Define First Ionization Enthalpy. Why is the first IE of Nitrogen higher than Oxygen? Explain the trend of IE down a group.

📚 Chapter: Classification of Elements and Periodicity in Properties

📘 Definition

First Ionization Enthalpy is the minimum amount of energy required to remove the most loosely bound electron from one mole of isolated gaseous atoms in their ground state to form one mole of gaseous unipositive ions.

X(g) → X⁺(g) + e⁻

It is represented by Δ_iH₁ and is generally expressed in kJ mol⁻¹.

First Ionization Enthalpy = Energy required to remove the first electron from an isolated gaseous atom.

📘 Explanation

Although ionization enthalpy generally increases from left to right across a period, Nitrogen has a higher first ionization enthalpy than Oxygen. This is an important exception in the second period.

Element	Atomic Number	Electronic Configuration
N	7	1s² 2s² 2p³
O	8	1s² 2s² 2p⁴

🔹 Nitrogen

Nitrogen has a half-filled 2p³ configuration. Its three 2p orbitals contain one electron each.

2p³ → ↑   ↑   ↑

A half-filled subshell is relatively stable because of symmetrical distribution and exchange energy. Therefore, more energy is required to remove an electron from nitrogen.

🔹 Oxygen

Oxygen has the configuration 2p⁴. One of the 2p orbitals contains a pair of electrons.

2p⁴ → ↑↓   ↑   ↑

The two electrons in the same orbital experience greater electron-electron repulsion. Therefore, one electron can be removed comparatively easily from oxygen.

Therefore, IE₁(N) > IE₁(O)

Nitrogen has a stable half-filled 2p³ configuration, whereas oxygen has electron-electron repulsion in one paired 2p orbital.

📘 Trend Down a Group

The ionization enthalpy generally decreases as we move down a group in the periodic table.

Ionization Enthalpy ↓ down the group

🔹 Reasons

1. Increase in atomic size: A new electron shell is added at each step down the group.
2. Increase in shielding effect: The inner electrons shield the valence electron from the attraction of the nucleus.
3. Increase in distance: The outermost electron is farther from the nucleus.
4. Reduced effective nuclear attraction: The valence electron is held less strongly by the nucleus.
5. Therefore, less energy is required to remove the outermost electron.

Atomic size ↑   →   Shielding effect ↑   →   Ionization Enthalpy ↓

📊 Example: Group 1

Element	Trend in Atomic Size	Trend in IE
Li	Small	High
Na	↑	↓
K	↑	↓
Rb	↑	↓
Cs	Large	Low

Thus, ionization enthalpy generally decreases down a group.

📝 20 MCQs with Answers

1. First ionization enthalpy is the energy required to:

A) Add an electron to an atom

B) Remove the first electron from an isolated gaseous atom

C) Remove a proton

D) Add a proton

✅ Answer

B

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2. First ionization enthalpy is generally expressed in:

A) g mol⁻¹

B) L mol⁻¹

C) kJ mol⁻¹

D) mol L⁻¹

✅ Answer

C) kJ mol⁻¹

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3. The process X(g) → X⁺(g) + e⁻ represents:

A) Electron affinity

B) First ionization

C) Second ionization

D) Bond formation

✅ Answer

B) First ionization

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4. Which has higher first ionization enthalpy?

A) N

B) O

C) Na

D) K

✅ Answer

A) N

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5. Electronic configuration of nitrogen is:

A) 1s² 2s² 2p²

B) 1s² 2s² 2p³

C) 1s² 2s² 2p⁴

D) 1s² 2s² 2p⁵

✅ Answer

B

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6. Nitrogen has high IE because of its:

A) Fully filled p-subshell

B) Half-filled p-subshell

C) Large size

D) Low nuclear charge

✅ Answer

B) Half-filled p-subshell

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7. Oxygen has the configuration:

A) 2p²

B) 2p³

C) 2p⁴

D) 2p⁵

✅ Answer

C) 2p⁴

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8. In oxygen, electron-electron repulsion occurs due to:

A) Unpaired electrons only

B) Paired electrons in one 2p orbital

C) Protons

D) Neutrons

✅ Answer

B

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9. Ionization enthalpy generally decreases:

A) Across a period

B) Down a group

C) From right to left only

D) Randomly

✅ Answer

B) Down a group

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10. Which factor increases down a group?

A) Atomic size

B) Ionization enthalpy

C) Effective nuclear attraction on valence electron

D) Electronegativity always

✅ Answer

A) Atomic size

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11. Shielding effect down a group:

A) Decreases

B) Increases

C) Becomes zero

D) Remains unchanged

✅ Answer

B) Increases

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12. Greater atomic size generally causes ionization enthalpy to:

A) Increase

B) Decrease

C) Become zero

D) Remain constant

✅ Answer

B) Decrease

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13. Which is easier to ionize?

