Class 11 Notes

Class 11 Chemistry Notes Chapter 2 (Structure of atom) – Chemistry Part-I Book

Philoid

Philoid

— 11 min read

Chemistry Part-I
Alright class, let's get straight into Chapter 2: Structure of Atom. This is a fundamental chapter, and understanding it well is crucial not just for chemistry but also for many competitive government exams where basic science is tested. We'll cover the key concepts systematically.

Chapter 2: Structure of Atom - Detailed Notes for Government Exam Preparation

1. Introduction & Discovery of Sub-atomic Particles

  • Dalton's Atomic Theory (Limitations): While foundational, it couldn't explain the existence of sub-atomic particles, isotopes, isobars, or why atoms combine.
  • Discovery of Electron (J.J. Thomson, 1897):
    • Experiment: Cathode Ray Discharge Tube experiment.
    • Observations: Rays travel from cathode to anode, deflected by electric/magnetic fields, consist of negatively charged particles.
    • Properties of Cathode Rays (Electrons):
      • Negatively charged.
      • Travel in straight lines (in absence of fields).
      • Possess kinetic energy.
      • Produce fluorescence.
      • Produce X-rays when striking heavy metals.
      • Charge-to-Mass Ratio (e/m): Thomson determined e/m = 1.758820 × 10¹¹ C kg⁻¹. This ratio is constant regardless of the gas used or cathode material, proving electrons are fundamental constituents of all matter.
    • Charge on Electron (R.A. Millikan, 1909): Oil Drop experiment determined the charge (e) = -1.6022 × 10⁻¹⁹ C.
    • Mass of Electron (mₑ): Calculated using e and e/m ratio. mₑ = 9.1094 × 10⁻³¹ kg.
  • Discovery of Proton (E. Goldstein, 1886 / Rutherford):
    • Experiment: Modified Cathode Ray Tube (using perforated cathode) - observed Canal Rays or Anode Rays.
    • Properties: Positively charged particles originating from the anode (actually ionized gas atoms).
    • The e/m ratio depended on the gas used (highest for hydrogen).
    • The positively charged particle from hydrogen was named Proton (by Rutherford).
    • Charge (eₚ): +1.6022 × 10⁻¹⁹ C
    • Mass (mₚ): 1.6726 × 10⁻²⁷ kg (approx. 1837 times heavier than electron).
  • Discovery of Neutron (James Chadwick, 1932):
    • Experiment: Bombardment of Beryllium (Be) with alpha (α) particles.
    • Observation: Emission of neutral particles with mass slightly greater than protons.
    • Properties: No charge, Mass (mₙ) = 1.6749 × 10⁻²⁷ kg. Located in the nucleus. Essential for nuclear stability.

2. Atomic Models

  • Thomson's Model (1904) - "Plum Pudding" or "Watermelon" Model:
    • Postulate: Atom is a sphere of positive charge with electrons embedded in it, like plums in a pudding. The atom is electrically neutral.
    • Limitation: Failed to explain the results of Rutherford's alpha-scattering experiment.
  • Rutherford's Nuclear Model (1911) - Based on α-particle Scattering Experiment:
    • Experiment: Thin gold foil bombarded with fast-moving α-particles (He²⁺ ions).
    • Observations:
      1. Most α-particles passed straight through (atom is mostly empty space).
      2. Some were deflected by small angles (positive charge is concentrated).
      3. Very few (1 in 20,000) bounced back by 180° (positive charge and mass are concentrated in a very small volume - the nucleus).
    • Conclusions:
      1. Atom has a tiny, dense, positively charged nucleus at the center containing protons (and later found, neutrons).
      2. Electrons revolve around the nucleus in orbits, like planets around the sun.
      3. The size of the nucleus (~10⁻¹⁵ m) is extremely small compared to the atom (~10⁻¹⁰ m).
    • Limitations:
      1. Instability: According to classical electromagnetism, an accelerating charged particle (electron orbiting the nucleus) should radiate energy continuously, spiral into the nucleus, and collapse the atom. This doesn't happen.
      2. Inability to Explain Spectra: Couldn't explain the discrete line spectra observed for atoms (especially hydrogen).

