R. J. ATKINSON, A. M. POSNER, AND J. P. QUIRK
550
Adsorption of Potential-Determining Ions at the Ferric Oxide-Aqueous Electrolyte Interface
by R. J. Atkinson, A. M. Posner, and J. P. Quirk Department of Soil Science and Plant 'l',trz%on, Institute of Agriculture, University of Western Australia, Nedlands, Western Australia (Receired July 6 , 1966)
Potentiometric titrations in indifferent aqueous electrolyte were used to memure the surface charge characteristics of microcrystalline particles of ferric oxides. The net adsorption density (I"+ - I'oH- microequivalents per gram) is described by equations for the surface excesses of H+ and OH-, derived by equating the electrochemical potentials of potential-determining ions and indifferent counterions in solution and surface phases, and neglecting counterions in the diffuse layer. The equations contain interaction constants arising from the small but finite distances of separation between planes of adsorbed potential-determining ions and indifferent counterions under conditions of ferric ion solvation at the interface. The BET nitrogen area does not satisfactorily represent the dependence of parameters and interaction constants on particle surface area, for different preparations. The equations were interpreted as showing that adsorbed counterions remain solvated, forming electrostatically bound ion pairs with only a small proportion dissociating into the diffuse layer.
Solid oxide particles immersed in aqueous electroor liquid mercury in contact with aqueous electrolyte.3-5 lyte solutions develop surface electrical charges by adQuantitative interpretation has not been attempted, and in previous publications the ferric oxide system has sorption or desorption of potential-determining ions. For ferric these ions are H + and OH-, and been characterized only by measurements of the pH a t the zero point of charge. ferric hydroxo complex ions derived from the solid phase by dissolution, when any ions which are specifiIn this paper, equations to describe potential-determining ion binding are derived for ferric oxides in cally adsorbed a t the oxide surface are absent. The contribution made to surface charge by transfer of aqueous suspension under conditions of high ionic strength of indifferent electrolyte. The form of analymetal hydroxo complex ions across the interface is sis was suggested by the ideas of Gilbert and RideaP usually considered to be very small in relation to H + on the titration of fibrous proteins. Two simple and OH- ion transfer, since solubility data3-6 show that methods of eliminating electrical potential terms from either H+ or OH- ions always greatly exceed ferric ion-binding equations are examined, and it is found that hydroxo complex ion in solution concentration for the the more satisfactory method does not require one to pH range 3.5 to 11. Surface charge measurement thus requires only adsorption-desorption measurements for H f , which are readily carried out by the direct (1) G. A. Parks, Chem. Rev., 65, 177 (1965). method of potentiometric titration of microcrystalline (2) P. G. Johansen and A. S. Buchanan, Australian J . Chem., 10, particles of oxide in dilute aqueous suspension using 392 (1957). various ionic strengths of an indifferent e l e ~ t r o l y t e . ~ ~(3) ~ G. A. Parks and P. L. de Bruyn, J . Phys. Chem., 66, 967 (1962). (4) A. E. Albrethson, Ph.D. Thesis, Massachusetts Institute of From such measurements, it is known that ferric oxideTechnology, 1963. aqueous electrolyte interfaces have important differ(5) G.Y.Onoda and P. L. de Bruyn, Surface Sci., 4,48 (1966). ences in double-layer structure and the magnitude of (6)G. A. Gilbert and E. K. Rideal, Proc. Roy. SOC.(London), A182, electrical properties when compared with AgI, Ag$, 335 (1944). ~~
The Journal of Physical Chemistry
ADSORPTION AT THE FERRIC OXIDE-AQUEOUS ELECTROLYTE INTERFACE
551
Table I
Goethite (27.7.65)
Hematite ((1)22.7.64)
70.9 x 104 20.3 f 0 . 4
43.5