A) An atom with a tightly held electron

B) An atom with a loosely held valence electron

C) A nucleus

D) A proton

✅ Answer

B

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14. The unusual IE trend between N and O is due to:

A) Half-filled stability and electron repulsion

B) Nuclear fission

C) Equal atomic radii

D) Different periods

✅ Answer

A

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15. Which subshell is half-filled in nitrogen?

A) 2s

B) 2p

C) 3s

D) 3p

✅ Answer

B) 2p

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16. Which element has the lowest IE among Li, Na and K?

A) Li

B) Na

C) K

D) All equal

✅ Answer

C) K

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17. Down a group, the valence electron is:

A) Closer to nucleus

B) Farther from nucleus

C) Removed automatically

D) Converted into proton

✅ Answer

B

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18. Which generally has higher IE?

A) Smaller atom

B) Larger atom

C) Larger ion always

D) None

✅ Answer

A) Smaller atom

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19. Ionization enthalpy is a measure of:

A) Ease of removing an electron

B) Ease of adding a proton

C) Number of neutrons

D) Atomic mass only

✅ Answer

A

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20. Correct statement is:

A) IE always decreases across a period

B) IE generally increases across a period and decreases down a group

C) IE always increases down a group

D) IE is independent of atomic size

✅ Answer

B

Q1. Define first ionization enthalpy.

It is the minimum energy required to remove the first electron from one mole of isolated gaseous atoms in their ground state to form one mole of gaseous unipositive ions.

X(g) → X⁺(g) + e⁻

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Q2. Why is IE of nitrogen greater than oxygen?

Nitrogen has a stable half-filled 2p³ configuration, whereas oxygen has 2p⁴ configuration with one pair of electrons. Electron-electron repulsion in the paired orbital makes electron removal from oxygen easier.

IE₁(N) > IE₁(O)

Q1. Explain why ionization enthalpy decreases down a group.

1. Atomic size increases down the group.
2. Shielding effect increases due to addition of shells.
3. The valence electron is farther from the nucleus.
4. Nuclear attraction on the outer electron decreases.
5. Hence, less energy is required to remove it.

IE decreases down a group.

Q2. Why is nitrogen an exception in the IE trend between N and O?

Nitrogen has the particularly stable half-filled 2p³ configuration. Oxygen has 2p⁴ configuration, in which two electrons occupy the same 2p orbital and experience repulsion. Therefore, oxygen loses an electron more easily than nitrogen.

Q. Explain first ionization enthalpy and discuss its trend down a group.

First ionization enthalpy is the minimum energy required to remove the first electron from one mole of isolated gaseous atoms in their ground state.

X(g) → X⁺(g) + e⁻

Down a group:

• Number of shells increases.
• Atomic radius increases.
• Shielding effect increases.
• Valence electron becomes farther from the nucleus.
• Effective nuclear attraction decreases.

Therefore, first ionization enthalpy generally decreases down a group.

Q. Explain why the first ionization enthalpy of nitrogen is greater than that of oxygen.

The electronic configuration of nitrogen is 1s² 2s² 2p³. The three 2p electrons occupy the three 2p orbitals singly. This is a half-filled configuration, which possesses extra stability.

N: 2p³ → ↑   ↑   ↑

Oxygen has the configuration 1s² 2s² 2p⁴. One of its 2p orbitals contains a pair of electrons. The electrons in the same orbital experience repulsion.

O: 2p⁴ → ↑↓   ↑   ↑

Therefore, an electron can be removed more easily from oxygen than from nitrogen.

IE₁(N) > IE₁(O)

Q1. Define first ionization enthalpy and explain its periodic trend down a group.

Solution

First ionization enthalpy is the minimum energy required to remove the most loosely bound electron from one mole of isolated gaseous atoms in their ground state to form one mole of gaseous unipositive ions.

X(g) → X⁺(g) + e⁻

Trend Down a Group

Ionization enthalpy generally decreases down a group due to the following factors:

1. Increase in the number of electron shells.
2. Increase in atomic radius.
3. Increase in shielding effect.
4. Greater distance of the valence electron from the nucleus.
5. Decrease in effective nuclear attraction on the valence electron.
6. Consequently, less energy is needed to remove the electron.

Therefore, Ionization Enthalpy ↓ down a group.

Q2. Explain why the first ionization enthalpy of nitrogen is greater than that of oxygen.

Solution

Element	Configuration	Important Feature
N	1s² 2s² 2p³	Stable half-filled 2p subshell
O	1s² 2s² 2p⁴	Paired-electron repulsion

Nitrogen has a stable half-filled 2p³ configuration. In oxygen, one 2p orbital contains a pair of electrons. The mutual repulsion between these paired electrons makes removal of one electron easier.

Hence, the first ionization enthalpy of nitrogen is greater than that of oxygen.

IE₁(N) > IE₁(O)

🎯 Quick Revision

First IE: Energy required to remove first electron
General trend across period: IE ↑
General trend down group: IE ↓
Reason down group: Atomic size ↑ + Shielding ↑
Important exception: IE(N) > IE(O)
Reason: N has stable half-filled 2p³; O has paired-electron repulsion