3. Atomic Number (Z) and Mass Number (A)

  • Atomic Number (Z): Number of protons in the nucleus. Defines the element. For a neutral atom, Z = number of electrons.
  • Mass Number (A): Total number of protons (Z) + neutrons (N) in the nucleus. A = Z + N.
  • Representation: ᴬ<0xE2><0x82><0x99>X (e.g., ¹²₆C has Z=6, A=12, N=A-Z=6).
  • Isotopes: Atoms of the same element (same Z) but different mass numbers (different N). E.g., ¹H (Protium), ²H (Deuterium), ³H (Tritium). They have similar chemical properties (due to same electron count) but different physical properties (due to mass difference).
  • Isobars: Atoms of different elements (different Z) but the same mass number (A). E.g., ⁴⁰₁₈Ar and ⁴⁰₂₀Ca. They have different chemical properties.
  • Isotones: Atoms of different elements having the same number of neutrons (N). E.g., ³⁰₁₄Si and ³¹₁₅P (both have N=16).

4. Developments Leading to Bohr's Model

  • Dual Nature of Electromagnetic Radiation:
    • Wave Nature (Maxwell): EM radiation consists of oscillating electric and magnetic fields perpendicular to each other and to the direction of propagation. Characterized by:
      • Wavelength (λ): Distance between two consecutive crests or troughs. (Units: m, cm, nm, Å).
      • Frequency (ν): Number of waves passing a point per second. (Units: Hz or s⁻¹).
      • Velocity (c): Speed of light in vacuum = 3 × 10⁸ m/s. Relationship: c = νλ.
      • Wavenumber (ν̄): Number of wavelengths per unit length. ν̄ = 1/λ. (Units: m⁻¹, cm⁻¹).
    • Particle Nature (Planck's Quantum Theory):
      • Blackbody Radiation: An ideal body that emits and absorbs all frequencies. Planck explained its spectrum by proposing that energy is emitted or absorbed discontinuously in small packets called quanta (plural of quantum).
      • Planck's Equation: Energy of one quantum (photon for light) is proportional to its frequency: E = hν, where h is Planck's constant = 6.626 × 10⁻³⁴ J s.
      • Energy is quantized: E = nhν (n = integer).
      • Photoelectric Effect (Explained by Einstein using Planck's theory): Ejection of electrons from a metal surface when light of suitable frequency strikes it.
        • Observations: Threshold frequency (ν₀) required; KE of ejected electrons depends on ν, not intensity; number of electrons depends on intensity.
        • Einstein's Explanation: Light consists of photons (E=hν). One photon ejects one electron.
        • Equation: hν = hν₀ + KE or KE = hν - hν₀. hν₀ = W₀ (Work function - minimum energy required to eject electron).

5. Atomic Spectra

  • Emission Spectrum: Pattern of light emitted by a substance when energy is supplied (e.g., heating, electric discharge). Continuous (like rainbow) or Line spectrum (discrete lines).
  • Absorption Spectrum: Pattern obtained when light passes through a substance, showing dark lines corresponding to frequencies absorbed.
  • Line Spectrum of Hydrogen: When hydrogen gas is subjected to electric discharge, it emits radiation which, when passed through a prism, shows distinct lines. These lines form series:
    • Lyman Series: UV region (n₁=1, n₂=2, 3, ...)
    • Balmer Series: Visible region (n₁=2, n₂=3, 4, ...)
    • Paschen Series: Infrared region (n₁=3, n₂=4, 5, ...)
    • Brackett Series: Infrared region (n₁=4, n₂=5, 6, ...)
    • Pfund Series: Infrared region (n₁=5, n₂=6, 7, ...)
  • Rydberg Formula: Empirically described the wavenumbers (ν̄) of spectral lines:
    ν̄ = 1/λ = R<0xE2><0x82><0x9C> [1/n₁² - 1/n₂²] where n₂ > n₁.
    R<0xE2><0x82><0x9C> = Rydberg constant = 109677 cm⁻¹.

