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Lecture #18 The Hydrogen Atom

Lecture #18 The Hydrogen Atom. March, 30 th. J.J. Thomson’s Plum Pudding Model of the Atom (1897). Thomson proposed that the electrons are embedded in a positively charged ‘pudding’. Rutherford’s Scattering Experiment (1911).

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Lecture #18 The Hydrogen Atom

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  1. Lecture #18The Hydrogen Atom March, 30th

  2. J.J. Thomson’s Plum Pudding Model of the Atom (1897) Thomson proposed that the electrons are embedded in a positively charged ‘pudding’

  3. Rutherford’s Scattering Experiment (1911) He found that, once in a while, the a-particles were scattered backwards by the target

  4. Discovery of the Atomic Nucleus Rationalizing the experimental data To explain the backscattering, the positive charge must be concentrated in a small region

  5. Rutherford’s Solar System Model of the Atom Rutherford’s “planetary” model of the atom The atom consists of electrons orbiting around a small but dense central nucleus

  6. Hydrogen Atom is Unstable? • It is known that accelerating charges emit radiation • Thus, electron should emit radiation, lose energy and eventually fall into the nucleus! • Why doesn’t this happen? Shows that something was wrong with this model of the hydrogen atom

  7. Absorption Spectrum of a Gas Dark lines will appear in the light spectrum

  8. Absorption spectrum of Sun Emission spectra of various elements

  9. Balmer’s Formula for Hydrogen • Notice there are four bright lines in the hydrogen emission spectrum • Balmer guessed the following formula for the wavelength of these four lines:where n = 3, 4, 5 and 6

  10. Bohr’s Model of the Hydrogen Atom(1913) He proposed that only certain orbits for the electron are allowed

  11. Bohr’s Empirical Explanation • Electrons can only take discrete energies (energy is related to radius of the orbit) • Electrons can jump between different orbitsdue to the absorption or emission of photons • Dark lines in the absorption spectra are due to photons being absorbed • Bright lines in the emission spectra are due to photons being emitted

  12. Absorption / Emission of Photonsand Conservation of Energy Ef - Ei = hf Ei - Ef = hf

  13. Energy Levels of Hydrogen

  14. Electron jumping to a higher energy level E = 12.08 eV

  15. Spectrum of Hydrogen Bohr’s formula: R=Rydberg’s constant f=final state; i=initial state

  16. Hydrogen is therefore a fussy absorber / emitter of light It only absorbs or emits photons with precisely the right energies dictated by energy conservation

  17. This explains why some nebulae are red or pink in colour One of the transitions in the Balmer series corresponds to the emission of red light

  18. Schrödinger’s Improvement to Bohr’s Model • Showed how to obtain Bohr’s formula using the Schrödinger equation • Electron is described by a wave function y • Solved for y in the electric potential due to the nucleus of the hydrogen atom

  19. Vibrational Modes of a String fundamental 2nd harmonic 3rd harmonic 4th harmonic

  20. Quantum Numbers • Energy levels can only take discrete values • Labelled by a ‘quantum number’ n, which takes values 1, 2, 3, ... • Each level has energy that increases with n

  21. Ground State (n=1) • Lowest or ground-state energy is non-zero • Electron cannot sit still but must be forever ‘jiggling around’ • Expected from the Heisenberg uncertainty principle

  22. 1s Orbital

  23. Density ofthe cloud gives probability of where the electron islocated

  24. Summary • Electron does not fly round the nucleus like the Earth around the Sun (Rutherford, Bohr) • Depending on which energy level it is in, the electron can take one of a number of stationary probability cloud configurations (Schrödinger) • Textbook, • 27-10, pp. 769-770, • 27-11, pp. 771-773

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