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Chapter 4 Review

Chapter 4 Review. Test is Thursday, December 22nd. Explain the term wavelength ( λ ) using a diagram. Explain the relationship between Wavelength ( λ ) and the energy of light. Speed of light (c) 3.00 x 10 8 m/s 186,000 miles per second 671,000,000 miles per hour

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Chapter 4 Review

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  1. Chapter 4 Review Test is Thursday, December 22nd

  2. Explain the term wavelength (λ) using a diagram

  3. Explain the relationship between Wavelength (λ) and the energy of light Speed of light (c) 3.00 x 108 m/s 186,000 miles per second 671,000,000 miles per hour Energy is inversely proportional to wavelength

  4. Explain the relationship between Wavelength (λ) and the energy of light Red light 650 nm – 750nm Violet light 380nm – 430nm 1m= 1 x 10-9nm (0.000 000 001 meters)

  5. Explain the relationship between Wavelength (λ) and the energy of light

  6. Write the equation for calculating energy (E) of light from its wavelength E = hc/λ E is energy of the electromagnetic radiation H is Planck’s constant 6.626 x 10-34 J s c is the speed of light 3.00 x 108m/s Remember λ must be in meters

  7. Draw a diagram to illustrate what is meant by line spectrum, and explain what the presence of line Spectra means in terms of energy that is emitted by atoms The fact that only 4 distinct lines are seen in the visible spectrum means that these atoms only emit the 4 amounts of energy associated with these wavelengths

  8. Explain the major differences between Classical Physics • Works well for large objects • Objects can have any energy Quantum Physics • Explains the behavior of very small particles • Objects have only certain particular energies

  9. Explain the meaning of the term quantized Electrons can only have particular discrete amounts of energy.

  10. Explain how the Bohr model accounts for line spectra • Electrons can jump from one energy level to another but never posses any energy in between (Quantum Leap)

  11. Explain the meaning of the term Principal Quantum Number • Bohr assigned each of his allowed electron “orbits” which he called shells a principal quantum number (n). The first shell was n=1 (lowest energy) The second shell n=2 (more energy) The third shell n=3 (more energy)

  12. Explain the meaning of the term Principal Quantum Number Ideally all electrons would like to be in the n=1 shell But… electrons repel each other and crowding them all so close would increase energy too much. To solve this problem and explain line spectra Bohr maintained that each shell can hold a maximum number of electrons

  13. The Bohr Model in a nutshell • Orbits get larger as the principal quantum number increases. • The energy of an electron in an atom increases with (n) the principal quantum number • Each shell can hold a maximum of 2n2 electrons

  14. Describe how valance shell configurations are related to the chemical properties of the elements. The Bohr model helped explain the periodic nature of the elements

  15. Describe how valance shell configurations are related to the chemical properties of the elements. The outermost shell is called the valance shell.

  16. Describe how valance shell configurations are related to the chemical properties of the elements. These elements have similar chemical properties because they have similar valance shell configurations.

  17. The number of valance shell electrons in an atom is equal to the roman numeral group number for the representative (group A) elements.

  18. Ground state • Arrangement of electrons that has the lowest energy

  19. Excited State • When enough energy is added to the atom (heating or passing electric current through) the electron can jump into a high energy level. • So an electron absorbs energy to jump to the excited state and releases energy to return to the ground state. (emitting light) (flame test lab)

  20. Bohr’s model was only able to predict line spectra for only one valance electron • Quantized energy • Maximum electrons in shells

  21. Calculate energy and determine the color of light emitted when an electron is excited and returns to the ground state. • Flame test lab • E = hc/λ

  22. Start here day 2

  23. Start day 3

  24. 1ev = 1.602 x 10-19J 1m = 1 x 109 nm h = 6.626 x 10-34 J * s c = 3.00 x 108m/s

  25. Subshells s, p, d, f

  26. s, p, d orbitalsWhere you would find an electron 95% of the time. (probability)

  27. Write the complete electron configuration for the following: Carbon Sodium Chlorine

  28. Write the shorthand notation electron configuration for the following: Carbon Sodium Chlorine

  29. Potassium problem • 2e-, 8e-, 9e- What? It should have 1e- in the valance shell • 1s2, 2s2, 2p6, 3s2, 3p6, 4s1

  30. Write the complete electron configuration for the following: Bromine Rubidium Tin (Sn)

  31. Write the shorthand notation electron configuration for the following: Bromine Rubidium Tin (Sn)

  32. Aufbau Principle: Electrons occupy the orbitals of the lowest energy levels first • Pauli exclusion principle: An atomic orbital can hold only two electrons. Electrons must have opposite spins. • Hund’s rule: electrons occupy orbitals of the same energy in a way that make the electrons with the same spin direction as large as possible.

  33. Compound formation and the octet rule • elements react to form compounds in such away as to put 8 electrons in their outermost valance shell. Just like noble gasses. • Some exceptions H = 2 Be and B sometimes 4 and 6. • Transition metals don’t often obey the octet rule.

  34. What is the chemical formula? Mg2+ F-

  35. Day 4

  36. Metals vs. Non Metals • Metals tends to lose valance electrons in a chemical reaction (cation) • Non metals tend to gain electrons in a chemical reaction (Anion)

  37. Explain why the trends in atomic size are opposite of the trends in ionization energy • Each step down adds an energy level thus increasing size • Each step across adds an electron and a proton more attraction between nucleus and electrons

  38. Heisenberg Uncertainty Principle

  39. Particles Behaving Like Waves

  40. s, p, d orbitalsWhere you would find an electron 95% of the time. (probability)

  41. Modern Model of the Atom • Rutherford’s small but massive positively charge nucleus. • Bohr’s “quantized” energy of electrons • Orbits replaced with probability orbital that has a wave like motion that gives it its shape • A particular size based on principle quantum number

  42. Photon: Discrete bundle of electromagnetic energy • Frequency: Number of wave cycles per second • Spectrum(electromagnetic spectrum) separation of light into different wavelengths • n=1 principal quantum number (Shell) • S, p, d, f sub shells

  43. Homework Due Thursday, December 22, 2011 • 4.40, 4.43, 4.60, 4.62, 4.72*, 4.73*4.80, 4.83, 4.117, 4.121, 4.123, 4.125, 4.137, 4.139, 4.141, 4.143, 4.151, 4.153, 4.160, 4.165, 4.174 • Blue worth 10 points • Gold worth 10 points

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