Atomic Structure

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Pauli Exclusion Principle -

No two electrons in an atom can be in exactly the same state

They cannot be in the same place at the same time

No two electrons in a given atom can have the same four quantum numbers

Energy levels, Sub levels and orbitals

First energy level - one sub energy level, one "s" orbital ( spherical shape-it can hold two electrons of opposite spin)

Second energy level - Two sub energy levels, The "s" sub level has one s orbital with spherical symmetry and the "p" sub level has three p orbitals with dumbbell shape

- px, py, pz ( they all have the same energy under normal conditions and differ in that one is oriented along x axis, one along y and the other along z - each orbital can hold 2 electrons making a total of 6 electrons in the p sub level)

Owing to increased electron-electron repulsion the p orbitals are at a slightly higher energy than the s orbitals

Third energy levels - three sub levels, s-2 electrons, p-6 electrons + 5 d orbitals - 10 electrons

+ the 3d orbitals are at a higher energy level than the 3p orbitals and slightly higher that 4s orbitals

Aufbau Principle -

the electrons in atoms always adopt the lowest energy configuration possible by filling one sub level completely before starting to fill the cub level of the next highest energy.

Hunds rule-

Sub levels are singly occupied as far as possible by electrons with the same spin

This configuration minimises mutual repulsion between them

Half filled 3d orbital

chromium is [Ar] 3d5 4s1 rather than [Ar]3d4 4s2

Fully filled 3d orbital

copper is [Ar]3d10 4s1 rather than [Ar] 3d9 4s2

This arrangement provides additional stability

First row d - block elements

With both d and s electrons in their valence shell

when elements between scandium and zinc form cations the first electrons that they lose are the 4s electrons even though this orbital was filled before the 3d orbital. This is a consequence of the change in relative stabilities of the 3d and 4s orbitals, which occurs as the 3d orbital stars to fill. Which means that the ion with the most d electrons is the most stable

The 3d and 4s sub levels are close in energy so that once the 4s2 electrons are lost the 3d electrons also behave as valence electrons

Schrodinger wave equation

This involves four constants - Quantum numbers

Principle quantum number - n - must be a positive integer - main or principle energy level

Subsidiary quantum number - l - can have integer values from o to (n-1) - sub level

Magnetic Quantum number - m - integer values from -/ to +/ including 0 - particular orbital within sub level

Spin quantum number - s - can be ±1/2 - differentiates between two electrons in that orbital

Ionisation Energies

The ionisation energy of an atom is the minimum amount of energy required to remove a mole of electrons from a mole of gaseous atoms to form a mole of gaseous ions

Second ionisation energy - the energy required to remove a second mole of electrons from the ion produced by the loss of one electron

ENDOTHERMIC CHANGE - work has to be done to remove a negatively charged electron from the attraction of a positively charged nucleus

The magnitude of ionisation energy depends in the charge on the nucleus .This will be counteracted by the repulsion or " shielding" of electrons in filled inner orbitals

( each electron in a filled inner shell will cancel out one unit of nuclear charge and after these have been subtracted the remaining nuclear charge is referred to as EFFECTIVE NUCLEAR CHARGE)

The third factor that affects the ionisation energy is the repulsion electron experiences from other electrons within the same shell

The more electrons that have been removed from an atom the greater the energy required to remove the next electron

Variation within period - ionisation energy increases across the period

Variation within the group - ionisation energy decreases