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Chapter 11 and 12.2: Basic Review Worksheet
1. What is electromagnetic radiation?
All forms of light- either in wave or photon form
2. Give some examples of electromagnetic radiation.
Microwaves, visible light, radiation, radio waves, infrared, ultraviolet
3. Explain what it means for an atom to be in an excited state and what it means for an atom
to be in its ground state.
Ground state: atom is at rest
Excited state: energy has been applied to the atom (i.e. as heat or electricity) and this extra
energy gets the atom in an excited state.
Ground State
4. How does an excited atom return to its ground state?
Gives off photons that we see as visible light or heat
5. What is a photon?
Packet of energy
6. How is the wavelength (color) of light related to the energy of the photons being emitted by
an atom? How is the energy of the photons being emitted by an atom related to the energy
changes taking place within the atom?
As an atom returns from its excited state it gives off photons of light that we can see as
different colors. Each color we see depends on which element we are viewing. Each
element has different levels of energy to show different wavelength of color.
7. Describe Bohr’s model of the hydrogen atom.
(e-) traveled about the nucleus in an orbit and this was based on the idea of the line
emission spectrum
8. How do wave mechanical orbitals differ from Bohr’s orbits? What does it mean to say that
an orbital represents a probability map for an electron?
Mechanical orbitals give a more precise idea of where electrons are where as Bohr’s model
was based on only looking at energy levels. The Heisenberg uncertainty principle states that
we do not know the exact location of an electron, but we can get a rough estimate or find
areas where there is a higher probability of finding an electron.
9. Explain what is meant by the term orbital.
Orbitals are the directions a particular sublevel can go. Each sublevel has a specific number
of directions in which it can turn itself. Each orbital can hold up to 2 electrons.
10. What is the symbol for the lowest-energy hydrogen orbital?
S – or specifically 1s
11. Describe electron spin.
Used in quantum numbers, electron spin is the direction the electron is going. An electron is
either spinning in a clockwise motion (+½) or a counter clockwise motion (-½).
12. How does electron spin affect the total number of electrons that can be accommodated in
a given orbital?
In each orbital a total of 2 electrons can be present. Since each electron has to have a
different spinning direction, either clockwise or counter clockwise. This limits the number that
can be within the space of an orbital.
13. What does the Pauli exclusion principle tell us about electrons?
No two electrons will have the same quantum number . It may even come down to the
electron being in the same location, but having a different spin.
14. Draw the diagnol rule. Explain why this is an important concept.
15. How many electrons can be placed in a given s subshell? In a given p subshell? In a
specific p orbital?
S subshell- 2 electrons
P subshell- 6 electrons total
P orbital- 2 electrons in each orbital
16. How many electrons overall can be accommodated in the first and second principal
energy levels?
Energy level number 1- maximum of 2 electrons (1s2)
Energy level number 2 – maximum of 8 electrons (2s2 2p6)
17. Define the valence electrons and the core electrons in an atom.
Valence electrons- electrons in the outermost energy level, used to identify bonding patterns
Core Electrons- electrons where each energy level is completely filled
18. Sketch the overall shape of the periodic table and indicate the general regions of the table
that represent the various s, p, d, and f orbitals being filled.
19. Why do we place unpaired electrons in the 2p orbitals of carbon, nitrogen, and oxygen?
What rule does this follow?
This follows the aufbau principal that states that electrons must be written in according to the
diagnol rule. Each energy level gets progressively larger needing more sublevels and more
orbitals.
20. Write the electron configurations for the following atoms (unabbreviated):
a. Na 1s22s22p63s1
b. N 1s22s22p3
c. Be 1s22s2
d. Sr 1s22s22p63s23p64s2 3d104p65s2
21. Write the electron configurations diagrams for the following atoms: See Section 11.9
a. P
b. Se
c. Ca
d. Ce
22. What is the abbreviated electron configurations for the following elements?
a. K [Ar] 4s1
b. O [He] 2s22p4
c. Mg [Ne] 3s2
d. Pb [Xe] 6s24f145d106p2
23. Write the quantum number that follows each of these quantum numbers?
a. 4 2 1 + ½
4 2 2 +½
b. 5 3 -2 -½
5 3 -1 -½
c. 2 0 0 -½
2 1 -2 +½
24. Write the quantum number associated with each electron configuration.
a. 4s2
4 0 0 -½
9
b. 6d
6 2 1 -½
c. 2s1
2 0 0 +½
8
d. 5f
5 3 -3 -½
25. What are the representative elements? In what region(s) of the periodic table are these
elements found? In what general area of the periodic table are the metallic elements found?
In what general area of the table are the nonmetals found? Where in the table are the
metalloids located?
Representative elements- represent the general bonding trends. They can be found in the
first period of each family.
Metallic elements: on either side of the staircase on the periodic table
Nonmetals: to the right of the staircase
Metals: to the left of the staircase
26. Why are the valence electrons more important to the atom’s chemical properties than the
core electrons? How is the number of valence electrons in an atom related to the atom’s
position on the periodic table?
Valence electrons determine bonding patterns which helps in predicting how the element will
be behave chemically. The columns on the periodic table are arranged based upon the
valence electrons that element has.
27. Explain how the valence-electron configuration of most of the elements can be written just
by knowing the relative location of the element on the table. Give specific examples.
Because if know which column the element is in they will have similar characteristics as the
rest of the elements in the that column. Also, the column number/heading indicates the
number of valence electrons.
28. Define the terms ionization energy and atomic radius.
Ionization energy: energy required to remove an electron from an atom (in the gas phase).
Atomic radius: size of the atom determined by number of energy levels and how far away
from the nucleus the valence electrons are
29. How do the ionization energies and atomic sizes of elements vary, both within a vertical
group (family) of the periodic table and within a horizontal row (period)?
Ionization energy: increases across a period (left to right), decreases down a group/family
Atomic radius: increases down a group/family, decreases across a period (left to right)
30. Arrange the following atoms from largest to smallest atomic radius, and from highest to
lowest ionization energy.
a. Na, K, Rb
AR: Rb, K, Na
IE: Na, K, Rb
b. C, O, F
AR: C, O, F
IE: F, O, C
c. Na, Si, O
AR: O, Na, Si
IE: O, Si, Na
31. Arrange the following atoms from largest to smallest atomic radius, and from highest to
lowest ionization energy.
a. Na, K, P AR: P, K, Na
IE: Na, K, P
b. Rb, N, Al AR: Rb, Al, N
IE: N, Al, Rb
c. Cs, I, O
AR: Cs, I, O
IE: O, I, Cs
33. What pattern according to periods and families does the electronegativity show? What is
electronegativity?
Increases going up a family/group. Increases left to right on a period.
Electronegativity is a measure of the ability of an atom in a molecule to attract electrons to
itself.
34. Arrange these elements from highest to lowest electronegativity.
a. Na, Li, K
b. C, O, B
c. Br, P, B
d. Mg, K, F
Na, Li, K
O, C, B
Br, P, B
F, Mg, K