Electron Configuration and Chemical Periodicity PDF

Preview

Summary

Electron Configuration and Chemical Periodicity- Electron-spin Quantum number, Pauli's Exclusion principle, Factors affecting orbital energy, Hund's Rule, Trends in atomic properties

Notes made by Jamie-Lee McDonic based on lecture notes received from University of Cape Town Chemistry D...

Description

Chapter 8: Electron Configuration and Chemical Periodicity The Electron-Spin Quantum Number  

Electron-spin quantum number: ms Indicates the direction of the electron spin about its own axis



Can only be one of 2 numbers: +



Each electron in an atom is described completely by a set of four quantum numbers, describing its size (n), shape (l) and orientation (ml)

1 2

OR -

1 2

Pauli Exclusion Principle   

No two electrons in the same atom can have the same four quantum numbers. Each electron must have a unique identity as expressed by its quantum numbers. The major consequence of the exclusion principle is that an atomic orbital can hold a maximum of two electrons and they must have opposing spins.

Factors affecting atomic orbital energies 



The splitting of energy levels into sub-levels of differing energies : the energy of an orbital in a many-electron atom depends mostly on its n value (size) and to a lesser extent on its l value (shape). The electrons of an atom in its ground state occupy the orbitals of lowest energy .

1. The Effect of Nuclear Charge (Z) on an Orbital Energy o Higher nuclear charge lowers orbital energy (stabilizes the system) by increasing nucleus-electron attractions. o One electron in an orbital makes another electron in the orbital easier to remove. Therefore the orbital is less stable (eg: Li) o A single electron in an orbital is harder to remove as there is a stronger electronnucleus attraction. Therefore this orbital is more stable (eg:

2+¿ ) Li¿

2. The Effect of Electron Repulsions (shielding) o Additional electrons in the same orbital: An additional electron raises the orbital energy through electron-electron repulsions. o This repulsion counteracts the nuclear attraction, making each electron easier to remove by helping to ‘push’ it away. o Additional electrons in inner orbitals: Inner electrons shield outer electrons more effectively than electrons in the same sublevel do. o Shielding reduces the full nuclear charge by to an effective nuclear charge (Zeff) which is the nuclear charge the electron actually experiences. This lower nuclear charge makes the electron easier to remove. o Shielding by inner electrons greatly lowers the Zeff felt by outer electrons.

o

For example: It takes 520kJ/mol to remove an electron from the Li atom and 2954kJ/mol to remove an electron from

2+¿ ion because the inner electrons in Li¿

Li shield very effectively. 3. Penetration: The Effect of Orbital Shape on Orbital Energy o At first, we might expect that an electron would enter the 2p orbital because it is slightly closer to the nucleus than the major portion of the 2s orbital. But a minor portion of the 2s radial probability distribution appears within the 1s region. o As a result, an electron in the 2s orbital spends part of its time ‘penetrating ‘ very close to the nucleus. o Charges attract more strongly if they are near each other than far apart. o Therefore penetration by the 2s electron increases its overall attraction to the nucleus relative to that of a 2p electron. o At the same time, penetration into the 1s region decreases the shielding of the 2s electron by the 1s electrons. o The 2s orbital is therefore of lower energy than the 2p orbital because it takes more energy to remove a 2s electron than a 2p.

Order of sublevel energies:

s...