Periodicity - Ionisation Energy
General Trend: The first ionisation energy generally increases as one moves from left to right across a period in the periodic table. This trend indicates that it becomes progressively harder to remove an electron from atoms as their atomic number increases within the same period.
Increasing Nuclear Charge: Across a period, each successive element has one more proton in its nucleus, leading to a greater positive nuclear charge. This increased charge exerts a stronger pull on all electrons, including the outermost ones.
Decreasing Atomic Radius: Despite the addition of more electrons, they are being added to the same principal energy shell. The increasing nuclear charge pulls the outer electron shell closer to the nucleus, resulting in a decrease in atomic radius across the period.
Constant Shielding: As electrons are added to the same principal energy shell, the number of inner-shell electrons, which are primarily responsible for shielding, remains relatively constant. Therefore, the shielding effect does not significantly counteract the increasing nuclear charge.
Overall Effect: The combined effect of a stronger nuclear charge and a smaller atomic radius, with relatively constant shielding, leads to a stronger effective nuclear charge experienced by the outermost electrons. This increased attraction makes them more difficult to remove, hence the general increase in ionisation energy.
Dip between Group 2 and Group 13 (e.g., Beryllium to Boron, Magnesium to Aluminium): A slight decrease in ionisation energy is observed when transitioning from a Group 2 element to a Group 13 element. This anomaly occurs because the electron being removed from the Group 13 element is from a p-subshell, whereas for the Group 2 element, it is from an s-subshell.
Subshell Energy Difference: Electrons in a p-subshell are, on average, slightly further from the nucleus and experience more effective shielding from the inner s-electrons compared to electrons in the s-subshell of the same principal energy level. This makes the p-electron easier to remove, despite the increased nuclear charge in the Group 13 element.
Dip between Group 15 and Group 16 (e.g., Nitrogen to Oxygen, Phosphorus to Sulfur): Another slight decrease in ionisation energy is observed when moving from a Group 15 element to a Group 16 element. This particular anomaly is attributed to the effect of spin-pair repulsion.
Spin-Pair Repulsion: Group 15 elements have half-filled p-subshells (one electron in each p-orbital), which is a relatively stable configuration. Group 16 elements, however, have one p-orbital containing two paired electrons. The mutual repulsion between these paired electrons makes it energetically more favorable, and thus easier, to remove one of them, leading to a lower ionisation energy than expected.
General Trend: The first ionisation energy generally decreases as one moves down a group in the periodic table. This indicates that it becomes progressively easier to remove an electron from atoms as their atomic number increases within the same group.
Increasing Nuclear Charge: As one moves down a group, the number of protons in the nucleus increases significantly. In isolation, this would suggest a stronger attraction for electrons.
Increasing Atomic Radius: However, down a group, electrons are added to new, higher principal energy shells. This results in a substantial increase in the atomic radius, meaning the outermost electrons are much further from the nucleus.
Increasing Shielding: The addition of more inner electron shells significantly increases the shielding effect. These additional inner electrons effectively block a greater portion of the nuclear charge from the outermost electrons.
Overall Effect: The combined effects of a substantially increased atomic radius and significantly increased shielding outweigh the increase in nuclear charge. The outermost electrons experience a much weaker effective nuclear charge, making them easier to remove, and thus the ionisation energy decreases down a group.
Master the Three Core Factors: When explaining ionisation energy trends, always analyze the interplay of nuclear charge, atomic radius, and shielding effect. Clearly articulate how each factor changes and its direct consequence on the electron's attraction to the nucleus.
Memorize Anomalies and Explanations: Be prepared to explain the specific dips in ionisation energy trends across a period, as these are common exam questions. These explanations require precise details involving subshell energy differences (s vs. p orbitals) and the concept of spin-pair repulsion.
Relate to Electron Configuration: A strong understanding of electron configurations is fundamental to explaining ionisation energy trends, particularly the anomalies. For example, knowing that boron's outermost electron is in a orbital and oxygen has paired electrons in a orbital is key to accurate explanations.
Practice Comparative Questions: Many exam questions involve comparing the ionisation energies of two or more elements. Systematically apply the factors (nuclear charge, shielding, distance) to justify your answer, ensuring you identify the dominant factor for each comparison.
Overemphasis on Nuclear Charge: A common mistake is to assume that increasing nuclear charge always leads to higher ionisation energy. While true across a period, down a group, the increased shielding and atomic radius become the dominant factors, leading to a decrease in ionisation energy.
Confusing Shielding and Nuclear Charge: Students sometimes mix up these two distinct concepts. Remember that nuclear charge is the total positive charge of the nucleus, while shielding is the reduction of this charge's effect on outer electrons due to inner electrons.
Ignoring Subshell Effects: Failing to recognize that electrons in different subshells (s, p, d, f) have different average energies and distances from the nucleus can lead to incorrect explanations for the dips in ionisation energy, such as the one between Group 2 and Group 13 elements.
Misunderstanding Spin-Pair Repulsion: Simply stating 'repulsion' is often insufficient for explaining the Group 15 to Group 16 dip. The explanation must specifically refer to 'spin-pair repulsion' occurring within a paired orbital, which makes that particular electron easier to remove.