Test,. concn. of oxide, g in 25 ml Washing method Initial charge before titration, pequiv g-1 PH, ZPC Surface density of charge at p H 4 in 1 M KCl, pequiv cm-2
0.1376
kH+VH+
BET Ns area, cmz g-l Temp of measurement,
-
Ferric oxide
r
x
Hematite (24.3.65)
Hematite (2)
Hematite (10.2.65)
34.1
0,2728
0.1576
0.1375
0.0949
Dialysis 4 i 1.0
KOH 7&2
Dialysis 85 f 2
KOH 12 i 3
Dialysis 72 i 4
KOH
7.55i0.15 1.41 x
8.60f0.2 1.63 X
9.27i0.10 3.68 X
8.45i0.2 4.22 x
8.90&0.15 4.89 x lo-'
8.70 7.25
0.527 X lo6 0.74 0.03647 2.59 x 104 0.047 3 x 104 KC1
0.37 X lo6 0.85 0.05387 2.34 x 104
3.40 X 106 7.62 0.03089 1.38 x 104 0.0165 0.9 x 104 KC1
0.889 X 106 2.61 0.02596 0.89 x 104
1.17 X lo6 3.21 0.02081 0.76 x 104 0.011 0 . 4 x 104 KCl
2.26 X lo6 9.83 0.02765 0.64 x 104
104
...
36.4 x 104 20 i 2
"C
ka+VH+/A K1, pequiv-' g KIA K2, pequiv-1 g K2A Indifferent electrolyte used
Hematite (Albrethson)
44.6 x 104 19.5 f 0 . 6
104
...
x
7
... ...
KNOs
account for the small proportion of counterions which dissociate from the surface as free ions.
Experimental Materials and Methods (i) Preparation of Ferric Oxides. Microcrystalline precipitates of hematite cu-FezO~and goethite a-FeOOH were prepared from ferric nitrate of at least 98% purity manufactured by May and Baker Ltd., Dagenham, England. Solutions containing 100 g of Fe(NO&. 9Hz0/1. were prepared at room temperature and boiled a t 100" for 18 days under reflux conditions.3 Final pH values were between 0 and 1. Goethite was prepared by the addition of 200 ml of 2.5 N KOH to 50 g of Fe(NOa)3.9H~0 in 825 ml of double-distilled water to give a final pH near 12, followed by aging for 24 hr in a 60" oven. Pyrex glass vessels were used, and no precautions were taken to prevent silicate contamination, which is possible a t high pH. Goethite and hematite precipitates were dialyzed in cellulose tubing using double-distilled water changed twice daily until NOa- concentrations were M. Some hematite precipitates were washed by repeated centrifugation and decantation using KOH to raise the pH t o 9 for the initial washes (Table I) ; at this pH the oxide was flocculated. Suspensions contained 0.02 to 0.03 g of ferric oxide/ml and were stored in polythene bottles from which 5-ml samples were pipetted as required for titrations or weight delivery determinations after 24 hr drying at 110". X-Ray powder diffractom-
... ...
KNO8
23 25
x
104
Hematite (Parks and de Bruyn)
22 21
x
...
*.. KOH
...
...
x
... ...
NaCl
104
8.5 10-2
..
1.32 X lo6 6.00 0.02405 0.53 x 104 0.043 0.9 x 104 KNO8
eter traces for hematites generally showed a slight goethite peak at 4.18-A spacing. This is the most intense goethite peak,? and its presence could be due to a goethite-like surface remaining on hematite when dried a t 110". Hematite was not detected in goethite preparations. Surface area measurements were made by application of the BET equation to conventional volumetric data for nitrogen adsorption on compressed cores of ferric oxide at -195" after 100" outgassing. Preliminary measurements showed that K+, C1-, or NO3- were not adsorbed at the zero point of charge (zpc), and that KC1 and KNO3 were equally indifferent. (ii) Surface Charges in the Absence of Specijfc Adsorption. The method used was essentially that of Parks and de Bruyn.8 Titration data were obtained using recording potentiometric titration apparatus (Radiometer, Copenhagen) with Radiometer G202B and K401 glass and calomel electrodes. Polythene titration cells were used containing 0.1 to 0.15 g of ferric oxide in 25.0 ml of potassium chloride solution mixed with a Teflon stirrer and flushed with wet COzfree Nz gas. Blank titrations were similarly carried out using 25.0 ml of the appropriate KC1 solution. Uptake of H+ or OH- by the oxide was found by taking differences between test suspension and blank (7) H. P. Rooksby in "The X-Ray Identification and Crystal Structures of Clay Minerals," G. Brown, Ed., Mineralogical Society, London, 1961, Chapter 10.
