Chemical reactionChemical reactions such as combustion in the fire, fermentation
and the reduction of ores to metals were known since antiquity. Initial
theories of transformation of materials were developed by Greek
philosophers, such as the Four-Element Theory of Empedocles
stating that any substance is composed of the four basic elements –
fire, water, air and earth. In the Middle Ages, chemical
transformations were studied by Alchemist. They attempted, in particular, to convert lead into gold, for which purpose they used reactions of lead and lead-copper alloys with sulfur.
The production of chemical substances that do not normally occur in nature has long been tried, such as the synthesis of sulfuric and nitric acids attributed to the controversial alchemist Jābir ibn Hayyān. The process involved heating of sulfate and nitrate minerals such as copper sulfate, alum and saltpeter. In the 17th century, Johann Rudolph Glauber produced hydrochloric acid and sodium sulfate by reacting sulfuric acid and sodium chloride. With the development of the lead chamber processLeblanc process, allowing large-scale production of sulfuric acid and sodium carbonate,
respectively, chemical reactions became implemented into the industry.
Further optimization of sulfuric acid technology resulted in the contact process in 1880s, and the Haber process was developed in 1909–1910 for ammonia synthesis in 1746 and the
From the 16th century, researchers including Jan Baptist van Helmont, Robert BoyleIsaac Newton tried to establish theories of the experimentally observed chemical transformations. The phlogiston theory was proposed in 1667 by Johann Joachim Becher.
It postulated the existence of a fire-like element called "phlogiston",
which was contained within combustible bodies and released during combustion. This proved to be false in 1785 by Antoine Lavoisier who found the correct explanation of the combustion as reaction with oxygen from the air.and
Joseph Louis Gay-Lussac
recognized in 1808 that gases always react in a certain relationship
with each other. Based on this idea and the atomic theory of John Dalton, Joseph Proust had developed the law of definite proportions, which later resulted in the concepts of stoichiometry and chemical equations.
Regarding the organic chemistry, it was long believed that compounds obtained from living organisms were too complex to be obtained synthetically. According to the concept of vitalism,
organic matter was endowed with a "vital force" and distinguished from
inorganic materials. This separation was ended however by the synthesis
of ureaFriedrich Wöhler in 1828. Other chemists who brought major contributions to organic chemistry include Alexander William Williamson with his synthesis of ethers and Christopher Kelk Ingold, who, among many discoveries, established the mechanisms of substitution reactions. from inorganic precursors by
Chemical equations are used to graphically illustrate chemical reactions. They consist of chemical or structural formulas
of the reactants on the left and those of the products on the right.
They are separated by an arrow (→) which indicates the direction and
type of the reaction. The tip of the arrow points in the direction in
which the reaction proceeds. A double arrow () pointing in opposite directions is used for equilibrium reactions. Equations should be balanced according to the stoichiometry,
the number of atoms of each species should be the same on both sides of
the equation. This is achieved by scaling the number of involved
molecules (A, B, C and D in a schematic example below) by the appropriate integers a, b, c and d.
More complex reactions are represented by reaction schemes, which in
addition to starting materials and products show important
intermediates or transition states.
Also, some relatively minor additions to the reaction can be indicated
above the reaction arrow; examples of such additions are water, heat,
illumination, a catalyst, etc. Similarly, some minor products can be
placed below the arrow, often with a minus sign.
The elementary reaction is the smallest division into which a chemical reaction can be decomposed to, it has no intermediate products.
Most experimentally observed reactions are built up from many
elementary reactions that occur in parallel or sequentially. The actual
sequence of the individual elementary reactions is known as reaction mechanism.
An elementary reaction involves a few molecules, usually one or two,
because of the low probability for several molecules to meet at a
The most important elementary reactions are unimolecular and
bimolecular reactions. Only one molecule is involved in a unimolecular
reaction; it is transformed by an isomerization or a dissociation
in one or more other molecules. Such reaction requires addition of
energy in the form of heat or light. A typical example of a
unimolecular reaction is the cis–transisomerization, in which the cis-form of a compound converts to the trans-form or vice versa.
In a typical dissociation reaction, a bond in a molecule splits resulting in two molecular fragments. The splitting can be homolytic or heterolytic. In the first case, the bond is divided so that each product retains an electron and becomes a neutral radical. In the second case, both electrons of the chemical bond remain with one of the products, resulting in charged ions. Dissociation plays an important role in triggering chain reactions, such as hydrogen-oxygen or polymerization reactions.
- Dissociation of a molecule AB into fragments A and B
For bimolecular reactions, two molecules collide and react with each other. Their merger is called chemical synthesis or an addition reaction.
