Which of the following elements will have highest ionisation energy? (a) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{1}\) (b) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{3}\) (c) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{4}\) (d) \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{1}\)

Understanding electron configuration is essential for predicting how atoms will interact and how reactive they may be. At its core, electron configuration is the distribution of electrons in an atom's orbitals, which are layers of space that surround an atom's nucleus. Electrons fill these orbitals in a specific order, following the 'aufbau principle,' which tells us they occupy the lowest energy levels first. For instance, the first two electrons will fill the 1s orbital, designated as the 1s² configuration.

When all the orbitals in a shell (or a layer) are filled, such as the 2p⁶ configuration, it indicates a stable, less reactive atom, because it takes more energy to remove an electron from a completely filled shell. On the other hand, an element with one electron in a new shell, such as 3s¹, may lose that electron more easily, indicating it might have lower ionisation energy compared to more filled configurations.

The periodic table is not just a chart of elements; it's a map of periodic trends that dictate elemental properties, one of which is ionisation energy. Ionisation energy tends to increase across a period, or row, from left to right. This happens because as protons are added to the nucleus, the overall positive charge increases, pulling electrons closer.

Conversely, as you move down a group, or column, in the periodic table, ionisation energy decreases. This is because with each additional layer of electron orbitals, the outer electrons are further from the nucleus and feel a lesser attraction, thus requiring less energy to be removed. These trends help us predict which element in a given set might have the highest or lowest ionisation energy, greatly narrowing down the analysis needed when comparing similar elements.

The effective nuclear charge (Z_eff) is the net positive charge experienced by an electron in a multi-electron atom. The more closely an electron is held by the nucleus, the higher the effective nuclear charge. Z_eff increases across a period because each successive element has an additional proton in the nucleus, increasing the nuclear charge without a corresponding increase in electron shielding.

The concept of electron shielding is important here; inner electrons can shield outer electrons from the full charge of the nucleus. Thus, as the inner electrons 'shield' the outer electrons from the nucleus, the effective nuclear charge that these outer electrons experience is less than the actual charge of the nucleus. This concept is key in determining why some atoms have higher ionisation energies than others.

The periodic table is not merely a list of elements but a profound tool that encapsulates four key areas of chemical properties: atomic size, ionisation energy, electronegativity, and electron affinity. Relative positions of elements within this table give insight into how their electrons are configured and how they react with other substances.

For example, elements on the right side of the table generally have higher ionisation energies due to their near-complete electron shells, making them less eager to give an electron away. This is in stark contrast to elements on the left side, which tend to lose electrons more readily given their sparser electron shells. Understanding these nuances helps us to not only predict chemical reactivities but also to delve into the reasons behind these predictions, enabling better grasp of complex chemical phenomena.