Volume 71, Number S February 1057
R. J. ATKINSON,A. M. POSNER, AND J. P. QUIRK
552
titration volumes at particular pH values. Adsorption densities ( r H + - r O H - ) microequivalents per gram of oxide (110" dry weight) were calculated relative to the zpc located by the common point of intersection of titration curves a t different ionic strengths. The temperature for titrations was 20 i 1.5". The titrants used were 0.1 N HC1 and 0.1 N KOH which were standardized against 0.005 M potassium hydrogen phthalate and 0.001 M borax NazB407.10HzO; all solutions were prepared with double-distilled water. Additions of titrant did not exceed 0.25 ml in the pH range 3.5 to 10.5, so that ionic strengths were kept essentially constant except at the lowest ionic strength, 0.002 M . Each titration was carried out twice using two different speeds which corresponded approximately to 2.5 and 5 pH units/hr. Good agreement (d=3%) was found between these replicate titrations. The uncertainty in pH measurements was iO.05, although absolute errors were greater at pH extremes. Preliminary experiments showed that suspension effects were not significant and liquid junction potentials were taken as being close to zero. The ferric oxides used in the adsorption experiments reported here were not subjected to any drying treatments.
Results and Discussion Criteria for attainment of equilibria in reactions at hydrated solid surfaces are often difficult to establish. After the initially rapid adsorption or desorption of Hf occurring with pH change, ferric oxide surfaces continue to adsorb or desorb protons very slowly. The slow H f transfer has been attributed to the extension of the proton-excess charge (or proton-deficit charge) into hydrated surface layers of the oxide,6 or may arise from structural rearrangement of ferric hydroxo complex ions existing at the surface, to expose effectively a greater number of proton acceptor or proton donor sites at the interface. The slow H+ transfer in any event involves reactions different in form t o the initial rapid H + transfer occurring when solution pH is changed, and accordingly only the initial fast reaction at the phase boundary is considered here. Other experiments are being conducted in these laboratories to gain information on the structural bases for differences in ferric oxide adsorption capacities for H+, which cannot be simply related to BET nitrogen surface area or known differences in preparation methods. Preliminary experiments on the adsorption behavior of dried (110') finely ground hematite indicated that H + adsorption or desorption from solution was greatly reduced with respect to quantity and rate, while phosphate adsorption was only slightly depressed. This suggested that the degree of crystal dispersion was not The Journal of Phv.rical Chemistry
the important factor, and instead the possibility that slow insertion of water molecules into surface crystalline layers is a rate-determining step for dried ferric oxides is being investigated. The quantity obtained from potentiometric titration measurements is the net adsorption, r H + - FoH-, in microequivalents per gram. Thus u =
1 A
-(rH+
- roH-) pequiv cmW2
(1)
where u is the surface density of charge, r H + and r O H are the surface excesses, in microequivalents per gram, and A is the surface area in square centimeters per gram. I n the presence of an excess of indifferent electrolyte, the adsorbed potential-determining ions are assigned to the solid side of the i n t e r f a ~ e ;that ~ is, the only counterions present in the solution phase are solvated C1- (or K + if the surface is negatively charged). Two features of the ferric oxide titration curves (Figures 1 and 2 ) suggest a form for theoretical analyses. (i) The titration curves may be regarded as adsorption isotherms for H + or OH- by use of the approximations U A
=
(rH+
-
roH-)
'v r H +
on the acid side of the zpc (2) a-4 =
(rH+
- roH-)
'v --OH-
on the alkaline side of the zpc (3) Positive charges on the surface are thus considered to be due entirely to excess adsorbed H + with respect to the zpc. Negative charges are considered to arise entirely from a surface excess of OH-. (ii) The uptake of HC1 or KOH does not appear to reach any maximum value as pH is decreased or increased, respectively. Dissolution of the solid is not thought to be appreciable until pH (CI-) - F($H- h i )
(14)
Activity coefficients have been omitted in view of the approximation made in relation to use of the potential difference ($H - $cl).ll Equation 14 may be rewritten in the form
FIGURE
3A
20
60
,
o
40
ao 100
120
c+-c,,-
1
1
1
140 360 IBO 200
pes
per 9.
