The Atom

Atoms

A Brief Introduction to The Atom

The word "atom" comes from the Greek "atomos" (named by Democritus) which means "indivisible" as it was thought in ancient times that atoms were the fundamental "units" of nature- they could not be split up into smaller constituents. However, now we know that atoms are not indivisible and that they are comprised of 3 sub-atomic particles known as protons, neutrons and electrons; where the protons and neutrons are comprised of even smaller fundamental particles called quarks and electrons are a type of fundamental particle in a group known as leptons- this is to do with particle physics, however, and so is not too relevant in this section.

Atoms are tiny constituents of matter (of molecules of compounds) which make up the entirety of the matter world we can see- all the stars, planets, nebulae etc. They consist of 3 main sub-atomic particles known as protons, neutrons and electrons; as mentioned earlier. Together, these 3 particles identify an element. The number of protons the nucleus of an atom contains gives its atomic number (z) on the periodic table- for example, the element Hydrogen has just one proton in its nucleus (giving z = 1) and 0 neutrons. The number of protons + the number of neutrons in an atom's nucleus gives the mass number (A_r) of the element/atom:

$A_{r}=N_{p}+N_{n}=z+N_{n}$

An atom with a neutral resultant ionic charge will have an equal number of protons and electrons, when C = 0:

$z=N_{e}$

The Atomic Structure

Atoms are comprised of mainly "empty space"- i.e. most of the space of an atom is a vaccuum. The protons and neutrons exist in the centre of the atom in the atomic nuclei- which is highly dense due to the compactness of the protons and neutrons- whereas the electrons exist in quantised energy levels surrounding the nucleus.

Note: It is a common misconception that electrons orbit the nucleus, there is uncertainty in their pattern of movement and so we say that they move in quantised energy levels surrounding the nucleus. If you're interested, we cannot know whether electrons orbit the nucleus due to Heisenberg's Uncertainty Principle which states that "we cannot know the exact position and momentum of a particle at any one time", or written mathematically:

$\Delta x\Delta p\geq {\frac {\hbar }{2}}$

Anyway, the size of an atom is often measured in nanometres (10^-9 m) or even smaller. Even though they are so small, the whole atom is huge compared to the size of the nucleus situated at the centre of the atom, this analogy is a good one to use to put the size into context: If you were to expand an atom out to the size of Wembley Football Stadium, then the nucleus would be the size of a garden pea on the centre spot of the football pitch. This puts the structure of an atom into context and clearly shows that the nucleus of an atom is extremely dense seeing that, on average 99.9% of the mass of an atom is concentrated in the nucleus of the atom.

The Nucleus and Relative Masses

All the bound protons and neutrons in an atom make up a tiny atomic nucleus, and are collectively called nucleons. The radius of a nucleus is approximately equal to

$1.07{\sqrt[{3}]{A}}$

where A is the total number of nucleons, given by

$\displaystyle \sum _{n=1}^{k}A_{n}=N_{p}+N_{n}$

Atoms of the same element have the same number of protons, called the atomic number. Within a single element, the number of neutrons may vary, determining the isotope of that element. The total number of protons and neutrons determine the nuclide. The number of neutrons relative to the protons determines the stability of the nucleus, with certain isotopes undergoing radioactive decay.

The neutron and the proton are different types of fermions. The Pauli exclusion principle is a quantum mechanical effect that prohibits identical fermions, such as multiple protons, from occupying the same quantum physical state at the same time. Thus every proton in the nucleus must occupy a different state, with its own energy level, and the same rule applies to all of the neutrons. This prohibition does not apply to a proton and neutron occupying the same quantum state.

For atoms with low atomic numbers, a nucleus that has a different number of protons than neutrons can potentially drop to a lower energy state through a radioactive decay that causes the number of protons and neutrons to more closely match. As a result, atoms with roughly matching numbers of protons and neutrons are more stable against decay. However, with increasing atomic number, the mutual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of the nucleus, which modifies this trend.

The relative masses of the sub-atomic particles in an atom are as follows:

Proton = A_r = 1, Neutron = A_r = 1, Electron = A_r = ${\frac {1}{1870}}$

As the relative mass of an electron is so small, it is often ignored in calculations concerning the mass of an atom and the periodic table masses.

Relative Charges

The relative charges of the sub-atomic particles of an atom are also given here:

Proton = C = +1, Neutron = C = 0, Electron = C = -1.

As the neutron has a neutral charge it is ignored in ionic charge and oxidation/reduction calculations. As an aside, the protons and neutrons obtain their charges from the charges on the fundamental particles known as quarks- specifically of the flavours "up" and "down" quarks:

Up Quarks have an electric charge

$C=+{\frac {2}{3}}$

and Down Quarks have an electric charge

$C=-{\frac {1}{3}}$

from here it is easy to see what protons and neutrons are specifically comprised of. A proton is made up of 2 up quarks and 1 down quark:

$C={\frac {2}{3}}+{\frac {2}{3}}-{\frac {1}{3}}={\frac {3}{3}}=1$

and neutrons are comprised of 1 up quark and 2 down quarks:

$C={\frac {2}{3}}-{\frac {1}{3}}-{\frac {1}{3}}=0$

where C obviously represents the relative charge.