THE REACTION BETWEEN UREA AND FORMALDEHYDE IN CONCENTRATED SOLUTIONS

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• Collection: General CIA Records
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• Release Decision: RIPPUB
• Original Classification: K
• Document Page Count: 10
• Document Creation Date: November 9, 2016
• Document Release Date: February 22, 1999
• Sequence Number: 3
• Publication Date: August 1, 1952
• Content Type: STUDY

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RECUEIL DES TRAVAUX CHIMIQUES DES PAYS-BAS
publ. par la Soc. chim, neerl. ['s-Gravenhage (Hollande), Lange Voorhout 5]
T. 71, 8 (aout 1952).
Soc. anon, d'editions scientifiques D. B. CENTEN, Amsterdam.,
547.495.7
THE REACTION BETWEEN UREA AND FORMALDEHYDE
IN CONCENTRATED SOLUTIONS,
BY _
J. I. DE JONG and J. DE JONGE
(Philips Research Laboratories N.V. Philips' Gloeilampenfabrieken,
Eindhoven - Netherlands).
The nature and the velocities of the reactions occurring in concentrated
solutions of urea and formaldehyde may be very similar to what is found
in dilute systems.
A rapid initial reaction between urea and formaldehyde, as reported in
literature, was not observed and is shown to be due to the analytical
methods, used,
Introduction.
Smythe 1) found the reaction between urea (U) and formaldehyde
(F) in four molar solutions to show a very high initial velocity.
Ignoring this stage, the further decrease in F could be described by
a second order reaction equation as was affirmed by Bettelheim and
Cedwall2). At first it occurred to the present authors that, at least
partially, the rapid initial reaction might be due to a contamination
of urea by ammonium salts 3). Recently, however, Smythe 1) showed
that even in the absence of ammonium salts there is a rapid reaction.
According to a previous investigation 4) 5), the reaction between
urea and formaldehyde in dilute aqueous solution did not show a fast
initial reaction stage. This led us to the present paper dealing with
some experiments with more concentrated solutions.
Again the reaction was found to be "normal" from the very
beginning and therefore our analytical method was compared with
those used by the other investigators,
1) L. E. Smythe, J. Phys. Colloid Chem. 51, 369 (1947); J. Am. Chem. Soc. 73,
2735 (1951).
Q) L. Bettelheim and J. Cedwall, Svensk Kem. Tid. 60, 208 (1948).
$) J. 7. de long, Rec. tray. chim. 69, 1566 (1950).
4) J. I. de Jong and J. de Jonge, Rec. tray. chim. 71, 643 (1952).
6) J. 1. de long and J. de Jonge, Rec, tray, chim. 71, 661 (1952),
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App rnviPd For RPIPasp 1999109194 ? CID-RnPR3-fi(1d93Rfi(19flflf113flflf13-9
891 J. 1. de long and J. de Jonge.
1. The reactions between area and formaldehyde in concentrated
solution.
The analytical methods applied are essentially the same as pre-
viously described. Free F was estimated by the iodometric sulphite
method and the sum of free F and methylol groups was determined
with an alkaline iodine solution.
a) The reaction between urea and formaldehyde.
Urea and formaldehyde were brought together in 2 and 4 molar
aqueous solution at pH = 7 and the decrease in the concentration
of formaldehyde with time was followed.
Two experiments are shown graphically in fig. 1. It is clearly seen
that the reaction takes a normal course from the very beginning.
The reaction constants were calculated with aid of the well-known
equation for a bimolecular reaction.
