Atomic Structure and Electronic Configuration

For convenience this is split into 5 sections:

Atomic Structure

Atoms are composed of 3 basic fundamental particles, electrons, protons and neutrons.
The protons and neutrons together form the central mass of the atom, called the nucleus. Orbiting electrons surround the nucleus.

These fundamental particles are often called sub atomic particles and have the following properties:

This shows that the proton and neutrons together form the mass of the atom (remember the mass of the electron is negligible).
All atoms are neutral therefore the number of protons and electrons must be equal in order for the charges to balance.

Representation of atoms

Mass Number		for example:
Atomic Number

A = mass no = number of protons + neutrons
Z = atomic no = number of protons

With the example of chlorine we can determine the make up of the atom.

Number of protons = Z = 17
Number of neutrons = A – Z = 18
As the atom is neutral, the number of electrons = no of protons = Z = 17

The atomic no Z, can be determined from the position of the element on the periodic table, starting at H = 1, He = 2, Li = 3 and so on across the table.

Knowing the number of electrons present is very useful in understanding the structure and bonding, as you will see.

Isotopes

The atomic no Z, defines the element, for example, only chlorine has the atomic no 17. However, individual elements may have different mass numbers; that is they have different numbers of neutrons.

These are called isotopes.

Eg
				Z = 17 and A = 35 and 37
or
				Z = 6 and A = 12, 13 and 14

Relative Atomic Mass

This is the weighted average of the mass numbers of all the isotopes present. Taking chlorine as an example, the isotopic composition is:

75% ³⁵Cl and
25% ³⁷Cl.

So its relative atomic mass is given by:

Ar = ( 75% x 35 ) + ( 25% x 37 ) = 35.5

Atomic masses are measured on a scale where one atom of ¹²C has a mass of exactly 12.00 amu.

Relative Molecular Mass

This again is measured on the same scale relative to ¹²C and amounts to the sum of the individual atomic masses making up the molecule.

Orbitals

Electrons do not simply orbit the nucleus of an atom at a set distance as initially thought.

Advances in quantum theory now tell us that electrons are located around definite energy levels. Within these energy levels there are areas where the probability of finding an electron is maximised and these are called orbitals.

There are a number of different orbitals, each having a different shape and an odd number of sub-orbitals.

Their properties are as follows:

The s – orbital is spherical about the nucleus.
The three p – orbitals are dumb-bell shaped along the 3 coordinate axes.

2px	2py	2pz

Each energy level is defined by a principal quantum no. with the level closest to the nucleus having principal quantum no, n = 1

The number of orbitals in each level = n²

So in the first level:

n = 1, there are 1² orbitals = 1 x s
n = 2, there are 2² orbitals = 1 x s + 3 x p
n = 3, there are 3² orbitals = 1 x s + 3 x p + 5 x d

The following diagram illustrates this and also demonstrates the order in which the orbitals are filled with electrons.

Orbital Filling Diagram

The order is determined by moving as far as possible down the diagonal line and then starting at the top of the next one.

For example, from 1s to 2s to 2p to 3s to 3p to 4s to 3d to 4p to 5s to 4d.

Although an electron will occupy the 4d orbital after the 5s orbital, the first electrons to be removed ionised will come from the orbital with the highest quantum no, in the above example 5s.

As with all quantum theory, there are a lot of rules to obey but learning them helps you determine the correct electronic arrangement of any atom.

RULES

1. Aufbau principle: electrons enter the lowest level orbitals first – see the above diagram for the energy level order.
2. Pauli exclusion principle: each sub orbital can hold a max of 2 electrons.
3. hund's rule: electrons are antisocial and so will occupy all sub orbitals of the same energy singly before pairing up.
   

This is because electrons repel each other and the total energy is reduced by seperating themselves.

Note also that electrons have a property called spin and only opposing spins pair up.

When talking about electronic configurations you can display them either by s, p, d rotation or by the box display.

Example: Titanium has atomic no 22 and so has 22 electrons

Box Level Diagram for Titanium

s, p, d notation: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d² 4s²

Experimental Support for Electronic Arrangement

Ionisation Energies

Atoms can either lose or gain electrons to become ions, which simply means charged atoms.

Gaining electrons gives negtive ions
Losing electrons gives positive ions

The ionisation energy is a measure of the energy required to remove an electron and therefore form a positive ion.

The more strongly an electron is attracted by the positive nucleus the higher the ionisation energy will be.

Factors affecting the size of this attraction and therefore ionisation energy are:

We will reinforce these ideas when we look at periodic trends - they are very important.

Ionistaion Trends

Successive ionisation energies provide evidence for the existence of electronic shells.

Taking sodium as an example, the first 3 successive ionisation energies are described as follows:

1st
2nd
3rd

1st Ionisation energy across period provides evidence for the existence of sub-shells.

The overall trend is for the ionisation energy to increase across the period as the increased nuclear charge is the most important factor at play.

Electronic Configuration and the Periodic Table

As the electronic configuration defines much of the reactivity characteristics of the elements, it is not surprising that the periodic table is closely related.