Introduction
Electrons
Electrons were discovered by Sir John Joseph Thomson in 1897 in his cathode ray experiment. He subjected residual gas to a high potential difference at a very low pressure. He observed rays travelling in straight lines from the cathode. He called the rays cathode rays. After many experiments of cathode-rays, J.J. Thomson demonstrated the ratio of mass to charge of cathode-rays which is 1.76 x 1011 coulombs per kg. He confirmed that cathode-rays are fundamental particles that have a negative charge. Cathode-rays became known as electrons. Robert Millikan in 1910, through oil-drop experiments, found the value of the charge – 1.6 x 10–19 coulombs.
Proton
- J. Thompson repeated his experiment using a perforated cathode. He observed rays moving in the opposite direction to that of the cathode rays. Experiment also revealed that the rays have the following characteristics:
- J. Thompson also determined the charge/mass ratio, which is 9.85 x 107 coulomb per kg. He called it the proton. R. A. Millikan also found out that the charge is +1.6 x 10-19 coulombs which is the same as that of the electron but opposite charge. The mass of proton is calculated be 1.67 x 10-27 kg. The proton is approximately 1840 times heavier than the electron.
- They travel in straight lines
- They are positively charged
Neutrons
Neutrons were discovered by James Chadwick in 1932, when he demonstrated that penetrating radiation incorporated beams of neutral particles. Neutrons are located in the nucleus with the protons. Along with protons, they make up almost all of the mass of the atom. The number of neutrons is called the neutron number and can be found by subtracting the proton number from the atomic mass number. The number of neutrons does not have to equal that of the protons.
Orbitals and Electronic Configuration
Orbital is a region or space where the probability of finding an electron is high while shell is an imaginary line on which electron revolves. Each shell is divided into orbitals. These orbitals are called s, p, d and f orbitals. Each shell contains its orbitals e.g. K shell contains only s orbital, while L shell contains s and p orbital, M shell contains s, p and d orbitals, N shell contains s, p, d and f orbital.
Each orbital has a maximum number of electrons it can hold. The sub-level s, has a maximum of 2 electrons. The sub-level p, has a maximum of 6 electrons. The sub-level d, has a maximum of 10 electrons. The sub-level f, has a maximum of 14 electrons.
Orbital Types in shell
| Shell | Shell Number | Orbital Types |
| K | 1 | 1s |
| L | 2 | 2s, 2p |
| M | 3 | 3s, 3p, 3d |
| N | 4 | 4s, 4p, 4d, 4f |
The sequence of filling up the orbitals with electrons is as follows:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4s etc.
Quantum Numbers
Each electron in an atom is described by four different quantum numbers. The first three (n, l, m) specify the particular orbital of interest, and the fourth (ms) specifies how many electrons can occupy that orbital.
Principal Quantum Number (n): n = 1, 2, 3, …, ∞
Specifies the energy of an electron and the size of the orbital. All orbitals that have the same value of n are said to be in the same shell (level). For a hydrogen atom with n=1, the electron is in its ground state; if the electron is in the n=2, it is in an excited state.
Subsidiary or Azimuthal Quantum Number (l): l = 0 to (n-1).
Specifies the shape of an orbital with a particular principal quantum number. The subsidiary quantum number divides the shells into smaller groups of orbitals called subshells (sublevels). The electrons with subsidiary quantum numbers 0, 1, 2 and 3 are usually referred to as the s, p, d, and f electrons respectively. The subshell with n=2 and l=1 is the 2p subshell; if n=3 and l=0, it is the 3s subshell, and so on. The value of l also has a slight effect on the energy of the subshell.
Magnetic Quantum Number (m): m = –l, …, 0, …, +l.
This number divides the subshell into individual orbitals which hold the electrons; there are 2l+1 orbitals in each subshell. Thus the s subshell has only one orbital, the p subshell has three orbitals, and so on.
Spin Quantum Number (s): s = +½ or -½.
