2. Unit 2: Atoms and Bonding 2.62 Unit 2: Atoms and Bonding 2.62 Photons and Energy Textbook ch 6.2 and ch 6.3

3. Big Idea 1: The chemical elements are fundamental building materials of matter, and all matter can be understood in terms of arrangements of atoms. These atoms retain their identity in chemical reactions. Students will be able to demonstrate understanding by laboratory investigation, analysis of data and creation of models. SWBAT: Explain what photons are and be able to calculate their energies given either their frequency or wavelengthExplain what photons are and be able to calculate their energies given either their frequency or wavelength Explain how line spectra related to the idea of quantized energy states of electrons in atomsExplain how line spectra related to the idea of quantized energy states of electrons in atoms Learning Objectives :

4. Max Planck He assumed that energy can be released (or absorbed) by atoms only in “chunks”. He called these chunks “quantum”. Quantum is the smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation

5. Think of it like comparing a ramp to climbing stairs. If you use the ramp your PE (potential energy) is uniform. If you climb stairs, you can only step on individual steps not between them. Your PE is restricted in certain values and is “quantized”

6. E = h E = h E = Energy, in units of Joules h = Planck’s constant (6.626 x 10 -34 J·s) = frequency, in units of hertz (hz or sec -1 ) = frequency, in units of hertz (hz or sec -1 ) The energy (E ) of electromagnetic radiation is directly proportional to the frequency ( ) of the radiation.

7. A few years later, Einstein used Planck’s theory to explain photoelectric effect. Light shining on a metal surface cause electrons to be emitted. Einstein went on to think the light was a stream of packets of energy. He called these packets of energy photons. E = hEnergy of photon = E = h Photon is a “particle” of light with energy E= h Planck and Einstein at Nobel Conference

9. Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table

10. So… short wavelengths, (like X-Rays) …have high frequency and high Energy X - ray photons can cause tissue damage and even cancer

11. e - can only have specific (quantized) energy values The e-’s energy correspond to orbits around the nucleus. Outer orbits have higher energy light is emitted as e - moves from higher energy level to lower energy level Photon E= h Niels Bohr’s Model of the Atom (1913) Ground State: n = Excited State: n Ionized: n = 1 >1 ∞ Niels Bohr and Max Planck at MIT

12. Figure 7.10 Quantum staircase

13. E = h Spectroscopic analysis of the hydrogen spectrum… …produces a “bright line” spectrum

14. This produces bands of light with definite wavelengths. Electron transitions involve jumps of definite amounts of energy.

15. Flame Tests sodiumlithiumpotassiumcopper Many elements give off characteristic light which can be used to help identify them.

16. Let’s try… 1. 1.Calculate energy of one photon of yellow light at 589nm 2. 2.Calculate energy of one mole of yellow light at 589nm

17. References www.sciencegeek.net/Chemistry and Mr. Perekupa’s PPT Cinnaminson.com from Cinnaminson High school. I modified the original PPTs to fit our needs in AP Chemistry. Our textbook: Brown, Lemay et all. AP edition chemistry, 13 th edition, 2015