DOCSLIB.ORG

Effects of Molecular Orientation on Electron-Transfer Collisions Peter W

Total Page：16

File Type：pdf, Size：1020Kb

Download full-text PDF Read full-text
Effects of Molecular Orientation on Electron-Transfer Collisions Peter W Subscriber access provided by RICE UNIV Effects of molecular orientation on electron-transfer collisions Peter W. Harland, Howard S. Carman Jr., Leon F. Phillips, and Philip R. Brooks J. Phys. Chem., 1991, 95 (21), 8137-8142 • DOI: 10.1021/j100174a023 Downloaded from http://pubs.acs.org on December 9, 2008 More About This Article The permalink http://dx.doi.org/10.1021/j100174a023 provides access to: • Links to articles and content related to this article • Copyright permission to reproduce figures and/or text from this article The Journal of Physical Chemistry is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 J. Phys. Chem. 1991,95,8137-8142 8137 Effects of Molecular Orlentatlon on Electron-Transfer Colllslons Peter W.Harland,' Howard S,Carman, Jr.,* Leon F. Phillips: and Philip R. Brooks* Department of Chemistry & Rice Quantum Institute, Rice University, Houston, Texas 77251 (Received: February 12. 1991) K+ ions have been detected from the intersection of a beam of K atoms (5-30 eV) with beams of various simple molecules, such as CH3Br and CF,Br, which had been oriented prior to the collision. Production of ions in the collision is found to be highly dependent on orientation. The effect is most pronounced near threshold (4eV) and almost disappears at higher (30 eV) energies. Attack at the "reactive' halogen end produces the most ions, regardless of the polarity of that end. For each molecule, the reactive end seems to have the lower threshold energy. These observations may be a result of the electron being transferred to a specific end of the molecule, but the experiments measure only the net result of an electron transfer followed by the separation of the ions. Whether or not electron jump per se depends on orientation is still an open question, but we are able to qualitatively interpret the experimental results as being due to interactions between the ions as they separate in the exit channel. Most of the negative molecular ions dissociate, ejecting a halogen X- in the direction of the (oriented) molecular axis. If the X end is oriented away from the incoming K atom, the ejected X- will travel in the same direction as the K+,making the electron more likely to return to the K+ ion and reducing the K+ signal in this unfavorable orientation. I. Introduction equal, step 1 was assumed to be independent of orientation. The "Chemical intuition", as well as a growing amount of direct dependence on orientation was assumed to be in step 2 because experimental observation, tells us that chemical reaction depends the I- would be ejected in the known direction of the molecular on the orientation of the molecules involved. By "orientation" we dipole. After we accounted for the momentum of all the partners, mean a spatial configuration where one end of a molecule can this mechanism qualitatively reproduced the experimental angular be distinguished from another. Although reactive species dif- distribution, not only for heads and tails orientations, but also for ferentiate between ends by not reacting with the "wrong" end, sideways orientations.' we have little experimental guidance for what constitutes the For the K + CF3Br reaction the angular distributions are in 'right" end because most experiments include all possible orien- qualitative accord with this harpoon model, but the reactive tations during collision and isolation of orientation effects has not probability is less at the tails (CF,) end, leading us to speculate8 been possible. that the electron-transfer process might be orientation dependent. Several molecular beam techniques have arisen to prepare Electron-transfer processes are important to many different reagent molecules in known orientations.' The simplest (and most kinds of reactions as well as to the harpoon reaction, and we have recent) is the "brute force" method of merely applying a strong begun to investigate how orientation might affect electron transfer. electric field on extremely cold polar molecules,2 but this method We have now studied several reactions, including K + CF3X (X is highly re~trictive.