Download The Periodic Table of Elements

THE PERIODIC TABLE
How’d They Come Up With That?



Our current society takes for granted all of
the hard work, research, chance,
and luck that has gone into creating
and discovering the materials that are used
in the products we utilize every day.
For example, who was the first person to set
or find a random black rock (coal) on fire
and discover that it provided a good,
constant source of heat?
Who was the first person to discover that a
substance found in some rocks was capable
of being the ultimate explosive (uranium)?
Coal
Uranium
Organizing the elements
In nature and in the lab we have
discovered over 100 different
elements.
We’ve organized the elements into
a table based on their PHYSICAL
and CHEMICAL PROPERTIES
It took us almost 2000 years to
figure out the properties of the
elements currently in the Periodic
Table of Elements and arrange
them.
I. Developing the Periodic Table

In the early 1800s,
enough information was
known about the
elements that scientists
wanted an easy way to
categorize the Earth’s
ingredients.
So why don’t we just list the elements in a
straight line? Why make a table?


An analogy: When you go to the grocery
store and are looking for a specific
ingredient, you want things to be grouped
in a logical way so that you can easily find
what you need.
Scientists wanted the elements that they
were using in their experiments to also be
grouped so that they could easily find
elements with the characteristics that they
desired.
Developing the Periodic Table

Many methods of
organization were
tried before scientists
found the most
effective way of
grouping the
elements
“Mayan”
Periodic
Table
The Who’s Who of Periodic Table Development
Johann Dobereiner

In 1829, he classified some elements into groups of three,
which he called triads. The elements in a triad had similar
chemical properties and orderly physical properties.
Model of triads
1780 - 1849
(ex. Cl, Br, I and Ca, Sr, Ba)
John Newlands

In 1863, he suggested that elements be arranged in “octaves”
because he noticed (after arranging the elements in order of
increasing atomic mass) that certain properties repeated every
8th element.
Law of Octaves
1838-1898
Dmitri Mendeleev

In 1869 he published a table of the elements
organized by increasing atomic
1834-1907
mass.
Developing the Periodic Table



Mendeleev stated that
the properties of
elements are a periodic
function of their atomic
masses.
If we lay out all of the
elements in order by
atomic mass, we would
see the same properties
showing up periodically
For example, as you would begin
down the list of the first 18
elements, every 9th (or so) atom
would be super reactive and want
to gain 1 electron
Lothar Meyer

At the same time, he published his own table of
the elements organized by increasing
atomic mass.
1830-1895
•Both Mendeleev and Meyer
arranged the elements in order of
increasing atomic mass.
•Both left vacant spaces where
unknown elements should fit.
•So why is Mendeleev called the
“Father of the Modern Periodic
Table” and not Meyer, or both?
Father of the Modern Periodic Table
Mendeleev …
• used his periodic function theory
and corrected the atomic masses of
Be, In, and U.
• was so confident in his table that
he used it to predict the physical
properties of three elements that
were yet unknown (Sc, Ga & Ge).
• After the discovery of these
unknown elements between 1874
and 1885, and the fact that
Mendeleev’s predictions were
amazingly close to the actual
values, his table was generally
accepted.
Mendeleev’s Periodic Table
Mendeleev’s Periodic Table


However, in spite of
Mendeleev’s great
achievement, problems
arose when new elements
were discovered and more
accurate atomic weights
determined.
By looking at our modern
periodic table, we can
identify some problems
which might have caused
chemists a headache.
Atomic Masses
•
•
•
•
Ar and K
Co and Ni
Te and I
Th and Pa
(40-39)
(59-58)
(128-127)
(232-231)
Henry Moseley

In 1913, through his work with X-rays, he determined
the actual nuclear charge (atomic number) of the
elements*. He rearranged the elements in order of
increasing atomic number.
*There is in the atom a fundamental quantity
which increases by regular steps as we pass from
each element to the next. This quantity can only
be the charge on the central positive nucleus.
His research was halted when the British
government sent him to serve as a foot soldier in
WWI. He was killed in the fighting in Gallipoli by
a sniper’s bullet, at the age of 28. Because of this
loss, the British government later restricted its
scientists to noncombatant duties during WWII.
1887-1915
Periodic Law

This arrangement allowed elements
with similar properties to be lined
up into the same column in the
periodic table
 Ex.
All of the elements in group 1 are
extremely reactive when by themselves
and want to give away 1 electron
Glenn T. Seaborg


After co-discovering 10 new elements, in 1944 he
moved 14 elements out of the main body of the
periodic table to their current location below the
Lanthanide series.
These became known as the
Actinide series.
1912-1999
Glenn T. Seaborg

