Chem. 152
Dr. Saidane
The Modern Atom
and the
Periodic Table
1. How do electrons create the colors in a line-emission spectrum? When an electron falls
from higher energy level to a lower a quantum of energy in the form of light is emitted
form the atom.
2. List the four quantum numbers and describe what specific information is given by each of the four
quantum numbers. There are four quantum numbers n, l, m, and s. n is the principle that is
used to determine the energy level of the electron. It determines the distance from the
nucleus. l is the azimuthal quantum number. l is used to determine the shape of a orbital
(Lobe, torus, sphere…). m is the magnetic quantum, which determines the number of
orbitals for each sublevel. s is the spin quantum which determines the direction (clockwise
and couterclockwise) the electron is spinning.
3. Compare atomic orbitals s, p, d, and f in terms of their shape, size, and energy.
a) s orbital
sphere
1 orbital
b) p orbital
Figure eight
3 orbitals
c) d orbital
Double figure eight
5 orbitals
d) f orbital
Quadruple figure eight
7 orbitals
S < p< d < f
Increasing order of size and energy
4. What is an atomic orbital? An atomic orbital is shape of the path of an electron.
5. How many orbitals are in the following?
a) 3p-3
b) 2s;- 1
c) 4f;- 7
d) 4p- 3
e) 3d; 5
f) n=3; 9
g) n=4 16
6. What is the maximum number of electrons that can go into each of the following sublevels?
a) 2s; 2
b) 3p; 6
c) 4s; 2
d) 3d; 10
e) 4p; 6
f) 4f- 14
7. What is the maximum number of electrons that can go into the following energy levels?
a) n=2; 8
b) n=3; 18
c) n=4 32
8. In which member of each of the following pairs of sublevel have a higher energy?
a) 2s or 3s; 3s
b) 5 p or 7 s; 7s
c) 4 f or 4 s. 4f
9. Complete the following table:
Principal
# of sublevel
Type of orbital
# of orbitals
# of electrons
1
1
1s
1
2
2
2
2s, 2p
1+3=4
8
3
3
3s, 3p, 3d
1+3+5=9
18
4
4
4s, 4p, 4d, 4f
1+3+5+7=16
32
Quantum #
10. For each of these elements: Se, Si, S, Sn, and Sc answer the following questions:
a) Write the complete electron configuration.
i. Se- 1s22s22p63s23p64s23d104p4
ii. Si- 1s22s22p63s23p2
iii. S- 1s22s22p63s23p4
iv. Sn- 1s22s22p63s23p64s23d104p65s24d105p2
v. Sc- 1s22s22p63s23p64s23d1
b) Write the abbreviated electron configuration.
i. Se-[Ar]4s23d104p4
ii. Si-[Ne]3s23p2
iii. S-[Ne]3s23p4
iv. Sn-[Kr]5s24d105p2
v. Sc-[Ar]4s23d1
c) Write the orbital diagram for the abbreviated electron configuration.
Se:
[Ar]
Si:
[Ne]
S:
[Ne]
Sn:
[Kr]
Sc:
[Ar]
d) What is the highest energy level occupied by electrons?
i. Se- 4
ii. Si- 3
iii. S- 3
iv. Sn- 5
v. Sc- 4
e) Give the number of valence electrons.
i. Se -2
ii. Si- 4
iii. S- 6
iv. Sn- 4
v. Sc- 2
f) Give the number of paired and unpaired electrons.
i. Se 32 paired 2 unpaired
ii. Si 12 paired 2 unpaired
iii. S 14 paired 2 unpaired
iv. Sn 62 paired 2 unpaired
v. Sc 20 paired 1 unpaired
11. Complete the following tables
Symbol Name
Lithium
Li
Identity
State
Period Group
Family
Block
valence e-s
Metal
Solid
2
Alkali
S
1
1A
Metals
B
Boron
Metalloids
Solid
2
3A
None
P
3
Se
Selenium
Nonmetals
Solid
4
6A
None
P
4
V
Vanadium Metal
Solid
4
5B
Transition D
2
Metals
Argon
Ar
Nonmetal
Gas
3
8A
Noble
P
8
Family
Block
valence e-s
Alkaline
S orbital
2
Gases
e- Configuration
Name
[Ne] 3s2
Magnesium Mg
Symbol Period Group
3
2A
Earth
Metals
[He] 2s22p5
Fluorine
F
2
7A
Halogens
P orbital
6
[Ar] 4s23d6
Iron
Fe
4
8B
Transition D orbital
2
Metals
[Xe] 6s25d14f3
Neodymium Nd
6
None
LanthanidesF orbital
2
[Rn] 7s26d15f1
Thorium
7
None
Actinides
2
Th
F orbital
12. What trends are observed among the atomic radii across a period? How can this trend be explained?
The atomic radii get smaller across a period. This is caused by the increase of the number of
protons which increases the pull and make the atom smaller.
13. What trends are observed among the atomic radii down a group? How can this trend be explained?
The atomic radii get bigger down a group. This is caused by the increase of the energy levels
which cause the increase of the shielding effect.
14. How does the size of a cation and an anion compare to the size of the neutral atom from which it is
formed? Explain.
a. The cation is smaller than a neutral atom because of an increase of the pull of protons (more
protons than electrons). The anion on the other hand is bigger because it has more electrons than
protons, which creates more repulsion between electrons and a weaker pull of the protons.
15. For each of the following groups, indicate whether electrons are more likely to be lost or gained and give
the number of such electrons as well as the symbol for the resulting ion: Group 1, 2, 13, 16, 17 and 18.
a. Group 1, Lose 1 electron
b. Group 2, Lose 2 electron
c. Group 13, lose 3 electron except for Boron
d. Group 16, gain 2 electron
e. Group 17 gain 1 electron
f.
Group18. Not likely to gain or lose electron
16. Classify the following atoms in order of increasing sizes: Ca, Si, F, Be, Mg, Br, As, Zr, and He. Explain.
First consider the energy level and then for atoms in the same period we use the number of
protons (more protons, stronger pull and smaller atom).
He <
F <
Be <
Si < Mg < Br < As < Ca < Zr
Shielding Effect Period
1
2
2
3
3
4
4
Pull Effect
9
4
14
12
35
33
20
Protons
4
5
17. Classify the ions formed from each of the following atoms in order of increasing ionic sizes: Na, Al, N, S,
Ca, and Cl
Each of the elements above will have the following ions: Na+, Al3+, N3-, S2-, Ca2+, Cl-.
Na+, Al3+, N3- have the configuration of Neon. They all have the same shielding but they have
different pull. The more positive Al3+ has the strongest pull (3 more protons than electrons) making
it the smallest ion followed by Na+. N3- is the largest of the three (3 more electrons than protons).
S2-, Ca2+, Cl- have the configuration of Ar. They all have the same shielding but they have
different pull. The more positive Ca2+ has the strongest pull (2 more protons than electrons) making
it the smallest ion followed by Cl-. S2- is the largest of the three (2 more electrons than protons).
18. Which of these elements chlorine, selenium, and bromine is a) the smallest atom? b) the atom with the
highest ionization energy?
Large atoms have the most shielding. When atoms have same shielding, atoms with less protons are
larger, less pull effect. Large atoms have low ionization energy, because it is easier to remove an
electron that is far from the nucleus, than an electron that is close to the nucleus.
a) the smallest atom: Cl (period 3) < Br (period 4, 35 protons) < Se(period 4, 34 protons)
b) the atom with the highest ionization energy: Cl > Br > Se