Question

A sudden large difference between the values of second and third ionization energies of elements would be associated with which of the following electronic configuration? (a) \(1 \mathrm{~s}^{2} 2 \mathrm{~s}^{2} 2 \mathrm{p}^{6} 3 \mathrm{~s}^{4}\) (b) \(1 \mathrm{~s}^{2} 2 \mathrm{~s}^{2} 2 \mathrm{p}^{6} 3 \mathrm{~s}^{2}\) (c) \(1 \mathrm{~s}^{2} 2 \mathrm{~s}^{2} 2 \mathrm{p}^{6} 3 \mathrm{~s}^{2} 3 \mathrm{p}^{1}\) (d) \(1 \mathrm{~s}^{2} 2 \mathrm{~s}^{2} 2 \mathrm{p}^{6} 3 \mathrm{~s}^{2} 3 \mathrm{p}^{2}\)

Short answer

The correct configuration is (c) \(1s^2 2s^2 2p^6 3s^2 3p^1\).

Step-by-step solution

Understand Ionization Energy

Ionization energy is the energy required to remove an electron from an atom or ion. The second ionization energy refers to the energy required to remove a second electron after the first has been removed, and the third ionization energy is for the third electron. A large increase in ionization energy usually indicates the removal of an electron from a full or stable electronic shell.

Analyze Electron Configurations

Compare each electronic configuration to see how many electrons need to be removed to reach a stable or filled shell configuration. Typically, elements achieve a stable configuration by having filled or half-filled subshells indicating a noble gas configuration.

Check Configuration (a)

Configuration (a) is \(1s^2 2s^2 2p^6 3s^4\). This is incorrect because the '3s' subshell can hold only 2 electrons. Thus, this configuration is not valid.

Check Configuration (b)

Configuration (b) is \(1s^2 2s^2 2p^6 3s^2\). This suggests a noble gas-like configuration after the removal of 2 electrons, which likely means no large jump between the second and third ionization energies.

Check Configuration (c)

Configuration (c) is \(1s^2 2s^2 2p^6 3s^2 3p^1\). Removing two electrons here would result in \(1s^2 2s^2 2p^6 3s^1\). Removing the third requires breaking a full subshell, causing a significant jump in energy.

Check Configuration (d)

Configuration (d) is \(1s^2 2s^2 2p^6 3s^2 3p^2\). Removing three electrons achieves \(1s^2 2s^2 2p^6 3s^2\), already relatively stable, but not indicating a large difference between ionization energies.

Conclusion

Among the given options, configuration (c) has the greatest potential for a sudden large difference between the second and third ionization energies, as the third electron removal disrupts a stable noble gas core.

Key concepts

Electronic Configuration

The electronic configuration of an atom refers to the arrangement of electrons in its various orbitals. Every element has a unique electronic configuration that determines its chemical behavior.
Each electron occupies the lowest energy orbital available, following the "Aufbau principle." This principle explains the order in which orbitals are filled.

• s-orbitals can hold a maximum of 2 electrons.
• p-orbitals can hold up to 6 electrons.
• d-orbitals can accommodate up to 10 electrons.

Understanding electronic configuration helps predict an atom’s reactivity, ionization potential, and the formation of chemical bonds, as described in periodic trends.

Stable Electronic Shell

A stable electronic shell is achieved when an atom has a filled or half-filled subshell.

• Filled shells are energetically favorable for the atom.
• Atoms typically strive to fill their outermost shell, achieving maximum stability.
• Unstable configurations often result in chemical reactivity.

The concept of a stable electronic shell is critical to understanding chemical bonding and reactions. Atoms will often gain, lose, or share electrons to achieve this stable state, reminiscent of noble gases, through ionic or covalent bonding.

Noble Gas Configuration

Noble gas configuration is a term used to describe an atom that has the same electron arrangement as a noble gas, which is known for its stability. The noble gases have complete outer shells, making them remarkably unreactive. Their electronic configuration serves as a model for other elements that seek to achieve similar stability.
Natural elements do this in various ways:

• Metals tend to lose electrons to match the prior noble gas configuration.
• Nonmetals often gain electrons to reach the next noble gas configuration.
• Ionization energy increases significantly after the electronic configuration of a noble gas is mimicked.

Atoms strive for this stable form due to its low energy state, often leading to chemical reactions.

Periodic Trends

Periodic trends are patterns in the periodic table that show how different properties change across groups and periods. Recognizing these patterns is crucial for predicting elements’ behavior in chemical settings.
For instance, ionization energy is one such periodic trend:

• Ionization energy generally increases across a period due to increased nuclear charge attracting electrons closer.
• It decreases down a group as the outer electrons are further from the nucleus and thus less tightly bound.
• Elements with filled or half-filled electron shells exhibit notably different (often higher) ionization energies after specific subshell configurations are disrupted.

Understanding these trends provides insight into reactivity, electron affinity, and electronegativity, aiding in the mastery of chemical concepts.