Periodic Table discussion

1. PERIODIC TABLE OF ELEMENTS
and QUANTUM MECHANICS
Kristine P. Sumalinog
Teacher III
Babag National High School – Senior High School

2. Lesson Objectives
By the end of the lesson, students are expected to:
● Identify elements on the periodic table;
● Know the sub-groups of the periodic table;
● Describe the trends within the periodic table;
● Identify the atomic symbol, atomic number and the atomic mass of each
element
● Use quantum numbers to describe an electron in an atom;
● Determine the magnetic property of the atom based on its electronic
configuration;
● Filled in an empty Periodic Table correctly; and
● Draw an orbital diagram to represent the electronic configuration of atoms.

3. Serenity Prayer
God, grant me the serenity
To accept the things I cannot change;
Courage to change the things I can;
And wisdom to know the difference.
Living one day at a time;
Enjoying one moment at a time;
Accepting hardships as the pathway to peace;
Taking, as He did, this sinful world
As it is, not as I would have it;
Trusting that He will make things right
If I surrender to His Will;
So that I may be reasonably happy in this life
And supremely happy with Him
Forever and ever in the next.
Amen.

4. Attendance Check
• District 1
• District 2
• District 3
• District 4
• District 5
•District 6
• District 7
• District 8
• District 9
• District 10

6. Review
•What is matter?
•What are the different phases of matter?

7. Introductory Activity
•Classifying Matters
Instruction: Classify the following as a
mixture or pure substance. If its a mixture,
identify if its a solution or a heterogeneous
mixture. And if its a pure substance, identify
if its an atom, molecule, or compound.

9. Pure substance
Molecule
Mixture
Heterogeneous Mixture
Pure substance
Compound
Pure substance
Atom
Mixture
Heterogeneous Mixture
Mixture
Solution
Mixture
Heterogeneous Mixture
Mixture
Heterogeneous Mixture
Mixture
Heterogeneous Mixture
Mixture
Heterogeneous Mixture
Pure substance
Atom
Pure substance
Molecule

10. Processing question:
•Can you explain the differences of
atoms, compounds and molecules?

11. Activity
Collaborative Approach,
Inquiry Approach

12. •1. The teacher has divided the class into
groups with three (3) members each.
•2. Per group will be given a box that
contained all the materials they may
need (Worksheets, cartolina, masking
tape, scissors, glues, and a clear copy of
the Periodic table)
Activity

13. •3.They will be given 15 minutes to fill in the
blank periodic table.
•4. Each group will be given 3 minutes to
present their completed periodic table and
make a short discussion of the concepts
about elements and their physical
properties.

15. Lab Report Rubrics
Team Number
1 2 3 4 5 6 7 8
Scientific
accuracy
Aesthetic
appearance
Teamwork
Total
3 = excellent 2 = good 1 = satisfactory

18. Analysis
•Based on the activity, what big
concepts can you draw out?

19. The
History
of the
Periodic
Table

20. 1829
Johann Dobereiner
• He classified some elements into
groups of three (triads). The
elements in a triad had similar
chemical properties and orderly
physical properties - model of
triads

21. 1863
John Newlands
• In 1863, he suggested that
elements be arranged in
octaves because he noticed
(after ordering the elements in
increasing atomic mass) that
certain properties repeated
every 8th element - Law of
octaves.

22. 1869
Lothar Meyer
• By using atomic weights, Meyer
was able to arrange 28 elements
into 6 families that held similar
chemical and physical
characteristics. He was able to
create this version of the Periodic
Table by listing the elements by
their valence, which is the
combining power of the element.

23. 1869
Dmitri Mendeleev
• Mendeleev’s first Periodic Table
arranged the elements in order
of increasing atomic mass and
grouped them by similarity of
properties. When a gap existed
in the table, Mendeleev
predicted a new element would
one day be discovered and
assumed its properties.
Mendeleev arranged the 63
known elements into a Periodic
Table and recorded it in his book,
Principles of Chemistry, in 1869.

24. 1913
Henry Moseley
• In 1913, through his work with X-
rays, he determined the actual
nuclear charge (atomic number) of
the elements. Moseley’s discoveries
produced a more accurate
positioning of the elements in the
Periodic Table. Moseley reevaluated
Mendeleev’s Periodic Table and
made one important modification
before it became the modern
Periodic Table- the use of atomic
number as the organizing principal
for the periods.

