Collision Theory
Collision theory in chemistry explains that chemical reactions occur when reactant molecules collide with sufficient energy and proper orientation. The theory emphasizes the importance of collision frequency, energy, and orientation in determining the rate of a chemical reaction. It provides a framework for understanding the factors that influence reaction rates and the conditions necessary for successful collisions to occur.
5 Key excerpts on "Collision Theory"
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
...Collision Theory states that, in order for molecules to react, they must collide with each other. But, a collision between reactant molecules is not enough to cause a reaction. These collisions must also be energetic enough to be able to break and reform molecular bonds. In addition, the reactant atoms must be oriented in such a way that they can easily rearrange to form the products. A molecular collision that results in a chemical reaction between reactants is called an effective collision. So, according to Collision Theory the rate of a chemical reaction is equal to the frequency of the effective collisions. Consider the gas phase reaction between reactants A and B, which collide to form the product A − B. A + B → A − B The number of collisions between reactants A and B is dependent upon the number of molecules of A and B and so the number of collisions is proportional to the concentration of each reactant. If the concentration of either A or B is doubled, the frequency of collisions between A and B will double. As the frequency of collisions is increased, the frequency of effective collisions is also increased. So, increasing the concentration of the reactants will have the effect of increasing the reaction rate. For conditions of one atmosphere pressure and a temperature of 273 K, the collision frequency between the gas reactants A and B has been estimated to be 2 × 10 8 mol • L − 1 • s − 1. If every collision between A and B was effective, this would be a very fast reaction rate where the two gases would almost completely react in a billionth of a second (10 − 9 s) after mixing. While there are a few gas reactions that approach this rate, most chemical reaction rates are much slower. Common rates for second order chemical reactions are in the range of 10 − 2 to 10 − 3 mol • L − 1 • s − 1...
eBook - ePub- Gavin Whittaker, Andy Mount, Matthew Heal(Authors)
- 2000(Publication Date)
- Taylor & Francis(Publisher)
...They may be determined experimentally from a plot of lnk against 1/ T. Collision Theory This simple model to describe the rate of a bimolecular reaction assumes that reaction occurs when two reactant species collide with an energy along their line of centers greater than the activation energy for the reaction. The species are treated as hard, structureless spheres that only interact when the distance between their centers is less than the collision radius (the sum of the radii of the colliding reactants). The derived rate constant also includes a steric factor to account for the probability that molecules collide with the correct relative orientation to permit reaction. Activated complex (transition state) theory This theory interprets chemical reaction in terms of a loosely-bound activated complex which acts as if it is in equilibrium with the reactant species. The molecular configuration of the activated complex corresponding to the maximum energy along the reaction coordinate between breaking of old bonds and formation of new bonds is known as the transition state. The derived rate constant is given by K‡K‡ where K‡ is the equilibrium constant between reactants and activated complex and k‡ is the first order rate constant for decomposition of the activated complex into products. These parameters can be calculated from statistical mechanics given a postulated model of the activated complex. Catalysts A catalyst increases the rate of chemical reaction by providing an alternative reaction pathway with lower activation energy than the reaction pathway in its absence. A catalyst is not consumed and therefore does not appear in the chemical equation for the reaction...
eBook - ePub- Warren C. Strahle, William A. Sirignano, William A. Sirignano(Authors)
- 2020(Publication Date)
- Routledge(Publisher)
...3 CHEMICAL KINETICS 3.1 INTRODUCTION All of the chemical reactions shown as examples in previous chapters almost never really take place as written. The true path of product creation is usually a complex one involving many steps at the molecular or atomic level. From the standpoint of thermodynamics, it did not matter that the true reaction mechanism was in error, because only the initial and final states were of interest. But often we need the true reaction path to study flame phenomena. We also need information on the rate in proceeding from an initial state to a final state. Such rates of chemical reactions are furnished by the science of chemical kinetics. The determination of actual reaction paths involves very complex inference from usually incomplete data and theoretical argument. However, at this time, there are many reaction schemes that are quite well known; there are also many others which are known quite incompletely. From an engineering viewpoint, complete knowledge of a reaction scheme is often not necessary, and useful approximations can be made. Indeed, in this text we shall use mostly approximate approaches, since they can yield valuable insight into flame behavior. 3.2 REACTION RATE We take it as axiomatic that atoms or molecules cannot react unless they collide. 1 We also take it as intuitive that the rate of atomic or molecular collisions must be proportional to the number densities (number per unit volume) of the colliding species. For example, in the reaction the rate of reaction for the reactants must be proportional to c H c O. We read the reaction above as “one atom of oxygen collides with one atom of hydrogen to form one molecule of OH.” We do not use moles or have fractional stoichiometric coefficients when dealing with actual chemical reaction mechanisms. We define the reaction rate, R, through a coefficient, k, called the specific reaction rate constant, by for the above reaction...
