**ELECTRONIC STRUCTURE**

__i) Energy levels__

Electrons do not orbit the nucleus randomly; they occupy certain fixed energy levels. Each atom has its own unique set of energy levels, which are difficult to calculate but which depend on the number of protons and electrons in the atom.

Energy levels in an atom can be numbered 1,2,3,…. To infinity. 1 is the lowest energy level (closest to the nucleus) and energy level infinity corresponds to the energy of an electron which is not attracted to the nucleus at all. The energy levels thus converge as they approach infinity:

## ii) __Orbitals and sub-levels__

Electrons do not in fact orbit the nucleus in an orderly way. In fact they occupy areas of space known as **orbitals**. The exact position of an electron within an orbital is impossible to imagine; an orbital is simply an area of space in which there is a high probability of finding an electron.

Orbitals can have a number of different shapes, the most common of which are as follows:

s-**orbital**

s-orbitals: these are spherical.

Every energy level contains one s-orbital.

An s-orbital in the first energy level is a 1s orbital.

An s-orbital in the second energy level is a 2s orbital, etc

## p-orbital

p-orbitals: these are shaped like a 3D figure of eight. They exist in groups of three:

Every energy level except the first level contains three p-orbitals. Each p-orbital in the same energy level has the same energy but different orientations: x, y and z.

A p-orbital in the second energy level is a 2p orbital (2p_{x}, 2p_{y}, 2p_{z})

A p-orbital in the third energy level is a 3p orbital (3p_{x}, 3p_{y}, 3p_{z}), etc

In addition, the third and subsequent energy levels each contain five d-orbitals, the fourth and subsequent energy levels contain seven f-orbitals and so on. Each type of orbital has its own characteristic shape.

S, p and d orbitals do not all have the same energy. In any given energy level, s-orbitals have the lowest energy and the energy of the other orbitals increases in the order p < d < f etc. Thus each energy level must be divided into a number of different sub-levels, each of which has a slightly different energy.

The number and type of orbitals in each energy level can thus be summarised as follows:

Energy level
| Number and type of orbital | ||||

1^{st} sub-level | 2^{nd} sub-level | 3^{rd} sub-level | 4^{th} sub-level | 5^{th} sub-level | |

1 | 1 x 1s | ||||

2 | 1 x 2s | 3 x 2p | |||

3 | 1 x 3s | 3 x 3p | 5 x 3d | ||

4 | 1 x 4s | 3 x 4p | 5 x 4d | 7 x 4f | |

5 | 1 x 5s | 3 x 5p | 5 x 5d | 7 x 5f | 9 x 5g |

## iii) __Shells__

Since the different sub-levels have different energies, and the energies of the different levels get closer together with increasing energy level number, the high energy sub-levels of some energy levels soon overlap with the low energy sub-levels of higher energy levels, resulting in a more complex energy level diagram:

Starting with the lowest energy, the orbitals can thus be arranged as follows:

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

Many of these sub-levels have similar energy, and can be grouped together.

A collection of sub-levels of similar energy is called a **shell**.

1s│2s 2p│3s 3p│4s 3d 4p│5s 4d 5p│6s 4f 5d 6p

The arrangement of shells and the maximum number of electrons in each can be summarised as follows:

Shell number
| Orbitals in shell |

1 | 1 x1s |

2 | 1 x 2s, 3 x 2p |

3 | 1 x 3s, 3 x 3p |

4 | 1 x 4s, 5 x 3d, 3 x 4p |

5 | 1 x 5s, 5 x 4d, 3 x 5p |

6 | 1 x 6s, 7 x 4f, 5 x 5d, 3 x 6p |

## iv) __Electrons__

Electrons repel each other. In a small space such as an orbital, it is impossible to put more than two electrons.

Since electrons are charged particles, and moving charges create a magnetic field, it is possible to create a small magnetic attraction between two electrons if they are spinning in opposite directions in the same orbital. This is the reason two electrons, and not one, are permitted in the same orbital.

It is thus possible to calculate the maximum possible number of electrons in each sub-level, and thus in each energy level:

Shell | Number of electrons in each sub-level | Max. no of electrons |

1 | 2 x 1s | 2 |

2 | 2 x 2s, 6 x 2p | 8 |

3 | 2 x 3s, 6 x 3p | 8 |

4 | 2 x 4s, 10 x 3d, 6 x 4p | 18 |

5 | 2 x 5s, 10 x 4d, 6 x 5p | 18 |

6 | 2 x 6s, 14 x 4f, 10 x 5d, 6 x 6p | 32 |

__v) Electron arrangement in orbitals__

There are three rules which determine the way in which electrons fill the orbitals

- Aufbau/building principle: electrons always fill the lowest energy orbitals first.

