Which element in the periodic table has the highest IE?

📚 Chapter: Classification of Elements and Periodicity in Properties

📘 Answer

Helium (He)

Helium (He), with atomic number 2, has the highest first ionization enthalpy among all elements.

Highest First Ionization Enthalpy → Helium (He)

🔹 Why does Helium have the highest IE?

1. Helium has a very small atomic size.
2. Its electrons are very close to the nucleus.
3. It has a high effective nuclear attraction on its electrons.
4. Its electronic configuration is 1s², which is a stable and completely filled shell.
5. Therefore, a large amount of energy is required to remove an electron.

He = 1s²

📈 Periodic Trend

Ionization enthalpy generally:

Increases → Across a Period →

Decreases ↓ Down a Group ↓

Helium is located at the top-right corner of the periodic table, where the first ionization enthalpy is maximum.

General Trend: IE increases from left to right and decreases down a group.

⚛️ Comparison with Other Noble Gases

Element	Electronic Configuration	Relative IE
He	1s²	Highest
Ne	2s²2p⁶	Very High
Ar	3s²3p⁶	High

Although noble gases have high ionization enthalpies because of their stable closed-shell configurations, helium has the highest due to its very small size and strong nuclear attraction.

✅ Final Answer

Helium (He) has the highest first ionization enthalpy in the periodic table because it has a very small atomic size, high effective nuclear attraction, and a stable 1s² electronic configuration.