~Most experiments so far have been done = I, Br), at energies of a few (5-30) eV, so that the ions move with symmetric top molecules, such as CH3Br. The orientation fast enough that they can escape from their Coulomb attraction is defined by a weak electric field because these molecules precess and be detected directly. Collisional ionization has been studied in an applied electric field in the same way a child's top precesses in some detail previously9 (for unoriented molecules) and the in a gravitational field. Each molecule is oriented. Even though products of theselo fast-atom collisions are predominantly K+,CF,, all orientations are present in a beam, passing the beam through and X-. In the experiments reported here, we orient symmetric an inhomogeneous electric field filters out molecules in orientations top molecules prior to collision, and we detect the K+ ion. such that (cos 8) > 0, where 8 is the angle between the dipole moment and electric field. The molecules passed by the inho- (1) For reviews see: (a) Brooks, P. R. Science 1976,193, 11. (b) Bern- stein, R. B.; Herschbach, D. R.; Levine, R. D. J. Phys. Chem. 1987,91,5365. mogeneous field have (cos8) < 0, and are reacted with an atomic (c) Parker, D. H.; Bernstein, R. B. Annu. Rev. Phys. Chem. 1989, 40, 561. beam in a weak uniform electric field. The weak field is applied (d) Stolte, S. In Atomic and Molecular Beam Methods; Scoles, G., Ed.; parallel or antiparallel to the relative velocity and determines which Oxford: New York, 1988; Vol. 1, p 631. end of the molecule is presented to the atoms. (2) (a) Loescb, H. J.; Remscheid, A. J. Chem. Phys. 1990,93,4779. (b) Friedrich, B.; Herschbach, D. R. 2. Phys., in press. Orientation effects are large and varied. For example, at (3) This method can be used if E, << MEand requires either very cold thermal energies, CHJ reacts with K or (Rb4) preferentially on molecules or very high fields, but it is the only method available for orienting the I end, but CF31and CF,Br react at both ends, with different 2-state molecules. For these 2 molecules, various optical methods have bctn angular distributions. The "harpoon" modelS of electron transfer used to produce polarization, where the plane of rotation is selected but one end is not distinguished from the other (described in ref 1 b). has been invoked6 to describe the CF31reaction, and this nicely (4) Chakravorty, K. K.; Parker, D. H.; Bernstein, R. B. Chem. Phys. 1982, explains the different angular distributions observed in the "heads" 68, 1. These latter experiments have been reinterpreted in terms of a steric and "tails" orientations. In this model, the K is expected to donate cone of nonreaction of 43O: Choi, S. E.; Bernstein, R. B. J. Chem. Phys. an electron (the harpoon) to the CF,I, followed by explosive 1985, 83, 4463. (5) Herschbach, D. R. Adv. Chem. Phys. 1966, IO, 319. decomposition of the CFJ-, after which the Coulomb attraction (6) Brooks, P. R. Faraday Discuss. Chem. Soc. 1973,55, 299. causes the K+ and I- to combine: (7) Brooks, P. R.; McKillop, J.; Pippen, H. G. Chem. Phys. Lett. 1979,615, 144. K + CFJ --* K+ + CFJ electron transfer (1) (8) Carman, H. S.;Harland, P. W.; Brooks, P. R. J. Phys. Chem. 1986, 90,944. CF31- - CF3 + I- ion dissociation (2) (9) For reviews of collisional ionization, see: (a) Kleyn, A. W.; Los, J.; Gislason, E. A. Phys. Rep. 1982.90, 1. (b) Lacmann, K. I. Potential Energy Surfaces; Lawley, K. P.; Ed.; Wiley: New York, 1980, p 513. (c) Bade, A. I- K+ + ionic recombination (3 1 + KI P. M. Adv. Chem. Phys. 1975, 30, 463. (d) Los, J.; Kleyn, A. W. Alkali Because the cross section for reaction at the two ends was roughly Halide Vapors; P. Davidovits, P., McFadden, D. L., Eds.; Academic Press: New York, 1979; p 279. (10) For collision with K: (a) McNamee, P. E.; Lacman, K.; Herschbach, 'Permanent address: Chemistry Department, University of Canterbury, D. R. Faraday Discuss. Chem. Soc. 1973,55,318. (b) Kleyn, A. W.; Hubers, Christchurch. New Zealand. M. N.; La, J. 10th ICPEAC Abstracts of Papers, Paris, 1977; p 1162. For 'Permanent address: Chemical Physics Section, Health & Safety Research collision with other alkali metals: (c) Compton, R. N.; Reinhardt, P. W.; Division, Oak Ridge National Laboratory, P. 0. Box 2008, Oak Ridge, TN Cooper, C. D. J. Chem. Phys. 1978,68,4360. (d) Tang, S. Y.; Mathur, B. 3783 I. P.; Rothe, E. W.; Reck, G. P. J. Chem. Phys. 1976, 64, 1270. 0022-3654191 12095-8 131302.50/0 0 1991 American Chemical Society 8138 The Journal of Physical Chemistry, Vol. 95, No. 21, 1991 Harland et al. Preliminary results have been obtained for CH$ and CF31.11*'2 io00 L I I I I Here we report on a survey of related molecules, CH3X (X = F, CI, Br, and I); CF,Y (Y = C1, Br, and I); t-butyl Z (Z = C1, Br); and CX3H (X = C1, F). We find in all cam a strong dependence on orientation at energies near threshold, and this effect decreases as the energy is increased. Regardless of the polarity of the dipole, the end producing the most ions is the "heads" end (the end with the most labile halogen). These results may stll be understood in terms of the orientation of the dissociating molecular ion. 11. Experimental Section The experimental apparatus and procedure for data collection 5 10 5 10 5 10 have been previously discussed in paper I (ref 12), and we provide E (eV) here only a brief summary of the experiment.
