Exceptions in Periodic Properties: Important Comparisons and Their Reasons

Periodic properties such as atomic radius, ionisation enthalpy and electron gain enthalpy generally follow predictable trends across periods and down groups. However, some elements do not follow the expected pattern.

These unusual observations are called exceptions in periodic properties. They often occur because of differences in orbital energy, electron-electron repulsion, shielding and electronic configuration.

For example, beryllium has a higher first ionisation enthalpy than boron, while chlorine has a more negative electron gain enthalpy than fluorine. Understanding the reasons behind these comparisons makes periodic trends easier to learn and apply in examinations.

Explore: Chemical Exceptions in Inorganic Chemistry

What Are Exceptions in Periodic Properties?

A periodic-property exception occurs when an element does not follow the general trend expected for a property in the periodic table.

For example, first ionisation enthalpy generally increases from left to right across a period because effective nuclear charge usually increases. However, the first ionisation enthalpy decreases from beryllium to boron.

This does not mean that periodic trends are incorrect. It means that the general trend is influenced by additional factors such as:

• The subshell from which an electron is removed.
• The relative energy of s and p orbitals.
• Electron-electron repulsion.
• The stability and energy of the resulting ion.
• Shielding and orbital penetration.

Students should understand both the general trend and the reason for the deviation.

1. Exceptions in First Ionisation Enthalpy

First ionisation enthalpy is the energy required to remove the most loosely held electron from an isolated gaseous atom in its ground state.

For example, when a gaseous sodium atom loses an electron, it forms a positively charged sodium ion and releases an electron.

First ionisation enthalpy generally increases across a period and decreases down a group. However, important exceptions occur when the electron being removed belongs to a different subshell or when electron pairing changes the energy of the atom.

Be vs B: Why does beryllium have higher first ionisation enthalpy than boron?

Beryllium and boron belong to the same period. The general expectation is that boron should have a higher first ionisation enthalpy because it has a greater nuclear charge.

However, beryllium has a higher first ionisation enthalpy than boron.

Their electronic configurations explain this difference.

In beryllium, the outermost electrons occupy the 2s subshell. In boron, the electron removed during first ionisation belongs to the 2p subshell.

The 2p electron is higher in energy and less penetrating than a 2s electron in the same principal shell. It is therefore easier to remove.

Conclusion: Boron has a lower first ionisation enthalpy than beryllium because its outermost 2p electron is less strongly held.

Mg vs Al: Why does magnesium have higher first ionisation enthalpy than aluminium?

Magnesium and aluminium show a similar exception in the third period.

The expected order based only on increasing nuclear charge would suggest that aluminium should have a higher first ionisation enthalpy. However, magnesium has a higher first ionisation enthalpy than aluminium.

Magnesium has a completely occupied 3s subshell. Aluminium has one electron in the higher-energy 3p subshell.

The 3p electron in aluminium is less penetrating and more easily removed than a 3s electron from magnesium.

Conclusion: Aluminium has a lower first ionisation enthalpy than magnesium because the electron removed from aluminium belongs to a higher-energy 3p orbital.

N vs O: Why does nitrogen have higher first ionisation enthalpy than oxygen?

Nitrogen and oxygen occur next to each other in the second period.

Although first ionisation enthalpy generally increases across a period, oxygen has a lower first ionisation enthalpy than nitrogen.

Nitrogen has a half-filled 2p subshell with one electron in each of the three p orbitals.

In oxygen, one 2p orbital contains a pair of electrons. Repulsion between the paired electrons makes it comparatively easier to remove one of them.

Conclusion: Oxygen has a lower first ionisation enthalpy than nitrogen because removing a paired electron reduces the electron-electron repulsion in the atom.

P vs S: Why does phosphorus have higher first ionisation enthalpy than sulphur?

The same type of exception occurs in the third period between phosphorus and sulphur.

Phosphorus has a half-filled 3p subshell. Sulphur has one paired set of electrons in its 3p orbitals.

The repulsion between paired electrons in sulphur makes the removal of one electron comparatively easier.

Conclusion: Sulphur has a lower first ionisation enthalpy than phosphorus because of additional repulsion between paired 3p electrons.

Quick comparison of first ionisation enthalpy exceptions

Important: Half-filled subshells are a useful part of the explanation for the nitrogen and phosphorus comparisons. However, the complete explanation also considers electron-electron repulsion and the energy changes involved in ionisation.

2. Exceptions in Electron Gain Enthalpy

Electron gain enthalpy is the enthalpy change when an isolated gaseous atom accepts an electron to form a gaseous anion.

For example, a gaseous chlorine atom can accept an electron to form a chloride ion.

When energy is released, electron gain enthalpy is negative. A more negative value indicates a more exothermic electron-addition process.

Across a period, electron gain enthalpy often becomes more negative, but this is not a universal rule. Atomic size, electronic configuration and electron-electron repulsion influence the actual value.

Cl vs F: Why does chlorine have more negative electron gain enthalpy than fluorine?

Fluorine is more electronegative than chlorine, but chlorine has a more negative electron gain enthalpy.