20 MCQs with Answers

1. Which element has the highest first ionization enthalpy?

A) H

B) He

C) Ne

D) F

✅ Answer

B) He

---

2. The atomic number of helium is:

A) 1

B) 2

C) 3

D) 4

✅ Answer

B) 2

---

3. Electronic configuration of He is:

A) 1s¹

B) 1s²

C) 2s²

D) 2p⁶

✅ Answer

B) 1s²

---

4. Helium has high IE mainly because of its:

A) Large size

B) Small size

C) Low nuclear attraction

D) High shielding only

✅ Answer

B) Small size

---

5. Ionization enthalpy generally increases across a period from:

A) Right to left

B) Left to right

C) Top to bottom

D) Bottom to top only

✅ Answer

B) Left to right

---

6. Ionization enthalpy generally decreases:

A) Across a period

B) Down a group

C) From right to left only

D) Nowhere

✅ Answer

B) Down a group

---

7. Helium belongs to:

A) Group 1

B) Group 2

C) Group 17

D) Group 18

✅ Answer

D) Group 18

---

8. Helium is a:

A) Alkali metal

B) Halogen

C) Noble gas

D) Transition metal

✅ Answer

C) Noble gas

---

9. Which factor strongly contributes to high IE of He?

A) Large atomic radius

B) Small atomic radius

C) High shielding

D) Many shells

✅ Answer

B

---

10. The valence shell of helium contains:

A) 1 electron

B) 2 electrons

C) 6 electrons

D) 8 electrons

✅ Answer

B) 2 electrons

---

11. Which has higher first IE?

A) Li

B) He

C) Na

D) K

✅ Answer

B) He

---

12. Noble gases generally have:

A) Low IE

B) High IE

C) Zero IE

D) Negative IE

✅ Answer

B) High IE

---

13. The nucleus attracts electrons because the nucleus is:

A) Negatively charged

B) Positively charged

C) Neutral

D) Uncharged

✅ Answer

B) Positively charged

---

14. A smaller atomic radius generally results in:

A) Lower IE

B) Higher IE

C) Zero IE

D) No relationship

✅ Answer

B) Higher IE

---

15. Which element is at the top-right of the periodic table?

A) Cs

B) He

C) Fr

D) Na

✅ Answer

B) He

---

16. Stable configuration of He is:

A) Half-filled

B) Completely filled first shell

C) Partially filled p-shell

D) Empty shell

✅ Answer

B

---

17. Which has the smallest atomic size among the noble gases?

A) He

B) Ne

C) Ar

D) Kr

✅ Answer

A) He

---

18. High ionization enthalpy means an electron is:

A) Easy to remove

B) Difficult to remove

C) Already removed

D) A proton

✅ Answer

B) Difficult to remove

---

19. Which statement is correct?

A) He has the lowest first IE

B) He has the highest first IE

C) He has zero IE

D) He is a metal

✅ Answer

B

---

20. The highest first ionization enthalpy in the periodic table belongs to:

A) Hydrogen

B) Helium

C) Fluorine

D) Neon

✅ Answer

B) Helium

Q. Which element has the highest first ionization enthalpy? Give one reason.

Solution: Helium (He) has the highest first ionization enthalpy because it has a very small atomic size and its electrons are strongly attracted by the nucleus.

Answer: Helium (He)

Q. Why does helium have the highest first ionization enthalpy?

Solution:

1. Helium has a very small atomic radius.
2. Its electrons are very close to the nucleus.
3. Its stable configuration is 1s².
4. Therefore, its electron is held very strongly and requires maximum energy for removal.

Q. Explain the periodic trend of ionization enthalpy and identify the element having the highest first IE.

Solution:

• Ionization enthalpy generally increases from left to right across a period.
• It generally decreases from top to bottom in a group.
• The highest value occurs at the top-right of the periodic table.
• Helium (He) has the highest first ionization enthalpy.

Highest First IE = Helium (He)

Q. Explain why helium has the highest ionization enthalpy among all elements.

Solution:

1. Helium has atomic number 2 and electronic configuration 1s².
2. It has only one occupied electron shell.
3. The atomic size of helium is extremely small.
4. Its electrons experience strong attraction from the nucleus.
5. The 1s orbital is completely filled and highly stable.
6. Therefore, a very large amount of energy is required to remove an electron.

Hence, Helium (He) has the highest first ionization enthalpy.

Q. What is ionization enthalpy? Explain its periodic trend and state which element has the highest first ionization enthalpy.

Detailed Solution

Ionization enthalpy is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.

X(g) → X⁺(g) + e⁻

Trend Across a Period

From left to right, effective nuclear charge generally increases and atomic size decreases. Therefore, electrons are held more strongly and ionization enthalpy generally increases.

Across a period: IE ↑

Trend Down a Group

Down a group, new shells are added. Atomic size and shielding effect increase, so the outer electron is easier to remove. Hence, ionization enthalpy generally decreases.

Down a group: IE ↓

Highest Ionization Enthalpy

Helium has the highest first ionization enthalpy because of its very small size, high effective nuclear attraction, and stable completely filled 1s² configuration.

Final Answer:

Highest First Ionization Enthalpy = Helium (He)

🎯 Quick Revision

Highest First IE: He
Atomic Number: 2
Configuration: 1s²
Group: 18
Period: 1
Across Period: IE generally ↑
Down Group: IE generally ↓

How does Screening Effect (Shielding) affect Ionization Enthalpy?

📚 Chapter: Classification of Elements and Periodicity in Properties

📘 Screening Effect / Shielding Effect

The screening effect is the reduction in the attractive force of the nucleus on the outermost electron due to the presence of inner-shell electrons.

The inner electrons act as a shield between the positively charged nucleus and the valence electrons.

Nucleus → Inner Electrons → Valence Electron
Inner electrons shield the valence electron from the nucleus.

🔹 Effect on Ionization Enthalpy

When the screening effect increases, the outermost electron experiences less effective nuclear attraction. Therefore, the electron is held less strongly and can be removed more easily.

Shielding Effect ↑ → Nuclear Attraction ↓ → Ionization Enthalpy ↓

Thus, an increase in screening effect generally results in a decrease in ionization enthalpy.

📈 Trend Down a Group

As we move down a group, new electron shells are added. The number of inner-shell electrons increases, so the shielding effect also increases.

Down the Group	Effect
Number of shells ↑	Shielding effect ↑
Atomic size ↑	Distance of valence electron from nucleus ↑
Nuclear attraction on valence electron ↓	Ionization enthalpy ↓

Therefore, ionization enthalpy generally decreases down a group because the shielding effect increases.

