the H;01* ion in YH(C&lAHIO nrr n.lntm1 I J ~
a twofold axis perpendicular to the 0 . .O bund. This s y n merry places the proton nt thr center of the 0 . H .O bond.
However, the thermal ellipsoid of the hydrogen atom is
elonnatpd in the hond dirrctim (51.521.llrnce, the hwlrogen
atom may be centered in a shallow single minimum pbtential
surface or i t mav be non-centered with two staticallv or dvnamically relatid hydrogen atom sites on either side of the
symmetry center. These results indicate the difficulty enc k t e r e i in using neutrm diffraction and X - r n diffraction
~
studirs to distinguish between single and double minimum
hydrogen atom
surfaces. Careful analysis of thermal
parameters for the hydrogen atom in O...H...O and of
0...D. ..O bond distances in deuterium analoeues has heloed.
. ,
hut additional studies in this area are needei.
Hydrogen-bonded ferroelectrics (59) such as potassium
dihydrogen phosphate contain very short 0 . . .O distances
similar to those listed in Table 7. Ex~erimentalstudies of
these systems have been interpretedin terms of a double
minimum potential for 0.. .H.. .O or 0. . .D.. .0.
vibrational spectral studies of hydrogen dihalide ions have
also provided useful information about symmetric and
i ~ s ~ m m r r rhydrogen
ic
honds. Evnniand 1.0 (66,ohtnint.d twc,
different types of spectra for snlts of HCIz- and IC'IHHrI-.
Tvoe I was~attributedto the oresence of a stronelv
~." a&mmet&c
"
h&ogen bond while type h was interpreted in terms of a
symmetric hydrogen bond. Tetraethyl, tetrapropyl, and tetrapentyl salts of HC& gave type I1 spectra while tetrahutyl
Table 7. Neutron Dmractlon Data for Short 0 . . .H.. .O Hydrogen
Bonds
Compound
0- . . 0
O-H
W..O
Symmetry Referof H bond ence
gave type I spectra. All R4N+ salts of [CIHBrI- gave type I
mectra.
Ault and coworkers ( f i l 4 3 )have investigated a variety of
alkali metal saltsof HX2- and [XHX'I- in argon matrires to
reduce the environmental influencr of neiyht~uringioni which
are present in solution or solid state studies suchas those by
Evans and Lo. All infrared matrix spectra of HX2- salts are
characterized as type I1 with the hydrogen symmetrically located in X.. .H.. .X. However, infrared hands characteristic
of both type I and type I1 forms are present for alkali metal
salts of mixed [XHX'I- species such as CIHF-, BrHF-, [HF-1,
and ClHBr-. The data are interpreted as evidence that hoth
asymmetric and symmetric hyd&en bonds are present with
thc position of thr alkrtli metal cation in the ion pair controllingwhich type of anion is found in the matrix..
Infrared matrix isolation studies of H20-HCI (64) and
H3N-HCI (65) complexes in a nitrogen matrix illustrate the
influence of lattice or solvation energies on proton transfer.
No proton transfer takes place in the isolated complexes and
infrared spectral data are in agreement with an asymmetric
0 . . .H-CI bond in H20-HCI and a symmetric N.. .H. . C1
hydrogen bond in H3N-HC1.
Extent of Proton Transfer
The extent of proton transfer in the equilibria
A-H
+ B =A-H..
.B = A-. . .HBf e A-
+ HBf
(1)
denends on the aciditv of HA. the basicitv of B. and the solvent. Quantitative mensurt~mmtsof pai phase proton transfer
enereies (28.66.671 and determination uf orotonation constants in different solvents (68-71) have provided important
information about the effect of solvent on Droton transfer
reactions.
The equilibria in eqn. (1) will also be influenced by the
competition of A- with B to form hydrogen bonds with HA
or BH+. Both homoconjugate and heteroconjugate ions have
been studied and Table 8 lists examples of these. Kolthoff and
Chantooni (68,70,71) have reported formation constants for
several of the ions listed in Tahle 8 in their studies of the influence of aprotic solvents of varying hydrogen bonding ability
on the stabilitv of homo- and heteroconiueate ions.
The position of the proton transf; equilibrium in
A-H.. .B * A-. . .H-B+ has been followed bv IR snectroscopy (72-74), electronic spectroscopy (7), dfpole moment
studies (75), and NMR spectroscopy (15, 76). A paper by
Lindemann and Zundel (74) gives a detailed account of the
application of IR to studies of the extent of proton transfer
.HN+ for adducts of carboxylic acid dein OH.. .N e 0-..
rivatives with nitrogen bases. When the proton is on the carboxylic acid group, the C=O hand is observed a t 1715 cm-'.