He is the only person to have an element named
after him while still alive.
"This is the greatest honor ever
bestowed upon me - even
better, I think, than winning the
Nobel Prize."
1912-1999
Video - Review

Periodic Table History (6:33)
QOD

Mendeleev
published a table
of the elements
organized by
increasing what?
Atomic Mass

Which scientist
arranged
elements on the
periodic table
according to
atomic number?
Henry Moseley
PERIODIC
TABLE
GEOGRAPHY
Where are elements found?
The horizontal rows of the periodic table
are called PERIODS.
The elements in any group of
the periodic table have
similar physical and chemical
properties!
The vertical columns of the periodic table
are called GROUPS, or FAMILIES.
Organizing Information on
the Periodic Table
Highlight/darken the stair step line
starting between B and Al.
 Label the left side: metals
 Label the right side: nonmetals
 Shade B, Si, Ge, As, Sb, Te, At and key as
METALLOID.

General Properties of Metals
 Metals are good conductors of
heat and electricity
 Metals are malleable
oEasily bent and shaped
 Metals are ductile
oAble to be drawn out into a
wire
Metals have a shiny luster
Examples of Metals
Copper, Cu, is a relatively soft metal,
and a very good electrical conductor.
Zinc, Zn, is more
stable than
potassium
Mercury, Hg, is the only metal
that exists as a liquid at room
temperature
General Properties of
Metalloids




They have properties of both
metals and nonmetals.
Metalloids are more brittle
than metals, less brittle than
most nonmetallic solids
Metalloids are semiconductors
of electricity
Some metalloids possess
metallic luster
Ex: Silicon possesses a metallic luster, yet
it is an inefficient conductor and is brittle.
General Properties of Nonmetals




Nonmetals are poor conductors
of heat and electricity
Good insulators
Nonmetals tend to be brittle
when solids
Many nonmetals are gases at
room temperature
Carbon, the graphite in “pencil lead” is a great
example of a nonmetallic element.
Examples of Nonmetals
Sulfur, S, was
once known as
“brimstone”
Graphite is not the only pure
form of carbon, C. Diamond is
also carbon; the color comes
from impurities caught within
the crystal structure
Microspheres of
phosphorus, P, a
reactive
nonmetal
12
Mg
4
19
K
5
37
Rb
6
55
Cs
7
87
Fr
20
Ca
5
B
3
4
5
6
9
M
13
Al
6
C
7
N
8
O
14
Si
15
P
16
S
7
8
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga
32
Ge
33
As
34
Se
E
10 11 12
9
F
10
Ne
17
Cl
18
Ar
21
Sc
22
Ti
23
V
24
Cr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
56
Ba
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
88
Ra
104
Rf
105
Db
106
Sg
107
Bh
109
Mt
110
Uun
111
Uuu
112
Uub
57
La
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
89
Ac
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
38
Sr
Transition Metals
Transition Metals
Transition Metals
10
8
Transition
Hs
Metals
Lanthanides
Actinides
LO
T
I
A
D
35
Br
Halogens
3
11
Na
13 14 15 16 17
Chalcogens
P
E
R
I
O
D
4
Be
Alkali Earth Metals
2
3
Li
2
He
Write the group names!
2
Alkali Metals
1
1
H
18
The Periodic Table of Elements
L
36
Kr
Noble Gases
1
Alkali Metals
Alkaline Earth Metals
Transition Metals
These elements are also called
the rare-earth elements.
Notice that it
continues from
period 6 & 7
InnerTransition Metals
Halogens
Noble Gases
Last slide
QOD

More than twothirds of the
elements on the
periodic table are
what?
Metals

If solid, non metals
tend to be what?
Good
insulators
III. Periodic Law

Before we can discuss the characteristics of
each group, we must introduce the Period Law.
Periodic Law

When elements are
arranged in order of
increasing atomic
number, there is a
periodic pattern in
their physical and
chemical properties.
Periodic Table Review
Elements with similar e- configurations are placed
in the same group. (same number of e- in last
energy level)
 The number of electrons in the outermost energy
level determines how reactive an element will be
 These electrons in the highest energy level are
called “Valence Electrons”

 Therefore
elements in the same group have the same
number of valence e-
Group 1


EX:
H = 1s1
Li = 1s22s1
Na = 1s22s22p63s1
All elements in group 1 have 1 valence
electron (1 electron in highest energy level)
Group 17

Ex:
F = 1s22s22p5
Cl = 1s22s22p63s23p5
Br = 1s22s22p63s23p64s23d104p5

All elements in group 17 have 7 valence
electrons (7 electrons in highest energy level)
Valence Electrons

Highlight the valence electrons on your periodic table
Octet Rule


Octet Rule = Eight electrons in an outer level render
an atom unreactive (stable). Therefore all atoms are
trying to get to this state.
Atoms obtain an octet of electrons in their valence
orbital by
 taking
electrons from another atom,
 giving their valence electrons away,
 or sharing electrons with another atom.
Octet Rule

Whether an atom gains, loses or shares e- to form an
octet depends on how close the atom is to having the 8
e-.
 If an atom has less than 4 valence e- then it will try
to lose them so that the energy level below (which is
full w/ an octet) will then be considered the
outermost energy level.
 Ex: Na = 1s22s22p63s1
 If Sodium loses 1 electron in the 3s energy level,
then energy level 2 will be considered the
outermost energy level. Since it has an octet the
atom will be stable.
Octet Rule
If
an atom has more than 4 valence electrons
it will try to steal electrons from other atoms
to form an octet
Ex: Cl = 1s22s22p63s23p5
If Chlorine gains 1 electron it will have
an octet and be stable.
Octet Rule
Atoms
with 4 electrons in
their valence shell will
share 4 electrons with
other atoms to feel as
though they have an
octet.
Last slide
QOD

The element
found in Group
14 and in
Period 2 is?
Carbon

The number of
valence electrons
in an atom with
an electron
configuration of
1s22s22p63s23p3
is? 5
IV. Oxidation Numbers
Our knowledge of e- configurations and the stability
of noble gases allows us to predict oxidation numbers
for elements.
 Oxidation numbers represent the charge an ion
obtains after losing or gaining valence electrons.
Ex. Ca = 1s22s22p63s23p64s2
2 valence e-, will want to get rid of them so
that the full 3rd energy level (8 electrons) will then be
it’s outermost shell
Ca2+=1s22s22p63s23p6

Oxidation Numbers
Ex. Ca = 1s22s22p63s23p64s2
Ca2+=1s22s22p63s23p6
*** Octet Rule: atoms want to have a full highest
energy level (8 electrons)
** If an atom has less than 4 valence electrons
– they will lose their valence electrons so that the
full energy level below will then be considered the
outermost shell, and become a positive ion (cation)
Oxidation Numbers
Ex.
F = 1s22s22p5
F- = 1s22s22p6
*** Octet Rule: atoms want to have a full highest
energy level (8 electrons)
** If an atom has more than 4 valence
electrons – they will gain (steal) electrons until they
have 8 electrons (full energy level – octet), and
become a negative ion (anion)
Oxidation Numbers
*on your periodic table, write in the oxidation
number of each group
1+
2+
Tend to have more
than one oxidation
number
3+
3+ or 4+
3+
0
3-
2-
1-
0
QOD



What is the element
with the electron
configuration of:
1s22s22p63s2
How many valence
electrons?
What would be the
oxidation number?
Magnesium
2 valence eMg+2



What is the element
with the electron
configuration of:
[Ar] 4s23d104p3
How many valence
electrons?
What would be the
oxidation number?
Arsenic
5 valence eAs-3
V. Periodic Groups & Their Properties
Alkali Metals
•
•
•
•
•
Group 1 Metals
Soft silver metals
Less dense than other metals so melt and
boil easily
All alkali metals have 1 valence electron
• This makes them very reactive because
they want to get rid of it to be stable
(octet in energy level below)
NEVER found alone in nature – are
always bonded to something else; they
are too reactive
Potassium, K
reacts with
water and
must be stored
in kerosene
Alkaline Earth Metals






Group 2 Metals
Shiny silvery - white
All alkaline earth metals have 2 valence electrons
 This makes them very reactive because they
want to get rid of it to be stable (octet in
energy level below)
Alkaline earth metals are less reactive than alkali
metals
Alkaline earth metals are not found pure in
nature; they are too reactive
Found in Earth’s crust in mineral form.
Alkaline Earth Metals
Halogens
•
•
Group 17 Nonmetals
Halogens all have 7 valence electrons
•
•
•
•
This makes them very reactive because they are SO CLOSE
to having an octet, just need 1 more e-
Halogens are never found unbonded in nature; they
are too reactive
Halogens in their pure form are diatomic molecules
(F2, Cl2, Br2, and I2)
Most important group to be used in industry
Chlorine is a yellow-green poisonous gas
Chalogens
• Group 16 nonmetals
• Have 6 valence
electrons
– Semi-reactive
• Diverse group that
includes nonmetals,
metalloids, and metals
Noble Gases
•Group 18 Nonmetals
•Noble gases have 8 valence electrons
• (except helium, which has only 2)
•This makes them extremely stable and unreactive
because they have filled valence shells. They don’t
bond/react with other elements because they don’t need
to gain or lose e-)
• Noble gases are ONLY found alone in
• nature – monatomic gasses
• Colorless, odorless; they were among the last of the
natural elements to be discovered
Noble Gases
QOD

