Atomic structure and chemical bonding: Bohr model, spectrum of hydrogen atom, quantum numbers; Wave-particle duality, de Broglie hypothesis; Uncertainty principle; Quantum mechanical picture of hydrogen atom (qualitative treatment), shapes of s, p and d orbitals; Electronic configurations of elements (up to atomic number 36); Aufbau principle; Pauli's exclusion principle and Hund's rule; Orbital overlap and covalent bond; Hybridisation involving s, p and d orbitals only; Orbital energy diagrams for homonuclear diatomic species; Hydrogen bond; Polarity in molecules, dipole moment (qualitative aspects only); VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral).
The lecture is based on chapter 4 : the Structure of Atom in Fundamentals of Chemistry by Fred H. Redmore, Prentice Hall 1979.
In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics. The energy levels proposed by Bohr were calculated and it was further modeled that each energy level or orbit has sublevels and each sublevels has one or more orbitals.
Orbit - sublevel – Orbital
Electrons are moving around the nucleus and orbitals represent a space where a pair of electrons is most likely to be found. This means that electrons may sometimes go out of the orbital but most of the time they are found in the orbital.
The quantum mechanics calculations gave out the result that there is a limit to the number of electrons that can occupy a given orbit or energy level.
The orbits are called as shells. The energy level of orbits or shells increases as they increase in distance from the nucleus of the atom. The orbits or shells are represented by numbers as 1,2,3,4,5,6 or 7. They are represented by letters as K,L,M,N,O,P,Q.
It is found that the maximum number of electrons in each energy level is equal to 2n2 where n is the number of energy level.
Therefore energy level 1 will have 2 electrons.
Energy level 2 will have 2*4 = 8 electrons
Energy level 3 will have 2*9 = 18 electrons.
Sublevel of an Orbit
The energy levels, or orbits or shells are further divided into sublevels, or subshells. These subshells are designated by letters: s for the first possible sublevel, p for the second possible sublevel, d for the third, f for the fourth, g for the fifth, and from here on they simply go in alphabets.
The number of sublevels of each energy level is equal to the number of the energy levels. This means energy level 1, the K shell will have only one sub levels – s sublevel. The energy level 2, the L shell will have 2 sub levels – s and p.
Sublevels have further divisions called orbitals. Electrons are found in these orbitals. Each orbital contains two electrons.
“s” sublevel has only one orbital. “p” sublevel has 3 orbitals. “d” sublevel has 5 orbitals. “f” sublevel has 7 orbitals.
As each orbital can hold two electrons, orbitals of s can hold two electrons. The orbitals are of p sublevel are named as px and py and pz. The orbitals of p contain 6 electrons. The orbitals of d are 5. The orbitals of d are named as dxy, dxz,dyz,dx2-y2 and dz2. The d sublevel orbitals contain 10 electrons.
The two electons in each orbital spin in different directions.
Shape of Orbitals
Each type of orbital( s, px and py and pz, dxy, dxz,dyz,dx2-y2 and dz2 ) has a unique shape.
1. Spherical shape for s.
2. Dumbbell shape for orbitals of p.
3. Four-lobed shape for orbitals of d.
4. Complex shape for all orbitals of higher sublevels.
Electron Structure or Configuration
In the orbital-sub level-orbit structure, energy level of orbit 1 is less than that of 2 and so on. In each orbit, the sublevel s if of lower energy than the p sublevel, and p is of lower energy than the d sublevel and so on. Orbital of a sublevel are all of equal energy.
Electrons occupy the lowest energy sublevels that are available. This is known as ‘aufbau’ order or principles. In the case of an atom having atomic number of 1, the lone electron occupied the s orbital of sublevel s of orbit 1(represented as 1s1). In case of an atom having atomic number 3 the electrons first occupy the sublevel of orbit 1(this can hold only two electrons) and then occupy p sublevels of orbit 2 (represented as 1s2,2s1).
Hund’s rule says that, for any set of orbitals of equal energy say p orbitals of orbit 2, there is one electron is each orbital before the second electron enters or occupies an orbital.
The energy level of some sublevels at higher orbits is less than the some sublevels at lower orbitals. This fact is to be kept in mind when electron configuration is determined for any atom. The increasing order of energy levels of sublevels is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f*, 5d, 6p, 7s, 5f*, 6d, 7p, 8s
* In this case one electron one electron goes into 5d and then 4f fills completely and then rest of 5d. Similar thing happens in 5f and 4d.