Bands for the carboxylate group appear a t about 1575 cm-1
(asym) and 1400 cm-I (sym) when proton transfer has occurred. The percentage proton transfer is calculated from
changes in either vc-0 or v,C02-. In pure 1:l carboxylic
acid-nitrogen base mixtures 50% proton transfer occurs when
PKBH+- PKAH(APK,,) = 2.3 if the nitrogen base has one
additional NH group (e.g., imidazole). When no additional NH
groups are available (e.g., pyridine), ApK. = 4. In the presence
Table 8.
(a)
Homoconjugate and Heteroconlugate Ions a
Homoconjugate Anions A-H. . .AA = RCOOH
A = X-CsH40H
A = HX
(b)
A- = RCOOA- = X-CeH40A- = X-
Heterocanjugate Anions A-H.. .X-
HA = HCI
X- = BT. F(c) HomoconjugateCations B-Ht. . .B
B = pyridine. RNH,. R2NH. RsN
(d)
Heteroconiugate Cations B-Ht. . .B'
B = R3N
B' = pyridine, Cmethylpyridine
R = dkyl Or aryl group. X = halogen
of water molecules ApK, is reduced from 2.3 to 0.9 for imidazole and from 4 to 2 for pyridine.
Other studies have verified that the value of ApK. for 50%
proton transfer is not zero but depends on the acid, base, and
solvent. For example, Pawlak et al. (76) reported ApK. values
of 2.5 for 50% proton transfer in 1:1 alkylcarhoxylic acidpentachlorophenolate adducta dissolved in nitrobenzene. The
ApK, value was lowered to 2 for corresponding arylcarhoxylic
acid adducts. They also found ApK, values increase as the
polarity of the solvent decreases.
H30+(H20). and OH-(H,O).
Several recent reviews have described the structural evidence for hydrated proton species in solids (52, 77). The
species containing hydrated H30+ are normally represented
as HsOz+, H703+, Hg04+ in the literature even though symmetric hvroaen
. .. bonds are the exceotion rather than the role
(see section nn synimetric and asymrnptric hyrirugen bunds).
Ciauere has itralted 178, that H O ' ~ H , O J ,where n = O 3 is
a more accura& representation$ince <he hydrogen atoms in
H30+ are equivalent while those in higher species are not.
Thus Hs02+. H703+. and Hg04+ should be represented as
H30+(H20), H30+(H20)2,and HsO+(HzO)n,respectively. His
arguments have merit since experimental evidence from hoth
solution and solid state studies indicates H30+ is hydrogen
bonded to 1. 2. or 3 water molecules. However. the solid
structures often have additional water mo~ecu~ks
present
which form weaker hydrogen bonds with the HsOz+, H703+,
or Hg04+ units. AS a result, the cluster formulas have been
favored in descriotions of solid state structures (52.77). The
present discussion will iorus on the H O ' , H:,O'rH:,O, or
lIt,Or'. and OH-(HzO) or O,H.,- ions.
"
Giguere (78) has reviewed the IR, NMR, and diffraction
studies which orovide experimental s u.~.o o rfor
t the presence
of H30+ in solution and in the solid state. X-ray and neutron
diffraction measurements (52,771 have verified the presence
of HsO+ in solid HaO+CI-, H30+C104-, HsO+NOa-,
HqO+HSOa-. ( H"? O +.-M O.i - .. HIO+CF?SO?-.
..
. and
*
H~o+cH~~,H,so~-.
A recent ' 7 0 NMR studv has orovided information about
the structure of HnO+ in"solution (79). At -15°C the 170
NMR spectrum for 170-enriched water in SO2 (containing a
slight excess of HF-SbFd consists of a quartet of peaks with
a 170-H coupling constant of 106 Hz. This verifies the
equivalence of the protons in HeO+, and the coupling constant
is in the range expected for s p 2 hybridization of oxygen.
0 H ( H 2 0 ) or H302The isolation and structural characterization of the simplest
hydroxide hydrate, H302-, was reported recently by Raymond
and coworkers (80). The crystal structure contains tridthiobenzohydroximato)chromate(III) anions, hydrated sodium
cations, HsOz- anions, and water of crystallization which are
all linked by hydrogen bonds. The H302- anion lies on a
crystallographic inversion center with an O...O distance of
2.29 A. This is the shortest 0 . . .O distance reported thus far,
and is similar to the F. . .F distance in HF2- (2.26 A).
Y 4+
The hydrogen bond in Hs02+ may be symmetric or asymmetric. For example, neutron diffraction measurements (52)
indicate pirryls~tlfnnicnctd trtrahydrirte, and nitrilnilic acid
hcxahvdmte contain asymmetric HSO1- while yttrium h y drogen oxalate trihydrate contains symmetric-~s02+ion.
