1.1 The Structure of an Atom

An atom consists of a tiny dense nucleus surrounded by electrons. The nucleus contains positively charged protons and neutral neutrons. A neutral atom has an equal number of protons and electrons. Atoms can gain electrons and thereby become negatively charged, or they can lose electrons and become positively charged. However, there won't be any change in the number of protons. Most of the mass of an atom is in its nucleus and most of the volume of an atom is occupied by its electrons. The atomic number of an atom equals to the number of protons and the number of electrons. The mass number of an atom is the sum of its protons and neutrons. The number of neutrons of an atom can be varied, so an atom can have different mass number.

1.2 The Distribution of Electrons in an Atom

The electrons are the greatest importance in organic chemistry. A neutral atom of each element contains an equal number of protons and electrons. According to quantum mechanics, the electrons in an atom can be thought of as occupying a set of shells. The first shell is the closest to the nucleus and subsequent shells lies farther from the nucleus. Each shell contains subshells known as atomic orbitals that have a characteristic shape and energy. They occupy a characteristic volume of the space. The atomic orbital that is close to the nucleus, is lower in its energy.

The first shell consists of only s atomic orbital. The second shell consists of s and three degenerate p atomic orbitals and the third shell contains, in addition, five degenerate d atomic orbitals. The fourth and higher shells contain, in addition, seven degenerate f atomic orbitals. Degenerate orbitals are orbitals that have the same energy. Each atomic orbital can have maximum of two electrons. There are only two electrons in the first shell as it has only s atomic orbitals. The second shell can have a total of eight electrons for one s and three p atomic orbitals. The third shell has nine atomic orbitals, one s , three p , and five d atomic orbitals, so eighteen electrons can occupy these nine atomic orbitals. Thirty two electrons can be occupied by the sixteen atomic orbitals of the fourth shell. When the electrons are in the available orbitals with the lowest energy, we call it as the ground-state electronic configuration of the atom. One or more electrons can jump into a higher energy orbital, if energy is applied to the atom in the ground state. We call it as an excited-state electronic configuration.

The following principles are used to determine which orbitals occupy the electrons:

• According to the aufbau principle , an electron always goes into the available lower energy orbital. The relative energies of the atomic orbitals follow:

1 s < 2 s < 2 p < 3 s < 3 p < 4 s < 3 d < 4 p < 5 s < 4 d < 5 p < 6 s < 4 f < 5 d < 6 p < 7 s < 5 f

1s atomic orbital is closer to the nucleus and lower in energy than 2s atomic orbital, which is lower in energy and closer to the nucleus than 3s atomic orbital. While comparing atomic orbitals in the same shell, s atomic orbital is lower in energy than p atomic orbital, and p atomic orbital is lower in energy than d atomic orbital.

• According to the Pauli Exclusion Principle, only two electrons can occupy each atomic orbital, and the two electrons must be of opposite spin. The single electron of a hydrogen atom occupies 1s atomic orbital, the second electron of a helium atom fills the 1s atomic orbital, the third electron of a lithium atom occupies 2s atomic orbital, the fourth electron of a beryllium atom fills the 2s atomic orbital. The fifth electron of a boron atom occupies any one of the three degenerate 2p atomic orbitals.
• According to the Hund's rule, when there are degenerate orbitals, an electron will occupy an empty orbital before it starts to pair up.