Chemical kinetics, also known as reaction kinetics,
is the study of rates of chemical processes. Chemical kinetics
includes investigations of how different experimental conditions can influence
the speed of a chemical reaction and yield information about the reaction's mechanism and transition
states, as well as the construction of mathematical models that can
describe the characteristics of a chemical reaction. In 1864, Peter Waage
and Cato
Guldberg pioneered the development of chemical kinetics by formulating the law of mass action, which states that the speed
of a chemical reaction is proportional to the quantity of the reacting
substances.
Chemical kinetics deals with the experimental determination
of reaction
rates from which rate laws and rate constants are derived. Relatively
simple rate
laws exist for zero-order reactions (for which reaction rates are
independent of concentration), first-order reactions, and second-order reactions, and can be derived
for others. In consecutive reactions, the rate-determining step often determines the
kinetics. In consecutive first-order reactions, a steady state approximation can simplify
the rate law.
The activation energy for a reaction is
experimentally determined through the Arrhenius equation and the Eyring
equation. The main factors that influence the reaction
rate include: the physical state of the reactants, the concentrations
of the reactants, the temperature at which the reaction occurs, and whether or
not any catalysts
are present in the reaction.
Factors affecting reaction rate
Nature of the reactants
Depending upon what substances are reacting, the reaction
rate varies. Acid/base reactions, the formation of salts, and ion
exchange are fast reactions. When covalent bond formation takes place
between the molecules and when large molecules are formed, the reactions tend
to be very slow. Nature and strength of bonds in reactant molecules greatly
influence the rate of its transformation into products.
Physical state
The physical state (solid, liquid, or gas) of a reactant is
also an important factor of the rate of change. When reactants are in the same phase,
as in aqueous solution,
thermal motion brings them into contact. However, when they are in different
phases, the reaction is limited to the interface between the reactants.
Reaction can occur only at their area of contact; in the case of a liquid and a
gas, at the surface of the fluid. Vigorous shaking and stirring may be needed
to bring the reaction to completion. This means that the more finely divided a
solid or liquid reactant the greater its surface
area per unit volume
and the more contact it makes with the other reactant, thus the faster the
reaction. To make an analogy, for example, when one starts a fire, one uses
wood chips and small branches — one does not start with large logs right away.
In organic chemistry, on water reactions are the exception to the rule
that homogeneous reactions take place faster than heterogeneous reactions.
Concentration
The reactions are due to collisions of reactant species. The
frequency with which the molecules or ions collide depends upon their concentrations.
The more crowded the molecules are, the more likely they are to collide and
react with one another. Thus, an increase in the concentrations of the
reactants will result in the corresponding increase in the reaction rate, while
a decrease in the concentrations will have a reverse effect. For example, combustion
that occurs in air (21% oxygen) will occur more rapidly in pure oxygen.
Temperature
Temperature usually has a major effect on the rate of a
chemical reaction. Molecules at a higher temperature have more thermal
energy. Although collision frequency is greater at higher temperatures,
this alone contributes only a very small proportion to the increase in rate of
reaction. Much more important is the fact that the proportion of reactant
molecules with sufficient energy to react (energy greater than activation
energy: E > Ea) is significantly
higher and is explained in detail by the Maxwell–Boltzmann distribution of
molecular energies.
The 'rule of thumb' that the rate of chemical reactions
doubles for every 10 °C temperature rise is a common misconception. This
may have been generalized from the special case of biological systems, where
the α (temperature coefficient) is
often between 1.5 and 2.5.
A reaction's kinetics can also be studied with a temperature
jump approach. This involves using a sharp rise in temperature and
observing the relaxation time of the return to equilibrium. A
particularly useful form of temperature
jump apparatus is a shock tube, which can rapidly jump a gas's temperature
by more than 1000 degrees.
Catalysts
Generic potential energy diagram showing the effect of a
catalyst in a hypothetical endothermic chemical reaction. The presence of the
catalyst opens a different reaction pathway (shown in red) with a lower
activation energy. The final result and the overall thermodynamics are the
same.