Fluorine has a very small atomic size. When an additional electron enters its compact 2p subshell, electron-electron repulsion is relatively high.

In chlorine, the incoming electron enters the larger 3p subshell. The additional repulsion is comparatively lower, so more energy is released.

Conclusion: Chlorine has a more negative electron gain enthalpy than fluorine because the compact 2p orbitals of fluorine experience greater electron-electron repulsion.

S vs O: Why does sulphur have more negative electron gain enthalpy than oxygen?

Sulphur also has a more negative electron gain enthalpy than oxygen.

Oxygen is smaller than sulphur. The incoming electron enters the compact 2p subshell of oxygen, where electron-electron repulsion is relatively high.

In sulphur, the incoming electron enters the larger 3p subshell. The comparatively lower repulsion makes electron addition more exothermic.

Conclusion: Sulphur has a more negative electron gain enthalpy than oxygen because its larger 3p orbitals have comparatively lower repulsion for the incoming electron.

Electron gain enthalpy exceptions: Important comparisons

Electron gain enthalpy vs electronegativity

These two concepts are related to an atom's tendency to attract electrons, but they are not the same.

Therefore, fluorine can be more electronegative than chlorine while chlorine has a more negative electron gain enthalpy.

3. Ga vs Al: An Important Atomic Radius Comparison

Atomic radius generally decreases across a period because effective nuclear charge tends to increase. Down a group, atomic radius generally increases because additional electron shells are occupied.

However, the comparison between gallium and aluminium is unusual.

Gallium is often reported as having an atomic radius comparable to, or slightly smaller than, aluminium, depending on the radius definition and the dataset being used.

Why does gallium show this unusual size?

Gallium has electrons occupying the 3d subshell before its outer 4p electron.

The 3d electrons do not shield the nuclear charge as effectively as might be expected. This imperfect shielding contributes to a greater effective nuclear attraction on the outer electrons.

As a result, gallium does not show the simple increase in size that students might expect when moving down from aluminium.

This behaviour is associated with the influence of the filled 3d subshell and is commonly discussed in connection with d-block contraction effects.

Important qualification: Atomic radius is not one single directly measured quantity. Covalent radius, metallic radius and van der Waals radius are defined differently. Comparisons must use the same radius definition and compatible data.

Do not memorise a numerical Ga–Al radius comparison without checking which radius type is being used.

4. Electron Gain Enthalpy of Noble Gases

Noble gases have completely filled valence shells in their ground-state configurations.

For example:

• Helium: 1s²
• Neon: [He] 2s² 2p⁶
• Argon: [Ne] 3s² 3p⁶

Their stable electronic configurations make the addition of another electron unfavourable under ordinary gaseous-atom conditions.

An incoming electron cannot simply enter the already filled valence shell. It must occupy a higher-energy orbital or shell.

Therefore, electron addition to noble-gas atoms is generally endothermic, meaning that energy must be supplied.

This is different from saying that noble gases can never form compounds. Their compound formation depends on the chemical species and conditions involved. The electron gain enthalpy statement refers specifically to adding an electron to an isolated gaseous atom.

5. Scientific Principles Behind Periodic Property Exceptions

The exceptions discussed above can be understood through a small group of connected atomic principles.

These principles should not be learned as independent memorisation points. A single comparison can involve more than one factor.

For example, the Be–B and Mg–Al comparisons involve both subshell energy and orbital penetration. The N–O and P–S comparisons involve electron pairing and repulsion.

6. Quick Revision Table: Periodic Property Exceptions

7. How to Prepare Periodic Property Exceptions for JEE and NEET

Use the following approach to revise periodic-property exceptions effectively.

1. Learn the general periodic trends first. Understand the normal changes in atomic radius, ionisation enthalpy and electron gain enthalpy.
2. Practise electronic configurations. Identify whether the outermost electron belongs to an s or p subshell.
3. Compare elements in pairs. Revise Be–B, Mg–Al, N–O and P–S together rather than as unrelated facts.
4. Understand electron pairing. Draw simple p-orbital box diagrams to understand the N–O and P–S comparisons.
5. Remember the difference between electron gain enthalpy and electronegativity. The Cl–F comparison is a useful test of this distinction.
6. Use consistent definitions for atomic radius. Do not compare values taken from different radius conventions without checking their meaning.
7. Practise explanation-based questions. Learn to explain why an observation occurs instead of only recalling which element has the higher value.
8. Revise with a comparison table. Keep the actual observation and the main reason together in your notes.

For JEE Main and NEET, focus on the standard comparisons and their conceptual explanations. For JEE Advanced, be prepared to apply orbital, shielding and electron-interaction concepts to unfamiliar comparisons.

Conclusion

Exceptions in periodic properties show that periodic trends are general patterns, not rigid rules. Subshell energy, orbital penetration, electron pairing, shielding and atomic size can change the expected behaviour of an element.

To understand these comparisons, learn the general trend first and then connect each exception with its scientific reason. Continue with the other guides in the Chemical Exceptions in Inorganic Chemistry series to study electronic configurations, chemical bonding and block-wise anomalies in greater detail.

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