Atoms, the building block of elements, consist of a nucleus surrounded by a cloud of orbiting electrons. The nucleus consists of positively charged protons and neutral neutrons, and so has a net positive charge that holds the negatively charged electrons, which revolve around it, in position by an electrostatic attraction.
The charges on the proton and electron are equal and opposite (1.602 × 10−19 coulombs) and the number of electrons and protons are equal and so the atom overall is electrically neutral. Protons and neutrons have approximately the same mass, 1.67 × 10−27 kg, whereas an electron has a mass of 9.11 × 10−31 kg, nearly 2000 times less.
These relative densities mean that the size of the nucleus is very small compared to the size of the atom. Although the nature of the electron cloud makes it difficult to define the size of atoms precisely, helium has the smallest atom, with a radius of about 0.03 nanometers, while caesium has one of the largest, with a radius of about 0.3 nanometres.
An element is characterised by:
- the atomic number, which is the number of protons in the nucleus, and hence is also the number of electrons in orbit;
- the mass number, which is sum of the number of protons and neutrons. For many of the lighter elements these numbers are similar and so the mass number is approximately twice the atomic number, though this relationship breaks down with increasing atomic number. In some elements the number of neutrons can vary, leading to isotopes; the atomic weight is the weighted average of the atomic masses of an element’s naturally occurring isotopes.
Another useful quantity when we come to consider compounds and chemical reactions is the mole, which is the amount of a substance that contains 6.023 × 1023 atoms of an element or molecules of a compound (Avogadro’s number).
This number has been chosen because it is the number of atoms that is contained in the atomic mass (or weight) expressed in grams. For example, carbon has an atomic weight of 12.011, and so 12.011 grams of carbon contain 6.023 × 1023 atoms.
The manner in which the orbits of the electrons are distributed around the nucleus controls the characteristics of the element and the way in which atoms bond with other atoms of the same element and with atoms from different elements.
For our purposes it will be sufficient to describe the structure of the so-called Bohr atom, which arose from developments in quantum mechanics in the early part of the 20th century. This overcame the problem of explaining why negatively charged electrons would not collapse into the positively charged nucleus by proposing that electrons revolve around the nucleus in one of a number of discrete orbitals or shells, each with a defined or quantised energy level.
Any electron moving between energy levels or orbitals would make a quantum jump with either emission or absorption of a discrete amount or quantum of energy.
Each electron is characterised by four quantum numbers:
- the principal quantum number (n = 1, 2, 3, 4 . . .), which is the quantum shell to which the electron belongs, also denoted by K, L, M, N . . . , corresponding to n = 1, 2, 3, 4 . . . ;
- the secondary quantum number (l = 0, 1, 2 . . . n − 1), which is the sub-shell to which the electron belongs, denoted by s, p, d, f, g, h for l = 1, 2, 3, 4, 5, 6, according to its shape;
- the third quantum number (ml ), which is the number of energy states within each sub-shell, the total number of which is 2l +1;
- the fourth quantum number (ms) which describes the electron’s direction of spin and is either +1/2 or −1/2.
The number of sub-shells that occur within each shell therefore increases with an increase in the principal quantum number (n), and the number of energy states within each sub-shell (ml ) increases with an increase in the secondary quantum number (l).
Table 1 shows how this leads to the maximum number electrons in each shell for the first four shells. Each electron has a unique set of quantum numbers and with increasing atomic number, and hence increasing number of electrons, the shells and sub-shells fill up progressively, starting with the lowest energy state.
Atomic Structure of Elements
The one electron of hydrogen is therefore in the only sub-shell in the K shell (denoted as 1s1 ), the two electrons of helium are both in this same shell (denoted as 1s2 ) and in lithium, which has three electrons, two are in the 1s1 shell and the third is in the 2s1 shell.
By convention, the configuration of lithium is written as 1s2 2s1 . The configuration of subsequent elements follows logically (for example, sodium with 11 electrons is 1s2 2s2 2p6 3s1 ). The structures of these elements are illustrated in Fig. 1.
An extremely important factor governing the properties of an element is the number of electrons in the outermost shell (known as the valence electrons), since it is these that are most readily available to form bonds with other atoms.
Groups of elements with similar properties are obtained with varying atomic number but with the same number of outer shell electrons. For example, the ‘alkali metals’ lithium, sodium, potassium, rubidium and caesium all have one electron in their outermost shell, and all are capable of forming strong alkalis.
A further factor relating to this is that when the outermost electron shell is completely filled the electron configuration is stable. This normally corresponds to the s and p states in the outermost shell being filled by a total of eight electrons; such octets are found in neon, argon, krypton, xenon etc., and these ‘noble gases’ form very few chemical compounds for this reason.
The exception to the octet rule for stability is helium; the outermost (K) shell only has room for its two electrons.
The listing of the elements in order of increasing atomic number and arranging them into groups of the same valence is the basis of the periodic table of the elements, which is an extremely convenient way of categorising the elements and predicting their likely properties and behaviour.
As we will see in the next article, the number of valence electrons strongly influences the nature of the interatomic bonds.
Thanks for reading about “atomic structure of elements.”