What is the trend of Ionic Radii down a group?

📚 Chapter: Classification of Elements and Periodicity in Properties

📘 Detailed Explanation

🔵 Trend Down a Group

The ionic radius generally increases down a group in the periodic table.

⬇️ Down the Group → Ionic Radius Increases ⬆️

This happens because a new electron shell is added as we move from one period to the next. The outermost electrons are therefore located farther from the nucleus.

🟣 Main Reasons

1. Increase in number of shells: Each successive element down a group has an additional principal energy level.
2. Increase in shielding effect: Inner electrons shield the valence electrons from the attraction of the nucleus.
3. Greater distance from nucleus: The outermost electron shell lies farther away from the nucleus.

More shells + More shielding → Larger ionic radius

🟢 Example: Group 1 Cations

Consider the alkali metal ions:

Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺

Ionic radius increases from Li⁺ → Cs⁺.

⚡ Important Point: Cations and Anions

Type	Effect
Cation (+)	Smaller than its parent atom
Anion (−)	Larger than its parent atom

However, within a particular group, both cationic and anionic radii generally increase as we move downward because additional electron shells are added.

✅ Final Answer

The ionic radius increases down a group because the number of electron shells increases, resulting in greater shielding and a greater distance between the nucleus and the outermost electron shell.

Down the group ↓ → Ionic radius ↑

📝 20 MCQs with Answers

1. What happens to ionic radius down a group?

A) Decreases

B) Increases

C) Remains constant

D) Becomes zero

✅ Answer

B) Increases

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2. The main reason for increasing ionic radius down a group is:

A) Decrease in shells

B) Addition of electron shells

C) Decrease in nuclear charge

D) Loss of protons

✅ Answer

B) Addition of electron shells

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3. Shielding effect down a group generally:

A) Decreases

B) Increases

C) Becomes zero

D) Remains exactly constant

✅ Answer

B) Increases

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4. Which has the larger ionic radius?

A) Li⁺

B) Na⁺

C) Both are equal

D) Cannot be compared

✅ Answer

B) Na⁺

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5. Correct order of Group 1 ionic radii is:

A) Li⁺ > Na⁺ > K⁺

B) Li⁺ < Na⁺ < K⁺

C) K⁺ < Na⁺ < Li⁺

D) Na⁺ < Li⁺ < K⁺

✅ Answer

B) Li⁺ < Na⁺ < K⁺

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6. Down a group, the number of occupied shells:

A) Decreases

B) Increases

C) Remains constant

D) Becomes zero

✅ Answer

B) Increases

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7. Which factor increases the distance of the outermost electrons from the nucleus?

A) Addition of electron shells

B) Removal of shells

C) Decrease in atomic number

D) Nuclear contraction only

✅ Answer

A) Addition of electron shells

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8. Which statement is correct?

A) Ionic radius always decreases down every group

B) Ionic radius generally increases down a group

C) Ionic radius is independent of shells

D) Ionic radius is always constant

✅ Answer

B) Ionic radius generally increases down a group

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9. A cation is generally ______ than its parent atom.

A) Larger

B) Smaller

C) Equal

D) Heavier only

✅ Answer

B) Smaller

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10. An anion is generally ______ than its parent atom.

A) Smaller

B) Larger

C) Equal

D) Without electrons

✅ Answer

B) Larger

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11. Which ion is expected to have the largest radius?

A) Li⁺

B) Na⁺

C) K⁺

D) Rb⁺

✅ Answer

D) Rb⁺

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12. What is the effect of inner electrons on outer electrons?

A) Shielding

B) Bonding

C) Ionisation

D) Fusion

✅ Answer

A) Shielding

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13. Which quantity generally increases down a group?

A) Ionic radius

B) Nuclear attraction on outermost shell

C) Electronegativity

D) Ionisation energy

✅ Answer

A) Ionic radius

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14. The outermost electron shell becomes:

A) Closer to nucleus down a group

B) Farther from nucleus down a group

C) Absent

D) Identical for all elements

✅ Answer

B) Farther from nucleus down a group

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15. Which pair represents a correct trend?

A) Down group → radius decreases

B) Down group → radius increases

C) Down group → shells decrease

D) Down group → shielding decreases

✅ Answer

B) Down group → radius increases

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16. Ionic radius is a measure of:

A) Size of an ion

B) Mass of an ion

C) Number of neutrons only

D) Melting point

✅ Answer

A) Size of an ion

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17. Which has a greater ionic radius?

