6 Most FAQ’s of Structure of Atom Chapter in Class 10th Physical Science (TS/AP)

8 Marks

LAQ-1 : Explain the significance of Quantum numbers in predicting the positions of an electron in an atom.

Introduction

Understanding the position of an electron within an atom is fundamental for grasping various concepts in physics and chemistry. Quantum numbers play a pivotal role in this understanding, providing detailed insights into the electron’s position, energy level, and orientation within an atom. These numbers are integral to the study of atomic structure and quantum mechanics, and their significance cannot be overstated. This guide aims to explain the role and importance of each type of quantum number in simple, straightforward terms.

Significance of Quantum Numbers in Predicting Electron Positions in an Atom

1. Principal Quantum Number (n)
   • Represented as n.
   • Indicates the main energy level or shell in an atom.
2. Angular-Momentum Quantum Number (l)
   • Represented as l.
   • Shows the shape of the sub-shell within a main shell.
3. Magnetic Quantum Number (ml​)
   • Represented as ml​.
   • Outlines the orientation of an orbital in space.
4. Spin Quantum Number (ms​)
   • Represented as ms​.
   • Details the spin orientation of electrons.

Summary

In summary, quantum numbers are integral in determining and predicting the precise position, orientation, and types of energy levels that electrons occupy within an atom. The principal quantum number (n) reveals the main energy shell, the angular-momentum quantum number (l) shows the shape of the sub-shell, the magnetic quantum number (ml​) details the orbital orientation, and the spin quantum number (ms​) indicates the electron’s spin direction. A thorough understanding of these quantum numbers not only deepens the knowledge of atomic theory but also lays the groundwork for advanced studies in quantum mechanics and chemical bonding.

---

LAQ-2 : Draw the shape of s, p and d orbitals.

Understanding the Shapes of s, p, and d Orbitals in Atomic Physics

S Orbital

1. Shape: Spherical
2. Details: The s orbital is a sphere with the nucleus at its center. In the hydrogen atom, the 1s electron is found within a spherical space surrounding the nucleus.

P Orbital

1. Shape: Dumbbell or double-lobed
2. Details: The p orbital consists of two lobes on either side of the nucleus. Each of the three p orbitals lies along one of the three Cartesian axes (x, y, and z), known as px​, py​, and pz​.

D Orbital

1. Shape: Cloverleaf and donut-shaped ring (toroidal)
2. Details: Four of the five d orbitals have a cloverleaf shape, while the fifth resembles a dumbbell with a donut-shaped ring around the middle. The d orbitals have more complex shapes with four lobes, except for one which has two lobes and a toroidal ring.

Summary

Understanding the different shapes of s, p, and d orbitals is fundamental for visualizing electron locations and behaviors in atoms. The s orbital is spherical, indicating a centered zone for electron probability. P orbitals, with their dumbbell shape, delineate a bi-lobed region. D orbitals, with their more complex cloverleaf and toroidal configurations, provide intricate patterns for electron distribution. These shapes are crucial for scientists and students in visualizing and comprehending electron placements and movements within the atom.

---

LAQ-3 : Write postulates and limitations of Bohr Hydrogen atomic model. (OR) Write the main features of Bohr’s model of hydrogen atom and mention any two limitations.

Introduction

The Bohr Model of the Hydrogen Atom The Bohr model of the hydrogen atom represents a significant milestone in the development of atomic theory. Proposed by Niels Bohr in 1913, this model was pivotal in explaining the stability of electron orbits and the emission spectrum of hydrogen. It bridged classical and quantum physics, providing a foundation for the quantum mechanical models that followed. Despite its historical importance and initial success, the Bohr model had its limitations, which became apparent with further scientific advancements.

Postulates of Bohr Hydrogen Atomic Model

1. Stationary Orbits:
   Electrons revolve in fixed orbits without losing energy.
2. Quantized Energy Levels:
   Electrons in an atom have specific energy levels.

Limitations of Bohr Hydrogen Atomic Model

1. Unable to Explain Line Spectra:
   The model cannot account for the fine splitting of spectral lines.
2. Doesn’t Explain Zeeman and Stark Effect:
   Fails to explain the Zeeman and Stark effects.

Summary

The Bohr model, while a critical step in atomic theory, faced limitations, notably in explaining complex atomic phenomena. It laid the groundwork for the quantum mechanical models, highlighting the evolving nature of scientific understanding in atomic physics. This comprehensive view of the Bohr model offers valuable insights into both its contributions and shortcomings within the field of atomic theory.

---

LAQ-4 : How we are using the nl^x method in writing electronic configuration?

