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KITED STATES URTMENT OF MMERCE

LICATION

IfaflonaJ Bureau of Standard',

Library, E-01 Admin. Bldg.

MAR 2 8 1970

NBS TECHNICAL NOTE 513

8E~ -, I<^sk? Hydrogen Fluoride

and the

Thermochemistry

of Fluorine

U.S.

:partment

OF

COMMERCE

National

Bureau

of

Standards
NATIONAL BUREAU OF STANDARDS

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UNITED STATES DEPARTMENT OF COMMERCE

Maurice H. Starts, Secretary

NATIONAL BUREAU OF STANDARDS • Lewis M. Branscomb, Director

NBS TECHNICAL NOTE 513

ISSUED FEBRUARY 1970

Nat. Bur. Stand. (U.S.), Tech. Note 513, 21 pages (Feb. 1970)

CODEN: NBTNA

Hydrogen Fluoride and the

Thermochemistry of Fluorine

George T. Armstrong

Physical Chemistry Division

Institute for Materials Research

National Bureau of Standards

Washington, D.C. 20234

Research supported by the

Air Force Office of Scientific Research

Office of Aerospace Research, U.S. Air Force

under AFOSR Contract No. ISSA. 68-0004

NBS Technical Notes are designed to supplement the Bureau's regular publications program. They provide a means for making available scientific data that are of transient or limited interest. Technical Notes may be

listed or referred to in the open literature.

For sale by the Superintendent of Documents, U.S. Government Printing Office, Washington, D.C, 20402.

(Order by SD Catalog No. C 13.46:513). Price 30 cents.

Table of Contents

Page

I. Introduction 1

Table 1 - Accuracy of Estimates of Heats

of Formation of Some Fluorine Compounds 3

II. Flame Calorimetry of Reactions Forming HF(g) 4

Table II - Flame Calorimetry Studies of the

Heat of Formation of HF(g) 6

Table III - Some Tabulated Values for

AH [HF] 7

III. Evidence Drawn from Other Reaction Processes 7

Table IV 7

Table V - A Series of Reactions Leading to

AH [HF(g)] by Hess 1 Law 8

Table VI - The Heat of Formation of HF(g)

from the Dissociation Energies of H_, F

and HF 9

IV. Recent Work on the Energy of Vaporization

from Water 10

Table VII - Cox and Harrop Study of HF(aq) 11

Table VIII - Vapor Pressure of HF Over LiHF (c)

.

Effect of a Zero-Point Entropy of LiHF (cj ... 12

V. Heat of Formation of Aqueous HF - Recent

Studies 12

VI. New Work on the Dissociation of F and HF 13

VII. References 14

VIII. Figures 17

in

Hydrogen Fluoride and the Thermochemistry of Fluorine

George T. Armstrong

The thermochemistry of hydrogen fluoride is

reviewed. The principal emphasis is on the ex- perimental basis of the heats of formation of

HF(g) and of HF(aq) . Data from flame calorimetry

in which HF(g) and HF(aq) are formed, evidence

from processes involving other fluorine compounds,

and some evidence from spectroscopic and photodissociation studies are brought together.

Keywords: Fluorides; fluorine dissociation

energy; fluorine thermochemistry; heats of formation; hydrofluoric acid; hydrogen fluoride.

I. Introduction

While it is possible for us to find certain unique features of

most elements, in this note we wish only to examine some features of the

element fluorine. We may list certain features which make it of special

interest

:

(1) It is a cornerstone element. With respect to the periodic

system it is in the extreme upper right of the non-noble

elements. As a result schemes that correlate properties

within period or within group find that in both cases fluorine

is the end element. Interpolation is rarely possible; - extrapolation is required in order to obtain properties by

estimation.

(2) It is the most electronegative element and generally conceded

to be the most reactive element. Thus bonds of unusual

strength are formed to other elements; and most elements are found in compound with it (even some of the inert gases).

The chemistry and many physical properties of fluorine compounds are,

therefore, frequently of interest as illustrations (paradigms) of these

properties as found among all compounds. The thermochemistry of fluorine

surely falls in this category.

The chemistry of fluorine showed a rather slow growth prior to

World War II. The primary problems were related to its reactivity and

to the need for proper materials for handling it. As a result of

technological development during the war, however, this barrier was

overcome, and the field of fluorine chemistry started very rapidly

growing and became one of the most active fields in both organic and

inorganic chemistry. This growth is symbolized in part by the formation
of a division of fluorine chemistry in the American Chemical Society - one of the few divisions which is built around an element. Such great

growth is probably due to the uniqueness of properties which fluroine

compounds tend to display.

