Chapter 5-2, 5-3 Study Guide- Chemistry
1. What is an atomic orbital?
An atomic orbital is a three-dimensional region around the nucleus that describes the electron’s
probable location.
2. What is the difference between a ground state electron configuration and excited state
electron configuration?
Electrons can become excited due to the addition of energy, this causes the electron(s) to go
from the lowest possible energy state (the ground state) to a higher energy level (the excited
state).
3. Study Figure 5-10 on pg 128. What happens when an electron absorbs energy? What
happens when an electron drops from a higher energy orbit to a lower energy orbit in a
hydrogen atom?
The electron moves to a higher energy level. It releases energy in the form of waves, specifically
colors.
4. Describe the relationship between the wavelength and energy of light?
They are inversely related, the higher the energy the shorter the wavelength and vice versa.
5. The n=3 energy level is of considerably greater energy than n=2 energy level. As an
electron falls from these two-energy level, use the Bohr model to describe:
a. How the wavelengths of light compare, b. How the energies of light compare
The wavelengths become longer again. The energy drops and is released as waves.
6. Based on the de Broglie equation, why is the wavelength of a moving soccer ball not
detectable to the naked eye?
The wavelength of the moving soccer ball is much smaller than the wavelengths of visible light.
The mass must be extremely tiny.
7. What is the Heisenberg uncertainty principle?
It states that it is fundamentally impossible to know precisely both the velocity and position of a
particle at the same time.
8. Study the graphic organizers for electron configuration and orbital diagrams.
9. Know the 4 quantum numbers of the quantum mechanical model (see your graphic
organizer)
Principle quantum, secondary quantum, magnetic quantum, & spin magnetic quantum
10. How do 1s and 2s orbital differ? How are they alike?
The 2s orbital has a higher energy level, but they both can only hold two electrons.
11. What letters are used to designate the different types of orbitals? Study Table 5-2 pg 134
S, p, d, & f
12. How many sublevels are found in each principal energy level?
It is dependent on the energy level. The 1st principal energy has one sublevel, the 2nd has two
sublevels, the 3rd three and so on so forth.
13. How many orbitals are in each subshell? For example: How many orbitals in 4d?
1, 3, 5, 7
14. What are the maximum number of electrons that can be found in the s, p, d, or f orbital?
2, 6, 10, 14
15. Sketch the s and p orbitals (see Figure 5-18 pg 137)
16. In the designation 3d7, what does the 3, d, and 7 represent?
3 = principal energy level, d = sublevel, 7 = # of electrons
17. Study Figure 5-19 on pg 138. This will help you remember the order of the lowest energy
orbitals.
18. Know the 3 rules, or principles of how electrons are arranged in an atom’s orbital.
Aufbau principle, Hund’s rule, and Pauli exclusion principle
19. Write electron configurations and orbital diagrams for different elements (or atoms). I
will give you a periodic table.
20. Know the electron configurations of Cr and Cu.
1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^5; 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^10
21. Identify elements from their electron configurations or noble gas notations. Example: pg
147 #80
22. Identify the valence electrons.
23. Draw electron dot structures. (See Table 5-5 pg 140)
24. Write the noble gas electron configurations.
25. Review Planck’s equation: E = h x frequency. Remember the units of frequency are Hz or
s-1
H = 6.626 x 10^-34 Jxs
26. Review c = wavelength x frequency. The units of wavelength are m, cm, or nm. Know
how to convert these units. Review your notes and homework.
C = 3.00 x 10^8 m/s
27. Review vocabulary words, your notes, homework and classwork.