- Hund’s rule: electrons never pair up in the same orbital until all orbitals of the same energy are singly occupied, and all unpaired electrons have parallel spin.

- Pauli exclusion principle: only two electrons may occupy the same orbital, and they must do so with opposite spin.

The arrangement of electrons in an atom is known as its **electronic configuration**. It can be represented in two ways:

The **arrow and box method** represents each orbital as a box and each electron as an arrow. The direction of spin is shown by the orientation of the arrow.

The electronic configuration of the first 18 elements using the arrow in box method is as follows:

1s 2s 2p 3s 3p

H | ↑ |

He | ↑↓ |

Li | ↑↓ | ↑ |

Be | ↑↓ | ↑↓ |

B | ↑↓ | ↑↓ | ↑ |

C | ↑↓ | ↑↓ | ↑ | ↑ |

N | ↑↓ | ↑↓ | ↑ | ↑ | ↑ |

O | ↑↓ | ↑↓ | ↑↓ | ↑ | ↑ |

F | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑ |

Ne | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ |

Na | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑ |

Mg | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ |

Al | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑ |

Si | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑ | ↑ |

P | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑ | ↑ | ↑ |

S | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑ | ↑ |

Cl | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑ |

Ar | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ | ↑↓ |

The **orbital method** indicates the number of electrons in each orbital with a superscript written immediately after the orbital.

The electronic configurations of the first eighteen elements can be shown with the orbital method as follows:

H: 1s^{1}

He: 1s^{2}

Li: 1s^{2}2s^{1}

Be: 1s^{2}2s^{2}

B: 1s^{2}2s^{2}2p^{1}

C: 1s^{2}2s^{2}2p^{2} or 1s^{2}2s^{2}2p^{6}3s^{2}3p_{x}^{1}3p_{y}^{1}

N: 1s^{2}2s^{2}2p^{3} or 1s^{2}2s^{2}2p^{6}3s^{2}3p_{x}^{1}3p_{y}^{1}3p_{z}^{1}

O: 1s^{2}2s^{2}2p^{4} or 1s^{2}2s^{2}2p^{6}3s^{2}3p^{2}3p_{x}^{2}3p_{y}^{1}3p_{z}^{1}

F: 1s^{2}2s^{2}2p^{5}

Ne: 1s^{2}2s^{2}2p^{6}

Na: 1s^{2}2s^{2}2p^{6}3s^{1}

Mg: 1s^{2}2s^{2}2p^{6}3s^{2}

Al: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{1}

Si: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{2} or 1s^{2}2s^{2}2p^{6}3s^{2}3p_{x}^{1}3p_{y}^{1}

P: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{3} or 1s^{2}2s^{2}2p^{6}3s^{2}3p_{x}^{1}3p_{y}^{1}3p_{z}^{1}

S: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{4} or 1s^{2}2s^{2}2p^{6}3s^{2}3p_{x}^{2}3p_{y}^{1}3p_{z}^{1}

Cl: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}

Ar: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}

A shorthand form is often used for both the above methods. Full shells are not written in full but represented by the symbol of the element to which they correspond, written in square brackets.

Eg. 1s^{2}2s^{2}2p^{6} is represented as [Ne] and 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6} is represented as [Ar].

The shorthand electronic configuration of the elements with atomic numbers 18 to 36 can be written as follows:

Note the unusual structures of chromium and copper.

The difference in energy between the 3d and 4s electrons is very small, and in chromium the energy required to promote and electron from 4s to 3d is recovered in the reduced repulsion which results from the fact that they are no longer paired. Thus the 4s^{1}3d^{5} structure in Cr is preferred.

In copper, the 3d orbitals are actually lower in energy than the 4s orbital, so the 4s^{1}3d^{10} structure in Cu is preferred.

## v) __Electron arrangement in ions__

The electronic configuration of ions can be deduced by simply adding or removing the appropriate number of electrons. The order in which electrons are to be removed can be deduced from the following rules:

- remove outer shell electrons first
- remove p-electrons first, then s-electrons and then d-electrons
- remove paired electrons before unpaired electrons in the same sub-level

__ __vi) __Effect of electronic configuration on chemical properties__

The chemical properties of an atom depend on the strength of the attraction between the outer electrons and the nucleus. These in turn depend on the number of protons and on the electronic configuration, and so it follows that these two factors are instrumental in determining the chemical properties of an atom.

This is in contrast with the neutron number however, which has no effect on the chemical properties of an atom. Neutrons have no charge and hence exert no attractive force on the nucleus.

**Isotopes, therefore, tend to have very similar chemical properties since they have the same atomic number and the same electronic configuration. They differ only in number of neutrons, which do not directly influence the chemical properties of an element.**