The thermochemistry of fluorine showed a parallel rapid growth with

a number of laboratories setting up facilities for handling elemental

fluorine as a reactant in the calorimetry; and others studying its

thermochemistry, but limiting themselves to reactions involving only the

combined element. We may refer to the Argonne National Laboratory; the

Fulmer Research Institute in England, John Margrave's laboratory at

Rice University; and the NBS laboratories, as examples of places where

elemental fluorine could be handled. Both flow calorimetry and bomb

calorimetry can be carried out in the NBS laboratories.

The use of reactions of fluorine compounds in which the element is never free has a longer history in calorimetry; and in fact the use of

hydrofluoric acid solution calorimetry has played an important role in

the thermochemistry of silcates and other mineral products for many

years. However, application of bomb calorimetry to reactions involving

aqueous HF has awaited the development of the rotating-bomb calorimeter.

Among the fluorine compounds there are many that have interesting

features. There is one, however, that is almost omnipresent in the

thermochemistry of fluorine - that is, hydrogen fluoride, either as the

gas or as its aqueous solution, hydrofluoric acid. This compound has a complexity of its character, in particular the large degree of hydrogen

bonding that it exhibits, which renders it a fit subject for investigation of many of its properties; but the feature which attracts our attention is its occurrence in so many reactions of fluorine compounds.

For the chemistry of fluorine it has an importance analogous to the

importance of water in the chemistry of oxygen.

In the thermochemistry of fluorine as in the chemistry of fluorine

the acid or its aqueous solution occur in many processes that have been

studied. As a result, the actual values that are derived for the heats

of formation of various compounds are dependent on the value assigned to

the heat of formation of hydrogen fluoride.

A few examples are shown below of processes typical of those by

which the thermochemistry of fluorine has been derived. These should be

multiplied a hundredfold to give the true importance of the compound in

this area.

In this note thermochemical values are given in kcal mol as taken

from the original literature or recent reviews in which the kcal is

equivalent to 4184 J (exactly).
Typical Processes By Which the Thermochemistry

of Fluorine Compounds Has Been Derived

(1) C

2 F

4

(solid polymer) +

2 (g) + 42^0(1) = 2C0

2 (g) +

4[HF-10H

2 0] (1)

(2) BeO(C) + 2HF(aq) = BeF^aq) + ^0(1)

(3) NaF(c) = Na

+

(aq) + F~(aq)

(4) SiF

4

(g) + 2H

2 0(g) = Si0

2 (cristobalite) + 4HF(g)

We might ask ourselves the question: How precise and accurate do

thermochemical values such as heats of formation have to be? This

is a question that really has no exact answer. The answers depend

upon the purposes to which the data are to be put.

We can give one answer which shows some inadequate values. See

Table I.

Table I. Accuracy of Estimates of Heats of

Formation of Some Fluorine Compounds

AH_ kcal mol

Estimate Later Measurement

Compound

NbF

5 (c)

PF

5

(g)

RuF

5 (c)

SF

4

(g)

TaF

5

(c)

-342

-315

-420

-300

-156

-300

-360

-432

-381.4

-213.4

-171.7

-455

Average error

Average bias

Error

+90

+66

-39

-87

+16

+155

+95

44 kcal mol

26 kcal mol

Estimates are selected from a review by Armstrong and Krieger [1].

-1
In a recent review of the thermochemistry of fluorine compounds [1]

we came across a number of compounds for which estimates of heat of

formation had been made at a time there was no available measured value.

The values that have been measured thereafter in some thirty cases were

compared with the estimates. A sampling of the results, recently

updated, is shown in Table I. The average error was 44 kcal mol; the

average bias was 26 kcal mol . While it is not easy to say what is

adequate accuracy; we would contend that this is not accurate enough

for anything but the roughest work.

If one is trying to make equilibrium calculations, one should

realize that an uncertainty of 1 kcal mol" in a heat of formation can

lead, at room temperature, to an uncertainty of a factor more than 5 in its equilibrium concentration.