A variation in the initial concentration of U or F appeared to have
no influence on the calculated reaction constant, indicating that the
rate is indeed proportional to the concentration of urea and formalde-
hyde (bimolecular reaction). Some experiments are summarized below:
Experiments at 35? C and pH = 7, using 0.0013
moles/litre KH=PO4 and 0.0009 moles/litre Na,HPO4
as a buffer solution:
4.0 2,0
2.0 2.0
Experiments at 25? C at pH = 7 with the same
buffer concentration:
Initial concentration
In moles per litre of
ki
in litre/mole second
urea and formaldehyde
2.0 2.0
0.50 X 10-4
2.0 4.0
0.44 X 10-{
4.0 4.0
0.49 X 10-4
The reaction constant at 350 C and with pH = 7 in these experi-
with concentrated solutions, is found to be k1 = 1.1 X 10_4
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The reaction between area and formaldehyde, etc. 71 (1952 -RECUEIL a92
F reacted in m7
0.1 N 12~0.6m1
T=35-? py_7
? 2
U+F
(UJO(F)O2M/L
U+F
.r
x
x
Fig. 1.
litre/moles second, while with about 0.1 molar solution, using the same
buffer concentration, a value of 0.9 X 10-4 was obtained (see fig. 4
in 4). The value for the rate constant at 25? C, obtained in concen-
trated solutions, viz. k1 = 0.5 X 10-4, is also very close to the result
kl = 0.6 X 10-4 recently found in dilute solution under similar
conditions of pH and buffer concentration.
b) The reaction between monomethylolurea and formaldehyde
was studied in a solution containing 2 moles/litre OF and
2 moles/litre F. At 35? C and pH = 7, using 0.0026 moles/litre
KH2PO4 and 0.0018 moles/litre Na2HPO4, the reaction rate constant
appeared to be 0.35 X 10-4 litre/moles second, while in about 0.1
molar solution a value of 0.37 X 10-4 litre/moles second was found.
c) The measurement of the decomposition of LIP in concentrated
.solution appeared to be rather difficult as the forward bimolecular
reaction between U and F prevails already at a very low decom-
position (see below). Though the result may not be expected to be
very reliable a run was made with initial concentrations of 1r1F, of
1.28 and 2.56 moles/litre (,T = 35? C, pH = 7, the solution contained
.0.0013 moles/litre KH2PO4 and 0.0009 moles/litre Na2HPO4). From
the initial slope of the decompositon curve (F - t plot) was found:
k2 = 2.7 X 106 sec -1 and 2.6 X 10-6 sec-1. Recently with a diluted
solution and under similar conditions the value was found to be
k2 = 2.5 X 10-6 sec-1.
d) Reversibility. When studying the reaction between U and, F
in concentrated solutions the reverse reaction may be ignored up to
a considerable conversion, as the following calculation shows: -
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893 4" 1. 1. de long and 1. de Jonge,
When in a solution originally containing 2 moles/litre LI and
2 moles/litre F. 0.7- moles/litre OF is formed the velocity of the
LIF formation is roughly: 1.1 X 10 1 X L3 X 1.3 2 X 10-4
moles/litre second and the velocity of the reverse reaction is,
2.7. X 10- 6 X 0.7. 2 X 10- a moles/litre second, where 1.1 X 10_4
and 2.7 X 10-6 are the values for k1 and k2 respectively at 35? C,
in dilute phosphate buffers at pH = 7.
Thus the decomposition velocity at 35 % conversion is only I % of
the OF formation velocity. This explains why previous investigators
did not find indications for the reversibility of the reaction in
concentrated solution.
e) Equilibria. From the foregoing experiments it is reasonable
to assume that the rate constants at 35? C and pH = 7 in dilute and
concentrated solutions will have the same value. Consequently the
equilibrium constants K = k2/kt and K* = k2*/k1* for the equilibria
k1_4k1*
U -}- F .-- LIF and LIF + F f-kzr' LIF2
2
as calculated from the rate constants in dilute systems, will also hold
in concentrated solutions.
In that case it is possible to calculate the amount of U, F, UP and
LIFq present at equilibrium conditions in a solution with given initial
concentrations of U and F. In equilibrium the following relations
hold:
[LI]eg [F]cq = K and [LIF]eq [, _ K.
[UF].q [UF21eq
with K = 0.036 moles/litre and K* = 0.22 moles/litre at 35' C. The
result of the calculations is as follows:
Initial concentrations Calc. equilibrium conc. in Conc. of F in moles,/litre
UF2 found exp.
of U and F in moles,/litre u Fles/IUF of
0.4
1.0
2.0
0.12
0.06
0.22
0.06
0.060
0.26
0.08
0.56
0.18
0.080
0.49
0.09
1.11
C.40
0.085
The calculated equilibrium concentrations of free F have been
checked experimentally in solutions of U and F left standing for
2-3 days, after which time the concentrations of F were constant.
