Ionisation energy is a fundamental periodic property that measures the energy required to remove electrons from gaseous atoms or ions. Its trends across the periodic table provide direct evidence for the existence of electron shells, subshells, and the effects of nuclear charge and shielding.
1. Definition & Core Concepts
Graph showing the general increase of ionisation energy across a period with characteristic dips at Group 3 and Group 6.
2. Underlying Principles: Factors Affecting IE
Nuclear Charge: As the number of protons in the nucleus increases, the positive charge increases, exerting a stronger electrostatic pull on the outer electrons. This generally increases the energy required to remove an electron.
Atomic Radius: The distance between the nucleus and the outer electrons significantly impacts attraction. Electrons further from the nucleus experience a weaker pull due to the inverse square relationship of electrostatic force.
Shielding: Inner shell electrons repel outer electrons, effectively 'shielding' them from the full positive charge of the nucleus. More inner shells result in lower ionisation energy.
Spin-Pair Repulsion: Within an orbital, two electrons with opposite spins repel each other. This repulsion makes it slightly easier to remove one of the electrons compared to an electron in a singly occupied orbital.
3. Periodic Trends
Across a Period
General Trend: Ionisation energy increases across a period because the nuclear charge increases while the shielding remains relatively constant. This results in a smaller atomic radius and a stronger pull on the valence electrons.
Group 2 to 3 Dip: There is a slight decrease because the outer electron in Group 3 is in a p-subshell, which is higher in energy and further from the nucleus than the s-subshell of Group 2.
Group 5 to 6 Dip: A decrease occurs because Group 6 elements have a paired electron in a p-orbital. The spin-pair repulsion between these two electrons makes the first electron easier to remove.
Down a Group
General Trend: Ionisation energy decreases down a group. Although nuclear charge increases, the effect is outweighed by the increase in atomic radius and the addition of more shielding shells.
4. Key Distinctions
5. Exam Strategy & Tips
First Ionisation Energy (IE1): This is defined as the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. It is measured in units of kJmol−1 and must always be represented with gaseous state symbols (g).
Standard Conditions: Ionisation energies are typically measured under standard conditions of 298K and 100kPa to ensure consistency across data sets.
Successive Ionisation Energies: These refer to the energy required to remove further electrons from increasingly positive ions (e.g., 2nd, 3rd, 4th IE). These values always increase because it is harder to remove a negative electron from a more positive nucleus.
Feature
Across a Period
Down a Group
Nuclear Charge
Increases
Increases
Shielding
Constant
Increases
Atomic Radius
Decreases
Increases
IE Trend
Generally Increases
Decreases
Big Jumps vs. Small Jumps: In successive ionisation energy data, a 'big jump' indicates a change in the principal quantum shell (moving closer to the nucleus). A 'small jump' or dip in the periodic trend indicates a change in subshell or the onset of electron pairing.
State Symbols: Always include (g) for both the atom and the ion in ionisation energy equations; omitting these is a frequent cause of lost marks.
Successive IE Identification: To determine the group of an unknown element, count the number of electrons removed before the first massive jump in energy. For example, if the jump is between the 3rd and 4th IE, the element is in Group 3.
Comparison Phrasing: When explaining trends, always mention three specific factors: Nuclear Charge, Shielding, and Distance/Radius. Use the phrase 'outweighs' to show which factor is dominant (e.g., 'increased shielding outweighs increased nuclear charge down a group').
Graph Reading: Ensure you read successive IE graphs from left to right (1st to last electron) and look for the 'steps' that represent different shells.