Which elements
have the most
similar chemical
properties…
Those in the same
family or those in
the same period
Same Family

Which of the
groups on the
periodic table are
too reactive to be
found as
monoatomic atoms
in nature?
Alkaline metals &
alkaline earth metals
VI. Periodic Trends
Using the Periodic Table to
Predict Properties of Elements



The basis of the periodic table is the atomic structure
(number of protons, neutrons and electrons) of the
elements.
Position on the table and properties of these elements
arise from the e- configurations of the atoms.
Properties such as density, atomic radius, oxidation
numbers, ionization energy, and e- affinity can be
predicted.
A. Atomic Radius (size of atoms)


As principal quantum number increases (you go down a
group), the size of the electron cloud increases (more
orbitals).
Size of atoms increase moving down Per. Table.
Atomic Radius

Atoms in the same period have the same quantum number;
however, positive charge on the nucleus increases by one
proton for each element in a period. This pulls the e- cloud
in tighter, decreasing atomic radius.
Decreasing Atomic Size Across a Period
Li
Be
B
++
+
++
+
+
+
++
++
Predicting Atomic Radius

General rule: atomic size increases as you move
diagonally from top right corner to bottom left
corner.
B. Ionization Energy
The energy required to remove an e- from an
atom.
 Remove the most loosely held e- is first
ionization energy.
 Measured in kilojoules per mole (kJ/mol)

Ionization Energy
•
As you move down the group, the ionization
goes from high to low ionization energy. The
larger the atom, the less energy is required
because the e- are farther from the positive
center.
Ionization Energy
•
As you move from left to right across a period,
the ionization goes from low to high ionization
energy.
Classification based on
First Ionization Energy
METAL

NONMETAL
1. Low 1st ionization
energy.
2. Located on left side of
Periodic Table.
3. Form positive ions
4. Will give away their
electrons.
4s2
Ca = [Ar]
2 valence epositive ion: Ca+2

1. High 1st ionization
energy.
2. Located on the right side
of Periodic Table.
3. Form negative ions
4. Will NOT give away their
electrons, they want to gain
electrons.
F = [He] 2s2 2p5
7 valence enegative ion: F-
Multiple Ionization Energies

Additional e- can be lost from an atom and the
ionization energies can be measured. Notice that the
energy increases greatly to remove more e-
IONIZATION ENERGIES (kilojoules per mole)
Element
1st
2nd
3rd
4th
H
1312.0
He
2372.3
5220
Li
520.2
7300
11750
Be
899.5
1760
14850
20900
B
800.6
2420
3660
25020
5th
32660
C. Ionization Energies

It increases from left to right, bottom to top and
bottom left to top right corners.
Video:

Ionization Energy (1:52)
C. Electronegativity

.
Electronegativity is the ability of an atom to
capture an electron.
Electronegativity
•
•
As you move from left to right across a period,
the electronegativity increases because
elements on the left side want to lose electrons
to be like the nearest noble gas (octet rule) so
do not want to grab electrons (low
electronegativity)
Elements on the right side want to gain electrons
to be like the nearest noble gas, so the will
grab electrons a lot. (high electronegativity)
Electronegativity
•
THE EXCEPTION TO THIS RULE:
•
NOBLE GASES
They have no electronegativity at
all because they are noble gases,
and don’t need any more
electrons. (No electronegativity)
•
Electronegativity
•
•
As you move from higher to lower in a group,
the electronegativity decreases due to the
shielding effect (which states that electrons in
outer energy levels are held less tightly than
those in lower energy levels).
Generally, if an atom doesn’t hold the electrons
it has very much, it won’t grab electrons from
other atoms much, either.
C. Electronegativity

It increases from left to right, bottom to top and
bottom left to top right corners.
Review

Based on our trends:
The most reactive
metal element would
be
Francium
The most reactive
nonmetal element
would be
Fluorine
Last slide
In Summary
Periodic table is a chart of elements in
which the elements are arranged based
on their e- configurations which dictates
their properties.
 Moving down a group in the periodic
table, atomic radii becomes larger
because more energy levels are needed
for more e-.

In Summary
As the size becomes larger, the e- are
located farther away from the positive
center.
 This decreases the affinity of that atom to
hold on to these outer e-, thus decreasing
e- affinity.
 Ionization energy is low because it is easy
for the atom to lose these outer e-.

In Summary
Moving across a period in the periodic table,
atomic radii becomes smaller because the
energy levels of periods are the same but
the positive centers of atoms increase. This
pulls the e- cloud closer to the nucleus,
making the atom smaller.
 Ionization energy increases for these smaller
atoms.

THE END