CP Chemistry
Unit 2 Review
Concepts:
Notes 2.1
 Electromagnetic radiation
o Photon
o Wavelength
o Frequency
o Wave speed
o Speed of light
o Energy of EM radiation
 Spectra
o Continuous spectrum
o Emission spectrum
o Visible light spectrum
o Electromagnetic spectrum
 Bohr model/planetary orbital model
o Absorption/emission of a photon
o Energy levels
o Quanta/quantized
o Ground state
o Excited state
o Energy of an electron
o Energy of transition + energy of
a photon
Notes 2.2
 Quantum Mechanical Model
o Quick history!
o Principle energy levels (n)
o Sublevel
o Atomic orbitals
o Electron spin
o Orbital diagrams
o Electron configuration (ground
state)
 Aufbau’s principle
 Pauli exclusion principle
 Hund’s rule
o Noble gas notation
o 4 blocks of periodic table
o Exception to expected electron
configuration
o Valence electrons
o Electron-dot diagrams
Practice problems:
1. What equations do you need to know for this test? What do all of the symbols stand for?
What are the units for each factor?
2. Which equation relates the frequency and wavelength of EM radiation?
3. Which equation allows you to find the energy of transition for an electron using the Bohr
model?
4. What is the energy of EM radiation with a wavelength of 450 nm?
5. What is the frequency of EM radiation with a wavelength of 900nm.
6. What color light has the most energy?
7. Compared to green light, UV light has more/less energy, a longer/shorter wavelength,
and a higher/lower frequency.
8. What is the frequency of a photon that has 4.50 x 10-19 J of energy?
9. What is the wavelength of a photon that has 6.78 x 10-20 J of energy?
10.What is the difference between continuous and emission spectra? What light source
would give you each type of spectra?
11.Draw a sketch of the Bohr model of the atom and label the parts. Draw an arrow that
indicates an electron transition where the electron loses energy and label it A. Draw
another arrow that indicates an electron gaining energy and label it B. Put a star next to
the arrow that indicates the atom emitting light.
12.When an atom becomes excited, has it gained or lost energy? Did it absorb or emit a
photon?
13.Explain the emission spectra for hydrogen.
14.What is the energy of an electron at n=1? How much energy does the electron gain
when it transitions from n=1 to n=3?
15.What is the wavelength of the photon emitted when an electron transitions from n=6 to
n=4?
16.If an electron transitions from n=2 to n=4, will ΔE be positive or negative? Why?
17.What is different between Bohr’s model of the atom and the Quantum Mechanical
Model?
18.How do we know that the energy levels in a hydrogen atom are not continuous?
19.Complete the chart
Energy
# of
orbital # of orbitals in
total # of
Level orbitals letters
each type
e1
2
3
4
20.When filling a p orbital, why do you separate the electrons into different orbitals, before
pairing them up. What is the name of this rule?
21.Identify the elements based on their configurations:
a. [Ar]4s23d9
b. [Kr]5s24d4
c. [Xe]7s25f4
d. 1s22s22p63s23p1
22.Give the electron configuration and noble gas notation for C, S, Si, Mg, Na, and Cu.
23.How many valence electrons do the elements in group 16 have? Group 13?
24.What information does an electron dot diagram give you? Draw the electron dot diagram
for F, Sr, Zn, C, P, Ne, H, V, U and Ag.
25.What are the elements that have an electron configuration that is different than what is
predicted by the three rules? Why do their configurations deviate?