If one is trying to correlate bond lengths and strengths or determine changes in strain energies, a kilocalorie may be adequate for some cases, but for others a much smaller uncertainty is needed. If one is

trying to carry out a Hess' law calculation that involves adding a number of reactions, one may want considerably better values.

We are dealing, in the case of hydrogen fluoride, with uncertainties of the order to 0.5 kcal mol" . To some who would be satisfied

with the 1 or 2 kcal mol" our discussion may seem to be nitpicking. On

the other hand, to those who are accustomed to the idea that the heat

of formation of water has an uncertainty of about 0.01 kcal mol this seems inadmissibly large.

II. Flame Calorimetry of Reactions Forming HF(g)

Several early calorimetric measurements of the energy of the

reaction of the elements were used to arrive at the heat of formation

of gaseous HF. A later reaction involving the combustion of ammonia

suffers the same problems. These are shown in Table II.

There are two problems. The first is the problem of corrosion and

accounting for all of the reaction products. The second which is a major one is the question of the non- ideality of gaseous HF has been the

subject of several studies none of which is definative in the lower

pressure region near room temperature. The gas undergoes polymerization,

forming significant quantities of H_F„, H_F_, H,F, , H,F,, even at pres- sures of less than one atmosphere.

Figure 1 shows the quantity Z as a function of pressure along

several isotherms. The quantity Z has the value 1 for an ideal gas.

Note in the figure that it deviates from unity even at very low pressures.

Figures in brackets indicate the literature references on pages 14-16,
We are not prepared here to discuss all the PVT data, not all of which

is shown here. Two or three sources but by no means all of them have

been used in deriving the graph. It will be noted that the data from

different sources are significantly different along several isotherms.

Also of some importance is the fact that the data are incomplete at the

very low pressure end, even though Z still deviates from unity. This

lack of completeness has caused some difficulty in extrapolation. A

model of the polymerization is necessary and the number of possible

polymers renders the interpretation of such a model in the light of the

available data somewhat ambiguous.

Figure 2 gives the parameter most interesting to us that can be

derived from the PVT data, H°-H

Remembering that H°

29g

= H

2gg

+ (H° 298

- H^)

we see from the figure that the enthalpy of the ideal gas differs from

that of the real gas by several tenths of a kilocalorie per mole in

this region. The solid lines represent values calculated from the work

of Long, Hildebrand, and Morell [8] . The dashed lines represent the

work of Strohmeier and Brieglib [9]. Note that these differ by a factor

of two in this region. Medvedev [7] in 1963 made a recalculation of the

correction for non-ideality based on the heat capacity of the real gas

studied by Franck and Meyer [10] and of the ideal gas compiled by

Gurevich, et al. [11]. This recalculation shows the non-ideality to

be less again by a factor of two than we derived from the work of

Breiglib and Strohmeier. In Table II we see several determinations of

the heat of formation of HF and effects of corrections based on different

interpretations of the non-ideality.

The value of -64.2 for NBS Circular 500 [12] was apparently derived

by using the value of Ruff and Laas (as corrected by Ruff and Menzel) or by averaging three of the values shown.

If we refer to figure 2, we see in the inset that at higher temperatures the differences between various non-ideality corrections approach

each other and tend to go to zero. Of the various measurements shown in

Table II, only the one by von Wartenberg and Schutza was done at a high

enough temperature that non-ideality corrections could be ignored. Thus

this value is the only really unambiguous value for gaseous HF obtained

by direct combination of the elements.

The measurement of the combustion of ammonia in fluorine suffers

the same ambiguity as the other measurements.

The value -64.2 kcal mol selected for NBS Circular 500 was

generally accepted and received wide usage. An analysis of the heat of

solution, some equilibrium constant data and heat of dilution data led

to the heat of formation at a series of aqueous solutions. These values

from Circular 500 are shown in Table III.
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Table III. Some Tabulated Values for AHf [HF]

Ideal gas

NBS Circular 500 [12] -64.2

NBS Tech Note 270-3 [13] -64.8

JANAF Tables [14] -64.8+0.3

Cox and Harrop [15]

based on -64.8 for the ideal gas rather than -64.92 as used

10 HO 03 H

2 -75.605 -78.66

-76.235 -79.50

-77. 247

-80.43''

by Cox and Harrop.

III. Evidence Drawn from Other Reaction Processes

The first inkling the author had of a possible error in the heat

of formation of HF came on examining the information collected in the

middle and late nineteen fifties on the heat of formation of CF, . 4

We may illustrate the nature of the clue by the following

reactions

Table IV.

