Anand Classes presents a comprehensive explanation of the Factors Governing Ionization Enthalpy — the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state. Ionization enthalpy is a key concept in chemistry that reflects how strongly an atom holds its electrons. Its magnitude is influenced by several factors, including the size of the atom, nuclear charge, screening effect, penetration power of electrons, and the stability of the atom’s electronic configuration. Understanding these factors is essential for mastering periodic trends and predicting the chemical reactivity of elements.
Table of Contents
Factors Affecting Ionization Enthalpy (Ionization Energy)
Ionization enthalpy is the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
The magnitude of ionization enthalpy depends on the following five major factors:
1. Size of the Atom
Trend:
- As atomic size increases, ionization enthalpy decreases.
- As atomic size decreases, ionization enthalpy increases.
Reason:
- The attractive force between the nucleus and the electron is given by Coulomb’s Law: $$F \propto \frac{Z_{\text{eff}}}{r^2}$$ where:
- Zeff = Effective nuclear charge
- r = Distance between the nucleus and the outermost electron
- Larger atomic radius → Outer electrons are farther from the nucleus → Weaker attraction → Easier to remove → Lower ionization enthalpy.
- Smaller atomic radius → Electrons are closer to the nucleus → Stronger attraction → Harder to remove → Higher ionization enthalpy.
Hence, as the atomic size increases, the outermost electrons are farther from the nucleus and experience less attraction. Consequently, these electrons are less tightly held and can be removed more easily, leading to lower ionization enthalpy.
Example:
- Compare Li and Cs:
- Li (small size) → High ionization enthalpy (~520 kJ/mol)
- Cs (large size) → Low ionization enthalpy (~375 kJ/mol)
2. Nuclear Charge (Z)
Trend: Ionization enthalpy increases with an increase in nuclear charge.
- Higher nuclear charge → Greater attraction → Higher ionization enthalpy.
- Lower nuclear charge → Lesser attraction → Lower ionization enthalpy.
Reason:
- The attractive force between the nucleus and an electron is directly proportional to the nuclear charge : F ∝ Z
- For the same shell and same screening effect, if Z increases, it becomes more difficult to remove the electron.
Therefore a greater nuclear charge makes electron removal more difficult, resulting in higher ionization enthalpy.
Example:
- Compare Boron (Z = 5) and Fluorine (Z = 9):
- Fluorine has a much higher nuclear charge → Its outer electrons are pulled more strongly → Higher ionization enthalpy.
3. Screening Effect (Shielding Effect)
Definition:
- In multi-electron atoms, inner-shell electrons repel outer-shell electrons, reducing the full nuclear charge felt by them.
- The effective nuclear charge is given by: Zeff = Z − S where S is the screening constant.
Effect on Ionization Enthalpy:
- High screening effect → Outer electrons feel less attraction → Easier to remove → Lower ionization enthalpy.
- Low screening effect → Outer electrons feel more attraction → Harder to remove → Higher ionization enthalpy.
Order of screening power of orbitals: s > p > d > f
This means electrons in s-orbitals are less shielded compared to those in d or f orbitals.
Example:
- In alkali metals, the outermost s-electron is shielded by all the inner-core electrons, making it easy to remove, hence low ionization enthalpy.
4. Penetration Effect of Electrons
Definition:
- Penetration refers to the extent to which an electron’s probability density is found close to the nucleus.
Order of penetration in the same shell: s > p > d > f
Reason:
- s-electrons have the highest probability of being near the nucleus, hence they experience the strongest attraction and are harder to remove.
- f-electrons are the farthest and most shielded, hence easiest to remove.
Effect on Ionization Enthalpy:
- Greater penetration → Higher attraction → Higher ionization enthalpy.
- Lesser penetration → Lower attraction → Lower ionization enthalpy.
Example:
- Removing a 2s electron requires more energy than removing a 2p electron from the same atom.
5. Electronic Configuration
Ionization enthalpy depends on the stability of the atom’s electronic configuration.
- Atoms with half-filled or fully filled subshells are exceptionally stable due to exchange energy and symmetry and making electron removal more difficult.
- These stable configurations have unusually high ionization enthalpies.
Examples:
- Noble gases: Noble gases have the stable configuration ns2np6 in each period and thus have very high ionization enthalpies.
- Be and Mg:
- Be: 1s22s2
- Mg: 1s22s22p63s2
- Be (1s² 2s²) and Mg (1s² 2s² 2p⁶ 3s²) have completely filled orbitals, leading to high ionization enthalpies.
- Completely filled s-subshell → High ionization enthalpy.
- N and P:
- N: 1s22s22px12py12pz1
- P: 1s22s22p63s23p63px13py13pz1
- N (1s² 2s² 2pₓ¹ 2pᵧ¹ 2p_z¹) and P (1s² 2s² 2p⁶ 3s² 3pₓ¹ 3pᵧ¹ 3p_z¹) have exactly half-filled p-subshells, which are stable and require large amounts of energy for electron removal.
- Exactly half-filled p-subshells → Extra stability → High ionization enthalpy.
Conclusion:
The more stable the electronic configuration of an atom, the greater its ionization enthalpy.