A) K⁺

B) Li⁺

C) Both equal

D) H⁺

✅ Answer

A) K⁺

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18. The shielding effect is caused mainly by:

A) Inner-shell electrons

B) Protons only

C) Neutrons only

D) Nucleus only

✅ Answer

A) Inner-shell electrons

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19. Which sequence shows increasing ionic radius?

A) Li⁺ < Na⁺ < K⁺

B) K⁺ < Na⁺ < Li⁺

C) Na⁺ < Li⁺ < K⁺

D) Li⁺ > K⁺ > Na⁺

✅ Answer

A) Li⁺ < Na⁺ < K⁺

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20. The best explanation for increasing ionic radius down a group is:

A) More electron shells and greater shielding

B) Fewer electron shells

C) Fewer protons

D) Complete removal of electrons

✅ Answer

A) More electron shells and greater shielding

Q. State the trend of ionic radii down a group.

Solution: Ionic radius generally increases down a group because a new electron shell is added at each successive period. The increased shielding effect also places the outermost electrons farther from the nucleus.

Down the group ↓ → Ionic radius ↑

Q. Why does ionic radius increase down a group?

Solution:

1. The number of occupied electron shells increases down the group.
2. The shielding effect of inner electrons increases.
3. The outermost electrons become farther from the nucleus, so the ionic radius increases.

Therefore: Ionic radius increases down a group.

Q. Explain the variation of ionic radius down a group with an example.

Solution:

Ion	Relative Size	Reason
Li⁺	Smallest	Fewer electron shells
Na⁺	↑	One additional shell
K⁺	↑↑	More shells and shielding
Rb⁺	↑↑↑	Still more shells

Li⁺ < Na⁺ < K⁺ < Rb⁺

Q. Discuss the trend of ionic radius down a group and explain the factors responsible for it.

Solution:

Ionic radius generally increases as we move down a group in the periodic table. The following factors are responsible:

• Increase in principal quantum number: Each step down adds a new electron shell.
• Increase in shielding effect: Inner electrons increasingly shield the outer electrons from the nucleus.
• Increase in distance: The outermost electron shell lies farther from the nucleus.
• Reduced effective attraction: Although nuclear charge increases, the additional shells and shielding cause the outer electrons to experience weaker effective attraction.
• Hence, the overall ionic size increases down the group.

General Trend:

Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺

Q. Explain in detail the trend of ionic radii down a group. Also explain the role of electron shells and shielding effect with a suitable example.

Detailed Solution:

The ionic radius is the effective size of an ion. It depends mainly on the number of occupied electron shells and the attraction between the nucleus and the electrons.

When we move down a group, the atomic number increases and a new principal energy level is added at each step. Therefore, the outermost electrons are located farther away from the nucleus.

Factor	Effect Down the Group
Number of shells	Increases
Shielding effect	Increases
Distance of outer shell from nucleus	Increases
Ionic radius	Increases

Example: Group 1 Cations

Li⁺ → Na⁺ → K⁺ → Rb⁺ → Cs⁺

Ionic radius increases from Li⁺ to Cs⁺.

For example, Li⁺ has electrons only in the first shell, while Na⁺ has electrons up to the second shell and K⁺ has electrons up to the third shell. Thus, the outer electron shell is progressively farther from the nucleus.

The inner-shell electrons also shield the outer electrons from the full positive charge of the nucleus. This shielding effect increases down the group and contributes to the increase in ionic size.

🎯 Final Answer:

The ionic radius generally increases down a group because successive elements have additional electron shells and greater shielding effect. These factors increase the distance of the outermost electrons from the nucleus, resulting in a larger ionic radius.

Down the group ↓ → Number of shells ↑ → Shielding ↑ → Ionic radius ↑

🎯 Quick Revision

Down a Group ↓
Electron shells ↑
Shielding effect ↑
Distance from nucleus ↑
Ionic Radius ↑

Example:
Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺

Arrange the following in increasing order of size: O 2− ,F − ,Na + ,Mg 2+ .

📚 Chapter: Classification of Elements and Periodicity in Properties

📘 Detailed Solution

🔵 Step 1: Count the electrons

First, determine the number of electrons in each ion:

Ion	Atomic Number (Z)	Electrons
O2−	8	8 + 2 = 10
F−	9	9 + 1 = 10
Na+	11	11 − 1 = 10
Mg2+	12	12 − 2 = 10

All four ions have 10 electrons.