The agreement with the calculated values is excellent, as is shown in
the table above, from which it may also be seen that a five-fold
dilution of a solution containing initially 2 moles LI and F per litre
has only a slight influence on the equilibrium concentration of F,
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The reaction between urea and formaldehyde, etc. 71 (1952) RECUEIL 894
In conclusion it may be said that the reactions between U and F
in dilute and in concentrated solutions will be very similar at 35? C
and pH = 7, and that the reaction rate constants will have nearly the
same value.
Experimental details.
The reaction vessel consisted of a platinum crucible, kept in a water bath of
constant temperature (within 0.1? C). Evaporation of water from the reaction
mixture was negligible during the short reaction periods. The formole was
previously neutralized with dilute sodium hydroxide solution.
Solutions containing U and F were adjusted to the reaction temperature chosen
and then added to the reaction vessel from a burette; the platinum allowed of
the temperature being well regulated.
In all experiments the temperature was 35? or 25? C and the pH was adjusted
to pH = 7 using 0.0013 moles/litre KH2PO4 and 0,0009 moles/litre Na2HPO4.
The total reaction volume was 25 ml. Samples of 1 or 2 ml were taken and run
into 6 ml ice water, the resulting solution then being diluted to 10.0 mI with
ice, water. From this solution 3 ml samples were analyzed.
Example of an experiment:
Initial concentrations of U and F: 2 moles/litre. Temperature 35? C, pH = 7.
2 ml samples were taken and diluted to exactly 10.0 ml. From the diluted solution
3 ml samples were taken for analysis on formaldehyde using the, method previously
described 4). The results are collected in the next table.
The concentration of F at zero time was known from the amount of formaldehyde
added to the .reaction mixture. Moreover F was estimated in the reaction mixture
by analyzing, according to the alkaline-iodine method, a sample a short time after
mixing. Even if some methylolurea should already be formed this method would
still give an accurate result since-with this method both free formaldehyde and
methylol groups are determined. The difference between the added amount and
estimated amount of formaldehyde in the reaction mixture never exceeded 1 %.
In the .experiment mentioned above 40,10 ml 0.0996 N 12 was added to a 3 ml,
sample. After reaction the amount of 12 not consumed appeared to be equivalent
to. 16.15 ml 0.0993 ml thiosulphate. Consequently the sample had consumed
24.0 ml 0.0996 N 12; calculated from the added amount of formaldehyde 24.0 ml
0.0996 N 12.-
11. Comparison of analytical methods for the urea-formaldehyde
solutions.
As the high initial reaction rate between urea and formaldehyde
i I
time in
seconds
ml 0.1 N
. IZ I
consumed F in
l moles/litre
kl in
litre/mole sec
0
24.00
0
720
20.85
0.26
1.05 X 104
1440
18.25
0.48
1.09 X 104
1980
16.90
0.59
1.06 X 10-4
2460
15.68
0.68
1.08 X 10-4
2880
19.85
0.76
1.07 X 104
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 CPYRGHT
895 1. 1. de long and 1. de Jonge,
as reported in literature 1) 1) was not found in the foregoing
investigation the present authors were led to suppose that the method
of analysis might be the cause of the differences in the results.
This was indeed found to be the case. The hydroxylamine method
as used by Smythe 1) the acidimetric sulphite method of Bettelheim
and Cedwall2) and the iodometric sulphite method as used by the
present authors were applied simultaneously in the same run (the
directions given by the various authors were followed as closely as
possible).
Initial concentration of' urea and formaldehyde 4 moles/litre;
temperature 30? C, pH = 7. The reaction mixture contained 0.0013
moles/litre K2HPO4 and 0.0009 moles/litre Na2HPO4.
I ml samples were analyzed. The results are given below:
time in
seconds
concentration of free F in i reaction rate
motes'titre using the method constant k in
of S. B. and C. de 1. and de 1. litre/mole sec.
