Physical Science
Chapter 4
Elements & The
Periodic Table
Today’s lesson……
The Modern Atomic
Model……
Objective of today’s lesson
-Describe the mass of an atom and its parts in
terms of amu
-Explain what the atomic number of an
element is
-Calculate atomic mass number or number of
neutrons
Read pages 106 to 108 and
complete the following idea
tree/concept map in your
notes……………
The modern atomic model has
three fundamental particles
The
Electron
Has a
negative
charge
Relative
mass is
1/1836
Found in
clouds
outside
the
nucleus
The
Proton
Has a
positive
charge
Is in the
nucleus
Has a
relative
mass of 1
(~1836X
more than
an electron)
The
Neutron
Has no
charge
Is in the
nucleus
Has about
the same
mass as a
proton
(1 amu)
Parts of an Atom






An atom consists of a nucleus surrounded by one or
more electrons
Atoms are electrically neutral w/ the same number
of protons as electrons.
Majority of the atom is empty space. If nucleus were
the size of a pencil eraser, the closest electron would be
100 yards away!
Subatomic Particles
– Protons
– Neutrons
– Electrons
Nucleus: Tightly packed
Protons & Neutrons
Electrons Orbiting nucleus
@ ~ speed of light!!
Atomic Mass






How much does an atom “weigh”?
What is the mass of an atom?
SI Unit for mass is the Gram…. Way toooo big to
accurately “mass” an atom
Came up w/ new unit, an AMU (atomic mass unit)
1 AMU = mass of 1 Proton
mass of subatomic particles
– Proton = 1 AMU
– Neutron = 1 AMU
– Electron = .0005 AMU

Atomic Mass = the total # of both Protons & Neutrons in
the atom
– ( we don’t worry about the mass of the electrons since they
have almost no mass)
Atomic Number

By definition:
– The Atomic Number = the number of Protons present
in the nucleus of an atom

Each Element in the Periodic Table has a
different number of Protons, therefore each
element has a different, unique, atomic number.
“Smaller” number is
always the atomic #,
therefore the number
of protons present
“Large” number is
always the Atomic
Mass which tells us the
total # of both Protons
& Neutrons present
When reading the Periodic table notice each
element has a unique 1 or 2 letter symbol and
“big” & “small” number listed
Electrical Atomic Charge

Electrical charge – all atoms have a neutral charge
– ( a zero net electrical charge)
Protons have a positive (+) electrical charge
Neutrons have a neutral (0) electrical charge
 Electrons have a negative (-) electrical charge
 Since the net electrical charge is 0 (neutral), if you have
10 Protons (10 “+” charges) then there must be 10 “-”
charges (10 electrons) present to balance out the atom.
 Therefore, as long as you know the Atomic #, you know
the # of Protons and also the # of Electrons!!


For example:
Carbon has an atomic # of 6, it
therefore has 6 Protons which has an
electrical charge of +6, to make the
atom neutral we need 6 negative
charges found in the 6 electrons
orbiting the nucleus.
How many Neutrons are there?

Remember:
– The Atomic # = the # of Protons
– The Atomic mass = The # of both Protons & Neutrons.
– Therefore, if you subtract the Atomic # (the number of Protons)
from the Atomic mass (the number of both Protons & Neutrons)
what is left over must be the number of Neutrons!!
For Example w/ Carbon:
Atomic Mass-Atomic # = # Neutrons
Atomic Mass = 12,
Atomic # 6
12 – 6 = 6 neutrons present in the Carbon nucleus
What makes
an Element?
What are isotopes?
What makes an Element?
Over 117 different elements
 http://www.privatehand.com/flash/elements.
html