⚛️ Example: Alkali Metals

Consider the Group 1 elements:

Li → Na → K → Rb → Cs

Down the group, more inner shells are added. These inner electrons shield the valence electron from the nucleus. Hence, the outermost electron is easier to remove.

Shielding effect increases → Ionization enthalpy decreases.

⚠️ Important Point

The shielding effect is one of the major factors affecting ionization enthalpy. However, ionization enthalpy also depends on atomic size, effective nuclear charge, and electronic configuration.

✅ Final Answer

The screening or shielding effect reduces the effective attraction of the nucleus on the outermost electron. As shielding increases, the valence electron is held less strongly and can be removed more easily. Therefore, ionization enthalpy decreases with an increase in shielding effect.

Shielding ↑ → Effective Nuclear Attraction ↓ → IE ↓

20 MCQs with Answers

1. Screening effect is caused mainly by:

A) Protons

B) Inner-shell electrons

C) Neutrons

D) Nucleus only

✅ Answer

B) Inner-shell electrons

---

2. Screening effect reduces:

A) Atomic mass

B) Nuclear charge

C) Effective nuclear attraction on valence electrons

D) Number of protons

✅ Answer

C

---

3. When shielding effect increases, ionization enthalpy generally:

A) Increases

B) Decreases

C) Becomes zero

D) Remains unchanged

✅ Answer

B) Decreases

---

4. Inner electrons act as a:

A) Magnet

B) Shield

C) Proton source

D) Neutron source

✅ Answer

B) Shield

---

5. Shielding effect generally increases:

A) Across a period strongly

B) Down a group

C) From right to left only

D) Nowhere

✅ Answer

B) Down a group

---

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

A) Decreases

B) Increases

C) Remains zero

D) Remains exactly one

✅ Answer

B) Increases

---

7. Greater shielding makes the valence electron:

A) More strongly held

B) Less strongly held

C) Converted into a proton

D) Closer to nucleus

✅ Answer

B

---

8. Greater shielding generally causes atomic radius to:

A) Increase

B) Decrease

C) Become zero

D) Remain unchanged

✅ Answer

A) Increase

---

9. Ionization enthalpy is the energy required to remove:

A) Proton

B) Neutron

C) Electron

D) Nucleus

✅ Answer

C) Electron

---

10. Which group shows a clear increase in shielding down the group?

A) Group 1

B) Only Group 18

C) No group

D) Only Group 17

✅ Answer

A) Group 1

---

11. Effective nuclear charge experienced by a valence electron is reduced by:

A) Shielding

B) Proton addition only

C) Neutron removal

D) Bond formation only

✅ Answer

A) Shielding

---

12. Which is easier when shielding is greater?

A) Removing a valence electron

B) Removing a proton

C) Removing a neutron

D) Removing the nucleus

✅ Answer

A

---

13. The outermost electron is also called:

A) Core electron

B) Valence electron

C) Nuclear electron

D) Proton electron

✅ Answer

B) Valence electron

---

14. Which factor generally decreases ionization enthalpy down a group?

A) Shielding effect

B) Nuclear charge

C) Number of protons

D) Atomic number

✅ Answer

A

---

15. Shielding is also called:

A) Screening effect

B) Magnetic effect

C) Photoelectric effect

D) Nuclear effect

✅ Answer

A) Screening effect

---

16. If shielding increases, the attraction between nucleus and valence electron:

A) Increases

B) Decreases

C) Becomes infinite

D) Remains unchanged

✅ Answer

B) Decreases

---

17. Which has generally greater shielding?

A) Li

B) Na

C) K

D) All are exactly equal

✅ Answer

C) K

---

18. Increased shielding is associated with:

A) More inner electron shells

B) Fewer shells

C) Fewer protons only

D) No electrons

✅ Answer

A

---

19. Which statement is correct?

A) Shielding increases IE

B) Shielding decreases IE

C) Shielding has no effect

D) Shielding removes protons

✅ Answer

B

---

20. Correct relationship is:

A) Shielding ↑ → IE ↑

B) Shielding ↑ → IE ↓

C) Shielding ↓ → Atomic size ↑ always

D) Shielding = IE

✅ Answer

B) Shielding ↑ → IE ↓

Q. How does shielding effect affect ionization enthalpy?

The inner-shell electrons shield the valence electron from the nucleus. Therefore, the effective nuclear attraction decreases and the valence electron becomes easier to remove.

