Question

The first ionisation enthalpy of the elements are in the order of (a) \(C

Short answer

The correct order of first ionisation enthalpy is Si < P < C < N.

Step-by-step solution

Understanding Ionisation Enthalpy

Ionisation enthalpy is the energy required to remove the most loosely bound electron from a gaseous atom in its ground state. The ionisation enthalpies generally increase across a period in the periodic table from left to right due to increasing nuclear charge, which attracts the electrons more strongly and decreases down a group due to increasing atomic radius which makes it easier to remove an electron.

Considering Periodic Trends

In the second period of the periodic table, the elements C, N, and Si, P follow the trend of increasing ionisation enthalpy from left to right with some exceptions due to electron repulsions in p-orbitals.

Analyzing Given Options

We know that in the second period, the trend is B < C < N < O < F. Barring the small disruption between N and O due to half-filled stability in N, the trend is increasing. Silicon and Phosphorus are from the third period and their ionisation energies are generally less than the elements directly above them in the second period.

Identifying the Correct Order

Given the trends, C should have a higher ionisation enthalpy than Si and P, but less than N. Therefore the correct order is Si < P < C < N, which matches option (c).

Key concepts

Understanding Periodic Trends

Periodic trends refer to patterns seen within the periodic table that provide insights into the properties of elements, including their ionisation enthalpies. As you move from left to right across a period (a row in the periodic table), the ionisation enthalpy generally increases. This is because atoms have a greater nuclear charge and a stronger attraction for their electrons, making it harder to remove an electron.

However, there are exceptions known as 'anomalies', mostly due to electron configuration and subshell stability. For instance, the halving or complete filling of s and p orbitals can cause elements like nitrogen or oxygen to have higher ionisation enthalpies than expected. These anomalies arise because electrons in these configurations are more stable and thus require more energy to be removed. Understanding these exceptions is crucial when analyzing patterns in ionisation enthalpy.

Deciphering Electron Orbitals

Electron orbitals are regions within an atom where electrons are most likely to be found. The configuration of these orbitals significantly influences the ionisation enthalpy. Electrons fill orbital spaces based on the principles of lowest energy and the Pauli exclusion, which states that no two electrons can have the same set of quantum numbers.

For example, the p-orbitals can hold six electrons but when they are exactly half-filled or completely filled (three or six electrons, respectively), they're extra stable. Hence, removing an electron from such a stable configuration requires more energy. This stability impacts the ionization enthalpy, as seen with nitrogen in the aforementioned exercise. Recognizing which electron configurations lead to stability can assist students in predicting periodic trends in ionisation energies.

Atomic Radius Influence

The atomic radius is the distance from the nucleus to the outermost stable electron orbital of an atom at equilibrium. There is an inverse relationship between the atomic radius and ionisation enthalpy; as the atomic radius increases, the ionisation enthalpy decreases.

This happens because electrons farther from the nucleus feel less nuclear attraction; hence it's easier for them to be removed, which equates to lower ionisation enthalpy. Down the groups of the periodic table, as the number of electron shells increases, the atomic radius increases, leading to lower ionisation enthalpies. On the other hand, across a period, the atomic radius decreases due to increasing nuclear charge pulling electrons closer, which increases ionisation enthalpy.