Atomic Structure & Nuclear Chemistry flash cards

1. 1.State the main postulates of Dalton's atomic theory.

   Matter is made of tiny indivisible atoms; atoms of an element are identical in mass and properties; atoms of different elements differ; atoms combine in small whole-number ratios to form compounds; atoms are neither created nor destroyed in reactions.

   Hint: Indivisible building blocks combining in whole-number ratios.

2. 2.Which observations could Dalton's theory NOT explain?

   It could not explain the existence of isotopes, isobars, subatomic particles (electron, proton, neutron), and the divisibility of the atom shown by discharge-tube and radioactivity experiments.

   Hint: Isotopes and subatomic particles broke 'indivisible'.

3. 3.What did the discharge-tube (cathode-ray) experiment establish?

   Cathode rays are streams of negatively charged particles (electrons) that travel from cathode to anode; they are identical regardless of gas or electrode, proving the electron is a universal constituent of all matter.

   Hint: Negative particles from the cathode, gas-independent.

4. 4.What is the charge-to-mass ratio of the electron found by Thomson?

   $e/m_e = 1.758 \times 10^{11}$ C/kg. Thomson measured this using deflection of cathode rays in electric and magnetic fields.

   Hint: Order $10^{11}$ C/kg, from field deflection.

5. 5.What did Millikan's oil-drop experiment determine?

   The magnitude of the electronic charge, $e = 1.602 \times 10^{-19}$ C. Combined with $e/m_e$ it gives the electron mass $m_e = 9.11 \times 10^{-31}$ kg.

   Hint: Charged oil droplets balanced against gravity.

6. 6.Describe Thomson's plum-pudding model of the atom.

   The atom is a uniform sphere of positive charge with electrons embedded in it like plums in a pudding (or seeds in a watermelon), making the atom electrically neutral overall.

   Hint: Electrons stuck in a positive sphere.

7. 7.What were the observations of Rutherford's gold-foil (alpha-scattering) experiment?

   Most alpha particles passed straight through; a few were deflected by large angles; very few (about 1 in 20000) bounced almost straight back.

   Hint: Mostly straight through, rare large-angle bounces.

8. 8.What conclusions did Rutherford draw from alpha scattering?

   The atom is mostly empty space; nearly all mass and the entire positive charge are concentrated in a tiny dense nucleus; electrons revolve around the nucleus at relatively large distances.

   Hint: Tiny dense positive nucleus, empty space around it.

9. 9.What were the two main drawbacks of Rutherford's nuclear model?

   (1) A revolving (accelerating) electron should continuously radiate energy and spiral into the nucleus, making the atom unstable. (2) It could not explain the discrete line spectrum of atoms.

   Hint: Spiral collapse and no explanation of line spectra.

10. 10.Give the discoverers and relative charges of the proton and neutron.

    Proton: discovered via anode/canal rays (Goldstein) and characterized by Rutherford; charge $+1$, mass $\approx 1$ u. Neutron: discovered by Chadwick (1932); charge $0$, mass $\approx 1$ u.

    Hint: Goldstein/Rutherford positive; Chadwick neutral.

11. 11.Define atomic number (Z) and mass number (A).

    Atomic number $Z$ = number of protons (= number of electrons in a neutral atom). Mass number $A$ = number of protons + number of neutrons = $Z + N$.

    Hint: $Z$ = protons; $A$ = protons + neutrons.

12. 12.Define isotopes, isobars, and isotones with one example each.

    Isotopes: same $Z$, different $A$ (e.g. $^{1}$H, $^{2}$H). Isobars: same $A$, different $Z$ (e.g. $^{40}$Ar, $^{40}$Ca). Isotones: same number of neutrons (e.g. $^{14}$C and $^{16}$O, both $N=8$).

    Hint: Same Z / same A / same N.

13. 13.What are isoelectronic species? Give examples.

    Species having the same number of electrons. Example: $\mathrm{N^{3-}}, \mathrm{O^{2-}}, \mathrm{F^-}, \mathrm{Ne}, \mathrm{Na^+}, \mathrm{Mg^{2+}}, \mathrm{Al^{3+}}$ all have 10 electrons.

    Hint: Same electron count, e.g. the 10-electron series.

14. 14.Relate wavelength, frequency, and wavenumber of electromagnetic radiation.

    $c = \nu \lambda$ where $c = 3 \times 10^{8}$ m/s. Wavenumber $\bar{\nu} = 1/\lambda = \nu/c$, usually in $\mathrm{cm^{-1}}$.

    Hint: $c=\nu\lambda$; wavenumber is reciprocal of $\lambda$.

15. 15.What is a black-body and what problem did its radiation pose for classical physics?

    A black-body is a perfect absorber and emitter of radiation. Classical theory predicted intensity rising without limit at short wavelengths (the 'ultraviolet catastrophe'), contradicting the observed peaked curve.

