Why do s-block elements have very low IE?

📚 Chapter: Classification of Elements and Periodicity in Properties

📘 Answer

s-block elements generally have low first ionization enthalpy because their outermost electron is present in an s-orbital and is relatively easy to remove.

s-block → ns¹ or ns²

The valence electrons are comparatively far from the nucleus and experience effective nuclear attraction that is relatively low. Hence, only a small amount of energy is required to remove them.

🔹 Main Reasons for Low Ionization Enthalpy

1. Large atomic size: s-block elements generally have a relatively large atomic radius, so the valence electron is farther from the nucleus.
2. Low effective nuclear attraction: Inner-shell electrons shield the valence electron from the nucleus.
3. High shielding effect: The inner electrons reduce the effective attraction experienced by the outermost electron.
4. Single or two valence electrons: Group 1 elements have ns¹ and Group 2 elements have ns² configurations. The outer electron can be removed comparatively easily.
5. Formation of stable ions: Removal of the valence electron often leads to a stable noble-gas configuration.

⚛️ Example: Sodium

Sodium has atomic number 11 and electronic configuration:

Na = 1s² 2s² 2p⁶ 3s¹

Its outermost electron is present in the 3s orbital. It is relatively far from the nucleus and is shielded by the inner 10 electrons. Therefore, it can be removed easily.

Na → Na⁺ + e⁻

After losing one electron, sodium forms Na⁺ with the stable noble-gas configuration of neon:

Na⁺ = 1s² 2s² 2p⁶

📊 Group 1 vs Group 2

Group	General Configuration	IE
Group 1	ns¹	Very Low
Group 2	ns²	Low, but higher than Group 1

Group 1 elements generally have lower first ionization enthalpy than Group 2 elements in the same period because removal of one electron from ns¹ gives a stable noble-gas configuration.

✅ Final Answer

s-block elements have very low ionization enthalpy because their valence electrons are relatively far from the nucleus and are strongly shielded by inner-shell electrons. Therefore, the effective nuclear attraction on the valence electron is low and the electron can be removed easily.

Large Size + High Shielding → Weak Attraction → Low IE

20 MCQs with Answers

1. s-block elements generally have:

A) Very high IE

B) Low IE

C) Zero IE

D) Negative IE

✅ Answer

B) Low IE

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2. The valence electron of Group 1 elements is present in:

A) p-orbital

B) d-orbital

C) s-orbital

D) f-orbital

✅ Answer

C) s-orbital

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3. General valence configuration of Group 1 elements is:

A) ns²

B) ns¹

C) np¹

D) nd¹

✅ Answer

B) ns¹

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4. General valence configuration of Group 2 elements is:

A) ns¹

B) ns²

C) np²

D) nd²

✅ Answer

B) ns²

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5. Which element belongs to the s-block?

A) Na

B) Cl

C) Fe

D) Br

✅ Answer

A) Na

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6. Which factor reduces nuclear attraction on outer electrons?

A) Shielding effect

B) Atomic number only

C) Proton mass

D) Neutron number

✅ Answer

A) Shielding effect

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7. Higher shielding generally causes ionization enthalpy to:

A) Increase

B) Decrease

C) Become infinite

D) Remain zero

✅ Answer

B) Decrease

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8. Sodium has which valence configuration?

A) 3s¹

B) 3s²

C) 3p¹

D) 2p⁶

✅ Answer

A) 3s¹

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9. Which has lower first IE in the same period?

A) Na

B) Mg

C) Al

D) Si

✅ Answer

A) Na

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10. Removal of one electron from Na produces:

A) Na⁻

B) Na⁺

C) Na²⁺

D) Na²⁻

✅ Answer

B) Na⁺

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11. Which configuration is obtained by Na⁺?

A) 2,8

B) 2,8,1

C) 2,7

D) 2,8,8

✅ Answer

A) 2,8

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12. Which group contains alkali metals?

A) Group 1

B) Group 2

C) Group 17

D) Group 18

✅ Answer

A) Group 1

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13. Which group contains alkaline earth metals?