6. Bohr's Model for Hydrogen Atom (1913)

  • Postulates:
    1. Electrons revolve around the nucleus in specific circular paths called orbits or stationary states, without radiating energy.
    2. The angular momentum of an electron in an orbit is quantized: mvr = nh / 2π (n = 1, 2, 3... called principal quantum number).
    3. Energy is emitted or absorbed only when an electron jumps from one orbit to another. ΔE = E₂ - E₁ = hν.
  • Key Results for Hydrogen-like species (1 electron, nucleus charge Ze):
    • Radius of nᵗʰ orbit (rₙ): rₙ = (0.529 Å) * n²/Z
    • Energy of nᵗʰ orbit (Eₙ): Eₙ = (-13.6 eV/atom) * Z²/n² = (-2.18 × 10⁻¹⁸ J/atom) * Z²/n²
    • Successfully explained the hydrogen spectrum using the energy difference formula and Rydberg formula.
  • Limitations:
    1. Failed for multi-electron atoms.
    2. Couldn't explain the splitting of spectral lines in magnetic (Zeeman effect) or electric (Stark effect) fields.
    3. Violated Heisenberg's Uncertainty Principle.
    4. Couldn't explain the ability of atoms to form molecules (chemical bonding).
    5. Did not consider the wave nature of the electron.

7. Towards Quantum Mechanical Model

  • Dual Behaviour of Matter (Louis de Broglie, 1924): Proposed that matter (like electrons) also exhibits wave-particle duality.
    • de Broglie Wavelength (λ): λ = h / mv = h / p (where p is momentum).
    • Significant only for microscopic particles (like electrons) due to their small mass.
  • Heisenberg's Uncertainty Principle (1927): It is impossible to determine simultaneously and precisely both the position (Δx) and the momentum (Δp) of a microscopic particle.
    • Mathematical Form: Δx * Δp ≥ h / 4π or Δx * mΔvₓ ≥ h / 4π.
    • Implication: Rules out the concept of fixed orbits (definite path and momentum) as proposed by Bohr. Introduces the concept of probability.

8. Quantum Mechanical Model of Atom

  • Based on wave nature of electron and probability. Does not specify exact path.

  • Schrödinger Wave Equation (Ĥψ = Eψ): A complex mathematical equation describing the wave behaviour of electrons in atoms. (No need to solve it for general exams).

    • ψ (Psi): Wave function. Represents the amplitude of the electron wave. Contains all information about the electron.
    • ψ²: Probability density. Gives the probability of finding the electron in a small region of space around the nucleus.

  • Orbitals: The region in space around the nucleus where the probability of finding the electron (ψ²) is maximum (typically > 90%). Orbitals replace the concept of orbits.

  • Quantum Numbers: Address/description of an electron in an atom, derived from the solution of the Schrödinger equation.

    1. Principal Quantum Number (n):
      • Determines the main energy level or shell.
      • Values: n = 1, 2, 3, ... (K, L, M, ... shells).
      • Indicates the average distance from the nucleus (size of orbital). Higher n = higher energy, larger size.
      • Maximum electrons in a shell = 2n².
    2. Azimuthal or Angular Momentum Quantum Number (l):
      • Determines the subshell within a shell.
      • Defines the shape of the orbital.
      • Values: l = 0 to (n-1).
      • l = 0 → s subshell (spherical)
      • l = 1 → p subshell (dumbbell)
      • l = 2 → d subshell (double dumbbell/complex)
      • l = 3 → f subshell (complex)
      • Number of subshells in a shell = n.
    3. Magnetic Quantum Number (m<0xE2><0x82><0x9C>):
      • Determines the orientation of the orbital in space within a subshell.
      • Values: m<0xE2><0x82><0x9C> = -l to +l, including 0. (Total 2l+1 values).
      • For l=0 (s), m<0xE2><0x82><0x9C>=0 (1 s orbital).
      • For l=1 (p), m<0xE2><0x82><0x9C>=-1, 0, +1 (3 p orbitals: pₓ, p<0xE1><0xB5><0xA7>, p<0xE1><0xB5><0xB3>).
      • For l=2 (d), m<0xE2><0x82><0x9C>=-2, -1, 0, +1, +2 (5 d orbitals: d<0xE1><0xB5><0xAB><0xE1><0xB5><0xA7>, d<0xE1><0xB5><0xA7><0xE1><0xB5><0xB3>, d<0xE1><0xB5><0xAB><0xE1><0xB5><0xB3>, d<0xE1><0xB5><0xAB>²-<0xE1><0xB5><0xA7>², d<0xE1><0xB5><0xB3>²).
      • Number of orbitals in a subshell = 2l+1.
      • Total number of orbitals in a shell = n².
    4. Spin Quantum Number (m<0xE2><0x82><0x9B>):
      • Describes the intrinsic angular momentum (spin) of the electron. Electron behaves like it's spinning on its axis.
      • Values: +1/2 (spin up, ↑) or -1/2 (spin down, ↓).
      • An orbital can hold a maximum of two electrons, and they must have opposite spins.