25. Parts of
the
Periodic
Table of
Elements

26. Groups
• Groups are the columns of the table. Atoms of
elements within a group have the same number of
valence electrons. These elements share many similar
properties and tend to act the same way as each other in
chemical reactions.

27. Periods • The rows in the periodic table are called periods. Atoms of these elements all
share the same highest electron energy level.

28. Metals
• With the exception of
hydrogen, the elements on the
left-hand side of the periodic
table are metals. The two rows
of elements below the body of
the periodic table are metals.
Specifically, they are a
collection of transition metals
that are called the lanthanides
and actinides or the rare earth
metals.

29. Metalloids
• There is a zig-zag line toward
the right side of the periodic
table that acts as a sort of
border between metals and
nonmetals. Elements on either
side of this line exhibit some
properties of metals and some
of the nonmetals. These
elements are the metalloids,
also called semimetals.

30. Nonmetals
• The elements on the right-
hand side of the periodic table
are the nonmetals. Nonmetals
properties are:
 usually poor conductors of
heat and electricity
 often liquids or gases at
room temperature and
pressure
 lack metallic luster
 readily gain electrons (high
electron affinity)
 high ionization energy

31. Periodic
Trends of
the
Periodic
Table of
Elements

32. Electronegativity Trends Electronegativity can be
understood as a chemical property describing an atom's ability to
attract and bind with electrons.

33. Electronegativity Trends
• Na 1s2 2s2 2p6 3s1 It is more energy efficient for sodium to lose an
electron rather than add seven more electrons to its outermost shell
• Na+ 1s2 2s2 2p6 Sodium ion complies with the octet rule.
• Fluorine 1s2 2s2 2p5 For fluorine, the atom needs only one electron to
comply with the octet rule.
• F- 1s2 2s2 2p6 Fluorine ion complies with the octet rule.

34. Ionization Energy Trends Ionization energy is the
energy required to remove an electron from a neutral atom in its
gaseous phase.

35. Electron Affinity Trends Electron affinity is
the ability of an atom to accept an electron.

36. Atomic Radius Trends The atomic radius is one-half
the distance between the nuclei of two atoms (just like a radius
is half the diameter of a circle).

37. Metallic Character Trends The metallic character
of an element can be defined as how readily an atom can lose an
electron.

38. Elements

39. ELEMENTS
 Are the simplest kind of pure substance.
 There are 118 elements, each made of a
different type of atom (periodic table).

40.  Each
element has
a name and
symbol.

41. Not all elements are equally common.

42. Most elements can exist as
a single atom (Fe, Cu):
But some elements
(mostly gases) usually
exist as diatomic
molecules (groups of 2
atoms).
Diatomic Gases
Hydrogen H2
Nitrogen N2
Oxygen O2
Fluorine F2
Chlorine Cl2
Bromine Br2
Iodine I2

43. Allotropes
Different structural forms of the same element.
O2
Diatomic
Oxygen
Molecule
O
Monatomic
Oxygen
(Single Oxygen
Atom)
O3
Ozone
Molecule
Oxygen has 3 allotropes:

44. Allotropes of Carbon All are pure carbon. Each has a different structure.

45. Atomic
Particles

46. Nucleus
• In the center of each atom is the nucleus. Within
the nucleus there are two kinds of particles. Positively-charged
particles called protons and particles with no charges
called neutrons. The protons give the nucleus a positive charge. For
example, a helium atom has 2 protons and 2 neutrons. It would have
a net charge of +2. A carbon atom has 6 protons and 6 neutrons and
a net charge of +6.

47. Protons
• The number of protons in an atom is unique to each element.
The number of protons in an atom is referred to as the atomic
number of that element.

48. Neutrons
• When you remove or add a neutron to the nucleus of an atom,
the resulting substance is a new type of the same element and
is called an isotope.

49. Electrons
• Moving around the nucleus are tiny, negatively-charged
particles called electrons. These particles are 2 000 times
smaller than protons and neutrons.