Load more

Recommended publications
  • Numerical Simulation of Plasma Kinetics in Low-Pressure Discharge in Mixtures of Helium and Xenon with Iodine Vapours
    12 Numerical Simulation of Plasma Kinetics in Low-Pressure Discharge in Mixtures of Helium and Xenon with Iodine Vapours Anatolii Shchedrin and Anna Kalyuzhnaya Institute of Physics, National Academy of Sciences of Ukraine Ukraine 1. Introduction Due to ecological problems related to the utilization of gas-discharge ultraviolet (UV) mercury vapour lamps widely used in lighting technology, photochemistry, and photomedicine, there arises a need for developing new mercury-free sources of UV radiation using electron bands of rare-gas monohalides and halogen molecules as well as spectral lines of their atoms (Lomaev et al., 2003). The most high-power mercury-free lamps are UV emitters with “chlorine - noble gas” active media emitting on transitions of the excimer molecules XeCl (308 nm) and KrCl (222 nm). The use of aggressive chlorine molecules in the gas mixtures of such emitters results in a comparatively low mixture life (1- 100 hours), which impedes their wide application in various optical technologies. It is mainly due to the absorption of chlorine by open metal electrodes (especially strongly heated cathode) and its heterophase chemical reaction with a quartz discharge tube accompanied by the formation of polymer compounds (chlorosiloxanes). That is why the replacement of chlorine molecules by less aggressive iodine ones in the working media of excilamps represents an urgent task. Ultraviolet radiation of glow discharge plasma in mixtures of helium with iodine vapours that is still transparent to air is mainly concentrated in the spectral ranges 175-210 nm and 320-360 nm. The mixture life of such lamps reaches 103 hours (Lomaev & Tarasenko, 2002; Shuaibov et al., 2005a, 2005b).
    [Show full text]
  • 3. Introduction to Reactive Molecular Collisions
    3. Introduction to reactive molecular collisions We examine in this chapter how the motion of the reactants as they approach each other governs chemical reactivity. This allows us to use a two-body point of view where the internal structure of the colliding specie is not explicitly recognized. All that we can do therefore is lead the reactants up to a reaction. But we will not be able to describe the chemical rearrangement itself nor to address such questions as energy disposal in the products. Chapter 5 takes up these themes. On the other hand, without the approach of the reactants there cannot be a bimolecular reaction. The tools already at our disposal are sufficient to discuss this approach motion. As expected, the striking distance that we have called the impact parameter will be a key player. We do all of this in section 3.2. What we will obtain is information about the dependence of the reaction cross section on the collision energy. In chemical kinetics one characterizes the role of energy in chemical reactivity by the temperature dependence of the reaction rate constant. In section 3.1 we review the input from chemical kinetics-the Arrhenius representation of the rate constant-then go from the rate constant to the reaction cross section. Next we go in the opposite direction, from the microscopic reaction cross section to the macroscopic rate constant. What we obtain thereby is the Tolman interpretation of the activation energy as the (mean) excess energy of those collisions that lead to reaction. MRD Chapter 3 page 1 © R D Levine (2003) 3.1.