CH

4

(g) + F

2 (g) = 4HF(g) + CF

4

(g)

Jessup, McCoskey, and Nelson [16]

Using -64.2 kcal for AH

f

[HF(g)] we get AH [CF, (g)] = -220.4

C

2

F

4

(polymer) +

2 (g) + 2^0(1) = 2C0

2 (g) + 4HF(aq- 10^0)

C

2F

4 (polymer) +

2 (g) = C02 (g) + CF

4

(g)

CF

4

(g) + 2H

2 0(1) = C02 (g) + 4HF(aq-10H2 O) AH = -41.5

Good, Scott , and Waddington [17]

Using -75.605 kcal for AH [HF'IOH 0] we get AH

f

[CF

4 (g)]= -218.3

C

2F

4

(g) = CF

4

(g) + C(c)

C

2

F

4

(g) + 2H

2 (g) + 4[118.2 H

2 0] (1) = 2C(c) + 4[HF- 118.2^0] (1)

CF

4

(g) + 2H

2

(g), + 4[118.2 H

2

0] (1)= C + 4 [HF- 118. 2^0] (1)

Neugebauer and Margrave [18]

HF-118.2H 0] weAH

f

CF

4

(g)

Using NBS 500 value for AH

f

[HF-118.2H 0] we get

-217.5
Note that CF, and HF appear on the same side in Jessup's reaction

and on opposite sides in the other two.

We suggested that the above values could be reconciled if AH [HF] was in error by a few tenths of a kilocalorie per mole.

A number of other reactions were also available and tended to

confirm this view.

One direction the above suggestion led was to a series of further

determinations on CF, (g) and another direction was to a new appraisal

of the heat of formation of HF(g) and HF(aq).

In every case, where the heat of formation of CF, and HF were tied

together, the above suspicion was further confirmed. It became apparent

at least to us, that so long as the uncertainty prevailed in the heat of

formation of HF, no reaction involving both HF and CF, would be of much

assistance in defining the heat of formation of CF, . Other methods were

sought and were brought successfully to a conclusion, we think.

With regard to the heat of formation of HF, a new series of studies was begun. The idea of trying to bring in the results of the relevant

data obtained by the new techniques of fluorine bomb calorimetry was seriously proposed by Feder, Hubbard, Wise and Margrave [19]. They

immediately derived a value of -64.92 kcal mol" for AH°F[HFg],

A typical series of reactions they introduced is shown in Table V.

Table V. A Series of Reactions Leading to AH [HF(g)] by Hess' Law

Si0

2

(a-quartz) + 2F

2

(g) = SiF^g) +

2 (g) (1)

AH*

98

(1) = -168.26 ± 0.28 [38]

SiO (crystobalite) = SiO (a-quartz) (2)

AH°

9g (2) = -0.35 ± 0.05 [39]

2H

2 (g) +

2 (g) = 2H

2 0(g) (3)

AH°

98

(3) = -115.598 ± 0.020 [12]

SiF

4

(g) + 2H

2 0(g) = Si02 (crystobalite) + 4HF(g) (4)

AH°

g8 (4) = +24.53 ± 0.36 [40,19]

2F

2 (g) + 2H2 (g) = 4HF(g) (5)

AH°

9g

(5) = -259.676

AH°

298

[HF(g)] = -64.92 ± 0.12
AH values are in kcal mol , uncertainties are from [19], and are twice the standard deviations of the means.

The revised NBS values (for TN 270-3) by Evans and Parker [13] were

prepared about this time and may have been influenced by this value.

There is available one completely different process for calculating

the heat of formation of HF. This is the spectroscopic observation of

the dissociation of HF.

