Chapter 6: Quantum Atom and Electron Configuration
OBJECTIVES
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
Give brief account of the early atomic theory.
State the Laws of Conservation of Mass, Constant composition, and Multiple Proportions and indicate how they are
related to the atomic theory.
Summarize the essential points of Dalton's Atomic Theory.
Summarize the experiments of Thomson, Rutherford, and Milliken.
Identify the three types of particles making up atoms; give the relative charges and masses of these particles.
Relate the nuclear symbol to the number of protons and neutrons in the nucleus.
Describe the Rutherford-Bohr model of the atom.
Define atomic mass, atomic number and relate it to properties of isotopes.
Calculate average atomic mass, and abundance of isotopes.
Given one of the three quantities: wavelength ( λ ),
frequency ( ν, or f), or Ehi - Elo , calculate the other two
quantities.
Discuss the significance of the line emission spectrum of hydrogen to the model of atomic structure.
Use the Bohr Theory to calculate the energy of an electron in a given
principal energy level of the hydrogen atom,
or the difference in energy between two levels.
Explain the wave-particle duality of the electron.
List the four quantum numbers and describe their significance and meaning.
Given the atomic number of an element, or its position in the
periodic table, describe the arrangement of electrons in
the atom using orbital notation, electron configuration, or electron dot notation (Lewis
structure).
State the Aufbau principle, Hund's rule, and the Pauli Exclusion Principle.
Using electron configuration of an element, discriminate between ground state and excited state of the atom.
Describe the historical background that led to the development of the
Periodic Table.
Use examples to describe the periodic properties of elements.
Predict trends in the Periodic Table with respect to atomic radius, ionization energy, electronegativity, electronaffinity,
and metallic character.
EQUATIONS:
E = hν
c= λv
c = 3.0 x 108 m/sec
E = mc2
Rh = 2.18 x 10-16 J
E=- 1 ( 1 )
2
ΔE = Ef - Ei
Rh n
h = 6.625 x 10-34 J∙sec
ASSIGNMENTS
READING:
Early Atomic Theory: Chapter 2, Review pages 29-35
Introduction to Periodic Table: Chapter 2, pages 35-42
Electron Structure: Chapter 6, pages 144 - 166
Periodic Trends: Chapter 6, pages 166-171
WRITTEN ASSIGNMENTS:
Chapter 6: page 174
Energy, Wavelength, and frequency: questions # 2, 4, 6, 8
Bohr Model: questions # 10, 14
Energy levels, Sublevels, and Quantum Numbers: questions # 18 through 38 even
Orbital Diagrams; Hund’s Rule: questions # 40 through 48 even
Electron Arrangement in Ions: questions # 50, 52
Trends in the Periodic Table: questions 54 through 60 even
Unclassified and Conceptual Questions: page 177, questions # 64, 66, 68, 70, 74
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1
Chapter 6: Quantum Atom and Electron Configuration
Quantum Mechanics and Electron Configurations - Vocabulary (1)
Choose words from the list to fill the blanks in the paragraphs.
Classical mechanics
De Broglie
Electron configuration
Energy level
Excited state
Ground state
Heisenberg
Newton
Orbitals
Principal quantum number
Probability
Quantum mechanics
The term ____(1)___ refers to the laws of motion that
were worked out by the 17th century English scientist
____(2)______. S, A mathematical description of the
behavior of electrons inside the atom, developed during
the 1920, is called ______(3)___. This theory of electron
behavior was based in part on the theory, proposed by
________(4)____, that particles can have wavelike
properties. Then, in 1927, ____(5)_____ proposed, in his
____(6_)___ , that it is impossible to know simultaneously
both the precise location and the precise velocity of an
electron. One can state only the ___(7) ___ or chance of
finding an electron at a particular time.
Radiation
Spectrum
Uncertainty Principle
Main Principal level
1. _________________________
2.
_________________________
3. _________________________
4. _________________________
5. _________________________
6. _________________________
7. _________________________
Letters and numbers are used to identify the energy level
of an electron in an atom. The symbol n stands for the
___(8)____. Each energy level represented by n has
sublevels indicated by the letters, s, p, d, and f. Each
sublevel contains one or more ___(9)____, which are
regions of space in which an electron of a certain may be
found.
8. _________________________
9. _________________________
10. _________________________
11. _________________________
The lowest energy state of an atom is called its
____(10)_____. The outermost principal energy level is
called the ___(11)_____. Atoms are said to be in a(n)
___(12)____ when one or more electrons are in a(n)
______(13)____ higher than the ground state. When an
electron in an excited atom falls to a lower energy level,
___(14)_____ is emitted whose energy is equal to the
energy difference between the higher and lower energy
levels. When the radiation emitted is visible, a line color
will appear in the ___(15)_____ of the element. The
arrangement of the electrons in the various levels and
sublevels of an atom may be written out in the form of
a(n) _____(16)____.
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12. _________________________
13. _________________________
14. _________________________
15. _________________________
16. _________________________
2
Chapter 6: Quantum Atom and Electron Configuration
Quantum Mechanics and Electron Configurations - Vocabulary (2)
Part I: In the space at the left, write the term that correctly completes each statement.
electron cloud
Newtonian mechanics
probability
sublevel
Heisenberg uncertainty principle
orbital
quantum mechanics
wave-particle duality of nature
Lewis electron dot structure
Pauli exclusion principle
quantum numbers
_________________________1. The volume of the region in space where an electron is located is referred to as
a(n) ___.
_________________________2. That there is always some uncertainty about the position and momentum of an
electron is a statement called the ___.
_________________________3. A depiction of the outer electrons around the symbol of an element is a notation
known as a(n) ___.
_________________________4. ___ refers to the concept of the two-sided nature of waves and particles.
_________________________5. The letter n represents the ___ which describes energy level.
_________________________6. Sometimes called classical mechanics, ___ describes the behavior of visible
objects traveling at ordinary velocities.
_________________________7. The ___ states that no two electrons in an atom can have the same set of four
quantum numbers.
_________________________8. There are four __ in Schrodinger’s equation that is used in describing electron
behavior.
_________________________9. The ____ for location is the mathematical chance of finding an electron at that
point in space.
_________________________10. A(n) ___ is one of many energy states grouped closely together to form an
energy level.
_________________________11. The behavior of extremely small particles traveling at speeds near that of light is
described by ___.
_________________________12. A(n) ____ is the space occupied by one pair of electrons.
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Chapter 6: Quantum Atom and Electron Configuration
13.
What is the maximum number of electrons that can be in the
_________________
a. second energy level?
_________________
b. third energy level?
_________________
c. fourth energy level?
_________________
14.
Which quantum number signifies the size of the electron cloud?
________________
15.
_________________
16.
The sublevel shape of the electron cloud is designated by which quantum
number?
The orbital describes the direction in space of the electron cloud. Which quantum
number is used to represent the orbital?
_________________
17.
The s quantum number is used to describe the clockwise and counterclockwise
rotation of the electrons in an orbital. What two numerical values can s have?
_________________
18.
When n has the numerical value 4, what values can l have?
_________________
19.
When l = 3, what values can m have?
_________________
20.
How many orbitals are contained inn the p sublevel?
_________________
21.
How many orbitals are contained in the d sublevel?
_________________
22.
How many electrons can be in one orbital?
_________________
23.
What is the maximum number of electrons that can be in the p sublevel?
_________________
24.
What is the maximum number of electrons that can be in the f sublevel?
_________________
25.
What is the maximum number of electrons that can be in a d sublevel?
_________________
26.
What names are used for the p orbitals?
_________________
27.
What is the name of the scientist who stated that no two electrons in the same
atom can have the same set of four quantum numbers?
_________________
28.
What is the name of the scientist who pointed out that it is impossible to know
both the exact position and the exact momentum of an electron at the same time?
________________
29.
What is the name of the scientist who treated the electron mathematically as a
wave?
_________________
30.
Mechanics is a word used to describe a system of mathematical equations.
Which system of mathematical equations is used to describe the behavior of
extremely small particles traveling at velocities near the speed of light?
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Chapter 6: Quantum Atom and Electron Configuration
Energy Levels and Quantum Numbers
1.
What experimental evidence requires that the emission of energy by an atom be quantized?
2.
What kind of evidence was used by physicists and chemists in establishing the order of energy levels?
3.
What is the maximum number of electrons that may be designated as
a. 2s _______b. 2p
______ c. 3s ________
d. 3d
________ e. 4f _ _____?
4.
What type of electron orbital ( s, p, d, f ) is designated by n=2, l=1, m l =1 ? _______________
Explain why!
5.
Match the orbital properties with the correct quantum number:
Quantum Number
Orbital Properties
1)
n
______ a. distance from the nucleus
2)
l
______ b. spin
3)
m
______ c. shape of orbital
4)
s
______ d. orientation in space
______ e. orbital
______ f. size of electron cloud
______ g. probability shape
______ h. direction of orbital in space
______ i. Shape of electron cloud
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Chapter 6: Quantum Atom and Electron Configuration
1.
List the four quantum numbers for 4px1 _______________________________________
2.
Which has higher energy, an electron in the 1s or 2p sublevel? ____________________
3.