Figure 3. Hematite no. 24.3.65. Figure 3A: plots of eq 10. Figure 3B: the reciprocal slopes of the straight lines in 3A plotted BS a function of dl(omitting activity coefficients).
The Journal of Physical Chemistry
(8) P. J. Anderson, Trans. Faraday SOC.,54, 130 (1958). (9) J. T. Davies, Proc. Roy. SOC.(London), A245, 417 (1958). (10) J. Th. G.Overbeek in “Colloid Science,” Vol. I, H. R. Kruyt, Ed., Amsterdam, 1952, Chapter IV. (11) E. A. Guggenheim, “Thermodynamics,” North-Holland Publishing Co., Amsterdam, 1959, Chapter 9.
ADSORPTION AT THE FERRIC OXIDE-AQUEOUS ELECTROLYTE INTERFACE
in which kHt =
[
exp -'AmR;
Arc"]
4'81
5.0 ~
Two further approximations are then made. If V H t , the maximum number of sites available for H+, is large compared with r H + , then B H / ~ - OH cv rHt/VHt. Secondly, it is assumed that r H t , the surface excess of H+, is a linear function of the potential difference ($H - $GI). The simplest approximation that +x = +GI? which was used by Gilbert and Rideal,6 does not prove useful in this case. The equation to be tested is r H t
555
A A
\
(16)
= kHtVH+d(H+)(Cl-) eXp(-KlrHt)
in which K1 is an interaction constant which is inversely proportional to surface area A . Dimensions are microequivalents-l gram for K1, and moles liter-' for kHc. The linear form for eq 16 is IOg
rH+
+ '/2pH
=
log
+
~H+VH+
3.4
KirH t log (Cl-) - - (17) 2.303
"2
which is tested in Figures 4 and 5 for hematite no.
......
I
.
.
.
,
.
.
.
.
.
I
20
1
1
I
I
40
60
80
xx)
TH.
I
120
I
140
pes perg.
Figure 5. Goethite no. 27.7.65. Plots of eq 17 on the acid side of the zpc.
.
5.4
24.3.65 and goethite no. 27.7.65, having approximately obtained r H t from r H t - r O H - data (Figures 1 and 2). The intercepts log k H tVH t l / 2 log (Cl-) should form a straight line of unit slope when plotted as a function of '/z log (Cl-) (see Figure 6). Both requirements appear to be satisfied. Some values of the constants K1 and k ~ t V H are t given in Table I. Recalculation of isotherms requires constants for the corresponding equation on the alkaline side of the zpc,
+
5.2
5.0 4.8
Ip 4.6
ViZ.
44
roH- = ~ o H - V O H - ~ ( K + ) ( O H eXp(-KzroH-) -)
F~~
+
i . E'
t
0
(18)
in which rOH-, the magnitude of the surface excess of OH-, is small compared with VoH-, the maximum number of sites for OH-. Other symbols are analogous to those in eq 16. The linear form is
4.2
40
IOg
rOH-
- '/zpH
'/z log Kw
+
3.8
3.6 .