🟣 Step 2: Identify the isoelectronic species

Since all four ions contain 10 electrons, they are called isoelectronic species.

O²⁻ F⁻ Na⁺ Mg²⁺

For isoelectronic species, the number of electrons is the same. Therefore, the size depends mainly on the nuclear charge.

🟢 Step 3: Compare nuclear charge

Ion	Number of Protons	Relative Nuclear Attraction
O2−	8	Lowest
F−	9	Higher
Na+	11	Still higher
Mg2+	12	Highest

As the number of protons increases, the same 10 electrons are attracted more strongly towards the nucleus. Therefore, the ionic radius decreases.

Higher nuclear charge → Stronger attraction → Smaller size

🎬 Size Trend

Increasing Nuclear Charge

O²⁻ → F⁻ → Na⁺ → Mg²⁺

Size decreases →

✅ Final Answer

Mg²⁺ < Na⁺ < F⁻ < O²⁻

All four species are isoelectronic with 10 electrons. As nuclear charge increases from Mg (Z = 12) to O (Z = 8), the attraction on the electrons decreases and the ionic size increases.

📝 20 MCQs with Answers

1. Which of the following species are isoelectronic?

A) O²⁻, F⁻, Na⁺, Mg²⁺

B) O, F, Na, Mg

C) Na, Mg, Al, Si

D) H₂, He, Li

✅ Answer

A) All contain 10 electrons.

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2. The number of electrons in O²⁻ is:

A) 6

B) 8

C) 10

D) 12

✅ Answer

C) 10

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3. The number of electrons in Mg²⁺ is:

A) 10

B) 12

C) 14

D) 8

✅ Answer

A) 10

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4. Isoelectronic species have the same number of:

A) Protons

B) Neutrons

C) Electrons

D) Nucleons

✅ Answer

C) Electrons

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5. Which ion has the highest nuclear charge?

A) O²⁻

B) F⁻

C) Na⁺

D) Mg²⁺

✅ Answer

D) Mg²⁺

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6. Among isoelectronic ions, size decreases with increase in:

A) Electron number

B) Nuclear charge

C) Mass number only

D) Neutron number

✅ Answer

B) Nuclear charge

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7. Which is the largest species?

A) O²⁻

B) F⁻

C) Na⁺

D) Mg²⁺

✅ Answer

A) O²⁻

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8. Which is the smallest species?

A) O²⁻

B) F⁻

C) Na⁺

D) Mg²⁺

✅ Answer

D) Mg²⁺

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9. Atomic number of oxygen is:

A) 6

B) 7

C) 8

D) 9

✅ Answer

C) 8

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10. Atomic number of fluorine is:

A) 8

B) 9

C) 10

D) 11

✅ Answer

B) 9

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11. Atomic number of sodium is:

A) 9

B) 10

C) 11

D) 12

✅ Answer

C) 11

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12. Atomic number of magnesium is:

A) 10

B) 11

C) 12

D) 13

✅ Answer

C) 12

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13. The electronic configuration of all four ions is:

A) 1s²2s²2p⁶

B) 1s²2s²2p⁴

C) 1s²2s²2p⁵

D) 1s²2s²2p⁶3s²

✅ Answer

A) 1s²2s²2p⁶

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14. Which factor mainly determines the size order of these ions?

A) Nuclear charge

B) Colour

C) Physical state

D) Melting point

✅ Answer

A) Nuclear charge

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15. Increasing nuclear charge causes the electron cloud to be:

A) More strongly attracted

B) Less strongly attracted

C) Completely removed

D) Unaffected

✅ Answer

A) More strongly attracted

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16. The correct decreasing order of size is:

A) Mg²⁺ > Na⁺ > F⁻ > O²⁻

B) O²⁻ > F⁻ > Na⁺ > Mg²⁺

C) F⁻ > O²⁻ > Mg²⁺ > Na⁺

D) Na⁺ > Mg²⁺ > O²⁻ > F⁻

✅ Answer

B)

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17. In an isoelectronic series, the ion with more protons is generally:

A) Larger

B) Smaller

C) Equal in size

D) Unstable always

✅ Answer

B) Smaller

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18. O²⁻ has how many more electrons than neutral O?