0 4.0 4.0 40 0.7 X 10 i
300 ' 2.43 3.00 3.70 0.7 X 10 4
720 2.05 2.86 3.30 0.7 X 10 S
1620 1.76 2.36 2.75 ' 0.7 X 10"-4
It may be seen that the hydroxylamine method and the acidimetric
sulphite method indeed suggest a rapid initial reaction with a
consumption of about 40 ;o and 25 % respectively, of the initial
amount of F, within the first five minutes. The iodometric sulphite
method appears to lead to a normal bimolecular reaction from the very
beginning.
As all these methods are known to give reliable results in the
determination of F in aqueous solution the reason for the differences
in the results must originate from the proporties of the U-F system
investigated.
From previous studies of the U-F system 4) 5) it may be taken that
the velocities of formation and dissociation of LIF and LIF2 are
subject to catalysis by hydrogen and by hydroxyl ions. Fig. 2 gives
the influence of pH on the rate constant of the reaction between U
and F at 35? C, clearly illustrating that at a pH higher than 9 and
lower than 4, the reaction becomes very rapid 4). It follows that
during the analytical procedure a high alkalinity or acidity must be
avoided in order to exclude the possibility that during the analysis of
a sample of F the amount of free F will decrease or increase owing
to formation or dissociation of LIF. It was found that the iodometric
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The reaction between urea and formaldehyde, etc. 71 (1952) RECU81L 896
pH7
0,24 6 e 10 12.
Fig. 2.
sulphite method as applied by the present authors gave good results
for F in the. presence of, U or UF (see below).
Following the acidimetric sulphite method to determine formalde-
hyde an excess of sulphite is added. to the sample and after one
minute the alkali liberated is titrated with HCI. A rough calculation
(see below) shows that the high alkalinity during one minute may
result in a considerable -decrease in free F due to reaction with U.
(The reason that Bettelheim and Cedwall obtained rather reproducible
results may be due to a careful standardisaton of the procedure.)
A similar complicaton is encountered when using the hydroxyl-
amine method. During the analysis the solution becomes strongly acid
through HCl liberated from NH2OH . HC1 by reaction with P. This
acidity appears to be sufficient to explain the disappearance of a
considerable portion of free F during the analysis.
Some more details as to the analytical methods and the calculations
are given below.
Comparison of the analytical methods.
a. The iodometric sulphite method.
A 1 ml or 2 ml sample from the reaction vessel is diluted to 10 ml with ice
water, 3 ml of the diluted solution, containing about 1.2 mmoles F, is made alkaline
with a few drops of 1 M Na2CO3 (pH = 10, blue colour of thymolphthalein) and
then mixed with 2-3 ml sulphite and 5-10 ice water: Indicator: three drops of a
0.1 % alcoholic thymolphthalein solution. Immediately the resulting deep blue colour
is reduced to a faint blue with 0.5 N HCI (pH 8.5). After being left for 7. minutes
the solution is acidified to about pH = 4 with 5 ml molar AcOH and the excess
of sulphite is exactly removed with iodine solution using starch as an indicator.
The formaldehyde bisulphite addition compound is decomposed with about 25 mI
2'N . Na2CO3 (pH 9-10) and the liberated sulphite is titrated with 0.1 1V I2 till
the blue colour of. starch is persistent. Only the normality of the iodine solution
has to be known, When during the 7 minutes waiting the faint blue colour intensifies
again (through the formation of alkali) the pH has to be readjusted with HCI,
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 CPYRGHT
Ap
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897 1, 1. de long and J. de Jonge,
The critical point in this analysis is the moment when the solution becomes
strongly alkaline through the formation of alkali by the reaction between F and
sulphite:
CH1O + Na5SO5 + H,O -- CH,OH. SO,Na + NaOH.
In strongly alkaline solution the OF formaton (k,) and decomposition (k,) become
rapid, as is clear from the following equations 4) holding at 35? C:
k,=0.56X10-4+ 0.14[HiI+1.7[OH'] and k,= 0.036 k,.