Table VI. The Heat of Formation of HF(g)

From the Dissociation Energies of H , F , and HF

Zffl298

J.15_

1 kcal mol

H(g) + F(g) = HF(g) -136.10 [20]

|F

2 (g) = F(g) +18.88 [13]

|H

2

(g) = H(g) +52.095 [13]

|H

2 (g) + iF2 (g) = HF(g) -65.125

The necessary series of reactions is shown in Table VI. Note that

the dissociation energy given by the work of Johns and Barrow [20] must

be combined with the dissociation energy of fluorine. This may be the

limiting step in the process. To illustrate the large amount of information available, we present in figure 3 the measurements uncovered in

a review [1] of the work reported on this limited subject in a period of some 10 years following the issuance of NBS Circular 500. Distilled out

of this mass of work, we find almost the only work upon which the current

best value rests follows the analysis of Stamper and Barrow [21] (Fig- ure 4), who obtained the dissociation energy from a 3rd law correlation

of the experimental work of Wise [22] and of Doescher [23], As you can

see from the boundary lines, almost full weight was given to the work

of Doescher. This has become a textbook example of the analysis of

equilibrium data by the third law of thermodynamics. The figure, from

Stamper and Barrow, was found in Pitzer and Brewer (Lewis and Randall)

[24].

Note that the uncertainty (0.01 ev. or 0.23 kcal) assigned by the

authors to the dissociation energy of HF would just about overlap the

value currently accepted for HF (TN 270-3) . In order for an error in

the value assigned to the dissociation of F to bring the AH into line

would require a slightly larger dissociation energy of F than we get

from Stamper and Barrow's analysis.
IV. Recent Work on the Energy of Vaporization from Water

One of the more important contributions that came out following the

revival of interest in HF was a series of reactions leading to the heat

of solution of HF in HO by Cox and Harrop [15]. This threw a real

monkey wrench into the situation. These reactions are found in Table VII.

The series of reactions was rather long, but each step was very

small (except for the vaporization step). The series was planned in

such a way as to allow no contribution from the non-ideality of the HF

gas.

The value shown in Table III for 10H was obtained from this sum

of reactions, by merely adding to the heat of formation of HF(g) as listed for TN 270-3 fl3].

The difficulty is that this is now a whole kilocalorie per mole

more negative than had been selected for TN 270-3 which was already

0.63 or more kcal more negative than had been listed in NBS Circular

500 [12].

We have looked carefully over the work of Cox and Harrop for a source of discrepancy and so far certainly nothing positive has been

discovered. We naturally turn to the largest term - the vapor pressure

point. To get one kcal mol" by an error here would require the pressure

to be in error by a factor of five, and even 0.4 kcal would require an

error of a factor of two.

One tantalizing prospect is that there is an error in the calculation of AS for the reaction. Now S for HF(g) can be calculated quite

accurately, and S 9QR -S~ for LiF and LiHF„ have both been measured in

very reliable laboratories. If S~ is zero for all substances, there

hardly seems to be a chance for an error in calculating AS.

The possibility of a residual entropy for LiF seems hardly con- ceivable. The possibility of a residual zero point entropy for LiHF (c" seems not so far fetched - and this reminds us of the classical thermodynamic problems in the verification of the third law of thermodynamics,

10
Table VII. Cox and Harrop Study of HF(aq)

(+) LiHF

2 (c) + 6017 H

2 0(1) = [LiHF

2 -6017 HO] (1)

(-) LiF(c) + 6008.67 H

2 0(1) [LiF-6008.67 H2 0] (1)

(+) [HF-8.33 H

2 0](1) + 6008.67 H^0(1) = [HF-6017 H^O] (1)

(-) [HF-8.33 H

2 0](1) + [LiF.6008.67 HO] (1) = [LiHF

2 -60l7 H 0](1)

(-) LlHF

2 (c) LiF(c) + HF(g)

^298 :ml

kcal mol

1.860 ± 0.038a

1.059 ± 0.028

-0.740 ± 0.006

-0.177 ± 0.006

13.410 ± 0.048

HF(g) + 6017 H

2 0(1) [HF-6017 H 0](1) -13.172 ± 0.068

Uncertainties are twice the standard deviation of the mean, assigned by

Cox and Harrop.

We have here a very strong hydrogen bond (FHF ). The question of

the zero point entropy in salts containing the bifluoride ion has been

the subject of numerous investigations. Westrum [25] did his thesis

with Pitzer on this very subject (KHF ) and satisfied himself by a thermodynamic cycle involving the heat capacity and vapor pressure of

KHF and the heat capacities of KF and HF that the proton is not

assymetric. Other evidence is from structural analysis (neutron diffraction) (F-F distance) and from the infrared spectra of the salts.

No positive evidence for an unsymmetrical hydrogen bond has been reported

for KHF ion the protons are assymetric as

2

, KDF , NaHF . would be expected.