An electron can spin in only one of two directions (sometimes called up and down).
| s orbitals | p orbitals | d orbitals | f orbitals | |
| L | 0 | 1 | 2 | 3 |
| M | 0 | -1, 0, +1 | -2, -1, 0, +1, +2 | -3, -2, -1, 0, +1, +2, +3 |
| Number of orbitals in designated subshell | 1 | 3 | 5 | 7 |
| n | l | M | Number of orbitals |
Orbital Name |
Number of electrons |
| 1 | 0 | 0 | 1 | 1s | 2 |
| 2 | 0 | 0 | 1 | 2s | 2 |
| 1 | -1, 0, +1 | 3 | 2p | 6 | |
| 3 | 0 | 0 | 1 | 3s | 2 |
| 1 | -1, 0, +1 | 3 | 3p | 6 | |
| 2 | -2, -1, 0, +1, +2 | 5 | 3d | 10 | |
| 4 | 0 | 0 | 1 | 4s | 2 |
| 1 | -1, 0, +1 | 3 | 4p | 6 | |
| 2 | -2, -1, 0, +1, +2 | 5 | 4d | 10 | |
| 3 | -3, -2, -1, 0, +1, +2, +3 | 7 | 4f | 14 |
Rules and Principles of Filling Electron
There are three basic principles or rules that govern electron filling into orbitals. These are:
Pauli Exclusion Principle: This principle state that two electrons in the same orbital of an atom cannot have same value for all quantum numbers. This implies that if two electrons are considered while they may have the same values of n, l and m, they will differ in s because while one is n = 1, l = 0, m = 0 and s = +1/2, and the other will be n = 1, l = 0, m = 0 and s = -1/2. This indicates that no two electrons in any atom can behave in an identical manner.
Hund’s Rule: This rule states that electrons occupy each orbital singly first before pairing takes place in a degenerate orbital.
Aufbau Principle: In the building up of atom, electrons enter into orbitals in order of increasing energy. This means the electron are fed into orbitals starting at the lowest energy level before filling higher energy level.
The table below shows the modern electronic configuration of the first twenty elements.
| Name | Atomic Number | Electron Configuration |
| Hydrogen | 1 | 1s1 |
| Helium | 2 | 1s2 |
| Lithium | 3 | 1s2 2s1 |
| Beryllium | 4 | 1s2 2s2 |
| Boron | 5 | 1s2 2s22p1 |
| Carbon | 6 | 1s2 2s22p2 |
| Nitrogen | 7 | 1s2 2s22p3 |
| Oxygen | 8 | 1s2 2s22p4 |
| Fluorine | 9 | 1s2 2s22p5 |
| Neon | 10 | 1s2 2s22p6 |
| Sodium | 11 | 1s2 2s22p63s1 |
| Magnesium | 12 | 1s2 2s22p63s2 |
| Aluminum | 13 | 1s2 2s22p63s23p1 |
| Silicon | 14 | 1s2 2s22p63s23p2 |
| Phosphorus | 15 | 1s2 2s22p63s23p3 |
| Sulphur | 16 | 1s2 2s22p63s23p4 |
| Chlorine | 17 | 1s2 2s22p63s23p5 |
| Argon | 18 | 1s2 2s22p63s23p6 |
| Potassium | 19 | 1s2 2s22p63s23p64s1 |
| Calcium | 20 | 1s2 2s22p63s23p64s2 |
EXERCISES
Lets see how much you’ve learnt, attach the following answers to the comment below
- Determine the maximum number of electrons that can occupy the principal energy level N of an atom. (a) 18 (b) 8 (c) 24 (d) 32
- Which element has an electronic configuration of 1s22s22p63s1? (a) chlorine (b) calcium (c) potassium (d) sodium
- There are _______ types of quantum numbers. (a) 3 (b) 4 (c) 5 (d) 6
- _______ states that electrons occupy each orbital singly first before pairing takes place in a degenerate orbital. (a) Hund’s rule (b) Pauli exclusive principle (c) Aufbau principle (d) hooke’s principle
- Which of these scientist discovered electron? (a) R. A. Millikan (b) neil bohr (c) J. J. Thompson (d) James Chadwick
