Electron Orbital Practice Worksheet
Name
Rania
Period
Unit 3 – Worksheet 1
Label the Orbital Blocks and Energy Levels
S-block
d-block
2s
3s
4s
5s
6s
<----------------3d ----------------->
<-------- 2p ---------->
<-------- 3p ---------->
<-------- 4p ---------->
<----------------4d ----------------->
<-------- 5p ---------->
7s
<----------------6d ----------------->
<----------------5d ----------------->
Number of Orbitals
1
S
P
D
F
f-block
<-------- 6p ---------->
<----------------4f ----------------->
<----------------5f ----------------->
Orbital Name
p-block
5d
6d
Maximum # of Electrons
2
6
3
5
10
7
14
Assigning Electrons: Examples
Element: Atomic #
H
1
Orbitals
1s1
Noble Gas Shortcut (abbreviated)
NA
He
2
1s2
NA
Li
3
1s22s1
[He]2s1
B
5
1s22s22p1
[He]2s22p1
Ne
10
1s22s22p6
[He]2s22p6
Mg
12
1s22s22p63s2
[Ne]3s2
S
16
1s22s22p63s23p4
[Ne]3s23p4
Ti
22
1s22s22p63s23p44s23d2
[Ar]4s23d2
Practice Problems:
Full Configuration
1. P
Noble Gas Shortcut (abbreviated)
15
2. Mn 25
3. Ga 31
4. O 8
5. Co 27
6. As 33
7. Ru 44
8. Cd 48
9. Nb 41
10. Sn 50
Challenge:
11. Eu 63
12. U 92
the following have F orbitals – there are differences of opinion on how to do these depending on
the periodic table you look at. You will not be asked to do the electron configuration of any
element with F orbitals on a test.
Atomic Theory and Orbitals
Name
Period
Unit 3 – Worksheet 2 (Goals 1-4)
1. Describe how the quantum theory of atomic structure differs from Bohr’s theory.
( Hint: Focus on electrons)
Quantum theory differs from Bohr’s model by saying that electrons don’t move in fixed orbits. Instead, they are
found in areas called orbitals, where their position is more about probability. It also states that we can’t know
both the exact position and speed of an electron at the same time.
2. Differentiate between an orbit and an orbital.
An orbit is a fixed path that an electron follows around the nucleus, as suggested in Bohr's model. An orbital, on
the other hand, is a region around the nucleus where there is a high probability of finding an electron, according
to quantum theory. Orbits are fixed, while orbitals are not.
3. What do electron configurations tell us about an atom?
Electron configurations tell us how electrons are arranged in an atom's orbitals. They show which energy levels and sublevels
electrons occupy, helping us understand the atom's chemical properties, such as its reactivity and bonding behavior.
4. How many orbitals are in the fourth energy level?
16 orbitals
5. How many orbitals of each type are there?
s- 1
p- 3
d- 5
f- 7
6. Draw the 2s and 2px orbital.
7. Write the electron configurations of the following elements using the shorthand notation for the noble
gas cores.
Full
Abbreviated
a. phosphorus
b. nickel
c. neon
d. potassium
e. titanium
8.
Which orbital is filled following these orbitals?
a. 3d 4p
c. 5p 6s
b. 4s 3d
d. 5f
6d
9. How many electrons can be accommodated in
a. a d subshell 10 electrons
d. the 7s orbital 10 electrons
b. a set of f orbitals 14 electrons
e. a px orbital 2 electrons
c. the n= 4 shell 32 electrons
f. a set of p orbitals 6 electrons
10. What is wrong, if anything, with the following ground state electron configurations? If there is a problem, how
should it be corrected?
1) Problem: The 3d orbital is completely empty, which is incorrect for most configurations
where the 3d orbital is lower in energy than 4s.
Correction: Populate the 3d orbital according to the Aufbau principle before moving to 4p.
2)Problem: The 4s orbital is empty, which would not typically happen in a ground-state
configuration where 4s would fill before 4p.
Correction: Place two electrons in the 4s orbital before placing any in the 4p orbitals.
Problem: One of the 3d orbitals has more than two electrons, violating the Pauli
exclusion principle, which states that an orbital can hold a maximum of two electrons
with opposite spins.
Correction: Distribute the electrons correctly in the 3d orbitals (two maximum per orbital).