Shielding effect ↑ → Ionization enthalpy ↓

Q. Explain the role of shielding effect in determining ionization enthalpy.

1. Inner-shell electrons shield the valence electrons from the nucleus.
2. Greater shielding reduces effective nuclear attraction.
3. The valence electron is held less strongly.
4. Hence, less energy is required for its removal.

Therefore, greater shielding generally results in lower ionization enthalpy.

Q. Explain the relationship between shielding effect and ionization enthalpy down a group.

As we move down a group, new electron shells are added. The number of inner-shell electrons increases and consequently the shielding effect increases.

1. Number of shells ↑
2. Shielding effect ↑
3. Effective nuclear attraction on valence electron ↓
4. Atomic size ↑
5. Ionization enthalpy ↓

Shielding ↑ → Attraction ↓ → Atomic Size ↑ → IE ↓

Q. What is shielding effect? Explain its effect on ionization enthalpy with a suitable example.

The shielding effect is the reduction in the effective nuclear attraction experienced by an outer electron due to the presence of inner-shell electrons.

For example, consider Li and Na. Sodium has an additional electron shell compared with lithium. The inner electrons in sodium shield its valence electron from the nucleus more effectively.

Li: 2,1
Na: 2,8,1

Therefore, the outer electron of Na is farther from the nucleus and is more strongly shielded. It is easier to remove than the valence electron of Li. Hence, sodium has lower first ionization enthalpy than lithium.

Greater shielding → weaker effective nuclear attraction → lower ionization enthalpy.

Q. Define screening effect and explain in detail how it affects ionization enthalpy.

Detailed Solution

The screening effect, also called the shielding effect, is the reduction in the attractive force of the nucleus on the outermost electron due to the repulsion and shielding provided by inner-shell electrons.

Effect on Ionization Enthalpy

1. Inner electrons lie between the nucleus and valence electrons.
2. They partially shield the valence electrons from nuclear attraction.
3. Thus, the effective nuclear attraction experienced by the valence electron decreases.
4. The valence electron becomes farther from and less strongly bound to the nucleus.
5. Consequently, less energy is required to remove the electron.
6. Therefore, ionization enthalpy decreases as the shielding effect increases.

Shielding Effect ↑
↓
Effective Nuclear Attraction ↓
↓
Valence Electron Held Less Strongly
↓
Ionization Enthalpy ↓

Example

Down Group 1, the shielding effect increases:

Li → Na → K → Rb → Cs

As a result, the outermost electron becomes progressively easier to remove, so the first ionization enthalpy generally decreases down the group.

Final Conclusion:

An increase in screening effect reduces the effective nuclear attraction on the valence electron and therefore lowers the ionization enthalpy.

Shielding ↑ → Effective Nuclear Charge ↓ → IE ↓

🎯 Quick Revision

Shielding Effect: Inner electrons protect outer electrons from nuclear attraction.
Down a group: Shielding ↑
Effective nuclear attraction: ↓
Atomic size: ↑
Ionization enthalpy: ↓

Shielding ↑ → IE ↓

Why is the second IE always higher than the first IE for any element?

📚 Chapter: Classification of Elements and Periodicity in Properties

📘 Detailed Explanation

The second ionization enthalpy (IE₂) is always greater than the first ionization enthalpy (IE₁) because the second electron is removed from a positively charged ion, whereas the first electron is removed from a neutral atom.

First Ionization:
X(g) → X⁺(g) + e⁻

Second Ionization:
X⁺(g) → X²⁺(g) + e⁻

🔹 Why does IE₂ > IE₁?