In the H F

The evidence actually put forward on LiHF is too limited to draw

any conclusions specifically for it, the only direct evidence being a measurement of the F-F distance, which turns out to be only slightly

greater than in KHF .

11
Table VIII. Vapor Pressure of HF Over LiHF (c)

Effect of a Zero-Point Entropy of LiHF (c)

LiHF

2

(c) - LiF(c) + HF(g) P

2gg

= 1.89 ± 0.14a Torr

AG = -RT In P = 3.553 ± 0.044 kcal mol"

1

AH = AG° + TAS°

With S °2QR an<J assuminS S

° or s

° n cal deg mol cal deg mol cal deg mol

HF(g) 41.51 ± 0.0^a

LiF(c) 8.52 ± 0.04

LiHF

2

(c) 16.97 ± 0.04 1.38 (R In 2)

AS AS

AS , kcal mol" 1 +33.06 31.68

TAS , kcal mol" 1 9.857 9.445

M° ,,, kcal mol" 1 13.410 12.998

298. 16'

Uncertainties given are estimated as twice the standard deviation of the mean.

If the proton were assymetric we would expect Rln 2 ± 1.38 for Sq

of LiHF and this would lead to a slightly lower enthalpy of vaporization and a slightly less negative heat of formation of HF(by 0.4

kcal mol" 1 ). See Table VIII.

The possibility of this is reduced further by the fact that a Clausius-Clapeyron calculation of the heat of vaporization of HF from

LiHF by Westrum and Burney [26] give 13.7 kcal mol" , which is even

larger than the vaporization energy given by Cox and Harrop. The work

on the (FHF") bond was recently summarized in an informal report by the

author [27].

V. Heat of Formation of Aqueous HF-Recent Studies

More recently evidence has begun to accumulate concerning the heat

of formation of aqueous HF, drawn from completely independent sources.

In Figure 5 we show as

- the lower line the heat of formation of

HF(aq) as a function of H 0/HF ratio, as listed in TN 270-3 [13]. In

the upper curve are shown the results of Cox and Harrop [15]. In

between are points representing the heats of formation as found from

more recent data. Typically these points are derived from the heat of

12
formation of a fluoride by direct combination of the elements or some other process; combined with a hydrolysis reaction in which HF(aq) is

formed. The points are located at the concentration of HF(aq) that was actually formed experimentally. We find definite evidence of a "concensus"

that the heat of formation of HF(aq) must lie between the two extremes

represented by the upper and lower curves.

These data come from several laboratories. The two BF points

(points 1 and 2) represent direct combinations of boron and fluorine as measured at NBS [28] and at Argonne [29], combined with a cycle for the

hydrolysis of BF proposed by Hubbard [29], The three CF, points

(points 3, 4, and 5) represent the recent determination of the heat of

formation NPL [31] and Wisconsin [18], This could be augmented by a very similar pair of points based on work on CF, at Argonne [32] . The

NF point was performed completely at Dow Thermal Laboratory [33]. The

HF point (point 7) was by directly forming the solution from the elements and water (by King) in our laboratories [34]. The latter point

if combined with the heat of solution of HF from TN 270-3 [13] leads to

£Hf (HF(g)] = -65.15 kcal mol . Similar values would also be found from

other data points in Figure 4.

One final piece of work may be mentioned in closing. Rodenberg

and Vanderzee [35] have recently completed a measurement of the heat

of solution of HF in NaOH solution leading to the heat of formation of

fluoride ion or of HF at °° dilution. By their calculation they find

AHF at °° HO is in agreement with the value in TN 270-3. Because the

extrapolation to ^ dilution is several kilocalories per mole this is

not necessarily inconsistent with the finding shown in the chart. If

confirmed, however, it could mean that the dilution energy, (to °° dilution)

is in error.

The recent review by D. R. Stull and collaborators [36] summarizes

in a systematic way much of the material presented above, possibly

including some additional material. The validity of such a summary,

and the enthalpy of formation derived from it depend strongly upon the

weights assigned to the various contributing reactions.

VI. New Work on the Dissociation of F and HF

An important very recent contribution to this subject is made by

Dibeler, and co-workers [37] who have measured the dissociation energy

of fluorine and of HF. Note that if their two values are taken together,

the enthalpy of formation of HF(g) can be calculated, in the same way

as in Table VT, and give a value in good agreement with values cited

above. However, if their F dissociation energy is combined with the

dissociation energy of HF reported by Johns and Barrow [20] a discrepancy of about 3 kcal mol' 1 is found. This is rather substantially

outside the limits suggested by the large aggregation of data on HF(g)

given above. A very real problem arises as a result of the new work.