However, the thermal ellipsoid representing the thermal
motion of the central hydrogen atom in the yttrium salt is
more elongated in the bond direction than the corresponding
hydrogen atom thermal ellipsoid in asymmetric H5O2+.The
shortest 0. . .H. . .O distance reported for HsOz+ is 2.336 A (81)
a
Volume 59 Number 5 May 1982
365
which is close to the shortest 0. . .H.. .O distance reported for
any hydrogen bond (2.31 A) (82).
Zundel and coworkers (72,83,84) have attributed the IR
continuum observed for acidic and basic aqueous solutions
to the fluctuation of the proton in H s 0 2 + or H 3 0 2 - groups.
Although there is disagreement about whether HsOz+ is a
distinct snecies in aoueous solution (78).
. .. the work of Zundel's
group provides useful information about the effect of concentration on the stahilitv of a articular H~O+(H?O),
species.
Conclusion
The present paper has focused on spectroscopic and diffraction studies which provide information about hydrogen
bonding and proton transfer. However, a number of important
tonics have not been discussed. These include hvdroeen
bonded structures of liquid water (85-89); kinetics d f p r 2 o n
transfer in solution (90,91) and in the gas phase (92); hydration energies of H 3 0 + and BH+ in solution (93) and in the gas
~ h a s (29.66.94):
e
hvdroeen hondine in excited states (95.96):
.
Hnd the theory of hydrogen bonding (97-99).
1361 Ref. 18). pp. 20&220.
1371 Sherry. A. D., ref 191. Chnpter 25. Vol. 3.
(381 Arnett. E. M.. Mitchell. E. J.,snd Murty.T. S. S. R., J. Ampr Chrm Soc.96.3875
119741.
1391 Drslg0.R. $.,end Ep1ey.T.D.. J. Amer Chem Sor.,91.2883 (19691.
I401 Sherry,A. D.,snd Pureell, K. F.. J. Phys Chsm.,74,3515 119701.
I411 Pur~ell.K. F.,Stikoloather, J. A..and Bmnk, S. D., J. Amer Chem Soc.. 91.4019
,.""",.
<,mo,
142) Draeo.R. S..O'Brvan. N..andVoeel.G. C.. J. Amer. Chem Soc.92.3924 119701.
I461 Ksmlaf. M, J., and Tsft, R. W., J. Amer. Chem. Soc.. 98,377 119761.
1471 Taft, R. \V.,and Kamlet, M. J.. J. Amrr. Chrm Soc..98,2E86 119761.
I481 See Emslev..l., Chem Soc R r a , 9.91 119801 for a recent review of wlyatmng hydrogen
h.nds~
~
I491 la1 McGaw, B.L.,andihers, J.A., J. Chem Phys.,39,2677 119681 lbl lbera, J. A,, J.
Chem Phye.. 40,402 119641.
(501 Williams, J.M.,and Sehneemeyer,L. F., J. Amw. Cham. Soc.,95.5780 (19731.
1511 Johnsun. C. K.. and Brunton. G. D.. A h t r a m of American Cmtalloersohic A~sociation
I541 Cur& M., and Speakman, J.C., J. Chem Soc.. A, 1923(19701.
I551 E1liron.R. D.,and Le\y.H. A..Arlo C~y81..19,260119651.
I561 Kostansek, E.C.,and Businp. W. R.,Acfo Cryat.. BZR.2454 119721.
1571 Lundgren. .I.-0.. and Tellaren. R.. Acln C~y$t.,B30.1937 119741.
1581 Kvick.A.. Kaetz1e.T. F..Thumns, R.,andTakuaaraws. F.. J. ('hem. Phw.60. 3866
~.~
Literature Cited
Ill Latimer W. M.,andRudebush, W.H., J. Amor Cham. Soc.42.1431 119201.
I21 h i s , G. N.. "Valence and the Structure of Atoms and Mnlecules? The Chemical
CelalugCo.,New York. 1923.p. 109.
I31 Huggina,M. L..Angew. Chsm..Int, Ed.. LO. 147 119711.
141 Hadri, D., Ed.."Hydrogen Bonding,"Pergamon Press, Oxford. England. 1959.
I51 Pimentsl. G. C., and MeClellsn. A. L.. "The Hydrugen Bcmd." W. H. Freeman, San
Francisco. 1960.
I61 Hamilton, W. C., and Ihors. J. A.,"Hydmen Bondinz inSolida." W.A. Renjsmin.San
Frsneiseu. 1968.
171 Vinogradou. S. N.,and Linnsll, a.A.,"Hydrugen Bonding."Van Nortrand~Reinhuld,
New Yurk, 1971.
181 Juerfen. M. D.,and Schsad. L.J., "Hydrogen Bonding," Marcel Dekker. New York.
I591 Schmidt,V. H., ref. 191,Chapter 23. Vol. 3.
(601 Euens.d. C..and 1.n.B. Y:S. J. Phva C h m . , 70.11. 20 119661.