50. Identifying
Sub-
atomic
Particles

54. Let’s complete the table. Use your periodic table to
complete the table.

55. Answers

56. Electron
Configuration

57. 1
2
3
4
5
6
7
6
7
1A
2A
3B 4B 5B 6B 7B 8B 8B 8B 1B 2B
3A 4A 5A 6A 7A
8A
group # = # valence (outside) e-
d p
f
s
Row
=
# shells

58. Electron Configuration
1s1
row #
shell #
possibilities are 1-7
7 rows
subshell
possibilities are
s, p, d, or f
4 subshells
group #
# valence e-
possibilities are:
s: 1 or 2
p: 1-6
d: 1-10
f: 1-14
Total e- should equal
Atomic #
What element has an electron configuration of 1s1?

59. Practice:
Ask these questions every time you have to write an electron
configuration
• Lithium:
1. find the element on the periodic table
2. what is the period number?
3. how many shells?
4. what is the group number?
5. how many valence electrons?
6. what subshell(s) does Li have?
7. what is the electron configuration?
atomic # = 3
2
2
1
1
s
1s2 2s1

60. Practice:
Ask these questions every time you have to write an electron
configuration
• Boron:
1. find the element on the periodic table
2. what is the row #?
3. how many shells?
4. what is the group #?
5. how many valence electrons?
6. what subshell(s) does B have?
7. what is the electron configuration?
atomic # = 5
2
2
3
3
p
1s2 2s2 2p1

61. Order of Electron Subshell Filling:
It does not go “in order”
1s2
2s2 2p6
3p6
4p6
5p6
6p6
7p6
3s2
4s2
5s2
6s2
7s2
3d10
4d10
5d10
6d10
4f14
5f14
1s2 2s2 2p6 3p6
3s2 4s2 4p6 5s2
3d10 5p6 6s2
4d10 6p6 7s2
5d10
4f14
7p6
6d10
5f14

62. 1
2
3
4
5
6
7
6
7
period
#
=
#
e-
shells
1A
2A
3B 4B 5B 6B 7B 8B 8B 8B 1B 2B
3A 4A 5A 6A 7A
8A
group # = # valence e-
d
f
3d
4d
5d
6d
4f
5f
Subshells d and f are “special”

65. Activity
Collaborative Approach,
Inquiry Approach

66. •1. The teacher has divided the class into
groups with three (3) members each.
•2. Per group will be given a box that
contained all the materials they may
need (Worksheets, cartolina, masking
tape, scissors, glues, and a clear copy of
the Periodic table)
Activity

67. •3.They will be given 15 minutes to
answer the worksheet entitled
“Electron Configuration PRACTICE”.
•4. Each group will be given 3 minutes
to present their completed worksheet
and make a short discussion of the
concepts about electron
configurations and orbital diagrams.

69. Lab Report Rubrics
Team Number
1 2 3 4 5 6 7 8
Scientific
accuracy
Aesthetic
appearance
Teamwork
Total
3 = excellent 2 = good 1 = satisfactory

73. Atomic
Theory
Changes over
time…

74. Atomic
Theory
• Because we can not see atoms, we use models to
teach and learn about atoms.
• The atomic theory has changed over time as new
technologies have become available.
• Remember: Scientific knowledge
builds on past research and
experimentation.

76. Who are these men?
In this part of the lesson, we’ll
learn about the men whose
quests for knowledge about the
fundamental nature of the
universe helped define our
views of the atom.

77. Democritus
• This is the Greek philosopher
Democritus who began the
search for a description of
matter more than 2400 years
ago.
• He asked: Could matter be
divided into smaller and
smaller pieces forever, or
was there a limit to the
number of times a piece of
matter could be divided?
400 BC

78. Atomos
•His theory: Matter could not be
divided into smaller and smaller
pieces forever, eventually the
smallest possible piece would be
obtained.
•This piece would be indivisible.
•He named the smallest piece of
matter “atomos,” meaning “not to
be cut.”

79. Because….
•The eminent
philosophers of the
time, Aristotle and
Plato, had a more
respected, (and
ultimately wrong)
theory.
Aristotle and Plato favored the earth, fire,
air and water approach to the nature of
matter. Their ideas held sway because of
their eminence as philosophers. The
atomos idea was buried for approximately
2000 years.

80. Dalton’s Model
• In the early 1800s, the
English Chemist John
Dalton performed a
number of experiments
that eventually led to the
acceptance of the idea of
atoms.

