GATE Life Sciences Atomic Structure and Periodicity Syllabus

Atomic Structure and Periodicity syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Atomic Structure and Periodicity in GATE Life Sciences, not a summary of it.

1. 1

   Quantum Theory and Atomic Models

   7 topics

   • Planck’s Quantum Theory
   • Wave-Particle Duality
   • Uncertainty Principle
   • Comparison between Bohr’s Model and Quantum Mechanical Model of Hydrogen Atom
   • Electronic Configuration of Atoms and Ions
   • Hund’s Rule
   • Pauli’s Exclusion Principle
2. 2

   Periodic Table and Periodic Properties

   4 topics

   • Ionization Energy
   • Electron Affinity
   • Electronegativity
   • Atomic Size

Atomic Structure and Periodicity flashcards for GATE Life Sciences

24 of 51 cards from the Atomic Structure and Periodicity deck — real questions with worked answers.

1. State Planck's quantum theory regarding the emission and absorption of radiation.

   Energy is emitted or absorbed not continuously but in discrete packets called quanta. The energy of each quantum is proportional to the frequency of radiation: $E = h\nu$.

2. Write the formula for the energy of one quantum (photon) of radiation and define each term.

   $E = h\nu = \dfrac{hc}{\lambda}$, where $h = 6.626 \times 10^{-34}\ \text{J s}$ (Planck's constant), $\nu$ is frequency, $c$ is the speed of light, and $\lambda$ is the wavelength.

3. What is the value of Planck's constant $h$ in SI units?

   $h = 6.626 \times 10^{-34}\ \text{J s}$.

4. According to Planck, how is the total energy radiated by a body related to the quantum energy?

   The total energy is an integral multiple of the quantum: $E = nh\nu$, where $n = 1, 2, 3, \dots$ is a positive integer.

5. How does the energy of a photon vary with wavelength?

   Energy is inversely proportional to wavelength: $E = \dfrac{hc}{\lambda}$, so shorter wavelength means higher energy.

6. What is wave-particle duality of matter, as proposed by de Broglie?

   Every moving particle of matter has an associated wave; matter exhibits both particle-like and wave-like properties.

7. Write the de Broglie equation for the wavelength of a moving particle.

   $\lambda = \dfrac{h}{p} = \dfrac{h}{mv}$, where $p$ is momentum, $m$ is mass, and $v$ is velocity.

8. Why is the wave nature of macroscopic objects not observable?

   Because their mass is large, the momentum $mv$ is large, making $\lambda = \dfrac{h}{mv}$ extremely small (negligible) and undetectable. Wave nature is significant only for very light particles like electrons.

9. Express the de Broglie wavelength of an electron in terms of its kinetic energy $KE$.

   $\lambda = \dfrac{h}{\sqrt{2m\,(KE)}}$, since $p = \sqrt{2m\,(KE)}$.

10. State Heisenberg's uncertainty principle.

    It is impossible to determine simultaneously, with absolute accuracy, both the exact position and the exact momentum of a microscopic particle such as an electron.

11. Write the mathematical expression of Heisenberg's uncertainty principle for position and momentum.

    $\Delta x \cdot \Delta p \geq \dfrac{h}{4\pi}$, where $\Delta x$ is uncertainty in position and $\Delta p$ is uncertainty in momentum.

12. Write Heisenberg's uncertainty relation in terms of velocity.

    $\Delta x \cdot \Delta v \geq \dfrac{h}{4\pi m}$, where $\Delta v$ is the uncertainty in velocity and $m$ is the mass of the particle.

13. What does the uncertainty principle imply about the concept of a fixed electron orbit?

    It rules out the existence of definite, fixed circular paths (Bohr orbits); we can only speak of the probability of finding an electron in a region of space (orbital).

14. In Bohr's model, what quantity is quantized for the electron's orbit?

    The angular momentum is quantized: $mvr = \dfrac{nh}{2\pi}$, where $n = 1, 2, 3, \dots$

15. How does Bohr's model describe the location of an electron compared to the quantum mechanical model?

    Bohr's model places the electron in fixed, well-defined circular orbits at exact distances. The quantum mechanical model describes the electron by a probability distribution (orbital) and gives no exact path, consistent with the uncertainty principle.

16. What is the key difference in dimensionality of motion between Bohr's orbits and quantum mechanical orbitals?

    Bohr orbits are two-dimensional fixed circular paths, whereas quantum mechanical orbitals are three-dimensional regions of probability around the nucleus.

17. How many quantum numbers describe an electron in Bohr's model versus the quantum mechanical model?

    Bohr's model uses only one quantum number $n$ (principal). The quantum mechanical model uses four quantum numbers: $n$, $l$, $m_l$, and $m_s$.

18. On what fundamental concept is the quantum mechanical model based that Bohr's model ignores?

    It is based on the wave nature of the electron (Schrödinger's wave equation) and the uncertainty principle, treating the electron as a standing wave rather than a particle in a fixed orbit.

19. Write the formula for the energy of the electron in the $n$-th orbit of a hydrogen atom (Bohr model).

    $E_n = -\dfrac{13.6}{n^{2}}\ \text{eV} = -\dfrac{2.18 \times 10^{-18}}{n^{2}}\ \text{J}$.

20. Write the expression for the radius of the $n$-th Bohr orbit of the hydrogen atom.

    $r_n = 0.529\,n^{2}\ \text{\AA} = 0.529 \times 10^{-10}\,n^{2}\ \text{m}$.

21. What is the physical meaning of the wavefunction $\psi$ and of $\psi^{2}$ in the quantum mechanical model?

    $\psi$ itself has no direct physical meaning, but $\psi^{2}$ gives the probability density of finding the electron at a given point in space.

22. Define an orbital in the quantum mechanical model.

    An orbital is a three-dimensional region around the nucleus where the probability of finding an electron is maximum (typically ~90–95%).

23. Write the ground-state electronic configuration of a neutral chlorine atom ($Z = 17$).

    $\ce{Cl}$: $1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{5}$.

24. State the Aufbau principle for filling atomic orbitals.

    Electrons occupy the available orbitals in order of increasing energy, filling the lowest-energy orbital first before moving to higher-energy ones.

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