Atomic Structure flash cards

1. 1.Who discovered the electron and by what experiment?

   J.J. Thomson discovered the electron through the study of cathode rays in a discharge tube.

   Hint: Cathode ray tube experiments.

2. 2.What is the charge and mass of an electron?

   Charge $= -1.602 \times 10^{-19}$ C; mass $= 9.109 \times 10^{-31}$ kg (about $\frac{1}{1837}$ of a proton).

   Hint: Very light, negative fundamental particle.

3. 3.Who discovered the proton and how?

   The proton was identified from anode (canal) rays; Goldstein observed canal rays and the lightest positive particle (from hydrogen) is the proton.

   Hint: Positively charged canal/anode rays.

4. 4.What are the charge and mass of a proton?

   Charge $= +1.602 \times 10^{-19}$ C; mass $= 1.672 \times 10^{-27}$ kg (about 1836 times the electron).

   Hint: Positive, nearly one atomic mass unit.

5. 5.Who discovered the neutron and in what year?

   James Chadwick discovered the neutron in 1932 by bombarding beryllium with alpha particles.

   Hint: Neutral particle, discovered later than proton/electron.

6. 6.What are the charge and mass of a neutron?

   Charge $= 0$ (neutral); mass $= 1.675 \times 10^{-27}$ kg, slightly greater than a proton.

   Hint: No charge, mass close to a proton.

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

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

   Hint: Z from protons; A from nucleons.

8. 8.How do you find the number of neutrons in an atom?

   Number of neutrons $= A - Z$ (mass number minus atomic number).

   Hint: Subtract protons from nucleons.

9. 9.Define isotopes, isobars, and isotones.

   Isotopes: same $Z$, different $A$. Isobars: same $A$, different $Z$. Isotones: same number of neutrons.

   Hint: Same protons / same mass number / same neutrons.

10. 10.What are isoelectronic species? Give an example.

    Species with the same number of electrons. Example: $Na^+$, $Mg^{2+}$, $F^-$, $O^{2-}$ and $Ne$ each have 10 electrons.

    Hint: Same electron count, different nuclei.

11. 11.Describe Thomson's model of the atom.

    The plum pudding (watermelon) model: positive charge spread uniformly over the atom with electrons embedded in it, like seeds in a watermelon.

    Hint: Uniform positive sphere with embedded electrons.

12. 12.Describe Rutherford's alpha-particle scattering experiment.

    A thin gold foil was bombarded with alpha particles. Most passed straight through, some deflected, and very few (about 1 in 20000) bounced back.

    Hint: Gold foil and alpha particles.

13. 13.What conclusions did Rutherford draw from the scattering experiment?

    Atom is mostly empty space; positive charge and nearly all mass are concentrated in a tiny dense nucleus; electrons revolve around it.

    Hint: Small dense central nucleus.

14. 14.What was the main drawback of Rutherford's model?

    According to classical electromagnetism, an orbiting (accelerating) electron should continuously radiate energy, spiral inward, and the atom would collapse. It also could not explain line spectra.

    Hint: Accelerating charge radiates; atom should collapse.

15. 15.State the relationship between frequency, wavelength, and wave number for electromagnetic radiation.

    $c = \nu \lambda$, and wave number $\bar{\nu} = \dfrac{1}{\lambda} = \dfrac{\nu}{c}$, where $c = 3 \times 10^{8}$ m/s.

    Hint: Speed = frequency times wavelength.

16. 16.What is Planck's quantum theory of radiation?

    Energy is emitted or absorbed in discrete packets called quanta (photons). Energy of one quantum $E = h\nu$.

    Hint: Energy comes in discrete packets.

17. 17.State the formula for the energy of a photon.

    $E = h\nu = \dfrac{hc}{\lambda} = hc\bar{\nu}$, where $h = 6.626 \times 10^{-34}$ J s.

    Hint: Involves Planck's constant and frequency.

18. 18.What is the photoelectric effect?

    The ejection of electrons from a metal surface when light of suitable frequency falls on it.

    Hint: Light knocking electrons out of a metal.

19. 19.Define threshold frequency and work function in the photoelectric effect.

    Threshold frequency $\nu_0$ is the minimum frequency needed to eject electrons. Work function $W_0 = h\nu_0$ is the minimum energy needed.

    Hint: Minimum frequency / minimum energy to emit.

20. 20.State Einstein's photoelectric equation.

    $h\nu = h\nu_0 + \dfrac{1}{2}m v^2$, i.e. photon energy = work function + kinetic energy of ejected electron.

    Hint: Photon energy splits into work function and KE.

21. 21.In the photoelectric effect, how does kinetic energy of ejected electrons depend on light intensity and frequency?

    Kinetic energy depends only on the frequency of light, not on intensity. Intensity affects only the number of electrons ejected.

    Hint: Frequency sets energy, intensity sets count.

22. 22.What is an atomic emission (line) spectrum?

    A spectrum of bright lines at specific wavelengths emitted when excited atoms release energy; each element has a characteristic set of lines.

    Hint: Discrete bright lines, a fingerprint of the element.

23. 23.State the postulates of Bohr's model of the hydrogen atom.

    Electrons revolve in fixed circular stationary orbits without radiating energy; angular momentum is quantized ($mvr = n\frac{h}{2\pi}$); energy is emitted/absorbed only when an electron jumps between orbits.

    Hint: Quantized orbits, quantized angular momentum, quantum jumps.

24. 24.State Bohr's quantization condition for angular momentum.

    $mvr = n\dfrac{h}{2\pi}$, where $n = 1, 2, 3, \dots$ is the principal quantum number.

    Hint: Angular momentum is an integer multiple of h/2 pi.

25. 25.Give the expression for the radius of the nth Bohr orbit of a hydrogen-like atom.

    $r_n = 0.529 \dfrac{n^2}{Z}$ angstrom (i.e. $r_n = 52.9 \dfrac{n^2}{Z}$ pm).

    Hint: Radius scales as n^2 / Z.

26. 26.What is the radius of the first Bohr orbit of hydrogen?

    $r_1 = 0.529$ angstrom $= 52.9$ pm $= 0.529 \times 10^{-10}$ m (the Bohr radius).

    Hint: Bohr radius, about half an angstrom.

27. 27.Give the formula for the energy of the electron in the nth orbit of 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.

    Hint: Negative, scales as Z^2 / n^2.

28. 28.Why is the energy of an electron in an atom negative?

    The zero of energy is taken at infinite separation (free electron). A bound electron has lower energy than this, so $E_n$ is negative; more negative means more tightly bound.

    Hint: Bound states lie below the free-electron reference.

29. 29.What is the ground state energy of the hydrogen atom?

    For hydrogen ($Z=1$, $n=1$): $E_1 = -13.6$ eV.

    Hint: Put Z=1, n=1 into the energy formula.

30. 30.State 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 $n_2 > n_1$.

    Hint: Difference of two inverse-square terms.

Open the interactive deck for the other 69 cards, with self-grading so the ones you keep missing come back.