Iodine-hydrogen reaction

The reaction with fluorine occurs spontaneously and explosively, even in the dark at low temperatures. This hydrogen—fluorine reaction is of interest in rocket propellant systems (99—102) (see Explosives and propellants, propellants). The reactions with chlorine and bromine are radical-chain reactions initiated by heat or radiation (103—105). The hydrogen-iodine reaction can be carried out thermally or catalyticaHy (106). 

Hydrogen-iodine reaction, 13 770 Hydrogen-ion activity, 14 23-34 nonaqueous solvents, 14 32 pH determination, 14 24-27 pH measurement systems, 14 27-31 Hydrogen ion concentration (total acidity), 14 23... 

The most frequently occurring simultaneous, interdependent reaction mechanism is the case in which the product, as its concentration is increased, begins to dissociate into the reactants. The classical example is the hydrogen-iodine reaction ... 

For many years the hydrogen-iodine reaction was quoted in textbooks as being virtually the only known example of a simple bimolecular reaction. There is now evidence 43 that in parallel to the main bimolecular transformation, some additional reactions involving iodine atoms do occur. 

Whereas in the hydrogen-iodine reaction, atomic iodine plays only a minor part, in the reaction between hydrogen and bromine, bromine and hydrogen atoms are the principal intermediates in the overall transformation. 

The kinetics of the reaction are quite different from those of the hydrogen-iodine reaction although the stoichiometric equation ... 

A classical example of a process that follows the mechanism of (21) is the hydrogen-iodine reaction H2 + l2 2HI. If = 0 and = = Cq at t = 0 for this bimolecular process, then equation (22) becomes... 

The hydrogen-iodine reaction is a classic in chemical kinetics. The work of Bodenstein on this reaction is one of the first systematic studies of the temperature dependence of reaction rates. For many years the formation of HI from H2 and I2 was regarded as the textbook example of a bimolecular four-center reaction as was its reverse. Recently, however, experimental results inconsistent with this interpretation have been obtained. ... 

While this work is requiring the revision of many textbooks which have used the hydrogen-iodine reaction as a classic example of a bimolecular reaction, it has also aroused interest in its implications for absolute reaction rate theory. Noyes has suggested that the results present a paradox in kinetics. In his discussion, he suggests that absolute rate theory as normally formulated fails to account for momentum effects which in some cases, namely the H2-I2 reaction, place severe restrictions on the path leading from reactants to products. In the... 

The hydrogen-bromine reaction is a classic example of a free radical reaction having been shown by Bodenstein et al. to have a much different kinetic behavior from the hydrogen-iodine reaction. Kassel has reviewed critically the early work. Other more recent reviews have also appeared ... 

For many years the reaction was thought to be a prime example of an elementary reaction. However, later work by J. H. Sullivan in 1967 (entitled. Mechanism of the bimolecular hydrogen-iodine reaction) showed that iodine atoms were involved in the reaction so that it must be composite. A two-step reaction mechanism can be proposed, in which molecular iodine dissociates, and the iodine atoms then react with a hydrogen molecule... 

For many years the hydrogen-iodine reaction had been the traditional example of opposing second-order reactions. Recent work by J. H. Sullivan indicates that the mechanism is not as simple as we have assumed here in fact, the mechanism now seems to be unresolved. For a discussion and references see R. M. Noyes, J.Chem. Phys. 48, 323 (1968). 

The kinetic law for the hydrogen-bromine reaction is considerably more complicated than that for the hydrogen-iodine reaction. The stoichiometry is the same. 

The values for (AA°) in Table 2.5e are obtained from pre-exponential factor measurements, values for p determined from equation (2-33a) or (2-34a), and the estimation procedure suggested in Illustration 2.5. In general, these are not elementary steps, although for many years it was believed that the hydrogen/iodine reaction was a true bimolecular reaction, since both collision theory and TST estimates of pre-exponential factors were in good agreement with experiment. This view, however, has changed (J.H. Sullivan, J. Chem. Phys., 46, 73 (1967)]. ... 