11. How many unpaired electrons are there in the following (consider orbital notation):
a. a nitrogen atom 3
b. an iodine atom 1
c. a nickel (Ni2+) cation 2
d. an oxide (O2-) ion? 0
12. Complete the chart below:
Element
Mg
Last Occupied
Orbital
3s
# of Valence
e- s
2
Li
U
2s
5f
1
2
O
Rn
2p
6p
6
6
F
As
2p
4p
7
5
13. Write out the full ground state electron configuration for an atom of aluminum.
14. Indicate all of the following for the electrons above:
Valance electrons, least tightly bound to the nucleus, most tightly bound to the nucleus, core electrons
15. The electron orbital notations above both show the filling patterns for 8 electrons. Why is “b” considered
incorrect? What principle does it violate?
B is considered in correct because it violates Hund's rule, which states that electrons will occupy separate orbitals of the
same sublevel with parallel spins before pairing up.
Electrons & Electromagnetic Radiation
Name
Period
Unit 3 – Worksheet 3 (Goals 2-5)
1.
Describe the difference between a neutral atom and an ion
A neutral atom has no charge because protons and electrons are equal. An ion has a charge because it gains or loses electrons.
Losing electrons makes a positive ion (cation). Gaining electrons makes a negative ion (anion).
2.
Indicate how many electrons each of the following neutral atoms will lose or gain to become ions:
a.
Na
Loses 1 electron to
form Na+
3.
b. Ca
Loses 2 electrons
to form Ca2+
c. O
Gains 2 electrons
to form O2−
d. Br
Gains 1 electron to
form Br−
Complete the information for each element listed below:
a. A neutral atom of sodium:
i. Full electron configuration –________________________________________________
ii. Noble gas elec config - ____________________________________________________
iii. Indicate which electrons above are the most energetic - ___________________________
iv. Indicate which electrons above are the least energetic - ___________________________
v. How many valence electrons are in the atom above and which are they? _______________
b.
A sulfide anion (S2-):
i. Full electron configuration –________________________________________________
ii. Noble gas elec config - ____________________________________________________
iii. Indicate which electrons above are the most energetic - ___________________________
iv. Indicate which electrons above are the least energetic - ___________________________
v. How many valence electrons are in the atom above and which are they? _______________
c.
A Barium cation (Ba2+):
i. Full electron configuration –________________________________________________
ii. Noble gas elec config - ____________________________________________________
iii. Indicate which electrons above are the most energetic - ___________________________
iv. Indicate which electrons above are the least energetic - ___________________________
v. How many valence electrons are in the atom above and which are they? _______________
4.
Describe what happens to an electron that absorbs energy?
When an electron absorbs energy, it jumps to a higher energy level, called an excited state. If it gains enough energy, it can leave the
atom entirely, causing ionization. Afterward, the electron usually releases the energy and falls back to a lower level.
5.
Will the electron above retain the extra energy? What happens?
No, the electron will not retain the extra energy permanently. After absorbing energy and moving to a higher energy level, it will eventually
release the extra energy.
6.
A valence electron goes from energy level n = 3 to n = 4 (n = energy level of orbital). At this point the atom is said to be in
an _________________________________________
state.
excited
7.
What is the reverse of the situation above? ______________________________________
state.
ground
8.
The following diagram shows 5 different energy levels (n) that electrons can occupy. Use the diagram to answer each of the
following questions:
a.
EXAMPLE Give a situation that could result in an electron making a quantum leap.
An electron gains energy and moves from n=1 to n= 3.
b.
Give a situation that would result in an electron giving off light.
An electron gives off light when it falls from a higher energy level (n=4 or n=5) to a lower energy level (n=3, n=2, n=1). The energy
difference is released as light.
c.
Give a situation that could produce red light
Red light is produced when an electron falls between two energy levels with a small energy difference, such as from
n=3 to n=2
d.
Give a situation that could produce blue light
Blue light is produced when an electron falls between two energy levels with a larger energy difference, such as from
n=5 to n=2.
9.
List all types of radiation in the Electromagnetic Spectrum. Label the following:
shortest wavelength, longest wavelength, lowest frequency, highest frequency, most energy, least energy
Gamma rays:
-shortest wavelength
-highest frequency
-most energy
X-rays:
slightly longer wavelength
than gamma rays
slightly lower frequency
very high energy
UV light
Visible light
Longer wavelength than X-rays Middle of the spectrum in
Lower frequency than X-rays
terms of wavelength,
High energy (but less than X-rays) frequency, and energy.