If the calorimetric thermochemistry be taken as conclusive evidence of

13
an enthalpy of formation of HF(g) of -64.8 ± 0.5 kcal mol~ , and the

work of Dibeler and co-workers be accepted, then an error must be present

in the analysis of Johns and Barrow, which is unexpected in view of the

very fine consistency of the results they obtained, or as suggested by

Dibeler, et al., they have measured vibrational levels above the thermodynamic dissociation level of HF. A further result of the work of

Dibeler, et al., is that bond energies calculated on the basis of the

dissociation energy of F will be appreciably lower than previously

calculated.

[1

[2

[3

[4

[5

[6

[7

[8

[9

VII. References

G. T. Armstrong and L. A. Krieger, Progr. Intern. Res. Thermodyn.

Transport Properties 1962 , 8-77.

H. von Wartenberg and 0. Fitzner, Z. anorg. u. allgem. Chem. 151 , 313-25 (1926).

0. Ruff and F. Laass, Z. anorg. u. allgem. Chem. JL83, 214-22 (1929).

H. von Wartenberg and H. Schuetza, Z. anorg. u. allgem. Chem. 206 , 65-72 (1932).

G. T. Armstrong and R. S. Jessup, J. Research Nat ' 1 Bur. Standards

64A, 49-59 (1960).

0. Ruff and W. Menzel, Z. anorg. u. allgem. Chem. JL98, 375-82

(1931).

V. A. Medvedev, Zh. Fiz. Khim. 37 (6), 1403-6 (1963).

L. H. Long, J. H. Hildebrand, and W. E. Morrell, J. Am. Chem.

Soc. 65, 182-7 (1943).

W. Strohmeier and G. Briegllb, Z. Elektrochem. _57, 662-7 (1953).

[10] E. U. Franck and F. Meyer, Z. Elektrochem. 63, 571-82 (1959).

[11] L. V. Gurvich, G. A. Khachkuruzov, and V. A. Medvedev, et al.,

Thermodynamic Properties of Individual Substances 1962 , Vols. I,

II.

[12] F. D. Rossini, D. D. Wagman, W. H. Evans, S. Levine, and I. Jaffe,

Selected Values of Chemical Thermodynamic Properties, Natl. Bur.

Standards Circular 500 (U. S. Government Printing Office, Washington, D. C, 1952).

[13] D. D. Wagman, W. H. Evans, V. B. Parker, I. Halow, S. M. Bailey,

and R. H. Schumm, Selected Values of Chemical Thermodynamic

Properties, Nat'l Bur. Standards Tech. Note 270-3, January 1968

(U. S. Government Printing Office, Washington, D. C. 20402).

14
[14] Dow Chemical Company. JANAF Thermochemical Tables. PB168370;

Clearinghouse for Federal Scientific and Technical Information,

Springfield, Virginia.

[15] J. D. Cox and D. Harrop, Trans. Faraday Soc. 61, 1328-37 (1965).

[16] R. S. Jessup, R. E. McCoskey, and R. A. Nelson, J. Am. Chem. Soc.

77., 244-5 (1955).

[17] W. D. Good, D. W. Scott, and G. Waddington, J. Phys. Chem. 60,

1080-9 (1956).

[18] C. A. Neugebauer and J. L. Margrave, J. Phys. Chem. j60, 1318-21

(1956).

[19] H. M. Feder, W. N. Hubbard, S. S. Wise, and J. L. Margrave, J. Phys. Chem. 67, 1148-50 (1963).

[20] J. W. C. Johns and R. F. Barrow, Proc. Roy. Soc. (London) A251 , 504-18 (1959).

[21] J. G. Stamper and R. F. Barrow, Trans. Faraday Soc. 54, 1592-4

(1958).

[22] H. Wise, J. Chem. Phys. 20, 927 (1952); J. Phys. Chem. 58, 389-91

(1954).

[23] R. N. Doescher, J. Chem. Phys. 19, 1070-1 (1951); 20, 330-4 (1952).

[24] K. S. Pitzer and L. Brewer (G. N. Lewis and M. Randall), Thermodynamics, (McGraw Hill Book Co., 1961).

