Download Ch 5 Electrons in Atoms

NA Chem F09 SAO Unit 3: (Atomic Structure & Electrons in Atoms & Quantum Mechanics)
Summative Assessment Objectives
Atomic Structure: OBJECTIVES
1. Identify three types of subatomic particles by charge, mass, location
2. Describe the structure of atoms according to the Rutherford atomic model
3. Compare each of the following scientists concept of the atom: Dalton, Thomson, Rutherford, Bohr
4. Of the three subatomic particles, identify which determines the element and which can be lost/gained and the atom still be of the same element
5. Of the three subatomic particles:
a. Name, charge, location, relative mass (in amu)
b. Explain why electron’s relative mass is considered to be 0 and whether an electron is actually matter
c. Since all atoms are electrically neutral, identify which subatomic particles must be in equal number in an atom
6. For ions…
a. Describe a formula to calculate charge of an ion if given the counts of each subatomic particle
b. Write a chemical symbol including ionic charge
7. For isotopes…
a. Describe how isotopes of the same element differ from one another
b. Do isotopes of the same element have the same chemical properties?
c. Explain how the number of neutrons affect the mass of an atom
8. Chemical properties of an atom are mostly determined by overall charge and total number of positive subatomic particles—based on this info,
explain which of the subatomic particles have the greatest affect on an atom’s chemical properties.
9. If given counts of each subatomic particle and periodic table, write a complete chemical symbol
10. If given “element name-mass number”, write a complete chemical symbol or vice versa
11. Explain how mass number differs from atomic mass in concept and in method to calculate
12. Calculate the number of neutrons in an atom
13. Calculate the atomic mass of an element if percent abundance is provided
14. Determine which isotope must have a greater percent abundance if atomic mass is provided
Electrons in Atoms: Objectives
15. Describe the Bohr model for electron location; contrast with quantum model
16. Contrast the concept of “electron cloud” vs. orbit or shell
17. Describe the atom as mostly empty space with an extremely small, dense nucleus consisting of the protons and neutrons and an electron
cloud surrounding the nucleus (C4.8B)
18. Explain the difference among energy levels, sublevels, and orbitals
19. Be able to recognize & sketch the shapes of each orbital (s, p, d, not responsible for f)
20. Describe shape and orientation of s and p orbitals (C4.8h)
21. Explain why as energy level (n) increases, the size of each orbital increases
22. Be able to calculate the number of electrons per sublevel, and the number of orbitals per energy level
23. Recognize which rule is being violated on incorrectly drawn orbital diagrams.
24. Using a periodic table, correctly write orbital diagrams and electron configurations for atoms and ions use Aufbau order.
25. Using a periodic table, be able to use orbital diagrams and electron configurations to identify an element.
26. Distinguish between an electron configuration and orbital diagram written in the element’s ground state (low energy) or excited state (high
energy)
27. Write the electron configuration of elements in the first four rows of the periodic table (C4.8e)
28. Write the kernel structures (aka “shorthand”) for main group elements (C4.8f)
29. Explain why the actual electron configuration for some elements (namely Chromium, copper, silver, and molybdenum) differ from those
predicted by the aufbau principal based on stability (*MMC—C4.8f))
30. Describe the wave characteristics of wavelength, frequency, amplitude, crest, and trough
31. Recognize units of wavelength and frequency
32. Organize types of radiation based on electromagnetic spectrum (p. 139, Figure 5.10)
33. Describe the relationship between the wavelength (λ) and frequency(ν) of light (speed of light = wavelength x frequency, or c = λν );
34. Compare various wavelengths of light (visible and nonvisible) in terms of frequency and relative energy (C2.4d)
35. Explain how electron movement can result in production of light
36. Describe the energy changes in flame tests of common elements in terms of the (characteristic) electron transitions (C2.4a)
37. Explain why an electron can only absorb certain wavelengths of light (C2.4 c)
38. Contrast the mechanism of energy changes and the appearance of absorption and emission spectra (C2.4 x)
39. Identify the source of atomic emission spectra and atomic absorption spectra
40. Contrast atomic emission spectra and atomic absorption spectra
41. Contrast atomic emissions spectrum (individual lines) vs. a continuous spectrum (blend of the lines of atomic emission spectra)
42. Explain how the frequencies (ν) of emitted light are related to changes in electron energies(E) (energy of photon = Planck’s constant x
frequency…or E = h ν );
43. Relate wavelength of light and energy of light
44. Explain the logic behind the Heisenberg Uncertainty Principle
45. Describe the fact that the electron location cannot be exactly determined at any given time (C4.8)