Infrared (IR)
Microwaves
Longer wavelength and lower
Longer wavelength and Radio waves
Longest wavelength,
frequency than visible light
lower frequency than
lowest frequency,
Emits heat, often associated with infrared
warm objects
Used in communication least energy
and heating food
10. Two lights are turned on. One glows with an orange color and the other glows with a purple color. Discuss the energy levels
of both lights. What is happening to create these different colors?
The orange light has lower energy because it has a longer wavelength and lower frequency. The purple light has higher energy because it has a shorter
wavelength and higher frequency. Different colors are created because electrons fall between energy levels with different energy gaps. A small gap produces
orange light, while a larger gap produces purple light.
11. An experiment is being performed with a spectral tube (light emitting device). The tube is energized slightly and gives off a
greenish light. Then the energy is increased and the green color ceases even though energy is still flowing through the tube
and electrons are still going through quantum leaps (energy transitions). What is occurring? (Hint – you would not want to
stay around the tube very long as it would be extremely unhealthy)
When the energy is increased, the green light stops because the electrons are now transitioning between much higher energy levels. This
produces higher-energy radiation, likely ultraviolet (UV) light or even X-rays, which are invisible but dangerous to humans.
Periodic Table and Trends
Name
Period
Unit 3 – Worksheet 4 (Goals 6 – 10)
1. What determines the vertical arrangement of the periodic table?
The vertical arrangement (groups or families) is determined by the number of valence electrons and similar chemical properties.
2. What determines the horizontal order of the periodic table?
The horizontal order (periods) is determined by the number of energy levels (electron shells) an atom has.
3. What determines the order of the periodic table?
The periodic table is ordered by increasing atomic number (number of protons).
4. Why did Mendeleev and other scientists of his time arrange elements in the periodic table in order of
atomic masses?
Mendeleev arranged elements by atomic masses because the atomic number was not yet discovered. Grouping by mass revealed repeating
patterns of chemical properties.
5. All halogens are highly reactive. What causes the similarity among the halogens?
Halogens are highly reactive because they have seven valence electrons, needing just one more to complete their outer shell, making them
eager to react.
6. What properties do Noble gasses share and how does this relate to their electron configuration?
Noble gases are inert and non-reactive because their outer electron shell is completely filled, making them stable and unlikely to gain or lose electrons.
7. For each element tell which period and group the element is in, identify the element, and state
whether it is a metal, nonmetal, or metalloid.
Element
#34
#40
#14
#56
#18
Group
Identity
4
16
Selenium (Se)
Nonmetal
Zirconium (Z)
Silicon (Si)
Metal
Metalloid
Barium (Ba)
Argon (Ar)
Metal
Non-metal
5
5
3
4
4
2
18
3
8. Classify the following as metals, nonmetals or metalloids:
a. manganese
b. arsenic
c. carbon
d. nitrogen
Metal
Metal, non,
metalloid
Period
Metalloid
Non-metal
Non-metal
e. niobium
Metal
f. radium
Metal
9. Iodine is used in many commercial chemicals and dyes. To what family does it belong? What are the
other members of this family? How many electrons are in the outermost energy level?
Iodine belongs to the halogen family (Group 17). Other members include fluorine, chlorine, bromine, astatine, and tennessine. It
has 7 valence electrons.
10. Define atomic size and describe its trend on the periodic table.
Atomic size is the distance from the nucleus to the outermost electron. It decreases across a period and increases
down a group.
11. Define ionization energy and describe its trend on the periodic table.
Ionization energy is the energy required to remove an electron. It increases across a period and decreases down a group.
12. Define electronegativity and describe its trend on the periodic table.
Electronegativity is an atom's ability to attract electrons. It increases across a period and decreases down a group.
13. Explain why as you go down the periodic table, within a group, the atoms get larger.
Atoms get larger down a group because they gain more energy levels, increasing the distance between the nucleus and
outermost electrons.
14. Explain why Cl # 17 is smaller than S #16 considering it has one more proton than sulfur.
Chlorine is smaller than sulfur because its higher nuclear charge pulls electrons closer, despite having the same number
of energy levels.
15. As an atom gets smaller the ionization energy becomes greater. Using atomic structure, explain this
statement.
In smaller atoms, the outer electrons are closer to the nucleus and experience a stronger attraction due to the nuclear charge.
This makes it harder to remove an electron, requiring more energy.