- Eugene Machlin(Author)
- 2010(Publication Date)
- Elsevier Science(Publisher)
...In gases, molecules come into contact by collision processes. In condensed matter, the atoms or molecules oscillate or rotate about a stable position or orientation, at any finite temperature. Even at 0°K they perform a zero point motion. These oscillations or rotations make possible the existence of kinetic phenomena in condensed matter. Each such phenomenon involves the motions of atoms or molecules in some particular way, which we will describe as a motion along a reaction path. If the energy of the system passes through a maximum as these particles move along the reaction path, then we describe this maximum energy as an “activation energy”, which the system must exceed for it to proceed from its initial state to a different state along this reaction path. Since there are many reaction paths from reactant state to product state, the path with the highest probability of being followed is that involving the least activation free energy, which is usually called the “saddle point” energy. We shall explore the deeper significance of the activation energy in this chapter. Kinetic phenomena that occur not too far from equilibrium conditions can be treated using linear approximations to rate relations. The linear realm including concepts from reaction rate theory and the thermodynamic theory of irreversible processes comprises the scope of this chapter. The non-linear regime applicable to far from equilibrium conditions and the production of spatial patterns is discussed in later chapters and Chapter XIII. We do not consider homogeneous chemical reaction kinetics, a subject outside the scope of this chapter – the kinetics of processes in condensed phases. 1 Activation energy Let us consider a process taking place in a condensed phase in which reactants existing in some metastable state proceed to a more stable product state. Atom or molecule arrangements define the reactant and product states...
eBook - ePub- Pieter Walstra(Author)
- 2002(Publication Date)
- CRC Press(Publisher)
...See further Section 5.2. 4.3.2 Activation Energy The first useful theory of reaction rates was due to Arrhenius, and it is easiest to envisage for a bimolecular reaction. It is assumed that the molecules have to overcome an energy barrier before they can react. This is depicted in Figure 4.4. The energy barrier per mole is called the activation energy, symbol E a. As mentioned, the average translational kinetic energy of a molecule is (3/2) k B T, and the average kinetic energy involved in a collision of two molecules is given by 2 times 1/3 of that value, i.e. k B T; the factor 1/3 arises because the molecules move in 3 dimensions and when they collide this happens in one dimension. The collision may now provide the activation energy needed for the molecules to react. From Eq. (4.9), the proportion of collisions of which the energy is higher than a given value U* can be derived to equal exp(- U*/kB T). By changing from molecules to moles and by putting U*N AV =E a, it follows that the rate constant would be (4.10)* where A often is called the frequency factor. It is also called the preexponential factor and denoted as k ‡ or k0, i.e., the value that k would attain at infinite temperature or zero activation energy, respectively. FIGURE 4.4 Arrhenius theory. Illustration of the energy state of the reactants, the reactants in the activated state, and the reaction product(s). E a represents the activation energy. Temperature Dependence. Eq. (4.10) predicts that log k is proportional to 1/ T, and this is indeed very often observed, especially for reactions of small molecules, involving breaking and formation of covalent bonds. The Arrhenius theory can thus be said to be very successful, and as a semiempirical relation Eq. (4.10) is indeed useful in many cases. Chemists generally express the temperature dependence of a reaction in Q 10, i.e., the factor by which a reaction is faster if one increases the temperature by 10 K...