0
20
40
60
80
1
,
,
*
*
,
*
'
.
xx) 120 140 l60 180
L+ peq per g.
Figure 4. Hematite no. 24.3.65. Plots of eq 17 on the acid side of the apc.
in which Kw is the appropriate dissociation constant for water. However, the approximations made to evaluate constants on the acid side of the zpc are not suitable for the restricted range of r O H - values on the alkaline side Volume 71, Number 8 February 1967
R. J. ATKINSON, A. M. POSNER, AND J. P. QUIRK
556
-
4.4
60
t
1 -1.5
-1.35
-1.0
- 0.5
0
112 log(C13
+
Figure 6. Plots of the intercepts log ~ H V H I/z log (Cl-) as a function of '/z log (Cl-) (see eq 17): 0,hematite hematite, Albrethsen, ref 4; 0 , hematite no. no. 24.3.65; 10.2.65; A, hematite no. 2; A, goethite no. 27.7.65; 0 , hematite no. 22.7.64. The dashed line gives the theoretical slope.
+,
4
5
6
7
8
9
10
11
PH Figure 7. Recalculation of adsorption density as a function of pH and molar concentration of KCl, for hematite no. 24.3.65. Dashed curves are the original experimental data. Solid curves have been calculated from eq 17 and 19 using the parameters: K1 = 0.0309; K Z = 0.0165; log ~ H V H = 6.53; '/z log K v f log ~ O H V O=H -2.95. 8.9 rOH-9 1.0 M KC1 (calcd); b, r H + - rOH-9 1.0 M KC1 (exptl); FOB-, 1.0 M KC1 (calcd); d, r H + , 1.0 M KC1 c, r H + (calcd); e, FOE-, 0.01 M KC1 (calcd); f, r H + - roH-, 0.01 kf KC1 (exptl); g, I"+- FOE-, 0.01 M KC1 (calcd); h, r H f , 0.01 M KC1 (calcd).
-
of the zpc. rOH- could not be satisfaGtorily approximated by the measured (rHt - roH-). The method chosen to obtain approximations for ~ o H - V O H - and K z was to plot log lI"+- rOH-1 - l/zpH as a function of IrHt - rOH-1 and fit straight lines having the intercepts predicted from the empirical equation for the acid side and the zero point of charge pH. Some values for Kz are given in Table I. The form of calculated isotherms is shown in Figure 7. These were obtained by solving the equations for pH, using the appropriate constants, at various values of rEtor roH-,and then obtaining the (rHt - rOH-) curve by graphical interpolation. Agreement between calculated isotherms and experiment (Figure 7) may be improved by introducing mean ionic activity coefficients for KC1I2 when evaluating the constant k ~ + v from ~ t the intercepts log kHtVHt '/z log (Cl-). For the present purposes, isotherms have been recalculated using arithmetic mean values of k H + V H +obtained without applying an activity coefficient to '/2 log (GI-). Under these cir-
+
The Journal of Physical Chemistry
cumstances, the agreement between recalculated isotherms and experiment is considered to be satisfactory. Model of the Ferric Oxide-Aqueous Electrolyte Interface. The assumption that rHt is a linear function of a potential difference ($H - $GI) made in obtaining eq 16 implies that the capacity of the double layer is a constant. Use of the relationship between Q and 6 the thickness of a Stern layer's
where E' is the dielectric constant within the Stern layer, should enable approximate calculation of the (12) R. A. Robinson and R. H. Stokes, "Electrolyte Solutions," 2nd ed, Butterworth and Co.Ltd., London, 1959. (13) J. T. Davies and E. K. Rideal, "Interfacial Phenomena," Academic Press, London, 1961, p 89.