OBJECTIVES:
–Explain how the atomic
number identifies an
element.
How Do Atoms of Different
Elements Differ?
• Starting Simply The
hydrogen atom has one
proton and one electron.
• Now for Some
Neutrons The helium
atom has two protons,
two neutrons, and two
electrons.
Elements Review………………
 The
number of protons determines the
element and it’s atomic number
– Atomic Number = the number of protons
– Atomic Mass Units (amu) = # of protons +
# of neutrons
A
neutral atom has the same number of
electrons as protons
Elements Review………………
 Atoms
have larger atomic numbers and
more mass as you add more protons
and neutrons
– Carbon (C): 6 protons + 6 neutrons = 12
amu
– Gold (Au): 79 protons + 118 neutrons =
197 amu
 An
atom does not have to have equal
numbers of protons and neutrons
Hydrogen
H
1
1
Protons: 1
Neutrons: 0
Electrons: 1
Sodium
Na
11
23
Protons: 11
Neutrons: 12
Electrons: 11
Rhenium
Re
186
75
Protons: 75
Neutrons: 111
Electrons: 75
EXAMPLE
How many protons, neutrons and electrons are
found in an atom of
Cs
55
133
Atomic number = protons and electrons
There are 55 protons and 55 electrons
Mass number = sum of protons and neutrons
133 – 55 = 78
There are 78 neutrons
Isotopes
What are they?
Isotopes Review…
 Isotopes
are atoms that have the same
number of protons but different numbers
of neutrons (they are still the same
element)
 Why? They have the same number of protons
 Which means? They have the same atomic
number
 Lets
review this more closely……….
Isotopes….......
Isotopes….......
Telling Isotopes Apart You can identify each isotope of an element
by its mass number. The mass number is the sum of the protons and
neutrons in an atom.
Isotopes
Atoms with the same number of protons
but different numbers of NEUTRONS
Carbon-12
N
NN
N N
N
Atomic # 6
Mass #12
?
6
6
12
N
Carbon-14
N
NNN N
N N
N
Atomic # 6
Mass #14
6
8
14
N
Name: ________________
Period : ______
Isotopes
Atoms with the same number of protons
but different numbers of NEUTRONS
B-11
Atomic #
Mass #
N
B-10
Atomic #
Mass #
N
Isotopes
Atoms of the same element can have different numbers of neutrons
The number of Neutrons in an atom will sometimes vary, that’s why the
atomic mass of the elements is not an even number. For Hydrogen, the
mass is 1.008. Most atoms of Hydrogen have 0 neutrons, but some have 1
neutron and a very very few will have 2 neutrons.
When you “weigh” trillions of Hydrogen atoms you find that almost all of them
will not have any Neutrons, & several of the atoms will have 1 neutron and
maybe 1 or 2 will have 2 Neutrons.
If you were to take an average of all of the Hydrogen atoms in your sample,
the atomic mass would reflect the different Isotopes present and be 1.008
AMU’s.
Organizing the Elements
Objectives:
• What is the periodic table ?
• What information can be obtained
from the table ?
• How the periodic table is organized
• How an elements properties can be
predicted
The scientists who organized it…..
•1829 J.W. Dobereiner (German) organized
elements into triads based on similar
properties and atomic mass
•1869 Dmitri Mendeleev (Russia) published a
classification schemes for all elements (in
order of atomic mass)
• Lothar Meyer (Germany) published a
nearly identical classification also in 1869
• Later, Henri Moseley
(England) 1887-1915,
established that each element
has a unique atomic number,
which is how the current
periodic table is organized
Organizing the Elements

1
The periodic table is laid out by increasing atomic number as you
go across and down the table
1
IA
1
H
Periodic Table
2
IIA
Atomic # increases 
1.00797
2
3
Li
13
IIIA
4
5
Be
B
6.939 9.0122
3
5
6
7
Atomic # increases 
4
11
Na
12
Mg
22.9898 24.305
19
20
K
Ca
14
IVA
6
C
15
VA
7
N
16
VIA
8
O
17
VIIA
9
F
18
VIIIA
2
He
4.0026
10
Ne
10.811 12.0112 14.0067 15.9994 18.9984 20.179
3
IIIB
21
Sc
4
IVB
22
Ti
5
VB
23
V
6
VIB
24
Cr
7
VIIB
25
Mn
8
26
Fe
9
VIIIB
27
Co
10
28
Ni
11
IB
29
Cu
39.102 40.08 44.956 47.90 50.942 51.996 54.9380 55.847 58.9332 58.71 63.54
37
38
39
40
41
42
43
44
45
46
47
17
Cl
18
Ar
26.9815 28.086 30.9738 32.064 35.453 39.948
31
32
33
34
35
36
Ga
65.37 65.37
48
49
As
Se
Br
Kr
72.59 74.9216 78.96 79.909 83.80
50
51
52
53
54
[99]
75
101.07 102.905 106.4 107.870 112.40 114.82 118.69 121.75 127.60 126.904 131.30
76
77
78
79
80
81
82
83
84
85
86
132.905 137.34 138.91 178.49 180.948 183.85 186.2
87
88
89
104
105
106
107
Fr
Ra
Ac
Ku
[223]
[226]
[227]
[260]
Os
Ir
Pt
Au
Hg
In
Ge
87.62 88.905 91.22 92.906 95.94
56
57
72
73
74
Re
Cd
S
85.47
55
W
Ag
16
Ru
Ta
Pd
P
Tc
Hf
Rh
15
Mo
La
Nb
Zn
14
Si
Sr
Ba
Zr
30
13
Al
Rb
Cs
Y
12
IIB
Tl
Sn
Pb
Sb
Bi
Te
Po
190.2 192.2 195.09 196.967 200.59 204.37 207.19 208.980 [210]
108
109
I
Xe
At
Rn
[210]
[222]
Across the Periodic Table
•
Periods: Are arranged horizontally across the
periodic table (rows 1-7)
•
All elements in a period (row) have the same number of
electron shells. (not the same as quantum orbitals)
1
IA
1
2
IIA
13
IIIA
2nd Period
2
3
3
IIIB
4
IVB
5
VB
4
5
6th Period
6
7
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
14
IVA
15
VA
16
VIA
18
VIIIA
17
VIIA
Down the Periodic Table
•Groups or Families: Are arranged vertically down the periodic table
(columns or groups, 1- 18 or I-VIII A,B)
•Elements in a group have the same number electrons in their outer most shells, the
valence shell.
1
IA
1
18
VIIIA
Alkali Family:
1 e- in the valence shell
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
2
3
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
4
5
6
7
Halogen Family:
7 e- in the valence shell
17
VIIA
When using a table with groups
3A to 8A Listed, the group #
equals the number of Valence
8A
electrons
(These are Groups 3A-8A)
1A
2A
3A
(Groups 3B-12B)
4A
5A
6A
7A
C. Bohr Models





Rings represent electron shells/energy levels
Dots represent the electrons in each shell (e-). Pair them!
Center circle is the nucleus with the number of protons and
neutrons
Example: Two elements in period 3 (three electron shells)
Sodium
1e-

7e-
8e-
8e-
2eP11
N12
Chlorine
2eP17
N18
*Add the electrons to each shell according to the number of
elements in a row up to the element you are modeling
C. Dot Diagrams
Dots represent the valence e-.
 EX: Sodium
 EX: Chlorine

C. Electron Shells & Energy Levels
When electrons in an atom are excited by electricity,
light or heat they can jump up to another energy level
 When they come down they emit a specific wave
length of light unique to that element!
 This is called an atomic emissions spectra

7e8e2eP17
N18
Online Animations:
http://einstein.byu.edu/~masong/htmstuff/Absorb2.html
http://ircamera.as.arizona.edu/NatSci102/NatSci102/lectures/spectroscopy.htm
“Need-to-Know Families
“Old Fashion Names” of certain Families
Alkali Metals
Alkaline Earth Metals
Noble Gases
Halogens
Chalogens
Transition Metals
The Alkali Metals – Group 1

Very reactive metals that have only one
valence electron in the outer orbit and
will freely give it away to become stable.
Very soft metal (you could cut it w/ a
plastic knife!). They form ionic bonds w/
Halogens and Chalogens. Examples
include Sodium and Potassium.
The Alkaline Earth Metals –
Group 2

not as reactive as Alkali Metals, but still
very reactive. They have two valence
electrons and generally give them up to
nonmetals to form ionic bonds. Examples
include Calcium and Magnesium
Transition Metals –
Groups 3 thru 12

These all vary dramatically in reactivity, Their oxidation
states (# of valence electrons) vary. They are a bridge
between the very reactive Alkali and Alkaline Earth
Metals and the nonmetals.
Chalogens AKA:
Oxygen Family – Group 16

nonmetals w/ 6 valence electrons, need 2
electrons to fill the outer shell. Most common
oxidation state is -2. Examples are Oxygen
(ozone is one of its allotropes), Sulfur
(responsible for that rotten egg smell when it
combines w/ oxygen to form sulfur dioxide)
and Selenium (one of the few non metals
that are also a good conductor of electricity).
Halogens – Group 17

Very reactive nonmetals w/ 7 valence
electrons. Need only one more electron to
fill their outer shell. Will steal an electron
from a reactive metal to form ionic bonds.
Examples include Chlorine (the most
abundant halogen), Iodine and Bromine
(found in Seawater).
Noble Gases - Group 18