Another possibility is that only a portion of one molecule is
transferred to the other molecule. This type of reaction occurs, for
example, in redox and acid-base reactions. In redox reactions, the
transferred particle is an electron, whereas in acid-base reactions it
is a proton. This type of reaction is also called metathesis.
- NaCl(aq) + AgNO3(aq) → NaNO3(aq) + AgCl(s)
Most chemical reactions are reversible, that is they can and do run
in both directions. The forward and reverse reactions are competing
with each other and differ in reaction rates.
These rates depend on the concentration and therefore change with time
of the reaction: the reverse rate gradually increases and becomes equal
to the rate of the forward reaction, establishing the so-called
chemical equilibrium. The time to reach equilibrium depends on such
parameters as temperature, pressure and the materials involved, and is
determined by the minimum free energy. In equilibrium, the Gibbs free energy must be zero. The pressure dependence can be explained with the Le Chatelier's principle.
For example, an increase in pressure due to decreasing volume causes
the reaction to shift to the side with the fewer moles of gas.
The reaction yield stabilized at equilibrium, but can be increased
by removing the product from the reaction mixture or increasing
temperature or pressure. Change in the initial concentrations of the
substances does not affect the equilibrium.
Chemical reactions are largely determined by the laws of thermodynamics. Reactions can proceed by themselves if they are exergonic,
that is if they release energy. The associated free energy of the
reaction is composed of two different thermodynamic quantities, enthalpy and entropy:
- G: free energy, H: enthalpy, T: temperature, S: entropy, Δ: difference
Reactions can be exothermic, where ΔH is negative and energy is released. Typical examples of exothermic reactions are precipitation and crystallization, in which ordered solids are formed from disordered gaseous or liquid phases. In contrast, in endothermic
reactions, heat is consumed from the environment. This can occur by
increasing the entropy of the system, often through the formation of
gaseous reaction products, which have high entropy. Since the entropy
increases with temperature, many endothermic reactions preferably take
place at high temperatures. On the contrary, many exothermic reactions
such as crystallization occur at low temperatures. Changes in
temperature can sometimes reverse the direction of a reaction, as in
the Boudouard reaction:
This reaction between carbon dioxide and carbon to form carbon monoxide is endothermic at temperatures above approximately 800 °C and is exothermic below this temperature.
Reactions can also be characterized with the internal energy which takes into account changes in the entropy, volume and chemical potential. The latter depends, among other things, on the activities of the involved substances.
- U: internal energy, S: entropy, p: pressure, μ: chemical potential, n: number of molecules, d: small change sign
The speed at which a reactions takes place is studied by reaction kinetics. The rate depends on various parameters, such as:
concentrations, which usually make the reaction happen at a faster rate
if raised through increased collisions per unit time. Some reactions,
however, have rates that are independent of reactant concentrations. These are called zero order reactions.
- Surface area
available for contact between the reactants, in particular solid ones
in heterogeneous systems. Larger surface areas lead to higher reaction
- Pressure –
increasing the pressure decreases the volume between molecules and
therefore increases the frequency of collisions between the molecules.
- Activation energy,
which is defined as the amount of energy required to make the reaction
start and carry on spontaneously. Higher activation energy implies that
the reactants need more energy to start than a reaction with a lower
which hastens reactions if raised, since higher temperature increases
the energy of the molecules, creating more collisions per unit time,
- The presence or absence of a catalyst.
Catalysts are substances which change the pathway (mechanism) of a
reaction which in turn increases the speed of a reaction by lowering
the activation energy needed for the reaction to take place. A catalyst is not destroyed or changed during a reaction, so it can be used again.
- For some reactions, the presence of electromagnetic radiation, most notably ultraviolet light, is needed to promote the breaking of bonds to start the reaction. This is particularly true for reactions involving radicals.
Several theories allow calculating the reaction rates at the
molecular level. This field is referred to as reaction dynamics. The
rate v of a first-order reaction, which could be disintegration of a substance A, is given by:
Its integration yields:
Here k is first-order rate constant having dimension 1/time, [A](t) is concentration at a time t and [A]0
is the initial concentration. The rate of a first-order reaction
depends only on the concentration and the properties of the involved
substance, and the reaction itself can be described with the
More than one time constant is needed when describing reactions of
higher order. The temperature dependence of the rate constant usually
follows the Arrhenius equation:
where Ea is the activation energy and kB is the Boltzmann constant. One of the simplest models of reaction rate is the collision theory. More realistic models are tailored to a specific problem and include the transition state theory, the calculation of the potential energy surface, the Marcus theory and the Rice–Ramsperger–Kassel–Marcus (RRKM) theory.