ADSORPTION AT
THE
FERRIC OXIDE-AQUEOUS ELECTROLYTE INTERFACE
interaction constants K1 and K2. $0 is identified with $a with and l?H+/Awith u. The expression obtained on evaluating constants at 293OK is
4 T ~ 6 2.16 x 1056 K=--pequiv-’ g (21) 2RTe‘A €‘A A is expressed in square centimeters per gram and 6 is expressed in Angstrom units. 6 is the average distance, in the direction normal to the surface plane, between the plane of surface charge and the plane of average approach of the centers of solvated counterions. Values for 6 and e‘ are not known; however, the correct magnitude for K is obtained on taking reasonable assumed values. For example, calculation using 6 = 0.5 A and a va1ue14 of e‘ = 6 for water at dielectric saturation, leads to K1 = 0.04 for hematite no. 24.3.65 of B E T nitrogen area 44.5 X lo4 em2 g-l, compared with the experimental value K1 = 0.03. It is possible that both e’ and 6 decrease with increasing surface charge while their ratio remains approximately constant. The hydrated ferric oxide surface is likely to have a sufficiently irregular configuration to permit small values of 6, of the order of 0.2 to 2 A, by “embedding” of the solvated counterions. The variation in values of the interaction constants K (Table I) for different hematite preparations may be understood qualitatively by consideration of the structural nature of the surface. I n the bulk of hematite crystals, each ferric ion is surrounded by a distorted octahedral arrangement of oxygen ions, but in the surface phase of crystals in aqueous suspension the coordinated species may be oxygen ions, hydroxyl ions, or water molecules. Ferric ions at the surface are partially solvated and the surface phase consists of polymerized derivatives of ferric hydroxo complex ions with coordinated water molecules and hydroxyl ions behaving as proton donors and acceptors. Although protons are probably relatively mobile within the surface phase, extension of the protonexcess (or proton-deficit) charge into the crystal lattice proper probably involves only a small fraction of the excess protons and would be limited to small distances of the order of 2 to 5 A. Little information is available on the effect of drying,15 but the surface phase found on precipitated hematite crystals is probably lost by mild drying to p / p , = 0.95. I n the preparation methods used here for hematite, the final stage of crystal growth occurs when the pH of the refluxed suspension is increased during washing to remove HK03, causing precipitation of the high solution content of ferric ions which exists at very acid pH. Washing by dialysis is thought to complete crystal growth under conditions of high surface charge
557
with slow attachment of ferric ions of the form Fe (OH)(HzO)52+or Fe(OH)2(Hz0)4+,resulting in a wellordered surface. Poorly ordered surfaces with relatively greater inclusion of water molecules result from KOH washing which completes crystal growth by rapid attachment of polymerized derivatives of ferric ions under conditions of slight surface charge. These considerations suggest the average distance of closest approach of counterions is likely to vary for different ferric oxide preparations. This factor is proposed as the cause of differences in values of the interaction constant K which remain after multiplication by surface area A . That is, dAK is proportional to 6 considering a’ to be a characteristic constant for all systems and considering relative values of A to be represented by the BET nitrogen area. Hematite preparations washed by KOH would thus be expected to have smaller values for AK1 than hematites washed by distilled water dialysis only, as closer penetration of C1counterions could occur into a more disordered surface structure compared with surface structures in dialyzed preparations. However, in Table I the relationship between the interaction constant K and the washing method is not clear. This suggests that the suitability of the BET nitrogen area for representing relative surface areas should be examined. The similarity of the interaction constants K I and Kz for positive and negative charges is considered to reflect similar distances of approach for C1- and K+ counterions. Under the pH conditions of these experiments, chloride ions or nitrate ions do not enter the first coordination shells of ferric ions, although it is possible that water molecules are shared between Fe3+ coordination shells and C1- hydration spheres. The binding of solvated K+, C1-, or X03- counterions is largely electrostatic in origin and arises from the requirements for electroneutrality in the surface phase. Anions such as phosphate which are desolvated and specifically adsorbed by ferric oxides are considered to form a new potential determining layer by ligand exchange reactions with OH- or H20 in the ferric first coordination shell. I n view of the approximations necessary to evaluate the constants of eq 14 and 16, the agreement of these equations with experiment is considered satisfactory. The approximations made to eliminate the electrical potential terms appear to be suitable for the ferric oxide-aqueous electrolyte interface. Since the equations neglect the presence of counterions in the diffuse (14) J. O’M. Bockris, M . A. V. Devanathan, and K. Muller, Proc. Roy. SOC.(London), A274, 55 (1963). (15) J. J. Jurinak, J. Colloid Sci., 19, 447 (1964).
