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The reaction rate (rate of reaction) or speed of reaction for a reactant or product in a particular reaction is intuitively defined as how fast or slow a reaction takes place. For example, the oxidative rusting of iron under Earth's atmosphere is a slow reaction that can take many years, but the combustion of cellulose in a fire is a reaction that takes place in fractions of a second.
Chemical kinetics is the part of physical chemistry that studies reaction rates. The concepts of chemical kinetics are applied in many disciplines, such as chemical engineering, enzymology and environmental engineering.
Consider a typical chemical reaction:
According to IUPAC's Gold Book definition the reaction rate r for a chemical reaction occurring in a closed system under isochoric conditions, without a build-up of reaction intermediates, is defined as:
where [X] denotes the concentration of the substance X. (Note: The rate of a reaction is always positive. A negative sign is present to indicate the reactant concentration is decreasing.) The IUPAC recommends that the unit of time should always be the second. In such a case the rate of reaction differs from the rate of increase of concentration of a product P by a constant factor (the reciprocal of its stoichiometric number) and for a reactant A by minus the reciprocal of the stoichiometric number. Reaction rate usually has the units of mol L−1 s−1. It is important to bear in mind that the previous definition is only valid for a single reaction, in a closed system of constant volume. This most usually implicit assumption must be stated explicitly, otherwise the definition is incorrect: If water is added to a pot containing salty water, the concentration of salt decreases, although there is no chemical reaction.
For any open system, the full mass balance must be taken into account: IN - OUT + GENERATION - CONSUMPTION = ACCUMULATION
where is the inflow rate of A in molecules per second, the outflow, and is the instantaneous reaction rate of A (in number concentration rather than molar) in a given differential volume, integrated over the entire system volume at a given moment. When applied to the closed system at constant volume considered previously, this equation reduces to: , where the concentration is related to the number of molecules by . Here is the Avogadro constant.
For a single reaction in a closed system of varying volume the so-called rate of conversion can be used, in order to avoid handling concentrations. It is defined as the derivative of the extent of reaction with respect to time.
Here is the stoichiometric coefficient for substance , equal to a, b, p, and q in the typical reaction above. Also is the volume of reaction and is the concentration of substance .
When side products or reaction intermediates are formed, the IUPAC recommends the use of the terms rate of appearance and rate of disappearance for products and reactants, properly.
Reaction rates may also be defined on a basis that is not the volume of the reactor. When a catalyst is used the reaction rate may be stated on a catalyst weight (mol g−1 s−1) or surface area (mol m−2 s−1) basis. If the basis is a specific catalyst site that may be rigorously counted by a specified method, the rate is given in units of s−1 and is called a turnover frequency.
For example, coal burns in a fireplace in the presence of oxygen, but it does not when it is stored at room temperature. The reaction is spontaneous at low and high temperatures but at room temperature its rate is so slow that it is negligible. The increase in temperature, as created by a match, allows the reaction to start and then it heats itself, because it is exothermic. That is valid for many other fuels, such as methane, butane, and hydrogen.
Reaction rates can be independent of temperature (non-Arrhenius) or decrease with increasing temperature (anti-Arrhenius). Reactions without an activation barrier (e.g., some radical reactions), tend to have anti Arrhenius temperature dependence: the rate constant decreases with increasing temperature.
For example, when methane reacts with chlorine in the dark, the reaction rate is very slow. It can be sped up when the mixture is put under diffused light. In bright sunlight, the reaction is explosive.
For example, platinum catalyzes the combustion of hydrogen with oxygen at room temperature.
All the factors that affect a reaction rate, except for concentration and reaction order, are taken into account in the rate equation of the reaction.
For a chemical reaction a A + b B → p P + q Q, the rate equation or rate law is a mathematical expression used in chemical kinetics to link the rate of a reaction to the concentration of each reactant. It is of the kind:
For gas phase reaction the rate is often alternatively expressed by partial pressures.