    Hint: Ideal emitter; classical theory failed in the UV.

16. 16.State Planck's quantum hypothesis and the energy of a quantum.

    Energy is absorbed or emitted only in discrete packets (quanta). Energy of one quantum: $E = h\nu = hc/\lambda$, with Planck's constant $h = 6.626 \times 10^{-34}$ J s.

    Hint: $E=h\nu$; energy comes in packets.

17. 17.State the key experimental observations of the photoelectric effect.

    Electrons are ejected only above a threshold frequency $\nu_0$; the effect is instantaneous; the number of electrons depends on intensity; the maximum kinetic energy depends on frequency, not intensity.

    Hint: Threshold frequency, KE from frequency, count from intensity.

18. 18.Write Einstein's photoelectric equation and define its terms.

    $h\nu = W_0 + \tfrac{1}{2}m v_{max}^2$, where $W_0 = h\nu_0$ is the work function (threshold energy) and $\tfrac{1}{2}m v_{max}^2$ is the maximum kinetic energy of the ejected electron.

    Hint: Photon energy = work function + max KE.

19. 19.How does the photoelectric effect contradict classical wave theory?

    Classically, energy should build up over time and any frequency should eject electrons if intense enough. Instead there is a sharp threshold frequency and instantaneous emission, showing light behaves as particles (photons).

    Hint: No threshold and time-lag expected classically; neither seen.

20. 20.What is the difference between emission and absorption spectra?

    Emission spectrum: bright coloured lines on a dark background emitted by excited atoms. Absorption spectrum: dark lines on a continuous bright background where a cool gas absorbs specific wavelengths. The line positions match.

    Hint: Bright lines emitted vs dark lines absorbed at same positions.

21. 21.Why is the atomic spectrum called a 'line spectrum'?

    Atoms emit or absorb only certain discrete wavelengths corresponding to allowed energy-level differences, giving sharp separate lines rather than a continuous band. This reflects quantized energy levels.

    Hint: Discrete lines = quantized energy jumps.

22. 22.State the postulates of the Bohr model of the hydrogen atom.

    (1) Electrons move in fixed circular orbits without radiating. (2) Angular momentum is quantized: $mvr = n h/2\pi$. (3) Energy is emitted/absorbed only when an electron jumps between orbits, with $\Delta E = h\nu$.

    Hint: Stationary orbits, quantized $mvr$, jumps give $\Delta E=h\nu$.

23. 23.Write Bohr's quantization condition for angular momentum.

    $m v r = \dfrac{n h}{2\pi}$, where $n = 1, 2, 3, \dots$ is the principal quantum number; angular momentum is an integral multiple of $h/2\pi$.

    Hint: $mvr = nh/2\pi$.

24. 24.Give the Bohr formula for the radius of the $n$th orbit of a hydrogen-like species.

    $r_n = 0.529 \, \dfrac{n^2}{Z}$ angstrom $= 0.529 \times 10^{-10}\dfrac{n^2}{Z}$ m. Radius increases as $n^2$ and decreases with $Z$.

    Hint: $r_n = 0.529\,n^2/Z$ angstrom.

25. 25.Give the Bohr expression for the velocity of an electron in the $n$th orbit.

    $v_n = 2.188 \times 10^{6} \, \dfrac{Z}{n}$ m/s. Velocity is proportional to $Z/n$.

    Hint: $v_n \propto Z/n$, about $2.19\times10^6$ m/s for H, n=1.

26. 26.State the Bohr energy expression $E_n$ for a hydrogen-like atom.

    $E_n = -13.6 \, \dfrac{Z^2}{n^2}$ eV $= -2.18 \times 10^{-18} \dfrac{Z^2}{n^2}$ J per atom. The negative sign indicates the electron is bound.

    Hint: $E_n = -13.6\,Z^2/n^2$ eV.

27. 27.Why is the energy of a bound electron negative?

    Zero energy is defined for a free electron at infinite separation (n = infinity). A bound electron has lower energy than this reference, so $E_n < 0$; energy must be supplied to remove it.

    Hint: Free electron at infinity is the zero; bound is lower.

28. 28.Define the ionization energy of hydrogen using Bohr theory.

    The energy to remove the electron from n = 1 to n = infinity: $E_\infty - E_1 = 0 - (-13.6) = 13.6$ eV for hydrogen ($Z=1$).

    Hint: From $n=1$ to infinity; 13.6 eV for H.

29. 29.Write the Rydberg formula for the hydrogen spectrum.

    $\bar{\nu} = \dfrac{1}{\lambda} = R_H Z^2 \left( \dfrac{1}{n_1^2} - \dfrac{1}{n_2^2} \right)$, with $R_H = 1.097 \times 10^{7}\ \mathrm{m^{-1}}$ and $n_2 > n_1$.

    Hint: $\bar\nu = R_H Z^2(1/n_1^2 - 1/n_2^2)$.

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