A) Group 1

B) Group 2

C) Group 16

D) Group 18

✅ Answer

B) Group 2

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14. In general, atomic size down Group 1:

A) Decreases

B) Increases

C) Remains unchanged

D) Becomes zero

✅ Answer

B) Increases

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15. As atomic size increases, ionization enthalpy generally:

A) Increases

B) Decreases

C) Becomes zero always

D) Doubles

✅ Answer

B) Decreases

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16. Which is generally more easily removed?

A) Inner-shell electron

B) Valence electron of an alkali metal

C) Proton

D) Neutron

✅ Answer

B

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17. s-block elements are mainly found in:

A) Groups 1 and 2

B) Groups 13–18

C) Groups 3–12

D) Only Group 18

✅ Answer

A) Groups 1 and 2

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18. Which generally has lower first IE?

A) K

B) Na

C) Li

D) Mg

✅ Answer

A) K

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19. The outermost electron in s-block elements is generally:

A) Easy to remove

B) Impossible to remove

C) In the nucleus

D) A proton

✅ Answer

A) Easy to remove

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20. Correct relationship is:

A) Shielding ↑ → IE ↑

B) Atomic size ↑ → IE ↓

C) Atomic size ↑ → IE ↑ always

D) Shielding has no effect

✅ Answer

B) Atomic size ↑ → IE ↓

Q. Why do s-block elements have low ionization enthalpy?

Solution:

s-block elements have relatively large atomic size and significant shielding of their valence electrons by inner-shell electrons. Therefore, the outermost electron experiences weaker effective nuclear attraction and can be removed easily.

Hence, s-block elements generally have low ionization enthalpy.

Q. Give three reasons for the low ionization enthalpy of s-block elements.

Solution:

1. They generally have relatively large atomic size.
2. Their valence electrons experience considerable shielding by inner electrons.
3. The effective nuclear attraction on the outer electron is relatively low.

Q. Explain why Group 1 elements have very low first ionization enthalpy.

Solution:

1. Group 1 elements have the general configuration ns¹.
2. The single valence electron is relatively far from the nucleus.
3. It is shielded by the inner-shell electrons.
4. Removal of this electron gives a stable noble-gas configuration.

Therefore, Group 1 elements have very low first ionization enthalpy.

Q. Explain the factors responsible for low ionization enthalpy of s-block elements with an example.

Solution:

The low ionization enthalpy of s-block elements is mainly due to their relatively large atomic size, high shielding effect and weak effective nuclear attraction on the valence electron.

For example, sodium has configuration:

Na = 2,8,1

Its single 3s electron is far from the nucleus and is shielded by the inner 10 electrons. Therefore, it can be removed easily:

Na → Na⁺ + e⁻

Thus, s-block elements generally have low ionization enthalpy.

Q. Explain in detail why s-block elements have low ionization enthalpy. Discuss the effect of atomic size, shielding effect and electronic configuration.

Detailed Solution

s-block elements are those elements in which the last electron enters an s-orbital. They mainly include Groups 1 and 2. Their general outer electronic configurations are ns¹ and ns².

1. Large Atomic Size

The valence electron is relatively far from the nucleus. Greater distance reduces the attractive force between the nucleus and the valence electron, making its removal easier.

2. Shielding Effect

The inner-shell electrons shield the valence electron from the attraction of the nucleus. Therefore, the effective nuclear attraction experienced by the outer electron is reduced.

Shielding ↑ → Effective Attraction ↓ → IE ↓

3. Electronic Configuration

Group 1 elements have ns¹ configuration. Removal of this single electron gives a stable noble-gas configuration. Therefore, the first electron can be removed comparatively easily.

Example: Sodium

Na = 1s² 2s² 2p⁶ 3s¹

The 3s electron is the valence electron. It is relatively far from the nucleus and is shielded by the inner electrons. On removal:

Na → Na⁺ + e⁻

Na⁺ obtains the stable configuration of neon: 1s² 2s² 2p⁶.

Final Conclusion:

s-block elements have low ionization enthalpy because their valence electrons are relatively far from the nucleus, strongly shielded by inner electrons, and are held comparatively weakly.

Large Atomic Size + Shielding + ns¹/ns² → Low IE

🎯 Quick Revision

s-block: Groups 1 & 2
Outer Configuration: ns¹ / ns²
Atomic Size: Relatively large
Shielding: High
Valence Electron: Relatively easy to remove
Ionization Enthalpy: Low

Shielding ↑ + Size ↑ → IE ↓