[25] E. F. Westrum, Jr., and K. S. Pitzer, J. Am. Chem. 71_, 1940-48

(1949).

[26] E. F. Westrum, Jr., and G. A. Burney, J. Phys. Chem. 65, 344-8

(1961).

[27] G. T. Armstrong, Nat'l Bur. Standards, Informal Report, July 1966.

[28] E. S. Domalski and G. T. Armstrong, J. Research Nat'l Bur. Standards

71A , 195-201 (1967).

[29] G. K. Johnson, H. M. Feder, and W. N. Hubbard, J. Phys. Chem. 70,

1-6 (1966).

[30] E. S. Domalski and G. T. Armstrong, J. Research Nat'l Bur. Standards

71A , 105-117 (1967).

[31] J. D. Cox, H. A. Gundry and A. J. Head, Trans. Faraday Soc. 61 , 1594-1600 (1965).

15
[32] E. Greenberg and W. N. Hubbard, J. Phys. Chem. 72, 222-227 (1968).

[33a]G. C. Sinke, J. Chem. Eng. Data J-0, 295-6 (1965).

[33b]L. C. Walker, J. Phys. Chem. 11, 361-3 (1967).

[34] R. C. King and G. T. Armstrong, 72A , 113-131 (1968).

[35] C. E. Vanderzee, Private communication, July 17, 1968.

[36] A. N. Syverud, AFRPL-TR-69-70, January 1969.

[37a] V. H. Dibeler and J. A. Walker, Inorg. Chem., Accepted for

publication (1969).

[37b] V. H. Dibeler and J. A. Walker, Private communication, April 1969.

[38] S. S. Wise, J. L. Margrave, H. M. Feder, J. Phys. Chem. 65, 2157-9

(1961).

[39] J. P. Coughlin, U. S. Bur. Mines Bulletin 542, (1954).

[40] F. A. Lenfesty, T. D. Farr, and J. C. Brosheer, Ind. Eng. Chem.

44, 1448-50 (1952).

16
3H SHOW hid 1V0M

\T> O

CM

(UM

•H

Fn

JH 310W y3d P* "H-oH

0)

60

•H

Ad/LU=z

17
TABLE 1. FLUORINE

Substance D or AH° (kcal/mole) Remarks

F(g) N.B.S. Circ. 500 [9] gives 18.3 for AHf° . See F

2

(g) for extensive later information.

F,(g)

F

2

(t) F2

(c)

Experimental measurements

leading to D or AH° for dissociation have been reported as follows:

66[54], 50 ±6[55], < 45[56], 37.7 at 759 to 1115°K or 38.9

at 1000°K[57], > 45 [58], 40-45[59], 37.7[60], 39.9 ±0.8[6l], 37.0±2[62], 31-5 ±0.9[63], 37.6 ±1.6[64],

37.6 ±0.8[65], 31[66,67],

38 ±0.4[68, 69], 32 ±3[70],

37.5[71], 41.3 ±0.5[73],

37.1[92], 31.6 ±4.3[606], < 39.0[606].

Discussions and calculations of the dissociation energy have been reported by [9, 11, 14, 18, 75, 76, 77, 78, 79, 80, 81, 82, 83, 84, 85, 86, 87, 89, 90, 91, 93, 94, 95, 96, 97, 98, 99, 100, 101, 102, 103, 104, 508]. [105] meas. v. p. and AH vap. [540] meas. ionization and dissociation by electron impact. [580] est. D values from -34 to +180.

[105] meas. AH fus.

[105] meas. AH trans. (c,l) (c,Il).

Data in tables 1-8 are citations from the literature and have not been critically reviewed for this paper. All numerical values are in kcal mole unless otherwise specified.

See [9] for additional references prior to 1949- Superscript a on AHf values for Br and I compounds indicates that such values are based on Br2 U or I~(g) as standard states.

Figure 3. Reproduction of a table listing work published on the

dissociation energy of fluorine in the interval about

1950 to 1960. The table is taken from reference [1]

and the reference numbers listed refer to the bibliography of reference [1].

18
1,200

Figure 4 Values of AHo of dissociation of F2 from combination of individual equilibrium measurements with (fq — n%)/T data. [J. G. Stamper and R. F. Barrow,

Trans. Faraday Soc, 54, 1592 (1958), fig. 1.]

78.0

o

00

0>

CM oX

77.0

< 76.0

75.0

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