1. After the first electron is removed, the atom becomes a positively charged ion.
2. The remaining electrons experience a stronger attraction towards the nucleus because the ion has fewer electrons but the same number of protons.
3. The electron-electron repulsion also decreases after removal of the first electron.
4. Therefore, the remaining electron is held more strongly by the nucleus.
5. Consequently, more energy is required to remove the second electron.

IE₂ > IE₁

⚛️ Example: Sodium (Na)

Sodium has atomic number 11 and electronic configuration:

Na = 2, 8, 1

The first electron is removed from the outermost 3s orbital:

Na → Na⁺ + e⁻
2,8,1 → 2,8

After this, Na⁺ has a stable noble-gas configuration of 2,8. The second electron must now be removed from the inner, completely filled shell. It is held very strongly by the nucleus.

Na⁺ → Na²⁺ + e⁻

For Na, the jump from IE₁ to IE₂ is especially large because the second electron must be removed from a stable inner shell.

📊 Comparison

Property	First IE	Second IE
Starting species	Neutral atom	Positive ion
Charge	0	+1
Electron removed	First electron	Second electron
Attraction by nucleus	Relatively lower	Stronger
Energy required	Lower	Higher

✅ Final Answer

The second ionization enthalpy is always higher than the first because after removal of the first electron, the atom becomes a positively charged ion. The remaining electrons are attracted more strongly by the nucleus. Therefore, more energy is required to remove the second electron.