Non reactive, have a full compliment of
valence electrons, 8 and are called the “Inert
Gases” because they do not react w/ other
elements. Examples include Helium (very low
mass and is used in filling children’s balloons
and even airships and the “Goodyear Blimp)
and Neon used in lighted bulbs to make a red
glowing light ( a neon light).
More Need-to-Knows
Transition Metals
Actinides
Lanthanides
Rare Earth Elements – AKA Inner Transition Metals
Properties of Metals





Metals are good conductors
of heat and electricity.
Metals are shiny.
Metals are ductile (can be
stretched into thin wires).
Metals are malleable (can
be pounded into thin
sheets).
A chemical property of
metal is its reaction with
water which results in
corrosion.
The Metals
Examples include: Iron, Tin, Sodium, Calcium, Gallium
Most of the elements are metals.
Metals tend to form positive (+) ions. (They lose electrons)
Physical Properties
– Such as hardness, shiny, malleability (pounded into shapes),
– ductility (stretched or pulled into a wire) electrical conductivity
and magnetic.
 Chemical Properties
– Form metallic bonds with metals and form Ionic bonds w/
nonmetals.




Properties of Non-Metals
Non-metals are poor
conductors of heat and
electricity.
 Non-metals are not
ductile or malleable.
 Solid non-metals are
brittle and break easily.
 They are dull.
 Many non-metals are
gases.

Sulfur
17 Nonmetals





a. There are 17 nonmetals, each are located to the right
of the zigzag line in the periodic table.
b. Non metals tend to steal electrons when they form
negative (-) ions.
c. Physical Properties – in general the physical properties
of nonmetals are opposite those of metals. non conductive,
dull, not ductile or malleable, brittle, gaseous
d. Chemical properties – usually form ionic bonds when
combined w/ metals (NaCl, FeO2, and CaCl2 ) and usually
form covalent bonds when combined w/ other nonmetals
(CO2, O2, C6H12O6)
e. Even though Hydrogen (H) is located in
Group 1, it is still a nonmetal and exhibits
Nonmetals
are the
oxidation states
of +1 and
-1.
light blue elements
Properties of Metalloids
Metalloids (metal-like) have
properties of both metals
and non-metals.
 They are solids that can be
shiny or dull.
 They conduct heat and
electricity better than nonmetals but not as well as
metals.
 They are ductile and
malleable.

Silicon
Metalloids




AKA “semi metals”
7 elements on the zigzag border between metals and the non
metals.
Their properties will sometimes make them act like a metal and
then sometimes act like a nonmetal.
Most important characteristic is their varying ability to conduct
electricity. Silicon is used to make Semiconductors which are used in
making computer chips.
Isotopes
And Radioactivity
Isotopes Review…
 Isotopes
are atoms that have the same
number of protons but different numbers
of neutrons (they are still the same
element)
 Why are they still the same element?
They have the same number of protons
 Which
 Lets
means?
They have the same atomic number
review this more closely……….
Isotopes….......
Isotopes of Carbon
What did we learn?…
 Isotopes
are atoms that have the same
number of protons but different numbers
of neutrons (they are still the same
element)
 Why? They have the same number of protons
 Which means?
They have the same Atomic Number but different
atomic mass (they can also be radioactive!)
 Lets
talk about radioactivity……….
What is Radioactivity?
Radioactivity occurs when certain
elements literally fall apart because their
nucleus is unstable
 Usually protons and neutrons are emitted
by the nucleus
 Sometimes an electron is emitted by the
nucleus, which changes a neutron to a
proton
 Sometimes just high energy gamma
radiation is emitted

Stopping Radiation
Half Life…….
The time it takes for ½ the mass of a radioactive
element to decay into another element

Some Half Lives
Carbon-14: 5,730 years
Uranium-235: 704 MY
Potassium-40: 1.3 BY
Uranium-238: 4.5 BY

Rubidium-87: 48.8 BY




Each half-life, the amount of
atoms gets cut in half.
One half-life.
Two half-lives.
Three half-lives.
Four half-lives.
Don’t worry about the last atom. You start with so
many trillions that you never really get there.
(It will just decay and then
they’re all gone.)
http://lectureonline.cl.msu.edu/~mmp/applist/decay/decay.htm
WE B
FINISHED !!
All Done !