12.1 - Electron Configuration
12.1.1 - Explain how evidence form first ionisation energies across periods accounts for the existence of main energy levels and sub-levels in atoms
First Ionisation Energy - The amount of energy required to remove one mole of electrons from one mole of an element in the gaseous state.
The graph shows peaks in the ionisation energy at each of the noble gases. These are the most stable elements because of their full electron shells, meaning that it takes a large amount of energy to remove an electron. These peaks give evidence of the main energy levels.
The troughs occur at the group 1 elements since they only have one outer shell electron, and it is easily removed. There is gradual increase along the graph, but there are still dips in between. These act as evidence for the presence of sub-levels. When all the subshells are half-full, they become more stable.
There are three factors affecting the force of attraction between the nucleus and electrons: 1. Size of the nucleus
2. Repulsion between electrons
12.1.2 - Explain how successive ionisation energy data is related to the electron configuration of an atom
The first electron removed during ionisation is the one at the highest energy level. This is because it has the least attraction to the nucleus and is in the valence shell. When there are multiple electrons in the shell, the ionisation energy will increase as they are subsequently removed.
Once the outer shell is empty, the next shell will be emptied. This will require more energy because it is still full and stable. This provides evidence for the existence of energy levels, allowing for the electron configuration to be determined.
Ionisation Ionisation energy (kJ mol-1)
1400 2856 4578 7475 9440 53266 64358
The large increase between the 5th and 6th ionisations indicates that we have moved to the next energy level. This is full, and closer to the nucleus. The electrons experience greater attraction to the nucleus
There is a clear pattern in the successive
ionisation energies and the number of electrons removed. Low ionisation energies occur when
12.1.3 - State the relative energies of s, p, d and f orbitals in a single energy level
A subshell is a group of orbitals with the same energy. The number of orbitals that make up a subshell increases as the energy of the orbitals increases.
The lowest energy electrons are found in an s
orbital. The next orbitals are the p orbitals, then
the d orbitals and finally the f orbitals.
12.1.4 - State the maximum number of orbitals in a given energy level
An energy level in an atom can have up to n2 orbitals, n = energy level number. Since each orbital can hold up to 2 electrons, this means that each energy level can contain up to 2n2 electrons.
The s orbitals on a single energy level can hold a total of 2 electrons. The p orbitals can hold up to a total of 6 electrons. The d orbitals can hold 10 electrons and the f orbitals can hold 14 electrons.
12.1.5 - Draw the shape of an s orbital and the shapes of the px, py and pz orbitals
Electrons do not have orbits due to the uncertainty principle. Instead, they occupy an orbital, which is a region of space in which an electron or electrons may be found
Since electrons occupy space in three dimensions, orbitals have three dimensional shapes. The uncertainty of their positions leads us to describe them as an electron cloud, an area where electrons are likely to be found.
S Orbitals
These have a spherical shape. There is one s orbital per energy level, and can hold up to two electrons. They are the lowest energy orbitals in a subshell.
P Orbitals
These are dumbbell or figure-8 shape. Each orbital is made of two lobes, and between the lobes there is a node where the probability of finding an electron is 0.
The orbitals are at right angles to each other
px orbital on the x-axis
py orbital on the y-axis
pz orbital on the z-axis
12.1.6 - Apply the Aufbau principle, Hund's rule and the Pauli exclusion principle to write electron configuration for atoms and ions up to Z = 54
Aufbau Principle
This is also called the building up principle. It states that electrons fill the lowest energy level first, and that the subshell with the lowest energy is filled before the next one.
Pauli Exclusion Principle
Orbitals can hold 0, 1 or 2 electrons, and no more.
Hund's Rule
For orbitals of the same subshell, lowest energy is achieved when there are an equal number of
electrons with the same spin for each orbital.
They will be filled with one electron per orbital first, and then the second electrons are added. A half-filled subshell is fairly stable.
An orbital containing electrons is referred to as being occupied. Orbitals still exist when they are not occupied, but are not included in the electron configuration. Electron configurations can be shown two different ways:
Orbital diagrams
These show electrons in boxes and show the spin with up or down arrows
Superscripts
Occupied orbitals are written in order of increasing energy with the number of electrons in each orbital is written as a superscript.
Example Configurations:
Fluorine
Calcium
Chromium
This arrangement occurs because the 4s and 3d orbitals are very close in energy. So, the configuration is the most stable with one electron in each orbital.
NOTE
Orbitals are filled according to the following order:
