The Electron Unit
Name______________________________
Period_______Date__________________
1.
2.
The Electromagnetic Spectrum
For hundreds of years, scientists believed that light energy was made up of tiny particles which
they call “corpuscles.” In the 1600s, researchers observed that light energy also had many
characteristics of waves. Modern scientists know that all energy is both particles, which they
call photons, and waves.
Photons travel in electromagnetic waves. These waves travel at different frequencies, but all
travel at the speed of light. The electromagnetic spectrum is the range of wave frequencies
from low frequencies (below visible light) to high frequencies (above visible light).
The radio wave category includes radio and television waves. These low-energy waves bounce
off many materials. AM waves bounce off the ionsphere and are reflected back to the Earth.
FM and television waves bounce off satellites for long-distance transmission.
Microwaves pass through some materials but are absorbed by others. In a microwave oven, the
energy passes through the glass and is absorbed by the moisture in the food. The food cooks,
but the glass container is not affected.
Like other wavelengths, infrared or heat waves are more readily absorbed by some materials
than by others. Dark materials absorb infrared waves while like materials reflect them. The
Sun emits infrared waves, heating the Earth and making plant and animal life possible.
Visible light waves are the very smallest part of the spectrum and are the only frequencies
visible to the human eye. Colors are different frequencies within this category, ranging from
red wavelengths, which are just above the visible infrared, to violet. Most of the Sun’s energy
is emitted as visible light.
The Sun also emits many ultraviolet waves. High-frequency ultraviolet wavelengths from the
Sun cause sunburn.
X-rays can penetrate muscle and tissue, but are blocked by bone, making medical and dental xray photographs possible.
Gamma-ray waves, the highest frequency waves, are more powerful than x-rays and are used
to kill cancerous cells.
The atmosphere protects Earth from dangerous ultraviolet, X-ray, and gamma-ray radiation.
Waves & the Electromagnetic Spectrum
3. The speed of light is _________________________ meters per second.
4. All waves can be described in terms of their amplitude, wavelength and ____________________.
5. The ___________________ of a wave is the number of complete waves passing a fixed point in a given
time.
6. The __________________________ __________________ is the range of wave frequencies from radio
waves to gamma waves.
7. The wavelength of microwave radiation is _______________ than the wavelength of visible light.
8. The color of visible light that has the longest wavelength is __________________.
9. A heat lamp produces _____________________ radiation.
10. A wave with a high frequency has a _______________ wavelength.
11. The brightness of light depends on the _____________________ of the light wave.
12. Microwaves are about one ________________________ long.
13. X-rays have ___________ wavelengths, ____________ frequency and ___________ energy.
14. Visible light is measured from ________nm (violet) to ________nm (red.)
15. Wavelengths are abbreviated ______ and are measured from _____________ to _____________.
16. If I lower the amplitude of a red wave, it becomes ____________ bright.
17. If I shorten the wavelength of a red wave it becomes _____________ light.
18. The relationship between frequency () and wavelength () is __________________ proportional.
Increasing Wavelength
Type of Electromagnetic
Radiation
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
Wave Description
These waves have a long wavelength
and a low frequency.
These are colors of the visible spectrum.
wavelengths between 750 nm and 400 nm.
These waves have a short wavelength
and a high frequency.
Homework I
Bohr reasoned that the obits of electrons
surrounding the nucleus must have a definite diameter.
He determined that an electron could emit energy of
one or two quanta, but not 1.5 or 3.2 quanta, as it fell
to a lower energy level.
A ball bouncing from one step to another
represents the motion of an electron as it falls from
one energy level to another. The bottom of the staircase
represents the lowest energy level in an atom.
The spectra produced by a compound can be
used to determine the elements in the compound. Each
line of the spectrum represents one frequency of light
and therefore a certain energy. This energy is determined
by the movement of electrons between energy levels
which are specific for each element. The same set of
energy levels will always produce the same spectrum.
When an electron falls from a higher to lower energy level,
a photon of light is emitted. The greater the difference in
energy levels is the greater the energy of the emitted light.
As an electron moves farther from the nucleus, it absorbs energy.
_______________________33. According to the formula c=, as the frequency gets larger, the
wavelength gets ____.