81. Dalton’s Atomic Theory
• He deduced that all elements are
composed of atoms. Atoms are
indivisible and indestructible particles.
• Atoms of the same element are
exactly alike.
• Atoms of different elements are
different.
• Compounds are formed by the joining
of atoms of two or more elements.

82. Thomson’s Plum Pudding Model
•In 1897, the English
scientist J.J. Thomson
provided the first hint
that an atom is made of
even smaller particles.

83. Thomson Model
• He proposed a model of the atom
that is sometimes called the “Plum
Pudding” model.
• Atoms were made from a
positively charged substance with
negatively charged electrons
scattered about, like raisins in a
pudding.

84. Thomson Model
•Thomson studied the
passage of an electric
current through a gas.
•As the current passed
through the gas, it gave off
rays of negatively charged
particles.
•Thomson called the
negatively charged
“corpuscles,” today known
as electrons.

85. Rutherford’s Gold Foil Experiment
• In 1908, the English physicist
Ernest Rutherford was hard
at work on an experiment
that seemed to have little to
do with unraveling the
mysteries of the atomic
structure.

86. • Most of the positively
charged “bullets” passed
right through the gold
atoms in the sheet of gold
foil without changing
course at all.
• Some of the positively
charged “bullets,” however,
did bounce away from the
gold sheet as if they had hit
something solid. He knew
that positive charges repel
positive charges.

89. Rutherford
• Rutherford reasoned that all of
an atom’s positively charged
particles were contained in the
nucleus. The negatively charged
particles were scattered
outside the nucleus around the
atom’s edge.

90. Bohr Model
• In 1913, the Danish scientist
Niels Bohr proposed an
improvement. In his model,
he placed each electron in a
specific energy level.

91. Bohr Model
• According to Bohr’s atomic
model, electrons move in
definite orbits around the
nucleus, much like planets
circle the sun. These orbits,
or energy levels, are
located at certain distances
from the nucleus.

92. The Wave Model
• Today’s atomic model is
based on the principles of
wave mechanics.
• According to the theory of
wave mechanics, electrons
DO NOT move about an
atom in a definite path, like
the planets around the
sun.

93. The Wave
Model
• In fact, it is impossible to determine the
exact location of an electron. The probable
location of an electron is based on how
much energy the electron has.
• According to the modern atomic model, an
atom has a small positively charged
nucleus surrounded by a large region in
which there are enough electrons to make
an atom neutral.

94. Electron Cloud:
• A space in which electrons
are likely to be found.
• Electrons whirl about the
nucleus billions of times in
one second
• They are not moving around
in random patterns.
• Location of electrons
depends upon how much
energy the electron has.

95. Indivisible Electron Nucleus Orbit Electron
Cloud
Greek X
Dalton X
Thomson X
Rutherford X X
Bohr X X X
Wave X X X

96. Quantum
Mechanical
Model of
the Atom

97. The Quantum Mechanical Model of the Atom
• The approach they developed became known as wave mechanics or,
more commonly, quantum mechanics.
Werner Heisenberg
(1901–1976)
Louis de Broglie
(1892–1987)
Erwin Schrödinger
(1887–1961)

98. Bohr
Model
• Bohr’s model gave hydrogen atom
energy levels consistent with the
hydrogen emission spectrum.
• Ground state – lowest possible
energy state (n = 1)
• Bohr’s model is incorrect. This
model only works for hydrogen.
• Electrons DO NOT move around
the nucleus in circular orbits.

99. Schrödinger's Theory
• He agreed that electrons have a
specific amount of energy
• He believed that the distance between
rungs on the ladder were not
consistent – they get closer together as
you move higher up
• Quantum – the amount of energy
needed to move from one energy level
to another
The electrons move in regions of
probability around the nucleus called
ORBITALS

100. Quantum theory, also called
wave mechanics, describes
the arrangement and space
occupied by electrons.
Orbitals refers to the three-
dimensional regions in space
where there is a high
probability of finding an
electron around an atom.

101. Characteristics of
Electrons
• Extremely small mass
• Located outside the
nucleus
• Moving at extremely
high speeds in a sphere
• Have specific energy
levels

102. Energy of
Electrons
• When atoms are heated, bright
lines appear called line spectra
• Electrons in atoms arranged in
discrete levels.
• An electron absorbs energy to
“jump” to a higher energy level.
• When an electron falls to a
lower energy level, energy is
emitted.

