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The Concept Of Atoms And Molecules

The Concept Of Atoms And Molecules Long ago, ancient Greek philosophers had pondered on the idea that matter could be composed of very small bits called atoms. They however did not go beyond this. In the early part of the nineteenth century (1808), an English chemist named John Dalton came up with a set of testable suggestions about the atom, and it became known as Dalton’s atomic theory.

The ideas about the theory states as follows: 1. Matter is made up of small, indivisible particles called atoms. 2. Atoms can neither be created nor destroyed. 3. Atoms of a particular element are all exactly alike in every way and are different from atoms of all other elements. 4. Chemical combination is the union of atoms in small whole numbers.

This theory stood for about a century before certain faults were discovered in most of its ideas in the early part of the twentieth century . We shall highlight these faults in this lecture. Distinction between Atoms and Molecules The atomistic philosophers made no clear distinction between atoms and molecules because they did not know about the chemical elements.

The distinction between atoms and molecules became important when the concept of chemical elements was accepted (this concept was proposed by Robert Boyle, and it states that all materials are composed of a small number of elements). Based on the later development of the molecular theory (also known as the kinetic theory), a clear distinction between atoms and molecules has been established.

An atom can be defined as the smallest particle of an element, which can take part in chemical change or reaction and still retain the identity of the element. Notice that the atom is not the smallest particle of an element, as there are electrons, protons and neutrons which are smaller, but, it is the smallest which can be involved in chemical change and still retain the chemical properties of the element.

Electrons, protons and neutrons are fundamental particles of matter, they are found alike in all substances. Hence, they do not retain the identities of their respective elements when they are involved in chemical change. A molecule can be defined as the smallest, uncharged particle of an element or compound, which moves about as a whole, such that its component parts do not become detached during the motion.

In other word, a molecule is the smallest part of a substance, which has a separate existence. A molecule of a compound consists of more than one kind of atom (this is because compounds are formed between two or more elements). E.g. a molecule of water consists of two atoms of hydrogen and one atom of oxygen. A molecule of an element consists of one or more atoms of the same kind.

E.g. a molecule of argon (Ar) and helium (He) consist of one atom of argon and one atom of helium respectively, while a molecule of hydrogen (H₂), chlorine (Cl₂) and ozone (O₃) consist of two, two and three atoms of hydrogen, chlorine and oxygen respectively.

Note: - All gases consist of molecules. This is why it is possible for a quantity of a gas to expand and occupy every space of the container in which it is kept (irrespective of the size of the container)

- Since every gas becomes a liquid when cooled sufficiently, and every liquid turns to solid when cooled to its freezing point, it is therefore not only gases, but at least some liquids and solids also consist of molecules. 

The electrons and protons haven been found as the fundamental particles of an atom, it was now left to describe their arrangement. J.J. Thomson described a model of an atom which suggested that an atom is a uniform sphere of about 10^-8 cm radius, made up of even distribution of protons, with the electrons embedded, such as to give the most stable electrostatic arrangement.

This arrangement was found to be Inadequate, Diagram and was totally discarded in the light of a classic experiment performed by Rutherford. Rutherford directed an alpha particle (α- Helium ion, He²⁺) onto a thin foil of metal. He observed that most of the particles passed through the foil either undeflected or deflected by small angles (30⁰). A few of the particles were however deflected by very large angles (up to 180⁰) along their original paths.

The explanation of this observation is given thus: The fact that most of the α-particles passed through the foil either undeflected or deflected by small angles, suggests that the mass and the protons of an atom are not uniformly distributed (as proposed by J.J. Thompson). Otherwise, the α-particles would have encountered large repulsions uniformly distributed throughout the metal.

The large deflections observed for a few of these very energetic α-particles suggests that there must be a seat in the atom of concentration of large electrical force, which must be produced by a body of same charge as the α-particles and which is of significant mass - not electrons (electrons are very light and would be swept away by the heavier α-particles).

For the fact that only a few of the α-particles are deflected suggest that the large electrical force must be confined to a very small region of space. Therefore, Rutherfold suggested a model for the atom as consisting of a dense central core or nucleus, in which all the positive charges, (i.e. protons) and almost all the mass of the atom are concentrated; and around which equal number of negatively charged electrons move in regular, defined orbitals, completely balancing the positive charges of the nucleus - this is the modern atomic model.

The atomic number of an element, i.e., the number of protons present in the atom was determined by H.G. Moseley. He compared the wavelenghts of the characteristic X-ray emitted by a number of elements and thus arrived at values for the nuclear charges (produced by the protons contained) of the atoms.

These values became the atomic number of the respective element. Moseley’s discovery provided proof that there is in the atom a fundamental quantity which increases by regular steps as we pass from one element in the periodic system to the next. This quantity is the charge of the nucleus, i.e., the number of protons in the nucleus.

Observing that a helium atom, with atomic number 2 (containing 2 protons) actually have an atomic weight of approximately 4 amu instead of 2 amu expected of a 2 proton atom. It was believed that there must be some other fundamental unit in the atom whose weight must be adding to the total atomic weight of the element, but do not contribute anything to the electric charge.

Rutherford suggested it to be an electron/proton pair in the nucleus that is very much joined together that it acts like a single particle of no charge – the particle was called a neutron.

However, James Chadwick, a young scientist working in Rutherford’s laboratory, later showed neutrons not to be as described by Rutherford, but as single discrete particles with no charge, but with a mass of 1.0087 amu, almost equal to that of the proton. Therefore, the mass of an atom is the sum of its number of protons and its number of neutrons.