    [Show full text]
  • Molecular Reaction Dynamics Raphael D
    Cambridge University Press 052184276X - Molecular Reaction Dynamics Raphael D. Levine Index More information Index 1, 1’ binaphtyl conformational twist 455 in the spectator limit 415 kinematic effects 402 1, 3 cyclohexadiene ring opening in real of reaction products in a polarization 402 time 386 chemiluminescent reaction 414 sense of rotation 411 analog computation 200 spherical basis vectors 408 + Ar +H2 stripping 23 angle of attack 14, 97 vector correlations 419, 422 collision-induced dissociation 23 angle of deflection 111 angular velocity 113 above-threshold ionization, ATI angle-dependent barrier to reaction 96, anisotropy parameter in 332 163 photodissociation 280 absorption, collision-induced 310 angle-dependent line-of-centers anti-Morse potential 198 abstraction reaction 14 model 97 Arrhenius activation energy 1, 75, accommodation to the surface 480 angle variable 198 79 action variable 198 angular distribution see also Tolman expression activated adsorption 482 classical divergence in the forward Arrhenius equation 1, 74 activation energy 1 direction 119 by saddle-point integration 101 definition 74 for surface scattering 479 historical background 25 in transition state theory 254 hard-sphere-like in the backward surface desorption 480 negative 196 direction 119 Arrhenius pre-exponential factor 1 Tolman expression 79, 103 in complex mode reactions 140 high/low 255 adiabatic limit 149, 268, 372 forward-backward symmetry 140 see also entropy of activation, steric adiabaticity parameter 372 in direct reactions 138 factor and momentum
    [Show full text]
  • Reaction Kinetics
    1 Reaction Kinetics Dr Claire Vallance First year, Hilary term Suggested Reading Physical Chemistry, P. W. Atkins Reaction Kinetics, M. J. Pilling and P. W. Seakins Chemical Kinetics, K. J. Laidler Modern Liquid Phase Kinetics, B. G. Cox Course synopsis 1. Introduction 2. Rate of reaction 3. Rate laws 4. The units of the rate constant 5. Integrated rate laws 6. Half lives 7. Determining the rate law from experimental data (i) Isolation method (ii) Differential methods (iii) Integral methods (iv) Half lives 8. Experimental techniques (i) Techniques for mixing the reactants and initiating reaction (ii) Techniques for monitoring concentrations as a function of time (iii) Temperature control and measurement 9. Complex reactions 10. Consecutive reactions 11. Pre-equilibria 12. The steady state approximation 13. ‘Unimolecular’ reactions – the Lindemann-Hinshelwood mechanism 14. Third order reactions 15. Enzyme reactions – the Michaelis-Menten mechanism 16. Chain reactions 17. Linear chain reactions The hydrogen – bromine reaction The hydrogen – chlorine reaction The hydrogen-iodine reaction Comparison of the hydrogen-halogen reactions 18. Explosions and branched chain reactions The hydrogen – oxygen reaction 19. Temperature dependence of reaction rates The Arrhenius equation and activation energies Overall activation energies for complex reactions Catalysis 20. Simple collision theory 2 1. Introduction Chemical reaction kinetics deals with the rates of chemical processes. Any chemical process may be broken down into a sequence of one or more single-step processes known either as elementary processes, elementary reactions, or elementary steps. Elementary reactions usually involve either a single reactive collision between two molecules, which we refer to as a a bimolecular step, or dissociation/isomerisation of a single reactant molecule, which we refer to as a unimolecular step.