Detailed studies have been made of many different examples of equilibria. We look at two cases. In case 1, the hydrogen-iodine reaction, we conhrm that fc(r) is independent of reactant concentrations. In case 2, an esterihcation reaction, we conhrm that the equilibrium concentrations of reactants and products (the equilibrium composition ) depends upon their starting concentrations. 

Most reactions are not elementary. The hydrogen-iodine reaction looks like an elementary bimolecular reaction it is first order in H2, first order in I2, and second order overall. However, this reaction is not elementary [425]. Reactions that behave kinetically just like a single-step, elementary reaction are called kinetically simple. The order of each reactant in such a reaction is equal to its stoichiometric coefficients. Such reactions are also said to obey the law of mass action or to have mass-action kinetics. 

Sullivan, J. H. (1967). "Mechanism of the bimolecular hydrogen-iodine reaction." J. Chem. Phys. 46, 73-78. 

In 1894 Max Bodenstein (1871-1942), at the University of Heidelberg, published a landmark paper [1] which has played an important role in the development of gas-phase chemical kinetics. In these investigations [1,2], he reported the rate measurements for the text-book example of the hydrogen-iodine reaction... 

Max Bodenstein opened the field of gas phase chemical kinetics in 1894 with his report of experimental studies of the hydrogen-iodine reaction H2 + I2 HI + HI and its reverse HI + HI H2 + l2- Bodenstein measured the rates of the forward and reverse reactions, their equilibria, and their temperature dependence. He found second order kinetic expressions and an Arrhenius temperature dependence for the rate constants. He suggested several mechanisms for these reactions. 

Max Bodenstein opened the field of gas phase chemical kinetics at Heidelberg one hundred years ago with the publication of his landmark paper [1] reporting experimental measurements of the rates of the hydrogen-iodine reaction, H2 + I2 —> HI 4- HI, and its reverse, HI + HI —> H2 + l2- In this first systematic study of the kinetics of chemical reactions in the gas phase he determined the eflFects of reactant concentrations and temperature for both reactions. He found that the reactions follow overall second-order kinetic expressions, that their rate constants have an Arrhenius temperature dependence, and that the equilibrium constant is given by the ratio of forward to reverse rate constants. 

Despite Bodenstein s considerations of alternative mechanisms the forward and reverse reactions were generally assumed to occur as direct bimolecular reactions and these reactions soon became textbook examples of bimolecular reactions. Lewis [3] used Bodenstein s data to demonstrate the applicability of Arrhenius concept of active molecules and the Arrhenius rate expression. Hin-shelwood [4] cited the hydrogen-iodine reaction and its reverse as evidence for... 

Figure 2 Symmetric trapezoidal and collinear configuations lying on prospective pathways for the hydrogen-iodine reaction. From Anderson [21].

Our recent electronic structure calculations 3deld a potential energy surface adequate to explain, at least qualitatively and within the uncertainties due to an incomplete knowledge of relaxation rates, the available experimental observations for the hydrogen-iodine reaction. The rate expressions, the rate constants, their temperature dependence, the vibrational excitation of HI products, the excitation and/or dissociation of reactant I2, the photochemical rates - all are compatible with the recent ab initio potential energy surface and with the classical trajectory calculations carried out with a similar surface. And all are compatible with either the bimolecular or termolecular mechanisms. It appears most likely that both mechanisms contribute, but the matter is not resolved as yet. 

The hydrogen-iodine reaction was mentioned in the previous chapter ... 

Sometimes the process of interest involves a series of collisions, and thus the initial conditions are obtained from the final conditions of trajectories. The earliest use of this approach was in a study of the termolecular hydrogen-iodine reaction. The termolecular process was treated as a two-step process. Trajectories were computed for collisions of I - - H2 to determine the initial conditions of the second step in which the complex [IH2] collided with the second iodine atom to form 2HI. More recently, this kind of approach has been used for sequential collisional energy transfer. This is discussed in more detail in section 4. 

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