103. The electrons move in
regions of probability around
the nucleus called ORBITALS
Defining these ORBITALS:
• Quantum Numbers are
used to define:
• The energy of the electron
• The electron’s relative
distance from the nucleus
• The size and shape of the
ORBITAL
• The pairings of the
electrons

104. Quantum Numbers
Principle Quantum Number (n) – define the
energy of the electron
n=1 is closest to the nucleus – low energy
n=2 is farther than n=1, slightly more energy
n=3 is farther than n=1 and n=2, still increasing in
energy
n=4 …..
Remember – The difference in energy between
energy levels decreases as “n” increases

105.  The angular momentum quantum number (ℓ) has
integral values from 0 to n-1 for each value of n. This
quantum number is related to the shape of atomic
orbitals. The value of ℓ for a particular orbital is
commonly assigned a letter: ℓ = 0 is called s; ℓ = 1 is called
p; ℓ = 2 is called d; ℓ = 3 is called f. This system arises from
early spectral studies and is summarized in the table
below.
The Angular Momentum Quantum Numbers and
Corresponding Letters Used to Designate Atomic Orbitals
Value of ℓ 0 1 2 3 4
Letter
Used
s p d f g

106. • The magnetic quantum number (mℓ) has integral values
between ℓ and -ℓ, including zero. The value of mℓ is
related to the orientation of the orbital in space relative
to the other orbitals in the atom.
Quantum Numbers for the First Four Levels of Orbitals in the Hydrogen Atom
n ℓ Orbital Designation mℓ
Number of
Orbitals
1 0 1s 0 1
2
0 2s 0 1
1 2p -1, 0, 1 3
3
0 3s 0 1
1 3p -1, 0, 1 3
2 3d -2, -1, 0, 1, 2 5
4
0 4s 0 1
1 4p -1, 0, 1 3
2 4d -2, -1, 0, 1, 2 5
3 4f -3, -2, -1, 0, 1, 2, 3 7

107. Electron Spin and the Pauli Principle
• Spinning in one direction, the
electron produces the magnetic
field oriented as shown in (a).
Spinning in the opposite direction,
it gives a magnetic field of the
opposite orientation, as shown in
(b).

108. Pauli exclusion principle
• In a given atom no two electrons can have the
same set of four quantum numbers (n, ℓ, mℓ ,
and ms). This is called the Pauli exclusion
principle. Since electrons in the same orbital
have the same values of n, ℓ, and mℓ this
postulate says that they must have different
values of ms. Then, since only two values of ms
are allowed, an orbital can hold only two
electrons, and they must have opposite spins.
Wolfgang Pauli
(1900–1958)

109. Orbital Shapes and
Energies
• Two representations of the hydrogen
1s, 2s, and 3s orbitals. (a) The
electron probability distribution. (b)
The surface that contains 90% of the
total electron probability (the size of
the orbital, by definition).

110. Representation of the 2p orbitals. (a) The electron probability distribution for
a 2p orbital. (Generated from a program by Robert Allendoerfer on Project
SERAPHIM disk PC 2402; reprinted with permission.) (b) The boundary
surface representations of all three 2p orbitals. Note that the signs inside the
surface indicate the phases (signs) of the orbital in that region of space.

111. Representation of the 3d orbitals. (a) Electron density plots of selected 3d orbitals. (b) The boundary surfaces of all five 3d
orbitals, with the signs (phases) indicated.

112. Representation of the 4f orbitals in terms of their boundary surfaces.

113. Orbital energy levels for the hydrogen atom.

114. Processing questions
What would happen if all 118 elements
were not categorized? Justify your
answer.

115. Processing
questions
Why was
Mendelev’s
Periodic table
considered to be
the most valid of
all proposed
periodic tables?

116. Processing
questions
How do you relate the
quantum numbers to the
position of an atom in
the periodic table?

117. Quiz:
The table below shows various combinations of the three quantum numbers.
Indicate which combinations are allowed and which are disallowed. For those that
are allowed, give the subshell notation (e.g., 2p).

118. Solutions:

119. Homework:

120. Concluding Activity
•Clean up the laboratory

121. Reflection