In these equations is the reaction rate coefficient or rate constant, although it is not really a constant, because it includes all the parameters that affect reaction rate, except for concentration, which is explicitly taken into account. Of all the parameters influencing reaction rates, temperature is normally the most important one and is accounted for by the Arrhenius equation.
Stoichiometry, molecularity (the actual number of molecules colliding), and reaction order coincide necessarily only in elementary reactions, that is, those reactions that take place in just one step. The reaction equation for elementary reactions coincides with the process taking place at the molecuar level, i.e. molecule A collides with molecule B. From collision theory follows that the likelihood of a collision of three molecules is highly unlikely. Therefore molecularity for elementary reactions is either one or two. Empirically, other values can be assigned to allow mathematical description of the rate. Then, positive rational numbers are not uncommon but should not be assigned physical meaning.
By using the mass balance for the system in which the reaction occurs, an expression for the rate of change in concentration can be derived. For a closed system with constant volume, such an expression can look like
Each reaction rate coefficient k has a temperature dependency, which is usually given by the Arrhenius equation:
Ea is the activation energy and R is the gas constant. Since at temperature T the molecules have energies given by a Boltzmann distribution, one can expect the number of collisions with energy greater than Ea to be proportional to . A is the pre-exponential factor or frequency factor.
The values for A and Ea are dependent on the reaction. There are also more complex equations possible, which describe temperature dependence of other rate constants that do not follow this pattern.
A chemical reaction takes place only when the reacting molecules collide. However, not all collisions are effective in causing the reaction. Products are formed only when the colliding molecules possess a certain minimum energy called threshold energy. Basically, the number of activated molecules nearly doubles for a temperature T+10 kelvin. For a given reaction, the ratio of its rate constant at a higher temperature to its rate constant at a lower temperature is known as its temperature coefficient (Q). Q10 is commonly used as the ratio of rate constants that are 10 °C apart. Q10 (temperature coefficient)
The pressure dependence of the rate constant for condensed-phase reactions (i.e., when reactants and products are solids or liquid) is usually sufficiently weak in the range of pressures normally encountered in industry that it is neglected in practice.
The pressure dependence of the rate constant is associated with the activation volume. For the reaction proceeding through an activation-state complex:
the activation volume, , is:
where denote the partial molar volumes of the reactants and products and indicates the activation-state complex.
For the above reaction, one can expect the change of the reaction rate constant (based either on mole-fraction or on molar-concentration) with pressure at constant temperature to be:
In practice, the matter can be complicated because the partial molar volumes and the activation volume can themselves be a function of pressure.
Reactions can increase or decrease their rates with pressure, depending on the value of . As an example of the possible magnitude of the pressure effect, some organic reactions were shown to double the reaction rate when the pressure was increased from atmospheric (0.1 MPa) to 50 MPa (which gives =-0.025 L/mol).
For the reaction
The observed rate equation (or rate expression) is:
As for many reactions, the rate equation does not simply reflect the stoichiometric coefficients in the overall reaction: It is third order overall: first order in H2 and second order in NO, although the stoichiometric coefficients of both reactants are equal to 2.
In chemical kinetics, the overall reaction rate is often explained using a mechanism consisting of a number of elementary steps. Not all of these steps affect the rate of reaction; normally the slowest elementary step controls the reaction rate. For this example, a possible mechanism is:
Reactions 1 and 3 are very rapid compared to the second, so the slow reaction 2 is the rate determining step. This is a bimolecular elementary reaction whose rate is given by the second order equation :, where k2 is the rate constant for the second step.
However N2O2 is an unstable intermediate whose concentration is determined by the fact that the first step is in equilibrium, so that :, where K1 is the equilibrium constant of the first step. Substitution of this equation in the previous equation leads to a rate equation expressed in terms of the original reactants
This agrees with the form of the observed rate equation if it is assumed that . In practice the rate equation is used to suggest possible mechanisms which predict a rate equation in agreement with experiment.
The second molecule of H2 does not appear in the rate equation because it reacts in the third step, which is a rapid step after the rate-determining step, so that it does not affect the overall reaction rate.