A) 1

B) 2

C) 3

D) 4

✅ Answer

B) 2

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19. Na⁺ is formed when sodium:

A) Gains one electron

B) Loses one electron

C) Gains two electrons

D) Loses two electrons

✅ Answer

B) Loses one electron

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20. The correct increasing order is:

A) O²⁻ < F⁻ < Na⁺ < Mg²⁺

B) Mg²⁺ < Na⁺ < F⁻ < O²⁻

C) Na⁺ < Mg²⁺ < F⁻ < O²⁻

D) F⁻ < Mg²⁺ < Na⁺ < O²⁻

✅ Answer

B) Mg²⁺ < Na⁺ < F⁻ < O²⁻

Q. Arrange O²⁻, F⁻, Na⁺ and Mg²⁺ in increasing order of size.

Solution: All four ions are isoelectronic with 10 electrons. In an isoelectronic series, ionic size decreases as nuclear charge increases.

Mg²⁺ < Na⁺ < F⁻ < O²⁻

Q. Explain the size order of O²⁻, F⁻, Na⁺ and Mg²⁺.

Solution:

1. O²⁻, F⁻, Na⁺ and Mg²⁺ each contain 10 electrons.
2. Hence, they form an isoelectronic series.
3. The nuclear charge increases from O (Z=8) to Mg (Z=12), so the attraction on the same number of electrons increases and the ionic radius decreases.

Increasing size:
Mg²⁺ < Na⁺ < F⁻ < O²⁻

Q. What are isoelectronic species? Arrange O²⁻, F⁻, Na⁺ and Mg²⁺ in increasing order of ionic size.

Solution:

Species having the same number of electrons are called isoelectronic species.

Ion	Electrons	Z
O2−	10	8
F−	10	9
Na+	10	11
Mg2+	10	12

Since all have the same number of electrons, the species with greater nuclear charge has the smaller radius.

Mg²⁺ < Na⁺ < F⁻ < O²⁻

Q. Explain the variation of ionic size in an isoelectronic series with a suitable example.

Solution:

• Isoelectronic species have the same number of electrons.
• For example, O²⁻, F⁻, Na⁺ and Mg²⁺ all have 10 electrons.
• The number of protons, however, is different.
• As the number of protons increases, the attractive force between the nucleus and electrons increases.
• Therefore, the ionic radius decreases with increasing nuclear charge.

Size order:

Mg²⁺ < Na⁺ < F⁻ < O²⁻

Q. Explain in detail why O²⁻, F⁻, Na⁺ and Mg²⁺ have different ionic sizes even though they are isoelectronic. Arrange them in increasing order of size.

Detailed Solution:

The ions O²⁻, F⁻, Na⁺ and Mg²⁺ are all isoelectronic because each contains 10 electrons. Their electronic configuration is:

1s² 2s² 2p⁶

Ion	Atomic Number	Electrons	Protons
O2−	8	10	8
F−	9	10	9
Na+	11	10	11
Mg2+	12	10	12

Because the number of electrons is identical, the main factor controlling the size difference is the nuclear charge.

O²⁻ has only 8 protons, so its 10 electrons experience the weakest nuclear attraction and remain relatively far from the nucleus. Hence, it has the largest radius.

Mg²⁺ has 12 protons attracting the same 10 electrons. It therefore has the strongest nuclear attraction and the smallest radius.

Mg²⁺ < Na⁺ < F⁻ < O²⁻

Increasing ionic size:
Mg²⁺ < Na⁺ < F⁻ < O²⁻

🎯 Final Concept:

For an isoelectronic series:
Same number of electrons + Higher nuclear charge = Smaller ionic radius.

🎯 Quick Revision

O²⁻ = 10 e⁻
F⁻ = 10 e⁻
Na⁺ = 10 e⁻
Mg²⁺ = 10 e⁻

Isoelectronic Series → Higher Z → Smaller Size

Increasing Size:
Mg²⁺ < Na⁺ < F⁻ < O²⁻

Why is the atomic radius of Noble Gases (Vander Waals radius) larger than the preceding Halogens?

📚 Chapter: Classification of Elements and Periodicity in Properties

📘 Detailed Solution

🔵 Step 1: Understand the important point

Noble gases such as Ne, Ar, Kr and Xe do not normally form covalent bonds with other atoms under ordinary conditions. Therefore, their atomic size is represented by their Van der Waals radius.