    [Show full text]
  • Investigation of Airborne Particles by On-Line Single-Particle Mass Spectrometry with Novel Ionization Approaches
    "#$%&'()*'(+#!+,!-(./+.#%!0*.'(12%&!/3!! 4#56(#%!7(#)2%50*.'(12%!8*&&!79%1'.+:%'.3! ;('<!=+$%2!"+#(>*'(+#!-99.+*1<%& ?@:@2*'($%!A(&&%.'*'(+#! >@.!B.2*#)@#)!C%&!*D*C%:(&1<%#!E.*C%&! !"#$"%&%'%()&*+$(%+,-()!FA.G!.%.G!#*'GH! *#!C%.! 8*'<%:*'(&1<5=*'@.;(&&%#&1<*,'2(1<%#!I*D@2'J'! C%.!K#($%.&('J'!L+&'+1D! $+.)%2%)'!$+#! M@2(*#!-2%N*#C%.!71<*C%! )%/+.%#!*:!OPGQPGRPPQ!(#!A+.':@#C! L+&'+1DS!=+$%:/%.!OQOQ! https://doi.org/10.18453rosdok_id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
    [Show full text]
  • Lecture Notes
    LECTURE NOTES Subject: Chemistry-III (Basic Physical-I) Faculty: Dr. Monalisa Mohapatra Subject Code: CH-211 3rd Semester 5 yr Int. M.Sc. Module-I: (10 Hours) Dilute Solutions: Vapour Pressure, Raoult’s Law, and Thermodynamic Derivation of Laws Relating to Elevation of Boiling Point, Depression of Freezing Point and Osmotic Pressure. Ideal and Non-ideal Solution Association and Dissociation. Module-II: (10 Hours) Homogenous Equilibrium: Law of Mass Action and the Thermodynamic Derivation of the Expression for Equilibrium constant, Different forms of Equilibrium Constants. Le Chatelier's Principle, Illustrations with some Gaseous Reactions. Effect of Temperature on Equilibrium van’t Hoff’s Equation and its Integration. Module-III: (10 Hours) Chemical Kinetics: Order and Molecularity, Kinetics of Ist and IInd Order Reactions, Simple Opposing (A & B) Reaction, Consecutive of Sequential (A-B-C) Reaction, Chain Reaction (H2+Br2). Effect of Temperature on Reaction Rate. Collision Theory of Reaction Rate, Qualitative Treatment of Transition Theory. Module-IV: (10 Hours) Thermodynamic Concept: Heat Content and Heat Capacity Isothermal and Adiabaic Change, Work Done for Ideal and van der Waal Gases. Thermochemistry: Heat Changes in Chemical Reactions, Hess’s Law, Kirchoff’s Equation. 2nd law of Thermodynamics: Spontaneous Process, Carnot’s Theorem and Carnot’s Cycle, Efficiency of Heat Engine, Entropy Changes in Reversible and Irreversible Processes, Free Energy and Work Function Condition for Equilibrium, Clapeyron and Clausius Equation, Gibb’s Helmholtz Equation. Prescribed Books: 1. P. Atkins and J. de Paula, Elements of Physical Chemistry, 6th Edition, Oxford University Press, 2013. 2. P. Atkins and J. de Paula, Physical Chemistry, 9th Edition, W.
    [Show full text]
  • MOLECULAR REACTION DYNAMICS Hilary Term 2021 4 Lectures Prof
    MOLECULAR REACTION DYNAMICS Hilary Term 2021 4 Lectures Prof Mark Brouard MOLECULAR REACTION DYNAMICS Hilary Term 2021 4 Lectures Prof Mark Brouard Aims: Understanding, predicting, and controlling chemical reactivity. What is the nature of the forces at play in chemical reactions? What is the nature of the potential energy surface(s)? Outline plan: 1. Introduction 2. Potential energy surfaces 3. Experimental techniques - initial states 4. Experimental techniques - final states 5. Elastic and Inelastic scattering 6. Reactive scattering 7. Controlling reagents and characterizing products 8. Probing the transition state. Types of encounter Cross-sections can be defined and measured for many different types of collision We will consider three categories of collisions (see Professor Vallance’s lectures): (a) Elastic scattering: conserves kinetic energy, leads to a change in direction. (b) Inelastic scattering: leads to a change in internal energy of the colliders. (c) Reactive scattering: leads to a rearrangement of bonds. Apart from today’s lecture we will focus on reactive collisions. The motivation is to understand the forces operating in molecular collisions. Elastic scattering Simple measurement of a cross-section Example of a ‘beam-gas’ experiment. l I 0 I detector molecular beamofA lowpressure ofB When molecules of A collide with B they are deflected. This leads to attenuation of the molecular beam of A. Use the equivalent of Beer-Lambert law for light −σc l I = I0 e B σ — the elastic scattering cross-section, cB — the concentration of B, and l — the path length of B through which A passes. Elastic scattering cross-section: Glory scattering Striking quantum effects are observed for atom-atom scattering (e.g., K + Rg).