_______________________34. Which light waves contain more energy, infrared or ultraviolet?
_______________________35. Do microwave ovens produce waves with short or long wavelengths?
_______________________36. Microwaves are high-energy waves that can be used for cooking.
What equation can be used to show why a microwave has high energy?
_______________________ 37. The smallest orbit of an electron is called its ____ state.
_______________________ 38. _____________ said it is impossible to state the exact position and
momentum of a moving object.
_______________________ 39. Light travels in waves, but it acts like matter when it interacts with
small matter like an electron. This shows the ____ nature of light.
_______________________ 40. An electron must _______ energy to move to a higher energy level.
_______________________ 41. Visible radiations are less destructive because they have less ______
than UV radiation.
_______________________ 42. An electron is in a(n) ____ state when it goes from a lower energy
level to a higher energy level.
Quantum Theory & The Atom
Part I: True and False - If the statement is true write “true.” If it is false, change the underline word to
make it correct.
____________________43. Planck proposed that the energy emitted or absorbed by any object is restricted
to quanta of particular sizes.
____________________ 44. We are not aware of quantum effects in the world around us because quanta
of energy are very large.
____________________45. Light consists of protons, which are quanta of energy behaves like tiny
particles.
____________________ 46. We are constantly surrounded by low frequency X-rays.
____________________ 47. Planck’s theory relates the frequency of radiation to its energy.
____________________ 48. The dual nature of light means that light has the properties of a charge and a
wave.
Part II: Completion – Complete the following sentences.
49a. Every element has a uniquely characteristic atomic ____________________ spectrum.
49b. An electron that absorbs a quantum of energy can jump to a level of _______________ energy, called
an excited state of the atom.
49c. _________________is emitted when an electron jumps from a higher energy level to a lower energy
level.
50. The Bohr model is an inaccurate model of the atom because there is no way to measure the exact
_________________ of an electron in an atom.
51. When radiation is absorbed by a hydrogen electron, the hydrogen atom changes its ground state to an
____________________ state.
Part III: Matching – On the line at the left, write the first letter of the scientist’s name who made the
contribution to the quantum theory listed below. Each name may be used more than once.
Planck
Bohr
Heisenberg
______ 52. Stated that energy is emitted or absorbed in discrete pieces called quanta.
______ 53. Used Planck’s idea of quantization to explain the line spectrum of hydrogen.
______ 54. Stated that the position and momentum of a moving object cannot be simultaneously
measured and known exactly.
______ 55. Labeled each energy level in his atomic model with the principal quantum number, n.
______ 56. Wrote the equation E=h.
______ 57. The constant, h, is named for this scientist.
______ 58. Stated that you cannot observe or measure the “orbit” of an electron.
Part IV: Drawing
59. Label the energy levels n=1,n=2, and n=3 on the Bohr atom.
60. Label the ground state in this atom.
61. Label an excited state of this atom.
62. Draw an arrow to show the direction an electron moves when it absorbs energy.
Electron Configuration & Orbital Box Diagrams
Part I: True and False - If the statement is true write “true.” If it is false, change the underline word or
words to make it correct.
_______________ 65. The Pauli Exclusion Principle states that an orbital can hold a maximum of two
electrons.
_______________ 66. The sum of the superscripts in an electron configuration represents the total number
of neutrons in the atom or ion.
_______________ 67. The Aufbau Principle states that electrons are added one at a time to the highest
energy orbitals available until all the electrons of the atom have been accounted for.
_______________ 68. An orbital box diagram uses arrows to represent the spin of the electrons
_______________ 69. The ground state is the least stable energy state of an atom.
_______________ 70. According to Hund’s Rule, electrons occupy equal energy orbitals so that a
maximum number of unpaired electrons result.
Part II: Completion – Identify the elements (atoms) that have the following electron configurations.
Write the chemical symbol for each element in the space provided.
_________ 71. 1s22s22p3
________ 76. 1s22s22p63s23p64s23d104p5
_________ 72. 1s22s22p63s23p2
________ 77. 1s22s22p6
_________ 73. 1s22s22p63s23p64s23d10
________ 78. 1s22s22p63s23p64s23d4
_________ 74. 1s22s22p63s23p64s2
________ 79. 1s22s22p63s23p64s23d104p65s24d10
_________ 75. 1s22s22p63s23p64s23d104p65s24d105p66s24f145d6
Part III: Short Answer
80. Explain the Aufbau Principle, Pauli Exclusion Principle and Hund’s Rule when predicting an electron
configuration.
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
Part IV: Write the Electron Configuration on the first line and Orbital Box Diagram on the second line.
x
81.magnesium
2
2
6
2
1s 2s 2p 3s
______
y
z
     