  • Shapes of Atomic Orbitals:

    • s orbitals: Spherically symmetrical. Size increases with n. Have (n-1) radial nodes.
    • p orbitals: Dumbbell shaped. Have 3 orientations (pₓ, p<0xE1><0xB5><0xA7>, p<0xE1><0xB5><0xB3>) along the axes. Have one nodal plane passing through the nucleus.
    • d orbitals: Mostly double dumbbell shape (except d<0xE1><0xB5><0xB3>² which is dumbbell with a donut/ring in xy plane). Have 5 orientations. Have two nodal planes/surfaces.
    • Nodes: Regions where probability of finding the electron (ψ²) is zero.
      • Radial nodes = n - l - 1
      • Angular nodes = l
      • Total nodes = n - 1

  • Energy of Orbitals:

    • Hydrogen atom: Energy depends only on n (orbitals within a shell are degenerate, i.e., have same energy: 1s < 2s=2p < 3s=3p=3d ...).
    • Multi-electron atoms: Energy depends on both n and l due to electron-electron repulsions and screening effect. Order generally follows the (n+l) rule:
      • Lower (n+l) value = lower energy.
      • If (n+l) is same, lower n value = lower energy.
      • General order: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s ...

9. Filling of Orbitals in Atoms (Electronic Configuration)

  • Aufbau Principle (German: 'building up'): Electrons first occupy the lowest energy orbital available. Filling occurs in order of increasing (n+l) values.
  • Pauli Exclusion Principle: No two electrons in an atom can have the same set of all four quantum numbers. This implies an orbital can hold a maximum of two electrons, and they must have opposite spins.
  • Hund's Rule of Maximum Multiplicity: Pairing of electrons in the orbitals belonging to the same subshell (degenerate orbitals) does not take place until each orbital belonging to that subshell has got one electron each (i.e., is singly occupied). The singly occupied orbitals must have parallel spins (same m<0xE2><0x82><0x9B> value).
  • Electronic Configuration: Distribution of electrons into various orbitals. E.g., Nitrogen (Z=7): 1s² 2s² 2p³ (or 1s² 2s² 2pₓ¹ 2p<0xE1><0xB5><0xA7>¹ 2p<0xE1><0xB5><0xB3>¹).
  • Stability of Completely Filled and Half-Filled Subshells: Atoms having completely filled (s², p⁶, d¹⁰, f¹⁴) or exactly half-filled (s¹, p³, d⁵, f⁷) subshells are relatively more stable due to:
    1. Symmetrical Distribution: Leads to greater stability.
    2. Exchange Energy: More exchanges possible when degenerate orbitals are similarly filled (parallel spins), releasing more energy and increasing stability.
  • Exceptions to Aufbau Principle: Chromium (Cr, Z=24) and Copper (Cu, Z=29).
    • Cr Expected: [Ar] 4s² 3d⁴ Actual: [Ar] 4s¹ 3d⁵ (to achieve stable half-filled d-subshell)
    • Cu Expected: [Ar] 4s² 3d⁹ Actual: [Ar] 4s¹ 3d¹⁰ (to achieve stable fully-filled d-subshell)

Multiple Choice Questions (MCQs)

  1. The charge to mass ratio (e/m) for cathode rays was found to be constant irrespective of the gas used in the discharge tube. This observation proved that:
    a) Atoms are electrically neutral.
    b) Electrons are fundamental particles present in all atoms.
    c) Protons exist within the nucleus.
    d) The mass of the atom is concentrated in the nucleus.