1611 Ault,R.S.,snd Andrewr,I.., J. Chem.Phyr.,63,2466l1975l:M,1986 (19761.
1621 E1lisun.C. M.,andAdt,B. S.,J. Phys. Chcm.,83,832 119791.
163) Aulf, B.S.. J. Phyhyr. Chem..R3,887 lL9791.
164) A u k R.S.,snd Pimente1.G. C., J. Phys. Chem.. 77.57 (19731.
I651 Ault,B.S..and Pimente1.C. C.. J. Phys. Chem.77, 1649 119711.
I661 111 Arnetf. E.M.. Accounts of Chem. R o s , 6.404 (19791. lh) Epshtoin. I,. M.. Ilsp.
Khim.. 48. 16M) 119791. English Tranrlation, p. 854.
1671 Keharla, P., in "lsnr and Ion Pain in Organic Roactiona? Szwarc, M., Ed., WileyIntemcience. New Yurk. 1972, Vol 1.
I681 Kdfhsff. I. M., in "Essays on Analytical Chemistry." Wanninen. E.. Ed., Pergamon
Preab. Oxford, Englsnd. 1977.
1691 David. M. M.. in"The Chemistry of NonaqueouaSnlvente." IEdifor: 1.ngowrki. J. .I).
Academic Press. New Yurk, 1970. Val. 111.
1701 Kolthoff. I. M., and Chsntooni. M. K..Jr.. J. Amer. Chem. Snc..95,86:19l1973l.
1711 Kolthuff, I. M., and Chantooni, M. K.,Jr., J. Phyr Ch?m.,83.468 119791.
1721 Zundel,G.,ref. 191. Chapter 15. Vol. 2.
I731 Zunde1.G.. J. MoCculorSlruetur~,45,5bl19781.
1741 Lindemann. R.,and Zundel, G.. J. Chem. Sot. Farodoy Tions. 2.79.788 I19771 and
references cited therein.
1751 Subczyk, L., Engelhardt, H.. and Runrl, K., ref. 19). Chapter 20, VoI. :3.
1761 Pawlab, 2.. Magonski, J.. and Jasinski, T., J. Mol~cularStructure, 47.329 (19781.
1771 Williams. J. M..ref. 191,Chapter 14. Vol. 2.
I781 Giguere. P. A , J.CHEM. EDUC..56.571 (19791.
1791 Materscu.1;. D..snd 8enedikt.G. M.. J. Am*,. Chpm. Snc. 101.1959 119791.
~ ~~ ~ .
~K. ~ . ~, ~ d
~ ~D.P., J~ ~A
~ ?hem.
~~ ~ sor.
~dh Inl,
, 1688
~
.
(80) A ~ V - D K..
119791.
I811 Rino.A.,andCntton. F. A,. J. Amw Chem. Srrc., 101.4160119791.
1821 Bertrand, J. A.. Black, T. D., Eller, P. G., Helm, F.'L and Mahmmd. M.. Imr& Chmt..
..
,.
... ... .. .
86 Y ( l R i l,l 9 7 R i,
(831 Leuchr. M.,and '7,u:undcl.G.. Can. J. Ch~m.,S1.487119791.
I841 Leucha,M.,sndZundol.G.. J. Phys. C h m , 82.1632 119781.
I851 Homing. N.I.. Chopin. G. R.. and Renouitch. C., Applr~dSprrtrorcopy Reuirma, 8.
"" ,
.",-",.",,~,,.
I901 Junes. F. M.. 111. EurLlre. D.. and Grunwald. E.. J. Amer. r h r m S o r . 94, 8941
,,W",
lL95ll.
1351 Badger. R. M., and Rnuer, S. H., J. Chem. Phys, 5.839 119371
366
Journal of Chemical Education
I911 Kmsze.A.J.. AcrounU?h~m. Re*.,R,:l51 (19751.
1921 Bohme, D. K., Macksy. G. I., and Tanner, S. D., J. .Arne?. Chrm. Sor., 101, a724
(19791.
1911 Heple.. I.. f:.. and Woolley, R. M.. ref. 1881, Chapter 1.Vol. 3.
1941 nnuidwn. W. ~..sunner,.r..snd ~ a h ~ rP..~ Je. ,~ m p rw m . SG In!, 1fi7s (19791.
1951 Lippert, E.,ref. 191. Chapter 1,Vnl. 1
1961 Bsum..l.C.,amd McClore. D. S.,J. Amer Chpm SO;., IOI.2340 119791.
72,283 119721.
I971 Kellman. P. A..alid Allen. L.C.,l'hem. RPUX.,
1981 Schnad. I..d..reL 181. Chapter 2.
1991 Srhurter. P.. ref. 191. Chapter 2.Vol. I.
~
~