The Particle Nature of Light and the Photoelectric Effect

1. The Particle Nature of Light and the Photoelectric Effect

2. Quantum Leap? • The Quantum Concept Describe what a burner on an electric stove looks like (color) when: It is turned off. It is turned up to the high setting. If that metal burner were placed in a 900 degree furnace.

3. Quantum Leap? • What you just described is an example of the Quantum Concept. • Temperature is just an average of the kinetic energy an object possesses. • The metal on the stove, as it gets hotter, can emit more energy as different colors of light. • The different colors correspond to different wavelengths and frequencies.

4. Max Planck • German physicist who realized this phenomenon could not be explained by the Wave Nature of Light • Planck introduced the Particle Nature of Light which states that: matter can gain or lose energy only in small quantized amounts, called quanta. Quantized is just a fancy word for a certain amount.

5. Max sets things straight • While many argued that Max didn’t know what he was talking about – Max’s new theory helped explain a commonly observed phenomenon known as the photoelectric effect. • This effect is seen when photo-electrons are emitted from the surface of a metal. This is how your solar powered calculators work!

6. The photoelectric effect – When light shines on metals, electrons (photoelectrons) are ejected from their surface. • A certain frequency has to be achieved or the effect does not work Red light will not cause electrons to eject!

7. The photoelectric effect has practical applications in photoelectrical cells used for solar powered cars, and solar powered calculators.

8. Don’t Worry Max, Einstein has your back! • In 1909, Einstein concluded that there was a dual nature to light (meaning two ways to look at it!) • A beam of light has many wavelike characteristics • But it can also be thought of as a bundle of tiny particles (bundles of energy) known as photons.

9. Niels Bohr & the Bohr Model • Bohr worked in Rutherford’s Lab (what was Rutherford famous for?) • In 1913 Bohr was the first to propose a quantum model – to predict the frequencies of light from the Hydrogen atom. • He stated that each element only has certain allowable energy states:

10. Niels Bohr & the Bohr Model • Lowest allowable is the ground state • When atoms gain energy, they move to an excited state • When the atom moves back down to its ground state, it emits the energy it gained to move up, as light.

11. Niels Bohr & the Bohr Model • How does an atom gain and lose energy? • Our friend, Mr. Photon – a bundle of particles of electromagnetic radiation hits the atom and excites the electron (wouldn’t you be excited too?) • To lose the energy, there must be an emission of the photon when the e- drops back to the ground state

13. Bohr Model of the Atom – Absorption of Light

14. Bohr Model of the Atom – Emission of Light

15. We can think of different colors of light like balls with different kinetic energies. • Blue light has a higher energy than green light, like the balls that make it into the top window. • Red light has the lowest energy, like the balls that can only make it to the lowest window.

16. Mathematical Relationship • Ephoton = h • Where E (measured in joules) is the quantum energy that can be emitted at a certain frequency, . • h is known as Planck’s constant and has a value of: 6.626 x 10-34 J s

17. Time for Practice! • What is the energy of radiation with a frequency of 6.32 x 10 20 Hz? • Formula: Equantum = h • E = (6.626 x 10-34 J s)(6.32 x 10 20 Hz) • E = 4.19 x 10 -13 J

18. Your Turn • Practice problems:

19. The Wave Nature of Light

20. Rutherford was great, but … • He did tell us that all of an atom’s positive charge and most of its mass was concentrated in the nucleus & that electrons orbited this nucleus • BUT he lacked detail on how exactly the electrons occupied the space around the nucleus. • Most importantly, why aren’t these negatively charged electrons pulled into the positive nucleus?

21. Way back in the 1900’s • Scientists observed that certain elements emitted visible light when heated in a flame. • This lead to the discovery that different elements emitted different colors of light and that this difference was due to the arrangement of electrons. • But to understand atoms – you must first understand light.

22. Parts of a Wave crest trough

23. Frequency

24. Basically states that light acts like a wave (picture a slinky!) Can mathematically be described by the equation: c = λ All light travels at the speed of light, c, which is 3.00 x 108 meters/second. Wave Nature of Light

25. Electromagnetic Spectrum

27. Visible light is a small portion of this spectrum. This is the only part of this energy range that our eyes can detect. What we see is a rainbow of colors. RedOrangeYellowGreenBlueIndigoViolet ROY G BIV

28. The radiation to which our eyes are most sensitive has a wavelength near the middle of this range, at about 5.5 x 10-7m (550 nm), in the yellow-green region of the spectrum.

29. It is no coincidence that this wavelength falls within the range of wavelengths at which the Sun emits most of its electromagnetic energy—our eyes have evolved to take greatest advantage of the available light.