1s
2s
2p
3s
82. oxygen
_____________________________
____________________________________
83. aluminum
_____________________________
____________________________________
84. argon
_____________________________
____________________________________
85. scandium
___________________________________________________________________
___________________________________________________________________
86. phosphorus
___________________________________________________________________
___________________________________________________________________
87. potassium
___________________________________________________________________
___________________________________________________________________
Part V: Write the electron configuration for each of the following elements and then write the number of
unpaired electrons each atom has. Write the electron-dot diagram around each element’s symbol.
88.
Na
_________________________________________________________________________
89.
As
_________________________________________________________________________
90.
Ir
_________________________________________________________________________
91.
Se
_________________________________________________________________________
92.
Bi
_________________________________________________________________________
93.
Hg
_________________________________________________________________________
The Atom
Part I: Completion
Pauli Exclusion Principle
electron spin
spherical
principal quantum number
quantum-mechanical model
electron density
orbital
___________________________ 94. A region in space where an electron with a particular energy is likely
to be found.
___________________________ 95. The density of an electron cloud.
___________________________ 96. Number designating the main energy level in an atom.
___________________________ 97. States that each orbital in an atom can hold at most two electrons and
that these electrons must have opposite spins.
___________________________ 98. Explains the properties of atoms by treating the electron as a wave
and quantizing its energy.
___________________________ 99. The clockwise or counterclockwise motion of an electron.
___________________________ 100. The shape of s orbitals.
Part II: Multiple Choice
______ 101. The electron cloud is least dense where the probability of finding an electron is
A. greatest
B. lowest
C. highly likely
D. nonexistent.
______ 102. The first principal energy level of the hydrogen atoms contains only a(n)
A.
B.
C.
D.
s orbital
p orbital
d orbital
f orbital
______ 103. All p orbitals are shaped like
A. spheres
B. doughnuts
C. dumbbells
D. footballs
______104. The 3s orbital differs from the 2s in that it is
A. smaller
B. larger
C. a different shape
D. more crowded
______ 105. The number of sublevels in each principal energy level equals the
A.
B.
C.
D.
mass of the atom.
electron density of the atom.
Principal quantum number for that energy level.
number of electrons in the atom.
______ 106. Which of the sublevels can be found in the fourth principal energy level of an atom?
A. s and p
B. s,p and d
C. s,p,d and f
D. s,p,d,f and g
I. Locating the Electron:
A. The maximum number of electrons possible in each principle energy level increases as
shown in the table below.
B. Not all energy levels contain all types of sublevels or orbitals.
C. Energy levels are subdivided into sublevels.
D. Each sublevel contains 1 or more orbitals. An orbital is the name given to a probability
diagram, which describes where an electron can be found.
E. Different orbitals have different shapes.
F. Complete the table below showing the number of orbitals and number of electrons in each
principal energy level.
Energy
Level
Sublevels
available
Number of
orbitals in
each sublevel
Maximum
Number of
electrons in each
sublevel
Maximum Number
of electrons in each
energy level
111.
1
s
1
2
2
112.
2
113.
3
114.
4
PROBLEMS:
A. Of the 1s, 2s and 2p sublevels:
115. In which is there no probability of an electron being found near the nucleus? _________
116. In which is there the greatest probability of finding an electron nearest the nucleus? _______
117. In which is there a different probability of finding an electron when going out in different
directions from the nucleus? ________
B. Of the energy level, sublevel or orbital:
118. Which best pinpoints the probability region in which you can find an electron? ___________
Why?
119. Which pinpoints it the least?_________________ Why?
C. From the table on the previous page, what generalizations can you make regarding:
120. The # of s orbitals in any s sublevel? _______
121. The # of p orbitals in any p sublevel? _______
122. The # of d orbitals in any d sublevel? _______
123. The # of f orbitals in any f sublevel? _______
II. Rules for Locating the Electron:
A. The rules for deciding where the electrons are in an atom are:
1. An electron will always occupy the lowest energy orbital available to it.
2. An orbital, regardless of shape, can hold a maximum of two electrons.
(Pauli Exclusion Principle)
incorrect: ↑↑
correct: ↑↓
3. Electrons will remain unpaired within equal energy orbitals as long as possible.
(Hund’s Rule)
incorrect: ↑↓ ↑ __
correct: ↑ ↑
↑
Px Py Pz
Px Py Pz
B. An electron configuration is a shorthand notation for indicating the number of electrons
in each sublevel. The notation can be broken down as shown below:
number of electrons in the sublevel
principle energy level
1s
2
sublevel (shape of orbital)
PROBLEMS:
A. For each of the following elements, state how many electrons are present in each energy level..
124. beryllium (4 electrons) ___________________________________________________
125. fluorine (9 electrons) ____________________________________________________
126. magnesium (12 electrons) ________________________________________________
127. calcium (20 electrons) ___________________________________________________
B. Write an orbital box diagram and electron configuration for each of the following elements.
128. boron (5 electrons)
129. silicon (14 electrons)
130. argon (18 electron)
131. arsenic (33 electrons)
132. sulfur (16 electrons)
133. magnesium (12 electrons)
C. Look at the following ground state orbital box diagrams, if there is a mistake, correct it.
y
x
z
    