What is the highest energy sublevel in the principal energy level for which n is a. 1 ______________
b. 3
4.
What is the total capacity for electrons in
a. an orbital
c. the
5.
________________
3rd
________________
b. a d sublevel ________________
principal energy level? ___________________________
How many orbitals are there in the following sublevels?
a. s _________
6.
b. f ________________
What is the total number of orbitals in the principal energy level for which n is:
a. 1 _________________
7.
What is the total number of orbitals totally filled in Zn, at no. = 30 ________________________
Example: Ne, at. no. = 10
8
b. 4 ________________________
5 orbitals.
Write the complete electron configuration of
a. B ________________________
b. Sc
_______________________________
c. Bi ___________________________________
9.
Identify the elements which have the following outer electron configuration:
a. 3s23p3 _________
10.
c. 6s26p5 ___________d. 5s25p4 _____________
Determine the atomic number and identify the element in which there is:
a. a singe 2s electron
11.
b. 4s1 ________
____________________
b. three 2p electrons ________________
In a certain atom, the first and second principal energy levels are completely filled, as is the 3s sublevel.
There are three electrons in the 3p sublevel.
a. How many electrons are there in the atom? _____________________________
b. Give the electron configuration of the atom. _____________________________
c. What is the element’s symbol? _______________________________________
12.
Give the total number of half-filled orbitals in atoms which have the following electron configurations:
a. [Ar]4s23d3 ___________
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b. [Ar]4s13d5 _______________c. [Ar]4s23d7
______________
6
Chapter 6: Quantum Atom and Electron Configuration
Quantum Model of the Atom - Review
1.
What is the sign of E when an electron moves from a lower to a higher energy level such as n=1 to n=2?
What is the sign of E when the electron moves in the reverse direction?
2.
In the Bohr model of the hydrogen atom, what is the value of n in the ground state? What is the value of n
when the electron is completely removed from the atom?
3.
Give the total capacity for electrons of the principal energy level for which a. n= 1
c. n=3
d. n=4
e. n=5
4.
Which has higher energy, an electron in the
a. 1st or 2nd energy level?
b. 1s or 2s sublevel?
c. 2s or 2p sublevel?
d. 3d or 4d sublevel?
5.
What is the highest energy sublevel in the principal energy level for which n is a. 1
c. 3
d. f
What is the capacity for electrons in each of the following sublevels?
a. s
b. p
c. d
d. f
What is the total capacity for electrons in
a. an orbital
b. an s sublevel
c. a p orbital
d. the 1st principal energy level
e. the 3rd principal energy level?
6.
7.
8.
9.
10.
11.
12.
13.
14.
b. 2
Which of the following sublevels do not exist?
a. 2p
b. 1p
c. 2d
d. 3d
e. 3f
Arrange the following sublevels in order of increasing energy:
1s, 3p, 2s, 2p, 3s
In the third principal energy level, what is the total electron capacity? How many different sublevels are in
this level? List them in order of increasing energy. One of these is actually higher in energy than the lowest
sublevel in the 4th principal energy level. Which one is it?
Draw an 1s orbital; a 2s orbital; a 2p orbital. how do you interpret the shapes of these orbitals?
How many orbitals are there in the following sublevels?
a. s
b. p
c. d
d. f
What is the total number of orbitals in the principal energy level for which n is: a. 1
b. 2
c. 3
d. 4
What is the total number of orbitals totally filled in
a. Ar, at no. = 18
b. Zn, at no. = 30
Example: Ne, at. no. = 10
15.
b. n=2
5 orbitals.
Give the outer electron configuration of the elements in the following groups of the Periodic Table:
a. 3
b. 5
c. 17
d. 18
e. 13
f. 8
Example: Group 2: ns2
16.
17.
18.
19
Which group in the Periodic Table has the outer configuration:
a. ns2
b. ns2np2
c ns2np5
Give the total number of outermost level electrons (valence electrons) of an element in Group a. 1
b. 13
c. 14
d. 16
e. 17
In what group of the Periodic Table do all the elements have:
a. 2 valence electrons
b. 5 valence electrons
c. 6 valence electrons?
Using Bohr’s equation calculate E, in Kcal/mol, for the transition ( 1 Kcal = 4.18 kJ) a. n=1 to n=3
b. n = 5 to n = 2
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Chapter 6: Quantum Atom and Electron Configuration
20
21
Write the electron configuration of
a. B
b. Ne
c. Al
g. Fe
h. Zn
I. As
d. Cl
j. I
e. Ca
k. Sn
f. Sc
22.
Identify the elements which have the following outer electron configuration:
a. 3s23p3
b. 4s1
c. 6s26p5
d. 5s25p4
Determine the atomic number of the element in which there is:
a. a singe 2s electron
b. two 2p electrons
c. five 3d and two 4s electrons
d. an equal number of 4s and 3d electrons
23.
