Download Chapter Three: Periodic Table

Chapter Three: Periodic Table
Understanding the Periodic Table
Atomic Mass
Atomic Number
Selected Oxidation
Numbers
Name
of
Element
Electron Configuration
Understanding the Periodic Table
For Example:
Atomic Mass
Atomic Number
12.011
6
-4
+2
+4
C
Selected
Oxidation
Numbers
2-4
Electron Configuration
SUMMARY: All the elements or atoms have the same pattern.
Parts of the Periodic Table:
Periods
Periods are horizontal rows on the
periodic table.
 There are seven periods on the periodic
table.
 Period one has one principle energy
level and as the period number
increases so does the number of
principle energy levels.

Parts of the Periodic Table:
Groups
Groups are vertical columns on the periodic table.They are
numbered from 1-18.
Elements in the same group have similar properties because
they have the same amount of valence electrons.
As you go down a group radius increase, electronegativity
decreases and ionization energy generally decreases (because
valence electrons are further away from positive protons and
more electrons are repelling the valence electrons). [Table S]



Group 1
Li
Na
K
Rb
Cs
Fr
1 valence
electron
Group 2
Be
Mg
Ca
Sr
Ba
Ra
2 valence
electrons
Group 17
F
Cl
Br
I
At
7 valence
electrons
Periodic Law

The Periodic Law states that the
properties of the elements are a period
function of their atomic number. This
means that, when the elements are
arranged by atomic number, those with
similar properties will be at regular
intervals; these elements will be in the
same group.
Metals


All elements to the left of the staircase (zigzag line) are metals.
They have low ionization energy (little energy is needed to
remove an electron) and low electronegativity (little attraction for
electrons). [Shown in Table S]
Other properties of metals:
– Tend to lose electrons to form positive ions
– Have metallic luster (shine)
– Malleable (can be made into sheets)
– Ductile (can be made into wires)
– Good conductor of heat and electricity
– At room temperature most metals are solids except Hg
which is a liquid.
Nonmetals

All elements to the right of the staircase
(zigzag line) except for Group 18. They have
high ionization energy (a lot of energy is
needed to remove an electron) and high
electronegativity (big attraction for electrons).
[Shown in Table S]
 Other Properties of Nonmetals
– Tend to gain electrons when reacting with metals,
forming positive ions
– They share electrons when combined with
nonmetals
– Lack metallic luster and are brittle in the solid
phase
– Poor conductors of heat and electricity.
Metalloids

All elements touching the staircase
(zigzag line) on two sides with the
exception of Al and Po. These elements
have properties of both metals and
nonmetals.
Noble Gases

All elements in Group 18. They have 8
valence electrons (with He as an exception,
which has 2) and because of this, are stable.
Noble gases can bond with only themselves
(He-He) and only Kr, Xe, and Rn are able to
bond with F and O. The forces of attraction
which bonds these elements are called
dispersion forces and are very weak.
 As you go down the group dispersion forces
increases; therefore, the boiling point
increases.
Special Types of Elements

Alkali Metals - located in Group 1 and are the
most reactive metals.
 Alkaline Earth Metals - located in Group 2
and are the second most reactive metals.
 Transition Metals - located from groups 3-11
and have a color in solution.
 Halogens - located in Group 17 and are the
most reactive non-metals.
Periodic Table: Most Active
Metals and Nonmetals

The most active metal is located in the bottom
left hand corner. It is Fr because it is easier to
lose one valence electron than two, if it was in
Group 2, or three, if it was in Group 3.
 The most active nonmetal is located in the top
right hand corner. It is F because it is easier
to gain one valence electron than it is to gain
more than one.
Ions of Metals and Nonmetals

When a metal (Na) loses and electron loses
an electron, a nonmetal (Cl) gains an
electrons.
 The Na loses an electron and becomes an
Na+ and then causes the Na to have a
complete outer shell.
 The Cl gains an electron and becomes Clcausing it to have a full outer shell.
 An ion and an atom only differ in the number
of electrons.
Allotropes
Allotropes are forms of the same
element that have molecular formulas.
 For example O2 and O3 or carbon which
differ in crystalline structure
(arrangement of atoms) such as a
diamond, graphite, or coal.
 Allotropes have different properties.