Noble Gas → Van der Waals Radius

On the other hand, the radius of a halogen such as F, Cl or Br is commonly considered as its covalent radius when discussing the periodic trend.

🟣 Step 2: Why is Noble Gas radius larger?

The Van der Waals radius represents the distance at which two non-bonded atoms can approach each other. It is therefore larger than the covalent radius, which represents half the distance between the nuclei of two bonded identical atoms.

Cl → Ar

Van der Waals radius > Covalent radius

Thus, although the atomic size generally decreases from left to right across a period due to increasing effective nuclear charge, the apparent radius of a noble gas is larger because a different type of radius is being used.

🟢 Step 3: Comparison

Halogen	Noble Gas
Usually considered using covalent radius	Usually considered using Van der Waals radius
Atoms form covalent bonds such as Cl₂	Atoms generally exist as isolated atoms
Covalent radius is smaller	Van der Waals radius is larger

🎬 Radius Concept

Halogen vs Noble Gas

Cl — Cl

Covalent radius

↓

Ar     Ar

Van der Waals radius

Van der Waals radius is larger

✅ Final Answer

The atomic radius of noble gases is apparently larger than that of the preceding halogens because noble gases are assigned their Van der Waals radius, whereas halogens are commonly compared using their covalent radius.

The Van der Waals radius is the distance associated with non-bonded atoms and is larger than the covalent radius. Therefore, the noble gas appears to have a larger atomic radius despite the general decrease in atomic size across a period.

📝 20 MCQs with Answers

1. The atomic radius generally assigned to noble gases is:

A) Ionic radius

B) Covalent radius

C) Van der Waals radius

D) Metallic radius

✅ Answer

C) Van der Waals radius

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2. Noble gases generally do not form:

A) Atomic nuclei

B) Ordinary covalent bonds

C) Electron shells

D) Isolated atoms

✅ Answer

B) Ordinary covalent bonds

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3. Which radius is larger?

A) Covalent radius

B) Van der Waals radius

C) Nuclear radius

D) None

✅ Answer

B) Van der Waals radius

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4. The preceding element to a noble gas in the same period is generally a:

A) Alkali metal

B) Halogen

C) Transition metal

D) Lanthanoid

✅ Answer

B) Halogen

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5. Chlorine is a member of:

A) Group 1

B) Group 2

C) Group 17

D) Group 18

✅ Answer

C) Group 17

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6. Argon belongs to:

A) Group 17

B) Group 18

C) Group 1

D) Group 16

✅ Answer

B) Group 18

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7. Van der Waals radius refers to:

A) Size of a bonded ion

B) Size associated with non-bonded atoms

C) Nuclear size

D) Proton size

✅ Answer

B) Size associated with non-bonded atoms

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8. Across a period, effective nuclear charge generally:

A) Decreases

B) Increases

C) Becomes zero

D) Remains exactly constant

✅ Answer

B) Increases

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9. The apparent increase in noble gas radius is mainly due to:

A) Increase in nuclear size

B) Use of Van der Waals radius

C) Loss of electrons

D) Formation of cations

✅ Answer

B) Use of Van der Waals radius

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10. Covalent radius is related to:

A) Bonded atoms

B) Free electrons

C) Nuclei only

D) Protons only

✅ Answer

A) Bonded atoms

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11. Which is generally greater?

A) Covalent radius of Cl

B) Van der Waals radius of Ar

C) Both are always equal

D) Nuclear radius of Cl

✅ Answer

B) Van der Waals radius of Ar

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12. Noble gases are chemically:

A) Highly reactive

B) Generally very unreactive

C) Strongly electropositive

D) Strong reducing metals

✅ Answer

B) Generally very unreactive

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13. The Van der Waals radius is associated with:

A) Non-bonded atoms

B) Only ionic compounds

C) Only metals

D) Only anions

✅ Answer

A) Non-bonded atoms

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14. Which pair is correctly matched?

A) Noble gases — Van der Waals radius

B) Noble gases — metallic radius

C) Halogens — nuclear radius

D) Halogens — proton radius

✅ Answer

A) Noble gases — Van der Waals radius

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15. Which element follows chlorine in the same period?

A) Na

B) Mg

C) Ar

D) K

✅ Answer

C) Ar

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16. The radius trend across a period normally shows:

A) General decrease

B) General increase

C) No change

D) Random change

✅ Answer

A) General decrease

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17. Why can noble gas radius not be directly compared with halogen covalent radius without qualification?

