Problem 79 - A Complete Resource Book in Chemistry for JEE Main

Question

Which one of the following arrangements represents the correct order of electron gain enthalpy (with negative sign) of the given atomic species? (a) \(\mathrm{S}<\mathrm{O}<\mathrm{Cl}<\mathrm{F}\) (b) \(\mathrm{Cl}<\mathrm{F}<\mathrm{S}<\mathrm{O}\) (c) \(\mathrm{F}<\mathrm{Cl}<\mathrm{O}<\mathrm{S}\) (d) \(\mathrm{O}<\mathrm{S}<\mathrm{F}<\mathrm{Cl}\)

Step-by-Step Solution

Answer

The correct order is (d) \( \mathrm{O

Step 1: Understanding Electron Gain Enthalpy

Electron gain enthalpy is the energy change when an electron is added to a gaseous atom. It is usually negative if energy is released during the addition. A more negative electron gain enthalpy means the atom has a stronger tendency to gain an electron.

Step 2: Electron Gain Enthalpy Trends in the Periodic Table

In general, electron gain enthalpy becomes more negative across a period from left to right, as the atoms have higher effective nuclear charge and a greater attraction for additional electrons. In a group, electron gain enthalpy becomes less negative downwards due to increased atomic size and electron shielding.

Step 3: Comparing Electron Gain Enthalpy of Given Elements

Among the elements given, Fluorine (F) and Chlorine (Cl) are in Group 17, while Oxygen (O) and Sulfur (S) are in Group 16. Chlorine has the most negative electron gain enthalpy due to optimal balance between nuclear charge and atomic size, even more negative than Fluorine. Oxygen and Sulfur, being in the same group, have less negative values compared to Group 17 elements.

Step 4: Arranging the Elements

Comparing individual electron gain enthalpy values: Chlorine generally has the most negative value, followed by Fluorine. Oxygen and Sulfur have less negative values, with Sulfur being less negative than Oxygen. Therefore, the order from most to least negative electron gain enthalpy is: Cl > F > O > S.

Key Concepts

Periodic Table Trends

Periodic table trends help us understand how different elements will react or interact under certain conditions. By recognizing these trends, we can predict the behaviors of elements, such as their electron gain enthalpy—the energy change that occurs when an atom gains an electron.
One crucial trend to remember is that as you move from left to right across a period, the electron gain enthalpy generally becomes more negative. This pattern is due to the increasing effective nuclear charge. Down a group, however, the electron gain enthalpy often becomes less negative because of increasing atomic size. The attraction between the nucleus and the incoming electron decreases due to increased distance and electron shielding.

• More negative values indicate a stronger attraction for gaining electrons.
• Less negative or positive values indicate weaker attraction.

This understanding helps us compare elements like Chlorine and Fluorine in Group 17, and Oxygen and Sulfur in Group 16, where Chlorine stands out for its high affinity for electrons.

Atomic Structure

The atomic structure of an element plays a critical role in determining its electron gain enthalpy. An atom consists of a nucleus containing protons and neutrons, surrounded by electrons in various energy levels or shells. The arrangement of these electrons defines how an atom interacts with others—particularly how it gains or loses electrons.
For example, in Fluorine, the high effective nuclear charge, with electrons held tightly, leads to a more negative electron gain enthalpy as compared to Oxygen or Sulfur. Moreover, in smaller atoms with a lesser number of electron shells, the added electron experiences a stronger pull from the nucleus.

• The lesser the number of shells, the stronger the pull of the electron toward the nucleus.
• Electrons in outer shells experience a weaker pull due to increased distance from the nucleus and electron shielding.

Understanding atomic structure details helps explain why elements gain or lose electrons in such specific manners.

Nuclear Charge

Nuclear charge refers to the total positive charge of the nucleus, which comes from the number of protons present. A higher nuclear charge indicates a stronger pull on electrons, influencing many of an atom’s properties, including its electron gain enthalpy.
As electrons are added to an atom, they are drawn closer to the nucleus if the nuclear charge is high. This close pulling makes it easier for an atom to gain additional electrons, reflecting in a more negative electron gain enthalpy. However, a delicate balance is necessary because too high of a nuclear charge can tightly bind electrons, potentially making electron addition less favorable.

