1. Mullis Chemistry Holt Ch.41 Atomic Structure Summary of Contributions Max Planck –A hot object emits energy in small, specific amounts called quanta. Albert Einstein –E = mc 2 –Light behaves as both a wave and a particle. –Each particle of light carries a photon ( a quantum of energy). Niels Bohr –Created a model of the atom that showed a single electron of hydrogen orbits the nucleus only in allowed orbits with a fixed energy.

2. Mullis Chemistry Holt Ch.42 Atomic Structure Summary of Contributions, continued Werner Heisenberg –Uncertainty principle: It is impossible to determine simultaneously both the position and velocity of an electron. Erwin Schrödinger –Helped lay the foundation for quantum theory with an equation that treats electrons like waves.

3. Electromagnetic Spectrum ©antonine-education.co.uk

4. Solar Radiation: Infrared 45% of solar radiation is in the infrared (IR) region of the electromagnetic (EM) spectrum. IR causes molecules to vibrate or rotate faster. Faster movement of molecules is perceived by us as higher temperature in any material that received the energy. CO 2 and H 2 O vapor absorb most of the sun’s IR radiation. Some of the higher-energy solar radiation will be absorbed and then reflected back—it is “re-radiated” away from the surface of the Earth and is captured in the Earth’s atmosphere.

5. Solar Radiation: Ultraviolet (UV) 9% of solar radiation is in the UV part of the EM spectrum. Based on the actual wavelength, UV is divided into 3 types. UV-A (>320 nm): Lowest energy and is most likely to be absorbed by a layer of ozone high in the atmosphere. UV-B (280-320 nm): Much of this is also absorbed by the ozone layer. Causes sunburn. UV-C (<280 nm): Highest energy and is absorbed in the atmosphere. Can be used to kill viruses and bacteria.

6. Mullis Chemistry Holt Ch.46 Wave Description of Light Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space Wavelength (λ) –Distance between corresponding points on adjacent waves. –Unit: nm,cm,m Frequency (ν) –Number of waves that pass a specific point in a given time –Unit: Hz or waves/sec Recall that Speed = Distance/time (m/sec) Speed of light (c) C = λ ν

7. Mullis Chemistry Holt Ch.47 Behavior of Light Photoelectric effect –The emission of electrons when light shines on the metal –Scientists found that below a certain frequency, no electrons were emitted. –Light also behaves as a particle: Since hot objects do not emit em energy continuously, they must emit energy in small chunks called quanta. Quantum –Minimum quantity of energy that can be gained or lost by an atom

8. Mullis Chemistry Holt Ch.48 Light as a particle and a wave Planck and Einstein Max Planck: Relationship between quantum of energy and wave frequency Planck’s constant h = 6.626 x 10 -34 J-s E = hνE is energy, ν is frequency Albert Einstein: Established dual wave-particle nature of light 1 st –Einstein explained PE effect by proposing that EM radiation is absorbed by matter only in whole numbers of photons. –Electron is knocked off metal surface only if struck by one photon with certain minimum energy.

9. Mullis Chemistry Holt Ch.49 Quantum Theory Ground state: An atom’s lowest energy state Excited state: Higher potential energy than ground state. Photon: A particle of electromagnetic radiation having zero mass and carrying a quantum of energy (i.e., packet of light)

10. Mullis Chemistry Holt Ch.410 Niels Bohr links hydrogen’s electron with photon emission Bohr proposed that an electron circles the nucleus in allowed orbits at specific energy levels. –Lowest energy is close to nucleus Bohr’s theory explained the spectral lines seen in hydrogen’s line emission spectrum, but it did not hold true for other elements.

11. Mullis Chemistry Holt Ch.411 Electron Configuration: The Rules Aufbau principle –An electron occupies the lowest energy orbital that will receive it. Pauli exclusion principle –No two electrons in the same atom can have the same set of 4 quantum numbers. –Therefore, electrons can pair in an orbital as long as their spins are opposite. Hund’s rule –Each of the orbitals at a particular level have one electron before any of them can have two electrons. –All single electrons in the orbitals at a particular level have the same spin.

12. Label and color the electromagnetic spectrum (See p.92). Position the energy and wavelength as shown. 12 Include the following: Visible spectrum (ROYGBIV) X- raysInfrared (IR) Gamma RaysMicrowavesUltraviolet (UV) Radio Waves Wavelength, (λ) (Larger wavelengths to left) Increasing Energy and Frequency(ν ) 10 4 m 10 cm 10 3 nm 400-700nm 10nm 1 nm 10 -2 nm 400 nm700 nm

13. Vocabulary Photon_______________________________ __________________________________ Quanta_______________________________ __________________________________ Excited state of an electron _____________________________________ Ground state of an electron _____________________________________ Mullis Chemistry Holt Ch.413

14. Indicate meaning (and value if a constant) by each symbol. λ____________________________________ ν____________________________________ c___________________________________ (This is a constant) h____________________________________ (This is a constant) Mullis Chemistry Holt Ch.414

15. Write the main contribution of each scientist for this topic. Bohr Aufbau Pauli Hund Einstein Heisenberg Planck 15