A) They have different atomic numbers

B) Different radius definitions are being used

C) Noble gases have no electrons

D) Halogens are metals

✅ Answer

B) Different radius definitions are being used

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18. The noble gases have a complete:

A) Valence shell

B) d-subshell

C) f-subshell

D) Nuclear shell

✅ Answer

A) Valence shell

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19. Which statement is correct?

A) Van der Waals radius is usually smaller than covalent radius

B) Van der Waals radius is generally larger than covalent radius

C) Both are always equal

D) Neither measures atomic size

✅ Answer

B) Van der Waals radius is generally larger than covalent radius

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20. The larger noble-gas radius does NOT mean that:

A) The radius definition has changed

B) Noble gas necessarily has a physically larger bonded atom

C) Van der Waals radius is larger

D) Comparison involves different radius definitions

✅ Answer

B) Noble gas necessarily has a physically larger bonded atom

Q. Why is the radius of noble gases larger than that of the preceding halogens?

Solution: Noble gases are assigned their Van der Waals radius, while halogens are commonly compared using their covalent radius. Since the Van der Waals radius is larger than the covalent radius, noble gases appear to have a larger atomic radius.

Q. Explain the apparent increase in atomic radius from a halogen to the noble gas in the same period.

Solution:

1. Atomic size generally decreases across a period because effective nuclear charge increases.
2. Halogens are commonly assigned their covalent radius.
3. Noble gases are assigned their larger Van der Waals radius because they do not normally form covalent bonds.

Therefore, Noble Gas Van der Waals Radius > Preceding Halogen Covalent Radius

Q. Differentiate between covalent radius and Van der Waals radius.

Covalent Radius	Van der Waals Radius
Used for bonded atoms	Used for non-bonded atoms
Half the internuclear distance between identical bonded atoms	Related to the closest approach of non-bonded atoms
Generally smaller	Generally larger
Commonly used for halogens when discussing covalent bonding	Used for noble gases

Q. Explain the periodic trend of atomic radius and the exception shown by noble gases.

Solution:

• Atomic radius generally decreases from left to right across a period.
• This is due to an increase in effective nuclear charge while electrons are added to the same principal shell.
• Halogens have their covalent radius considered because they form covalent molecules such as Cl₂.
• Noble gases generally do not form ordinary covalent molecules, so their Van der Waals radius is used.
• The Van der Waals radius is larger than the covalent radius, making the noble gas appear larger than the preceding halogen.

Key Point:
The apparent increase is mainly due to the different definitions of atomic radius, not a reversal of the normal periodic trend.

Q. Explain in detail why the atomic radius of noble gases is larger than that of the preceding halogens.

Detailed Solution:

Across a period, the atomic radius generally decreases from left to right. This happens because the nuclear charge increases while the added electrons enter the same principal energy level. Consequently, the effective nuclear charge increases and pulls the electron cloud closer to the nucleus.

However, an apparent exception is observed when the radius of a noble gas is compared with that of the preceding halogen. The reason is that different types of atomic radii are used.

Halogen	Noble Gas
Forms covalent molecules	Normally exists as individual atoms
Covalent radius is used	Van der Waals radius is used
Smaller radius	Larger radius

For example, chlorine forms the covalent molecule Cl₂. Therefore, its covalent radius can be obtained from the internuclear distance between the two bonded chlorine atoms.

Argon, however, does not normally form an Ar₂ covalent molecule under ordinary conditions. Hence, its atomic size is represented by its Van der Waals radius, based on the closest approach of non-bonded atoms.

Van der Waals Radius > Covalent Radius

Therefore, the noble gas appears to have a larger atomic radius than the preceding halogen even though the underlying periodic trend in atomic size across the period is a decrease.

🎯 Final Answer:

The atomic radius of noble gases appears larger than that of the preceding halogens because noble gases are assigned their Van der Waals radius, whereas halogens are commonly assigned their covalent radius. Since the Van der Waals radius is larger than the covalent radius, the noble gas appears larger. This is therefore mainly a consequence of using different definitions of atomic radius, rather than a reversal of the normal periodic trend.

🎯 Quick Revision

Halogen → Covalent Radius
Noble Gas → Van der Waals Radius
Van der Waals Radius > Covalent Radius
Reason → Different radius definitions
Across period → Atomic radius generally decreases