Introduction

Writing electronic configurations is a method to denote the arrangement of electrons in an atom’s atomic orbitals. The nlx method is a shorthand notation that makes this process more concise and easier to understand. It employs three variables: n, l, and x to depict the arrangement of electrons in an atom.

Understanding the nl^x Method

1. n – Principal Quantum Number:
   • Represents the main energy level.
   • Indicates which shell the electrons are in.
2. l – Angular Momentum Quantum Number:
   • Represents the sub-level or sub-shell.
   • Indicates the shape of the orbital (s, p, d, f).
3. x: Denotes the number of electrons in a specific orbital.

How to Use the nl^x Method

1. Example: Hydrogen Atom (H):
   • Atomic Number (Z) of Hydrogen: 1.
   • Number of Electrons: 1.
2. Steps:
   • The electron occupies the orbital of the first main shell, n=1 (K shell).
   • The sub-shell is s (since l is 0 in this case).
   • There is only one electron, so x=1.
3. Electronic Configuration:
   • The electronic configuration is written as 1s¹.
4. Set of Quantum Numbers for Hydrogen Atom:
   • n=1, l=0, ml​=0, ms​=+1/2.

Summary

The nl^x method is a straightforward and effective way to write electronic configurations and predict the quantum numbers of electrons in an atom. It provides a clear picture of electron arrangement and helps in understanding the properties and behavior of elements.

---

LAQ-5 : How orbital is different from Bohr’s orbit?

Introduction

In the study of atomic structure, the concepts of orbital and Bohr’s orbit are fundamental yet distinct. Understanding the differences between these two is crucial in grasping the advances in atomic theory from classical to modern quantum mechanics. This discussion aims to elucidate how an orbital differs from Bohr’s concept of an orbit.

Concept of Orbital

1. Definition: An orbital is a probabilistic region in space where there is a high likelihood of finding an electron.
2. Nature: Represents uncertain, probabilistic areas, unlike defined paths.
3. Shape and Size: Varies, including spherical (s), dumbbell-shaped (p), and more complex for d and f orbitals.

Bohr’s Orbit Concept

1. Definition: In Niels Bohr’s model, an orbit is a circular path where electrons revolve around the nucleus.
2. Nature: Features well-defined circular paths.
3. Fixed Energy Levels: Electrons have specific quantized energy levels.

Key Differences

1. Nature of Path: Orbitals are probability regions, while Bohr’s orbits are fixed paths.
2. Energy Emission: Movement between energy levels in orbitals, but fixed levels in Bohr’s orbits.
3. Acceptance: Orbitals align with modern quantum mechanics, whereas Bohr’s orbits are part of an older atomic model.

Summary

The distinction between an orbital and Bohr’s orbit reflects the evolution of atomic theory. Orbitals, with their probabilistic nature and varying shapes, align with the principles of quantum mechanics and provide a more accurate depiction of electron behavior than Bohr’s orbits, which depict electrons in fixed paths with quantized energy levels. This understanding is pivotal for students delving into the complexities of atomic structure and quantum theory, showcasing the shift from classical to modern interpretations in atomic physics.

---

LAQ-6 : How does Hund’s rule help in writing electronic configuration of an atom? Explain with a suitable example.

Introduction

Hund’s Rule is an essential principle in determining the electron configuration of atoms, crucial for understanding how electrons are arranged in an atom’s orbitals.

Hund’s Rule Explained

1. What is Hund’s Rule?
   • Electrons fill degenerate orbitals (those with the same energy level) singly before pairing up.
   • Electrons in the same subshell have the same spin when they are unpaired.
2. Why is it Important?
   • Ensures electrons occupy individual orbitals in a subshell before pairing, minimizing electron repulsion and lowering the energy of the atom.

Example: Carbon (Z=6)

1. Atomic Number and Configuration:
   • Atomic number of Carbon (C) is 6.
   • Basic electron configuration:
     $$1s^2 2s^2 2p^2$$
2. Applying Hund’s Rule:
   • In the 2p subshell, there are three orbitals (p_x​,p_y​,p_z​).
   • According to Hund’s Rule, each orbital is singly occupied before any pairing occurs.
   • For Carbon, two of the three 2p orbitals will each have one electron $$(p_x^1 p_y^1​)$$.
3. Result:
   • Electrons in 2p orbitals will have the same spin, showing as parallel spins.
   • Electron configuration with Hund’s rule:
     $$1s^2 2s^2 2p_x^1 2p_y^1$$

Summary

Hund’s Rule plays a crucial role in determining the arrangement of electrons in an atom, influencing the atom’s chemical properties and behavior. The example of Carbon illustrates how electrons are distributed in degenerate orbitals to minimize repulsion and energy, ensuring the stability of the atom. Understanding and applying Hund’s Rule is fundamental in writing the electron configuration of atoms.