    [Show full text]
  • Investigation of Monochromatization Light Effect at Molecular/Atomic Level in Electronegative-Electropositive Gas Mixtures Plasma L.C
    World Academy of Science, Engineering and Technology International Journal of Physical and Mathematical Sciences Vol:6, No:11, 2012 Investigation of Monochromatization Light Effect at Molecular/Atomic Level in Electronegative-Electropositive Gas Mixtures Plasma L.C. Ciobotaru + → 2e + X2 (4) Abstract—In electronegative-electropositive gas mixtures → X + X- (5) * plasma, at a total pressure varying in the range of ten to hundred → e + X2 (6) Torr, the appearance of a quasi-mochromatization effect of the At low pressures, (0.5-5) Torr, the major reaction for the emitted radiation was reported. This radiation could be the result of excimer formation is accepted to be the reverse Harpoon – the generating mechanisms at molecular level, which is the case of reaction between the ground state rare gas atom and the the excimer radiation but also at atomic level. Thus, in the last case, 3 1 in (Ne+1%Ar/Xe+H2) gas mixtures plasma in a dielectric barrier excited states ( P, D) of the molecular halogen: * discharge, this effect, called M-effect, consists in the reduction of the Rg +X2 → Rg X *(B) + X (7) discharge emission spectrum practice at one single, strong spectral This reaction is a two-body reaction and for this reason is line with λ = 585.3 nm. The present paper is concerned with the very attractive one for explaining the formation of excimer characteristics comparative investigation of the principal reaction molecules at low pressures. * mechanisms involved in the quasi-monochromatization effect At high pressure (10-760) Torr, the RgX molecules excimer existence in the case of the excimer radiation, respectively of the M- are generated either by the direct Harpoon - reaction or by the effect.
    [Show full text]
  • A Multichannel Harpoon Model for Reactive Quenching of Xe 5P56p by Cl2 *
    x, CONF-8705246--2 DE88 004263 A Multichannel Harpoon Model for Reactive Quenching of Xe 5p56p by Cl2 * M. R. Bruce, W. B. Layne, and J. W. Keto Physics Department The Univ. of Texas at Austin Austin, Tx. 78712 ABSTRACT A harpoon reaction model employing multiple crossings based on the formalism of Bauer et al * and Gislason and Sachs2 is developed to explain the large cross sections (a ~ 400 A2) measured for the reaction Xe 5p56p + Cl2- The model calculates the Landau-Zener transition probability for each intermediate ionic crossing with the covalent surface. The transition matrix elements are represented as a product of the electronic interaction (modeled by the empirical result of Olson et al)3 and a Frank-Condon factor for the CI2—>Cl2~ transition. The model predicts near unit probability for a transition to the ionic surface for impact parameters less than 20 Bohr. Once transfer occurs, the pair is captured by dissociation of Cl2~ to form XeCl*. The large temperature dependence observed qualitatively in the experiments is explained by the increased cross section for vibrationally excited G2. A simple model for orbiting collisions is used to estimate the effect of termolecular collisions. DISCLAIMER This report was prepared as an account of work sponsored by an agency of the United States Government. Neither the United States Government nor any agency thereof, nor any of their employees, makes any warranty, express or implied, or assumes any legal liability or responsi- bility for the accuracy, completeness, or usefulness of any information, apparatus, product, or process disclosed, or represents that its use would not infringe privately owned rights.