Volume 71,Number 9 February 1967
NORMAN COHENAND JULIAN HEICKLEN
558
layer or Gouy layer by equating OH and Ocl, it is suggested that under these conditions the proportion of counterions in the diffuse layer is small, and surface potentials are reduced as a result of ion binding between counterions and the surface.I6 Electrokinetic phenomena, determined by the small proportion of counterions in the diffuse layer, are not expected to give reliable calculations of amounts of adsorbed ions under the conditions of ionic strength used here. The considerable extent of ion-pair binding16 of counterions is suggested to be primarily related to the closeness of counterion approach which is possible under the conditions of ferric ion solvation at the interface. The electrostatic interaction is then large compared with the thermal energy of ions. An al-
ternative suggestione is that for large particles, of “microscopic” rather than “atomic” dimensions, only a very small fractional dissociation of counterions is required to give large electrostatic interactions relative to the thermal energies of the particle and counterions, when surface densities of charge are not low.
Acknowledgments. Scholarship assistance received by R. J. Atkinson from the Australian Agricultural Council and the Commonwealth Scientific and Industrial Research Organization is gratefully acknowledged. (16) D. G. Edwards, A. M. Posner, and J. P. Quirk, Trans. Faraday SOC.,61, 2808 (1965).
Reaction of NO(A2C+) with Carbon Dioxide
by Norman Cohen and Julian Heicklen Aerospace Corporation, El Seoundo, California (Received July 21, 1966)
I n the presence of COz, NO(A2Z+) was produced by irradiation of NO-COz mixtures with a cadmium arc. The products monitored were Nz,whose yield dropped as the (COz): (NO) ratio was enhanced, and CO, whose yield rose with the (COZ): (NO) ratio to a constant upper limit. Product formation was unaffected by temperature changes (23-300’) or by the addition of xenon or NOz. The results are explained by the simple competition NO(A2Z+) NO -+ Nz other products and NO(A2Z+) COZ+ NOz CO. Carbon dioxide is about 3.4 times more efficient than NO in the competition.
+
+
I. Introduction As part of a continuing program in our laboratory on the reactions of NO(A2Z+), we have examined its reaction with COz. Fluorescence quenching of the A2Z+ state of NO by COZhas been studied in a number of l a b o r a t o r i e ~ , ~and - ~ the quenching constant is reasonably well known. It is very large (-loll M-’ sec-9, and it has been suggested that chemical reaction may 0ccur.l However, as far as we know, no detailed study of product formation has been made prior to our work. The Journal of Physical Chemistry
+
+
II. Experimental Details The quartz reaction vessel was 10 cm long and 5 cm in diameter. It was jacketed in a wire-wound furnace (1) A. V. Kleinberg and A. N. Terenin, Dokl. Akad. Nauk S S S R , 101, 1031 (1955). (2) N. Basco, A. B. Callear, and R. G. W. Norrish, Proc. Roy. SOC. (London), A260, 459 (1961). (3) R. Young and R. Sharpless, Discussions Faraday SOC.,3 3 , 228 (1962). (4) A. B. Callear and I. W. M. Smith, Trans. Faraday SOC.,59, 1720 (1963). (5) H. Broida and T. Carrington, J . Chem. Phys., 3 8 , 136 (1963)