IE₂ > IE₁

20 MCQs with Answers

1. Which is generally greater?

A) IE₁

B) IE₂

C) Both are equal

D) Neither

✅ Answer

B) IE₂

---

2. The second electron is removed from:

A) Neutral atom

B) Negative ion

C) Positive ion

D) Molecule only

✅ Answer

C) Positive ion

---

3. First ionization process is:

A) X → X⁻ + e⁻

B) X → X⁺ + e⁻

C) X⁺ → X²⁺ + e⁻

D) X²⁺ → X³⁺ + e⁻

✅ Answer

B

---

4. Second ionization process is:

A) X → X⁺ + e⁻

B) X⁺ → X²⁺ + e⁻

C) X⁻ → X + e⁻

D) X²⁺ → X + e⁻

✅ Answer

B

---

5. After first ionization, the atom becomes:

A) Neutral

B) Negative

C) Positive

D) Unstable nucleus

✅ Answer

C) Positive

---

6. Why is the second electron held more strongly?

A) More protons are created

B) Remaining electrons experience stronger nuclear attraction

C) Nucleus disappears

D) Electron mass increases

✅ Answer

B

---

7. Which charge is present on the ion before second ionization?

A) 0

B) −1

C) +1

D) +2

✅ Answer

C) +1

---

8. Which has fewer electrons?

A) Neutral atom

B) X⁺

C) Both equal

D) X⁻

✅ Answer

B) X⁺

---

9. Removal of the first electron generally causes electron-electron repulsion to:

A) Increase

B) Decrease

C) Become infinite

D) Remain exactly unchanged

✅ Answer

B) Decrease

---

10. Sodium has electronic configuration:

A) 2,7,2

B) 2,8,1

C) 2,8,2

D) 2,8,8

✅ Answer

B) 2,8,1

---

11. Na⁺ has configuration:

A) 2,8,1

B) 2,8

C) 2,7

D) 2,8,8

✅ Answer

B) 2,8

---

12. For sodium, the second electron is removed from:

A) 3s orbital

B) 2p orbital

C) 1s orbital

D) 3p orbital

✅ Answer

B) 2p orbital

---

13. Which statement is correct?

A) IE₂ < IE₁

B) IE₂ = IE₁ always

C) IE₂ > IE₁

D) IE₂ = 0

✅ Answer

C) IE₂ > IE₁

---

14. The nucleus has:

A) Positive charge

B) Negative charge

C) Zero charge

D) Variable charge only

✅ Answer

A) Positive charge

---

15. Removal of an electron from a cation is:

A) Easier than from neutral atom

B) More difficult than from neutral atom

C) Always impossible

D) Independent of charge

✅ Answer

B

---

16. A very large jump between successive IEs may indicate:

A) Removal of an inner-shell electron after valence electrons are removed

B) Addition of neutrons

C) Formation of nucleus

D) Melting

✅ Answer

A

---

17. Which ion is formed after first ionization?

A) X⁻

B) X⁺

C) X²⁺

D) X²⁻

✅ Answer

B) X⁺

---

18. Ionization enthalpy involves removal of:

A) Proton

B) Neutron

C) Electron

D) Nucleus

✅ Answer

C) Electron

---

19. After one electron is removed, effective attraction on remaining electrons generally:

A) Becomes weaker

B) Becomes stronger

C) Becomes zero

D) Does not exist

✅ Answer

B) Becomes stronger

---

20. The main reason IE₂ is higher than IE₁ is:

A) Second electron is removed from a more positively charged species

B) Second electron has greater mass

C) Nucleus loses a proton

D) Atom becomes larger

✅ Answer

A

Q. Why is the second ionization enthalpy greater than the first ionization enthalpy?

Solution:

After the first electron is removed, the atom becomes a positively charged ion. The remaining electrons are attracted more strongly by the nucleus. Hence, greater energy is required to remove the second electron.

IE₂ > IE₁

Q. Explain why successive ionization enthalpies increase.

Solution:

1. Removal of the first electron converts the atom into a positive ion.
2. The remaining electrons are held more strongly by the nucleus.
3. Electron-electron repulsion also decreases.
4. Therefore, more energy is required for each successive electron.

IE₁ < IE₂ < IE₃ < IE₄ ...

Q. Differentiate between first and second ionization enthalpy.

First Ionization Enthalpy	Second Ionization Enthalpy
X(g) → X⁺(g) + e⁻	X⁺(g) → X²⁺(g) + e⁻
Electron removed from neutral atom	Electron removed from positive ion
Lower energy	Higher energy
IE₁	IE₂ > IE₁

The second electron is more strongly attracted by the nucleus, so its removal requires more energy.

Q. Explain why successive ionization enthalpies of an element increase. Illustrate with sodium.

Solution:

Sodium has electronic configuration 2,8,1. Its first electron is easily removed from the outermost shell:

Na → Na⁺ + e⁻

After this, Na⁺ has the stable configuration 2,8. The second electron must be removed from an inner shell and is therefore held much more strongly.

Na⁺ → Na²⁺ + e⁻

Thus, the second ionization enthalpy is much greater than the first. Similarly, successive ionization enthalpies generally increase.

IE₁ < IE₂ < IE₃ < ...

Q. Why is the second ionization enthalpy always higher than the first? Explain with a suitable example.

Detailed Solution

Ionization enthalpy is the energy required to remove an electron from an isolated gaseous species. The first ionization involves removal of an electron from a neutral atom, whereas the second ionization involves removal of an electron from a positively charged ion.

IE₁: X(g) → X⁺(g) + e⁻

IE₂: X⁺(g) → X²⁺(g) + e⁻

Reasons for IE₂ > IE₁

1. The species after first ionization is positively charged.
2. The number of electrons decreases, while the number of protons remains unchanged.
3. Effective nuclear attraction on the remaining electrons increases.
4. Electron-electron repulsion decreases.
5. The remaining electron is therefore more strongly bound.
6. Consequently, more energy is needed for its removal.

Example: Sodium

Na = 2,8,1
Na⁺ = 2,8

The first electron is removed from the third shell. After its removal, Na⁺ acquires a stable noble-gas configuration. The second electron has to be removed from the inner second shell, which is much more strongly held.

Therefore:

IE₂ > IE₁

and generally,

IE₁ < IE₂ < IE₃ < IE₄ ...

🎯 Quick Revision

IE₁: X → X⁺ + e⁻
IE₂: X⁺ → X²⁺ + e⁻
After IE₁: Positive ion is formed
Nuclear attraction: Increases for remaining electrons
Electron-electron repulsion: Decreases

IE₁ < IE₂ < IE₃ < IE₄ ...