    [Show full text]
  • REACTION DYNAMICS Hilary Term 2011 6 Lectures Prof M. Brouard
    REACTION DYNAMICS Hilary Term 2011 1. Introduction 6 Lectures Prof M. Brouard Thermally averaged rate coefficients The thermal rate coefficient, k(T ), is a highly averaged quantity. Understanding Look again at simple collision theory: Predicting chemical reactivity 1 • Reactions occur by collisions Controlling • Collision energy2 (E ) must exceed the barrier height (E ) t 0 At fixed relative velocity, crel (or collision energy, Et) 1. Introduction Reaction rate = Rate of reactive collisions at velocity crel × 2. Elastic and Inelastic scattering Fraction of collisions at velocity crel, f(crel) 3. Determining the angular distribution miss hit 2 sc = p d 4. Controlling reagents and characterizing products crel 5. Interpretation Rate of reactive collisions = k(crel) [A] [B] 6. Probing the transition state. 1 Reactants are structureless spheres, with radii rA and rB and collision diameter d = rA + rB 2 relative motion 1 2 The collision energy is the translational energy associated with Et = 2 µcrel with mAmB v v µ = mA+mB and crel is the relative speed (crel = | B − A|). 1 k(crel) defines the reactive collision volume swept out per unit time State-resolved properties k(crel)= crel σ(crel) , Example 1 where σ(crel) is the cross-section for the process of interest at velocity K + HCl(v) −→ KCl + H ∆H◦− = 4 kJ mol−1 crel. −1 f(crel)dcrel is the fraction of collisions with velocities in the range crel to HCl(v = 1) has an energy of 34.5kJmol . crel + dcrel: 3/2 µ 2 −µcrel/2RT 2 f(crel) dcrel = e 4πcreldcrel Put this energy into HCl(v = 1) and rate increases by ×100.
    [Show full text]
  • Understanding Chemical Reactions at the Molecular Level
    This page intentionally left blank Molecular Reaction Dynamics Molecular Reaction Dynamics is a brand new version of the text by Levine and Bernstein. The book delivers an updated treatment of this fundamental topic. An appreciation of how chemical reactions occur and their control is essential to chemists and to those in interdisciplinary fields such as materials and nanoscience, drug design, and astrochemistry. The first half of the book describes experimental techniques for initiating and probing reaction dynamics and the essential insights gained. The second part explores key areas including photoselective chemistry, stereochemistry, chemical reactions in real time, and chemical reaction dynamics in solutions and interfaces. Typical of the new challenges are molecular machines, enzyme action, and molecular control. With problem sets included, this book is aimed at advanced undergraduate and graduate students studying chemical reaction dynamics, as well as physical chemistry, biophysics, and materials science. R L is Max Born Professor of Natural Philosophy at the Hebrew University of Jerusalem and Distinguished Professor of Chemistry at the University of California, Los Angeles. He is active in the area of chemical reaction dynamics and his published scientific work has earnt the recognition of the Israel Prize and the Wolf Prize. He is a member of the Israel National Academy of Sciences and a foreign member of the National Academy of Sciences of the United States and of Academiae Europaeae. Molecular Reaction Dynamics Raphael D. Levine Cambridge, New York, Melbourne, Madrid, Cape Town, Singapore, São Paulo Cambridge University Press The Edinburgh Building, Cambridge , UK Published in the United States of America by Cambridge University Press, New York www.cambridge.org Information on this title: www.cambridge.org/9780521842761 © R.
    [Show full text]
  • FACTORS AFFECTING RATE of REACTION -.:: GEOCITIES.Ws
    FACTORS AFFECTING RATE OF REACTION 1)Chemical nature of the reactants a freshly exposed metallic sodium surface furnishes almost instantly when explored to air and moisture – very fast reaction . Reaction between iron and moisture (air) to form rust is slow. Reason- iron do not lose electron easily. It depends on the inherent tendencies of atoms ,molecules, or ions to undergo changes in chemical bonds. 2) Ability of the reactants to come in contact with each other. For reaction to occur the reactants must collide with each other: (a) homogeneous reaction –when reactants are in the same phase, e.g liquid, gas. (b) heterogeneous reaction –reactant only able to meet at the interface. Rate of reaction dictated by the area of contact between the phases i.e surface area plays an important role. Assoc. Prof. Dr Misni Misran SCE2230 Sesi 2005/2006 36 1 cm 0.01 cm 0.01 cm 0.01cm 1 cm 1 cm 1,000,000 cubes Total surface area = 6 cm2 total surface are = 600 cm2 3) Concentration of the Reactants The rates of both homogeneous and heterogeneous reactions are affected by the concentrations of the reactants. 4)Temperature of the system Chemical reaction will occur faster at higher temperature. e.g chemical metabolism slows down in cooler surrounding. 5) Presence of rate promoting agents or catalysts. Catalysts are substances that increase the rate of reactions without itself being used up. Best example is enzymes. ZERO –ORDER REACTION Say our reaction of interest is, aA + bB → products No need to specify the products since order of reaction is defined interms of the reactant concentration.
    [Show full text]