1. QUANTUM AND NUCLEAR PHYSICS

2. Wave Particle Duality In some situations light exhibits properties that are wave-like or particle like. Light does not show both properties at the same time and place. Wave packet or photon is particle like in its finite size but wave like in its varying amplitude. The amplitude is a measure of the probability of finding a photon at a particular place. A large amplitude is a high probability of finding the photon

3. A Comparison Wave – Like Diffraction – bending around obstacles Interference Polarisation Reflection Refraction Particle Line emission spectra Exerts radiation pressure or force Photoelectric effect Gravitational bending of light eg around the Sun

4. Planck’s Quantum Theory Classical Physics Describes electromagnetic energy as a wave. Quantum Mechanics Describes electromagnetic energy as a “particle” equivalent.

5. Planck’s Quantum Theory 1901 Planck proposed that electromagnetic energy was quantised and only occurred in discrete amounts. (not continuous) called quanta A photon is a monochromatic (single frequency) quantum of EM energy. Planck’s Law states that the energy of each photon is directly proportional to its frequency.

6. energy α frequency Constant of proportionality is h Planck’s constant 6.626 x 10 -34 J.s photon = Planck’s x photon energy constant frequency E = h. f (J)

7. Energy of a photon Since the speed of electromagnetic radiation is given by: where c = 3.00 x 10 8 ms -1 then This is the energy of a photon of a wavelength of electromagnetic radiation

8. 1.What is the energy of a quantum of green light of frequency 5.0 x 10 14 Hz? 2. How much energy is contained in a photon of UV of wavelength 1.0 x 10 -7 m 3. Calculate the frequency and wavelength of a photon of light with energy 4.0 x 10 -19 J

9. Classical Predictions Intensity– as light increases number of photo electrons increases Emission time – low intensity light longer emission time Frequency – emission is independent of frequency Energy – as light intensity increases the kinetic energies of the photo- electrons increases Experimental Results Intensity – number of photoelectrons increased with intensity Emission – instantaneous Frequency – emission is frequency dependent. Below cut-off frequency no electrons emitted Energy- as intensity incr KE max stayed constant. KE is frequency dependent with type of material

10. Photoelectric Effect metals bombarded with high frequency light emit electrons (photoelectrons) occurs in solid, liquids and gases first discovered by Hertz in 1887

11. Einstein’s Theory Based his explanation of the photoelectric effect on Planck’s quantum idea. He proposed that: energy was in a particle called a photon Each photon had an energy of E = hf A photon could give up all its energy to one electron but not part “all or nothing” The maximum KE of the emitted electron was equal to the initial photon minus the work done in overcoming the attractive forces near the metal surface.

12. Ф is the work function, the minimum energy needed to escape the nuclear attractive force on the electrons. Ф is given in electron volts Ф = h.f 0 where f 0 is the cut-off or threshold frequency KE MAX = h.f - Ф ½ mv 2 = h.f – Ф

13. Electron volt This is the energy gained by an electron as it moves through an electric field of 1V from E P = qV or eV charge on an electron is 1.6 x 10 -19 C 1eV = 1.6 x 10 -19 C x 1V = 1.6 x 10 -19 J

14. The threshold frequency of rubidium is 5.0 x 10 14 Hz. Find a)the work function and b) the max velocity of electrons ejected by light of a frequency of 8.0 x 10 14 Hz ( the mass of an e is 9.1 x 10 -31 kg

15. Do now Write down Einstein’s photoelectric equation and explain the meaning of each symbol What is meant by threshold frequency calculate the f 0 of a material having a work function of 4.5eV 1.09 x 10 15 Hz

16. The f 0 for a metal is 3 x 10 14 Hz. If blue radiation of frequency 7 x 10 14 Hz falls on the material, find (a) Energy of incident photon (b) work function (c) max KE of ejected photoelectrons (d) max velocity of ejected photoelectrons. (mass of electron = 9 x 10 -31 kg

17. answers (a) 4.64 x 10 -19 J (b) 1.99 x 10 -19 J (c) 2.65 x 10 -19 J (d) 7.67 x 10 5 ms -1

18. Photocell and Planck’s constant

19. Summary Photoelectric Effect Define work function Energy is α frequency E α f or = hf Define threshold f Photoelectron emission is instantaneous Increase Intensity = more e emitted Insufficient energy or Φ metal heats up but no photoemission Wave/ Particle Duality Particle like nature eg. line emission spectra, radiation pressure or force Wave like nature eg. Interference light passing through slits; and diffraction – laser light falling on a screen to produce a central bright spot

20. Atomic Spectra Newton identified components of white light.white Bunsen & Kirchoff identified elements by the spectra produced.Bunsen Gases can be excited by heating or by electrical discharge There are three types of spectra: 1.Continuous spectrum – all frequencies represented; usually a heated light source eg SunSun 2.Emission spectrum- when an excited atom returns to its ground state, it emits e.r. of particular frequencies only. The energy the electron loses jumping from a higher to a lower orbit is emitted as a photon. E=hfE=hf 3. Absorption spectrum – a photon of energy can raise an electron to a higher orbit. If white light is passed through a gas containing sodium, certain frequencies will be removedspectrum

21. Absorption and Emission

24. Atomic Line Spectra When excited by heat or an electrical discharge the atoms gives off particular frequencies of light. The Balmer formula can be used to predict the wavelengths of visible light. This was later modified by Rydberg to give: where S = Series number L= Line number (S+1) R= Rydberg constant = 1.1 x 10 7 m -1

25. Hydrogen Spectra Analysis Rydberg’s formula also applies to the series of lines in the infra-red and ultraviolet spectrum 1.Lyman formula (ultra-violet range) 2. Balmer formula (visible light range) 3. Paschen formula (infra-red

26. The lines of the hydrogen spectrum correspond to particular frequencies and energy levels given by E=hf These levels are not evenly spaced but vary by 1/n 2 according to Rydberg. When the electron is in the lowest level or ground state, then n=1. The highest level is called the ionisation level, n=α Hydrogen Atom

27. Bohr’s Hydrogen Atom (1913) 1.electrons have fixed amounts of energy (quanitsed); they revolve around the nucleus in allowed orbits or stationary states without radiating energy 2.Electrons can jump from one energy level to another by either absorbing or emitting a photon of light equal to the difference in energy levels ΔE=|E - E| = hf 3. Angular momentum (mvr) is quantised and can only take multiples : L = n where n is the principal quantum number

28. Calculating Energy Emissions/Absorptions What are the possible energies that an emitted photon could have? E photon =hf=E n - E m = hc/λ n=3 n=2 n=1 -1.5eV -3.4eV -13.6eV

29. For n=3 to n=2 E= E 3 – E 2 = -1.5 – (-3.4) =1.9eV For n=3 to n=1 E = E 3 – E 1 = -1.5 – (-13.6) = 12.1eV For n=2 to n=1 E = E 2 – E 1 = -3.4 – (-13.6) = 10.2eV values +ve for emission (b) E = hf = hc/λ so λ=hc/E E values are in eV; convert to Joules 1eV = 1.6 x 10 -19 J h = 6.62 x 10 -34 Js c = 3.0 x 10 8 ms -1

30. From Rydberg’s formula The energy of the photon emitted is This is the energy difference between 2 electron energy levels. Electrons can only exist with energy levels of Where n=1,2,3….. Where n is the quantum number

31. Limitations of Bohr’s Model Does not explain why electrons have fixed energy levels Why some spectral lines are more intense than others the presence of hyperfine spectral lines Spectra for atoms larger than the hydrogen atom – one electron