  2. Which observation from Rutherford's α-scattering experiment indicated that most of the space inside an atom is empty?
    a) Most α-particles passed straight through the gold foil.
    b) Some α-particles were deflected by small angles.
    c) A very small fraction of α-particles bounced back.
    d) α-particles caused fluorescence on the screen.

  3. According to Bohr's model, the angular momentum of an electron in the 3rd orbit is:
    a) h / 2π
    b) 2h / 2π
    c) 3h / 2π
    d) 9h / 2π

  4. The energy of an electron in the nᵗʰ Bohr orbit of a hydrogen atom is given by Eₙ = -13.6 / n² eV. What is the energy required to excite an electron from n=1 to n=2 state?
    a) 13.6 eV
    b) 3.4 eV
    c) 10.2 eV
    d) -10.2 eV

  5. Which principle states that it is impossible to determine simultaneously the exact position and exact momentum of an electron?
    a) Aufbau Principle
    b) Hund's Rule
    c) Pauli Exclusion Principle
    d) Heisenberg's Uncertainty Principle

  6. Which set of quantum numbers (n, l, m<0xE2><0x82><0x9C>, m<0xE2><0x82><0x9B>) is NOT permissible for an electron in an atom?
    a) (2, 1, 0, +1/2)
    b) (3, 2, -2, -1/2)
    c) (1, 0, 0, +1/2)
    d) (2, 2, 1, +1/2)

  7. The number of radial nodes and angular nodes for a 4d orbital are, respectively:
    a) 2, 2
    b) 1, 2
    c) 3, 1
    d) 2, 1

  8. Which rule is violated in the following electronic configuration: 1s² 2s² 2pₓ² 2p<0xE1><0xB5><0xA7>¹ 2p<0xE1><0xB5><0xB3>⁰ ?
    a) Aufbau Principle
    b) Pauli Exclusion Principle
    c) Hund's Rule of Maximum Multiplicity
    d) Heisenberg's Uncertainty Principle

  9. What is the correct electronic configuration of Chromium (Cr, Z=24)?
    a) [Ar] 4s² 3d⁴
    b) [Ar] 4s¹ 3d⁵
    c) [Ar] 4s⁰ 3d⁶
    d) [Ar] 3d⁶ 4s⁰

  10. If the wavelength of a photon is 2.2 × 10⁻¹¹ m, Planck's constant is 6.6 × 10⁻³⁴ J s, then the momentum of the photon is:
    a) 3 × 10⁻²³ kg m/s
    b) 3.33 × 10²² kg m/s
    c) 1.452 × 10⁻⁴⁴ kg m/s
    d) 6.89 × 10⁴³ kg m/s

Answer Key for MCQs:

  1. b
  2. a
  3. c (mvr = nh/2π, here n=3)
  4. c (ΔE = E₂ - E₁ = (-13.6/2²) - (-13.6/1²) = -3.4 - (-13.6) = 10.2 eV)
  5. d
  6. d (For n=2, l can only be 0 or 1. l=2 is not allowed.)
  7. b (For 4d: n=4, l=2. Radial nodes = n-l-1 = 4-2-1 = 1. Angular nodes = l = 2)
  8. c (Hund's rule requires filling each degenerate orbital singly before pairing.)
  9. b (Due to stability of half-filled d-subshell)
  10. a (For photon, p = h/λ = (6.6 × 10⁻³⁴ J s) / (2.2 × 10⁻¹¹ m) = 3 × 10⁻²³ kg m/s)

Make sure you understand the concepts behind each point and MCQ. Revise the formulas, definitions, principles, and especially the exceptions. Good luck with your preparation!

Read more