A catalyst is a substance that accelerates the rate of a
chemical reaction but remains chemically unchanged afterwards. The catalyst increases
rate reaction by providing a different reaction mechanism to occur with a lower activation
energy. In autocatalysis a reaction product is itself a catalyst
for that reaction leading to positive
feedback. Proteins that act as catalysts in biochemical reactions are
called enzymes. Michaelis–Menten kinetics describe the rate
of enzyme mediated reactions. A catalyst does not affect the position of
the equilibria, as the catalyst speeds up the backward and forward reactions
equally.
In certain organic molecules, specific substituents can have
an influence on reaction rate in neighbouring group participation.
Agitating or mixing a solution will also accelerate the rate
of a chemical reaction, as this gives the particles greater kinetic energy,
increasing the number of collisions between reactants and, therefore, the
possibility of successful collisions.
Pressure
Increasing the pressure in a gaseous reaction will increase
the number of collisions between reactants, increasing the rate of reaction.
This is because the activity of a gas is directly proportional to
the partial pressure of the gas. This is similar to the effect of increasing
the concentration of a solution.
In addition to this straightforward mass-action effect, the
rate coefficients themselves can change due to pressure. The rate coefficients
and products of many high-temperature gas-phase reactions change if an inert
gas is added to the mixture; variations on this effect are called fall-off and chemical activation.
These phenomena are due to exothermic or endothermic reactions occurring faster
than heat transfer, causing the reacting molecules to have non-thermal energy
distributions (non-Boltzmann distribution). Increasing the
pressure increases the heat transfer rate between the reacting molecules and
the rest of the system, reducing this effect.
Condensed-phase rate coefficients can also be affected by
(very high) pressure; this is a completely different effect than fall-off or
chemical-activation. It is often studied using diamond
anvils.
A reaction's kinetics can also be studied with a pressure
jump approach. This involves making fast changes in pressure and observing
the relaxation time of the return to equilibrium.
Equilibrium
While a chemical kinetics is concerned with the rate of a
chemical reaction, thermodynamics determines the extent to which
reactions occur. In a reversible reaction, chemical equilibrium is
reached when the rates of the forward and reverse reactions are equal (the
principle of detailed balance) and the concentrations of the reactants and products no longer change. This is demonstrated
by, for example, the Haber–Bosch process for combining nitrogen and
hydrogen to produce ammonia. Chemical
clock reactions such as the Belousov–Zhabotinsky reaction demonstrate
that component concentrations can oscillate for a long time before finally
attaining the equilibrium.
Free energy
In general terms, the free energy change (ΔG) of a reaction
determines whether a chemical change will take place, but kinetics describes
how fast the reaction is. A reaction can be very exothermic
and have a very positive entropy change but will not happen in practice if the
reaction is too slow. If a reactant can produce two different products, the
thermodynamically most stable one will in general form, except in special
circumstances when the reaction is said to be under kinetic reaction control. The Curtin–Hammett principle applies when
determining the product ratio for two reactants interconverting rapidly, each
going to a different product. It is possible to make predictions about reaction
rate constants for a reaction from free-energy relationships.
The kinetic isotope effect is the difference in
the rate of a chemical reaction when an atom in one of the reactants is
replaced by one of its isotopes.
Chemical kinetics provides information on residence
time and heat transfer in a chemical
reactor in chemical engineering and the molar mass distribution in polymer
chemistry.
Applications
The mathematical models that describe chemical reaction
kinetics provide chemists and chemical engineers with tools to better
understand and describe chemical processes such as food decomposition,
microorganism growth, stratospheric ozone decomposition, and the complex
chemistry of biological systems. These models can also be used in the design or
modification of chemical reactors to optimize product yield, more efficiently
separate products, and eliminate environmentally harmful by-products. When
performing catalytic cracking of heavy hydrocarbons into
gasoline and light gas, for example, kinetic models can be used to find the
temperature and pressure at which the highest yield of heavy hydrocarbons into
gasoline will occur. Kinetics is also a basic aspect of chemistry.
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