Summary Table: Factors Affecting Ionization Enthalpy
| Factor | Effect on Ionization Enthalpy | Reason |
|---|---|---|
| Atomic size ↑ | ↓ | Greater distance → Less attraction |
| Nuclear charge ↑ | ↑ | Stronger attraction to electrons |
| Screening effect ↑ | ↓ | Outer electrons feel less nuclear pull |
| Penetration power ↑ | ↑ | Electrons closer to nucleus → Held more tightly |
| Stability of configuration ↑ | ↑ | Half-filled/fully filled shells are more stable |
FAQs – Factors Governing Ionization Enthalpy
Q1. What is ionization enthalpy?
Answer:
Ionization enthalpy is the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
It is usually expressed in kJ/mol or electron volts (eV).
Example:
For hydrogen:
H (g) → H+(g) + e−
Ionization enthalpy = 1312 kJ/mol.
Q2. How does atomic size affect ionization enthalpy?
Answer:
- Larger atoms → Outer electrons are farther from the nucleus → Less attraction → Easier to remove → Lower ionization enthalpy.
- Smaller atoms → Electrons are closer → Strong attraction → Higher ionization enthalpy.
Reason: Based on Coulomb’s law, attraction decreases with distance: $$F \propto \frac{1}{r^2}$$
Example:
- Li : 520 kJ/mol
- Cs : 375 kJ/mol
Q3. Why does an increase in nuclear charge increase ionization enthalpy?
Answer:
- Higher nuclear charge (Z) → Stronger pull on electrons → Harder to remove → Higher ionization enthalpy.
- If the number of inner-shell electrons remains constant, increasing protons means greater positive pull.
Example:
- Boron (Z = 5) has a lower ionization enthalpy than Fluorine (Z = 9) because Fluorine’s nucleus pulls electrons much more strongly.
Q4. What is the screening (shielding) effect and how does it affect ionization enthalpy?
Answer:
- In multi-electron atoms, inner electrons repel outer electrons, reducing the full nuclear charge felt by the outer electrons.
- The effective nuclear charge is: Zeff = Z − S
- Where Z = Nuclear charge and S = Screening constant.
Effect:
- High screening → Less nuclear pull → Easier to remove electrons → Lower ionization enthalpy.
- Low screening → More nuclear pull → Higher ionization enthalpy.
- Order of shielding power : s > p > d > f
Q5. What is penetration effect and why does it matter?
Answer:
- Penetration effect refers to how close an electron can get to the nucleus on average.
- Electrons in s-orbitals have the greatest penetration, followed by p, d, and f.
Order of penetration power : s > p > d > f
Effect:
- Higher penetration → Electron experiences stronger attraction → Harder to remove → Higher ionization enthalpy.
Example:
- In the same shell (n=2), a 2s electron requires more energy to remove than a 2p electron.
Q6. How does electronic configuration influence ionization enthalpy?
Answer:
- Certain configurations are exceptionally stable, especially:
- Fully filled subshells
- Half-filled subshells
- Noble gas configurations
Reason: Stability arises from symmetry and exchange energy. Removing an electron from such configurations requires much more energy.
Examples:
- Noble gases: ns2np6 → Very high ionization enthalpy.
- Be (1s² 2s²) and Mg (1s² 2s² 2p⁶ 3s²) → Fully filled s-subshells → High ionization enthalpy.
- N(1s22s22px12py12pz1) → Exactly half-filled p-subshell → Extra stability → High ionization enthalpy.
Q7. Why does ionization enthalpy not follow a perfectly smooth trend in the periodic table?
Answer:
While ionization enthalpy generally increases across a period and decreases down a group, there are exceptions due to:
- Subshell stability (half-filled and fully filled)
- Changes in penetration power
- Differences in screening effect
Example:
- Be (1s22s2) has higher ionization enthalpy than B (1s22s2 2px1) even though B is to the right, because removing a p-electron from B is easier than removing a fully filled s-electron from Be.
Q8. What is the difference between first and successive ionization enthalpies?
Answer:
- First Ionization Enthalpy (IE1): Energy needed to remove the first electron.
- Second Ionization Enthalpy (IE2): Energy needed to remove the second electron from a positively charged ion.
Trend:
- Successive ionization enthalpies are much higher because removing an electron from a positively charged ion is more difficult due to increased nuclear pull.
Example (Oxygen):
- IE1 = 1314 kJ/mol
- IE2 = 3388 kJ/mol
Q9. Why do alkali metals have the lowest ionization enthalpy in their periods?
Answer:
- Large atomic size
- Low nuclear charge relative to size
- Only one loosely bound electron in the outermost s-orbital
- Strong screening effect from inner electrons
Example:
Cs has one of the lowest ionization enthalpies (~375 kJ/mol) among naturally occurring elements.
Q10. Why do noble gases have the highest ionization enthalpy?
Answer:
- Very small size (in their respective periods)
- Very high nuclear charge
- Extremely stable ns² np⁶ configuration
- No tendency to lose electrons
Example:
Helium has the highest ionization enthalpy of all elements (~2372 kJ/mol).
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