Atoms

An atom (Greek language|Greek άτομον from ά: non and τομον: divisible) is a submicroscopic structure found in all ordinary matter. It is the smallest unit of an chemical element to retain all the chemical properties of that element. The word atom originally meant a smallest possible particle of matter. They thought that you couldn't divide the atom any further. Later, the objects that had been called atoms were found to be further divisible into smaller subatomic particles, but the word atom nonetheless continues to refer to them.

Most atoms are composed of three types of massive subatomic particles which govern their external properties:

electrons, which have a negative electric charge and are the least massive of the three;
protons, which have a positive electric charge and are about 1836 times more massive than electrons; and
neutrons, which have no electric charge and are about 1838 times more massive than electrons.

Together, protons and neutrons form the atomic nucleus of an atom, which is surrounded by the electrons.

Atoms can differ in the number of each of the subatomic particles they contain. Atoms of the same element have the same number of protons, but can differ in the number of neutrons, in which case they are called isotopes of that element. Atoms are electrostatically neutral if they have an equal number of protons and electrons. Electrons that are furthest from the nucleus are less tightly bound to the atom and are relatively easily transferred to other nearby atoms or even shared between atoms. Atoms which have either given away or taken electrons from another atom are called ions. Particles in the nucleus also sometimes escape atoms, in a process known as radioactive decay.

Atoms are the fundamental building blocks of chemistry, and are conserved in chemical reactions. Atoms are able to chemically bond into molecules and other types of chemical compounds. Molecules are made up of multiple atoms; for example, a molecule of water is a combination of two hydrogen and one oxygen atom.

Bonding / Reactions

Exothermic reactions

According to energy balance criteria, that is, chemical reaction equilibria criteria, any closed system will tend to minimize its free energy. Without any outside influence, any reaction mixture, too, will try to do the same. For many cases, an analysis of the enthalpy of the system will give a decent account of the energetics of the reaction mixture. The enthalpy of a reaction is calculated using standard reaction enthalpy|reaction enthalpies and the Hess' law of constant heat summation. Many of these enthalpies may be found in beginners' books on thermodynamics. For example, consider the combustion of Organic Chemistry/Alkanes and cycloalkanes/Alkanes/Methane|methane in oxygen:

CH₄ + 2 O₂ → CO₂ + 2 H₂O

By calculating the amounts of energy required to break all the bonds on the left ("before") and right ("after") sides of the equation using collected data, it is possible to calculate the energy difference between the reactants and the products. This is referred to as ΔH, where Δ (Delta) means difference, and H stands for enthalpy, a measure of energy which is equal to the heat transferred at constant pressure. ΔH is usually given in units of Joule|kilojoules (kJ) or in Calorie|kilocalories (kcal).

If ΔH is negative for the reaction, then energy has been released often in the form of heat. This type of reaction is referred to as an exothermic reaction (literally, outside heat, or throwing off heat). An exothermic reaction is more favourable and thus more likely to occur. An example reaction is combustion, known from everyday experience, since burning gas in air produces heat.

Endothermic reactions

w:Image:Endothermic.png|400px|thumb|right|A sketch of an endothermic reaction

A reaction may have a positive ΔH. If a reaction has a positive ΔH, it consumes energy as the reaction moves towards completion. This type of reaction is called an endothermic reaction (literally, inside heat, or absorbing heat).

The above rule, "Exothermic reactions are favourable", is usually true. However, there may be situations where exothermic reactions may not be favourable. This happens when the stability obtained due to loss of enthalpy is off set by a corresponding decrease in entropy (a measure of disorder). The exact rule is that a reaction is favourable when the Organic Chemistry/Introduction to reactions/Gibbs free energy|Gibbs free energy of that reaction is negative where ΔG = ΔH − TΔS; ΔG being the change in Gibbs free energy, ΔH being the change in enthalpy, and ΔS is the change in entropy

A reaction is called spontaneous process|spontaneous if its thermodynamically favoured, by that meaning that it causes a net increase on entropy. Spontaneous reactions (in opposition to non-spontaneous reactions) do not need external perturbations (such as energy supplement) to happen. In a system at chemical equilibrium, it is expected to have larger concentrations of the substances formed by the spontaneous direction of the process.

Thus, in a global isolated system (which it strictly isn't, see entropy), spontaneous reactions may be understood to occur without Intelligence_%28trait%29|human interference. Most spontaneous reactions in this system are exothermic (such as rusting) or metamorphosis, thus increasing the global entropy, though photosynthesis is an important exception (in a global system).