Question

The electronic configuration of elements \(\mathrm{A}, \mathrm{B}\) and \(\mathrm{C}\) are \([\mathrm{He}] 2 \mathrm{~s}^{1},[\mathrm{Ne}] 3 \mathrm{~s}^{1}\) and \([\mathrm{Ar}] 4 \mathrm{~s}^{1}\) respectively. Which one of the following order is correct for the first ionization potentials (in \(\mathrm{kJ} \mathrm{mol}^{-}\) ) of \(\mathrm{A}, \mathrm{B}\) and \(\mathrm{C}\) ? (a) \(\mathrm{A}>\mathrm{B}>\mathrm{C}\) (b) \(C>B>A\) (c) \(\mathrm{B}>\mathrm{C}>\mathrm{A}\) (d) \(C>A>B\)

Short answer

The correct order is option (a) \(\mathrm{A} > \mathrm{B} > \mathrm{C}\).

Step-by-step solution

Identify Elements A, B, and C

Based on their electronic configurations, we identify the elements:- Element A: \([\text{He}] 2s^1\) is Lithium (Li).- Element B: \([\text{Ne}] 3s^1\) is Sodium (Na).- Element C: \([\text{Ar}] 4s^1\) is Potassium (K).

Understand Ionization Potential

Ionization potential is the energy required to remove an electron from an atom. Elements with electrons closer to the nucleus hold them tighter, resulting in higher ionization energies.

Determine General Periodic Trends

Ionization energy generally decreases down a group in the periodic table. This is because additional electron shells increase distance from the nucleus, reducing the electrostatic attraction for outer electrons.

Apply Trends to Elements A, B, and C

Since Elements A, B, and C (Li, Na, and K) belong to the same group in the periodic table, their ionization energy follows the trend: \(\text{Li} > \text{Na} > \text{K}\).

Analyze Ionization Energy Order

With the order of ionization energies understood as Lithium (A) > Sodium (B) > Potassium (C), we compare it to the given options.

Identify Correct Option

The correct order of first ionization potentials from highest to lowest is A > B > C. Therefore, the correct answer is option (a) \(\text{A} > \text{B} > \text{C}\).

Key concepts

Electronic Configuration

Electronic configuration is the detailed arrangement of electrons in an atom's energy levels. For each element, electrons fill from lower to higher energy levels, following a specific order based on the principles of quantum mechanics. This configuration not only helps identify the element but also predicts its chemical behavior.

• For Lithium ( ext{Li}), the configuration is \([ ext{He}] 2s^1\), meaning it has two electrons filling the first shell (similar to Helium) and one electron in the second main energy level.
• Sodium ( ext{Na}) has the configuration \([ ext{Ne}] 3s^1\), with two electrons in the first level, eight in the second (like Neon), and one in the third.
• Potassium ( ext{K}) is configured as \([ ext{Ar}] 4s^1\), which indicates that it has the electron configuration of Argon in the first three shells and an additional electron in the fourth.

Understanding electronic configurations allows us to determine placement on the periodic table and predict how an element will react chemically.

Periodic Trends

Trends across the periodic table provide insight into various atomic properties, such as size and ionization energy. Periodic trends are patterns that occur across different periods (rows) or groups (columns).

• Across a period from left to right, ionization energy tends to increase. This is because as you move across a period, protons are added to the nucleus increasing the positive charge which pulls electrons closer.
• Down a group, the trend is the opposite. Ionization energy generally decreases because adding electron shells increases the distance between the nucleus and the outermost electrons.
• This increased distance reduces the nuclear attraction felt by electrons, making them easier to remove as you move down a group.

Understanding these trends helps predict behavior and reactivity of elements, crucial for solving problems like comparing ionization energies.

Ionization Energy

Ionization energy is the measure of the energy required to remove an electron from an atom in the gaseous state. It's a crucial concept that explains an element's reactivity and stability.
Ionization energy is influenced by several factors:

• Nuclear Charge: A higher positive charge in the nucleus attracts electrons more strongly, increasing the ionization energy.
• Shielding Effect: Inner electrons can shield the outer electrons from the attractive force of the nucleus. With more shells, this effect increases, reducing ionization energy.
• Atomic Size: Larger atoms have lesser nuclear pull on valence electrons due to greater distances, leading to lower ionization values.

When comparing ionization energies within the same group, lithium, sodium, and potassium, for example, higher ionization energy corresponds to a smaller atomic size or fewer shielding shells.

Group Trends in Periodic Table

Group trends refer to how elements' properties change as you move up or down a column in the periodic table. Each group typically has elements with similar properties due to a similar valence electron configuration.

• As you move down a group, several trends emerge:
• Atomic Size Increases: New electron shells are added as you go down, expanding the size of the atoms.
• Ionization Energy Decreases: The increased distance between valence electrons and the nucleus reduces the energy required to remove an outer electron.
• Electronegativity Decreases: The ability of an atom to attract electrons decreases down a group due to increased atomic size.

These trends help explain the chemical behavior of elements in a group, such as why lithium has a higher ionization energy than sodium, and sodium is higher than potassium in terms of ionization energy. They also illustrate why option (a) is correct for the original exercise.