However, by analyzing free F only in diluted samples and immediately removing
excess alkali, an accurate determination of F in the presence of U and UP appears
to be possible, as is shown by the following experiments.
1. Determination of P in the presence of U:
a) 28 mmoles F and 83 mmoles U were mixed at pH = 7 in 500 ml water at
25? C. A 20 ml sample containing about 1.1 mmoles F was taken two minutes
after mixing.
The sample consumed 21.34 ml 0.1044 N I,. Calculated: 21.4 ml.
Checked with an alkaline iodine solution: 21.40 ml 0.1044 N I.
b) 20 mmoles U and 2.26 mmoles F and 5 ml molair sulphite were mixed with
20 ml ice water and Immediately neutralized on thymolphthalcin with HCI, and so on.
Found: 45.33 ml 0.0996 N I,,
Calculated and experimentally checked, using an alkaline iodine solution:
45.37 ml 0.0996 N I.
However, when waiting five minutes before neutralisation we found: 43.40 ml
0.0996 N I,! This clearly Indicates the critical point of the method.
2) Determination of F in the presence of UF:
I mmol UP, 1.13 mmoles F, I mmole U and 3 ml molair sulphite were dissolved
in 10 ml ice water and analysed for free F.
Found: 22,50 ml 0.1000 N I,
Calc.: 22.64 ml 0.1000 N I.
b. The acidimetric sulphite method.
An excess of sulphite solution is added, to the sample to be analyzed (containing
U, F and UP). After one minute the alkali liberated is titrated with HCI.
The alkali liberated, when adding sulphite, increases the reaction velocity
between U and F enormously, as is shown by the following rough calculation:
To I ml solution, originally containing 4 m moles U and 4 mmoles F, 8 ml molair
sulphite is added. In the resulting solution the U and F concentrations are about
0.4 molar. When half of the F has reacted with sulphite the OH- concentration
has risen to 0.2 N. At that moment k1 (at 35? C) becomes: 1.7 X 0.2 litre/moles
sec. At room temperature (20? C) k, will be roughly 0.7 X 0.2.
Consequently per second 0.7 X 0.2 X 0.2 X 0.4 = 0.01 moles/litre P will be
converted into UF This explains the deviating results obtained by this method,
c, The hydroxylamine method.
During this analysis the solution becomes strongly add through HCI liberated
by F from NH2OH . HCI. This acidity is sufficient to explain the disappearance
of a considerable portion of the free F through reaction with U during the analysis,
as is shown by the following experiment.
To 23 mmoles U dissolved in 20 ml 10 % NH1OH . FICA solution, 23 mmoles F
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The reaction between urea and formaldehyde, etc. 71 (1952) RECU91L 89!(
(2 ml) were added. After 30 seconds the solution was titrated rapidly (in 1-2
minutes) with 1.23 N NaOH, using bromo-phenol blue as an indicator till the
colour of the solution became purple. For a set of titrations the results were:
11.7, 11.9, 11.2 and 11.7 ml NaOH. Calculated 18.7 ml NaOH.
The latter value was found experimentally for a solution of F without U titrated
with NaOH 20 minutes after the addition of NHZOH . HCI. When the solution
was titrated 30 seconds after the addition of NHZOH . HC1 the values were 3 %
too low,
Consequently, the presence of urea leads to erroneous results when applying the
hydroxylamine method for the determination of formaldehyde. In that case the
reproducibility is rather bad too.
.The following rough calculation illustrates how the method is influenced by the
high acidity and high concentrations: When half of the F has reacted with
NH,OH. HCI in the experiment just given the H+-ion concentration has increased
to 0.5 N. At 35? kl can be calculated to be 0.14 X 0.5 = 0.07 litre/moles second.
At 20? k, will be about 0.03 litre/moles second.
The result will be that then per second 0.03 X I X 0.5 = 0.015 moles/litre F
will be converted into UF.
It will be clear that this.. velocity is sufficient to make the reaction between U
and F fast enough to compete successfully with the analogous reaction between F
and NHQOH. Moreover, in strongly acid solutions the latter reaction is retarded.
Eindhoven, August 1951.
(Received March 4th 195,0.
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