134. nitrogen
1s
2s
2p
x
y
z
    
135. sodium
1s
2s

2p
y
z
  
3s
x
3p
z
y
x
z
       
136. phosphorus
1s
2s
2p
y
x
3s
3p
z
y
x
z
         
137. potassium
1s
2s
x
138. gallium
y
x
2p
y
3s
z
x
3p
y
4s
3d
z
              
1s
2s
2p
3s
3p
4s
3d
D. For each of the following changes of orbitals for an electron, state whether the electron would have
to absorb energy or emit energy to complete that change.
139.
3s to 3p ___________________
142.
3p to 2s ___________________
140.
3d to 4s ___________________
143.
5s to 3p ___________________
141.
5s to 4d ___________________
144.
3p to 3d
___________________
E. From the information given, and the periodic table, identify each of the following elements by name.
145. An atom of this element contains 2 electrons in the first energy level, 8 electrons in the
second energy level and 3 electrons in the third energy level. _______________________
146. An atom of this element has the following electron configuration:
1s2 2s2 2p6 3s2 3p6 4s2 3d1
147. An atom of this element has an atomic number of 16.
_______________________
_______________________
F. For each of the following atoms, state whether the electron configuration is written in the ground
state or excited state.
148. oxygen: 1s2 2s2 2p3 3s1
_________________________
149. potassium: 1s2 2s2 2p6 3s2 3p6 3d1
_________________________
150. chlorine: 1s2 2s2 2p6 3s2 3p5
_________________________
151. krypton: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 5s1
_________________________
III. Valence Electrons / Electron Dot Diagrams:
A. The outtermost energy level is given the name valance energy level. The electrons within
that energy level are known as valence electrons.
B. Determine the number of valence electrons in an element from its electron configuration.
C. There can be a maximum of eight valance electrons and they will be found in the
outtermost s and p orbitals.
PROBLEMS:
A. For each of the following elements, write the electron configurations and then determine
the number of valence electrons.
152. carbon:
153. phosphorus:
154. arsenic:
155. fluorine:
156. silicon:
157. iodine:
158. neon:
159. krypton:
B. Based on your answers to question A, which of the elements would you expect to
find in the same groups in the periodic table? Why?
C. Write the electron-dot diagram for each of the following elements:
160. strontium
165. calcium
161. oxygen
166. gallium
162. phosphorus
167. aluminum
163. iodine
168. fluorine
164. germanium
169. boron