Give the complete electron configuration of atom which have the following “condensed configurations”.
[Ne} 3s13p1
b. [Ar] 4s1 3d5
24.
In a certain atom, the first and second principal energy levels are completely filled, as is the 3s sublevel.
There are three electrons in the 3p sublevel.
a. How many electrons are there in the atom?
b. Give the electron configuration of the atom.
c. What is the element’s symbol?
25.
a. Which of the following configurations correspond to an atom in its ground state? b. Which would be
formed if one or more electrons were excited to higher sublevels?
c. Which are impossible?
(1) 1s22s22p4
(4) 1s22s12p2
(2) 1s22s22p32d1
(5) 1s22s22p1
(3) 1s22s22p33s1
(6) 1s22s3
26.
Give the number of electrons in each of the 3p orbitals in each of the following elements:
a. Al
b. Br
c. Ca
d. O
e. Ar
f. Ti
g. Fe
h. Sb
I. Ge
j. C
27.
Give the total number of half-filled orbitals in atoms which have the following electron configurations:
[Ar]4s23d3
b. [Ar]4s13d5
c. [Ar]4s23d7
1
List the four quantum numbers for 5px
28.
a.
a.
Answers:
1. (+), (-) 2. n = 1, n =
3. A) 2; b) 8; c) 18; d) 32; 3) 50
4. A) 2nd b. 2s; c)2p; d) 4d
5. a.s; b. p;
c. d; d. f
6. A)2; b) 6 ; c) 10; d) 14
7. A) 2; b) 2; c) 6; d) 2
8. 1p, 2d, 3f;
9. 1s,2s, 2p, 3s, 3p
10. 18; 3; s, p, d; 3d
11. o, O, ; probability.
12. a) 1; b) 3; c) 5; d) 7
13. A) 1; b) 4; c) 9; d) 16
14. Ar - 9 orbitals; b. Zn - 15
15. a) ns2(n-1)d1; b) ns2(n-1)d3; c) ns2np5; d) ns2np6; e) ns2np1; f) ns5(n-1)d6
16. a)2; b) 14; c) 17
17. A) 1; b) 3; c) 4; d) 6; e) 7
18. A) 2; b) 15; c) 16
19. a) +347 Kcal/mol; b) -65.52 kcal/mol
20. Check notes
21. a) P; b) K; c) At; d) Te
22. A) Li; b) C; c) Mn; d) Te
23. Check notes
24. a. 15 b. 1s22s22p63s23p3 c. P 25. Impossible - (2), (6); ground - (1) (5); excited - (3), (4)
26. Al -1; Br-6; Ca-6; O-0; Ar - 6; Ti-6; Fe-6; Sb-6; Ge-6; C-0
27. a. - 3 b. 6
c. 3
28. n=5, l= 1, m= +1; s= ½
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Chapter 6: Quantum Atom and Electron Configuration
Answer Sheet
Page 2
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
classical mechanics
Newton
Quantum mechanics
de Broglie
Heisenberg
Uncertainty Principle
Probability
Principal quantum #
Orbitals
Ground state
Principal energy level
Excited state
Energy level
Radiation
Spectrum
Electron config.
Pages 3, 4
1. electron cloud
2. Heisenberg uncert.
3. Lewis-dot diagram
4. Wave-particle duality
5. Principal quantum #
6. Newtonian mechanics
7. Pauli exclusion principle
8. Quantum numbers
9. Probability
10. sublevel
11. quantum mechanics
12. orbital
13. 8, 18, 32
14. n
15. l
16. m
17.  1/2
18. 0, 1, 2,3
19.  3,  2,  1, 0
20. 20. 3
21. 5
22. 2
23. 6
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24.
25.
26.
27.
28.
29.
30.
14
10
pxpypz
Pauli
Heisenberg
Schroedinger
Wave mechanics
Page 5
3a. 2
b. 6
c. 2
d. 10
e.
4. 2px
5 a. n
b. s
c. l
d. m
e. m
f. n
g. l
h. m
i. l
Page 6
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
n = 4; l = 1; m= -1
2p
a) s; b) d
a) 2; b) 10; c) 18
a) 1; b) 7
a) 1; b) 16
15
a) 1s22s22p1
b) 1s22s22p6 3s23p64s23d1
c) 1s22s22p6 3s23p64s23d104p6
4d104f145s25p65d106s26p3
a) P; b) K; c) At; d) Te
a) li; b) N
a) 15; b) 1s22s22p6 3s23p3 c) p
a) 3; b) 5; c) 3
9