CHEMICAL
KINETICS
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.
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 influences the rate of its transformation into products. The reactions
which involve lesser bond rearrangement proceed faster than the reactions which
involve larger bond rearrangement.
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
only occur at their area of contact, in the case of a liquid and a gas, at the
surface of the liquid. 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 doesn't 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
Concentration
plays a very important role in reactions, because according to the collision theory of
chemical reactions, molecules must collide in order to react together. As the
concentration of the reactants increases, the frequency of
the molecules colliding increases, striking each other more frequently by being
in closer contact at any given point in time. Think of two reactants being in a
closed container. All the molecules contained within are colliding constantly.
By increasing the amount of one or more of the reactants it causes these
collisions to happen more often, increasing the reaction rate.
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 Q10
(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 rate of
an equilibrium process.
CATALYSTS
Generic potential energy diagram showing the effect of a catalyst in an 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.
EQUILIBRIUM
While 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 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 generally 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.
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