• In elements like Chlorine, optimal balance between nuclear charge and atomic size creates favorable electron gain enthalpy.
• Comparatively, Oxygen and Sulfur, with less optimal balances, have less negative electron gain enthalpies.

This understanding clarifies why Chlorine, despite being down a group compared to Fluorine, boasts a more negative electron gain enthalpy.

Electron Shielding

Electron shielding impacts how strongly outer electrons feel the nucleus's pull. As the number of electron shells, or energy levels, increases, inner electrons can "shield" the outer electrons from the nucleus, reducing the effective nuclear charge felt by outermost electrons.
This reduced pull makes it less favorable for an atom to gain additional electrons, thereby resulting in a less negative electron gain enthalpy. In elements that have more electron shells, such as down a group, this becomes crucial.

• For example, while Fluorine is highly electronegative, Chlorine with its extra shell surprisingly has a more negative electron gain enthalpy.
• Higher periods elements like Sulfur face more electron shielding than smaller periods like Oxygen.

Understanding electron shielding is vital in predicting electron gain tendencies and resolving cases where atomic size and charge balance might seem ambiguous.

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Key Concepts

Periodic Table Trends

Periodic table trends help us understand how different elements will react or interact under certain conditions. By recognizing these trends, we can predict the behaviors of elements, such as their electron gain enthalpy—the energy change that occurs when an atom gains an electron.
One crucial trend to remember is that as you move from left to right across a period, the electron gain enthalpy generally becomes more negative. This pattern is due to the increasing effective nuclear charge. Down a group, however, the electron gain enthalpy often becomes less negative because of increasing atomic size. The attraction between the nucleus and the incoming electron decreases due to increased distance and electron shielding.

• More negative values indicate a stronger attraction for gaining electrons.
• Less negative or positive values indicate weaker attraction.

This understanding helps us compare elements like Chlorine and Fluorine in Group 17, and Oxygen and Sulfur in Group 16, where Chlorine stands out for its high affinity for electrons.

Atomic Structure

The atomic structure of an element plays a critical role in determining its electron gain enthalpy. An atom consists of a nucleus containing protons and neutrons, surrounded by electrons in various energy levels or shells. The arrangement of these electrons defines how an atom interacts with others—particularly how it gains or loses electrons.
For example, in Fluorine, the high effective nuclear charge, with electrons held tightly, leads to a more negative electron gain enthalpy as compared to Oxygen or Sulfur. Moreover, in smaller atoms with a lesser number of electron shells, the added electron experiences a stronger pull from the nucleus.

• The lesser the number of shells, the stronger the pull of the electron toward the nucleus.
• Electrons in outer shells experience a weaker pull due to increased distance from the nucleus and electron shielding.

Understanding atomic structure details helps explain why elements gain or lose electrons in such specific manners.

Nuclear Charge

Nuclear charge refers to the total positive charge of the nucleus, which comes from the number of protons present. A higher nuclear charge indicates a stronger pull on electrons, influencing many of an atom’s properties, including its electron gain enthalpy.
As electrons are added to an atom, they are drawn closer to the nucleus if the nuclear charge is high. This close pulling makes it easier for an atom to gain additional electrons, reflecting in a more negative electron gain enthalpy. However, a delicate balance is necessary because too high of a nuclear charge can tightly bind electrons, potentially making electron addition less favorable.

• In elements like Chlorine, optimal balance between nuclear charge and atomic size creates favorable electron gain enthalpy.
• Comparatively, Oxygen and Sulfur, with less optimal balances, have less negative electron gain enthalpies.

This understanding clarifies why Chlorine, despite being down a group compared to Fluorine, boasts a more negative electron gain enthalpy.

Electron Shielding

Electron shielding impacts how strongly outer electrons feel the nucleus's pull. As the number of electron shells, or energy levels, increases, inner electrons can "shield" the outer electrons from the nucleus, reducing the effective nuclear charge felt by outermost electrons.
This reduced pull makes it less favorable for an atom to gain additional electrons, thereby resulting in a less negative electron gain enthalpy. In elements that have more electron shells, such as down a group, this becomes crucial.

• For example, while Fluorine is highly electronegative, Chlorine with its extra shell surprisingly has a more negative electron gain enthalpy.
• Higher periods elements like Sulfur face more electron shielding than smaller periods like Oxygen.

Understanding electron shielding is vital in predicting electron gain tendencies and resolving cases where atomic size and charge balance might seem ambiguous.