Physical Properties of Elements

A physical property is a characteristic of a
substance that you can notice without
changing the substance into anything else.
 Examples:
– Color
• Odor
– Solubility
» Density
– Boiling Point
• Melting Point
– Hardness
• Conductivity
– Phase
Chemical Properties Of Elements
Chemical properties of an element
explain how an element reacts in a
chemical reaction.
 Examples:

– Burning
• Mixing with acid or water
– Distillation
Lewis Dot Structure

Use to show how many valence
electrons an element has.
Group
1
2
13
14
15
16
17
18
Element
X
X
X
X
X
X
X
X
X means elements
Means 1 valence electron
Sample Questions
1.
2.
The chemical properties of the elements are
periodic functions of their atomic
(1) masses
(2) weights
(3) numbers
(4) radii
Boron and arsenic are similar in that they
both
(1) have the same ionization energy
(2) have the same covalent radius
(3) are in the same family of elements
(4) are metalloids
Sample Questions
3. Which of the following substances is the best
conductor of electricity?
(1) NaCl(s)
(2) Cu(s)
(3) H2O(l)
(4) Br2(l)
4. Which statement best describes Group 2 elements?
(1) they have one valence electron and they form
ions with a 1+ charge
(2) they have one valence electron and they form
ions with a 1- charge
(3) they have two valence electrons and they
form ions with a 2+ charge
(4) they have two valence electrons and they
form ions with a 2- charge
Sample Questions
5. Which electron configuration represent the
first two elements in Group 17 of the
Periodic Table?
(1) 2-1 and 2-2 (2) 2-2 and 2-3
(3) 2-7 and 2-8-7 (4) 2-8 and 2-8-7
6. Which element is a liquid at room
temperature?
(1) K (2) I2 (3)Hg (4) Mg
Ionization Energy

It is the amount of energy needed to remove
an electron. The smaller the amount of
ionization energy, the easier it is to lose an
electron.
 Nonmetals have high ionization energy
because they are trying to gain electrons to
become more like the noble gases.
 Metals have low ionization energy because
they are trying to lose electrons to become
more like the noble gases.
Ionization Down A Group


Look at Group 1. The elements Li, Na, K, and Rb are listed in
order of atomic number in table S. Look at how the ionization
energy decreases as you go down the group.
Ionization decreases as you go down a group because electrons
are negative and protons are positive and the further away the
electrons are from the nucleus, the less attraction it has to the
protons so, therefore, it takes less energy to remove an electron.
Atomi c # Symbol Ionization Energy
3
Li
520
11
Na
496
19
K
419
37
Rb
403
Ionization Across A Period


Look at Period 2. The elements Li, Be, B, C, N, O, F,
and Ne are listed in order of atomic number in Table
S. Look at how the ionization energy increases as
you go across the period.
Ionization increases because there are more protons
attracting the electrons, making it harder to remove
an electron. Atomi c # Symbol Ionization Energy
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
520
900
801
1086
1402
1314
1681
2081
Electronegativity
Electronegativity is the attraction (or
“desire”) for electrons. The larger the
electronegativity, the more the atom
attracts electrons.
 Electronegativity is a scale from 0-4, 0
being the least desire for electrons and
4 being the most.

Electronegativity Down A Group


Look at Group 1. The elements Li, Na, K, and Rb are
listed in order of atomic number in Table S. Look at
how the electronegativity decreases as you go down
the group.
Electronegativity decreases because the elements try
to lose electrons to become more like the noble
gases, or more stable. This is because noble gases
have a full (8 electrons) outer shell.
Atomic # Symbol Electronegativity
3
Li
1.0
11
Na
0.9
19
K
0.8
37
Rb
0.8
Electronegativity Across A Period


Look at Period 2. The elements Li, Be, B, C, N, O, F,
and Ne are listed in order of their atomic number in
Table S. Look at how the electronegativity increases,
with the exception of Ne which has no desire for
electrons.
This is because the elements want to be more like a
Group 18 element. Atomi c # Symbol Ekectronegativity
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
1.0
1.6
2.0
2.6
3.0
3.4
4.0
--
Atomic Radius
An atomic radius is the distance from
the nucleus to the outermost electron.
 The atomic radius for all elements on
the periodic table can be seen in Table
S.

Atomic Radius Down A Group


Look at the elements in Group 1. The elements Li,
Na, K, and Rb are listed in order by atomic number in
Table S. Look at how the atomic radius increases.
The atomic radius increases because there are more
levels of electrons, making the atom bigger.
Atomic # Symbol Atomi c Radius
3
Li
155
11
Na
190
19
K
235
37
Rb
248
Atomic Radius Across A Period


Look at the elements Li, Be, B, C, N, O, F, and Ne in
Period 2. Look at how the atomic radius generally
decreases.
This is because there are more protons which makes
there be more attraction, pulling the electrons closer
to the nucleus.
Atomi c # Symbol Atomi c Radius
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
155
112
98
91
92
65
57
51
Ionic Radius

For a metal, it loses electrons and becomes
smaller. They become positive and are called
cations. For example, Li loses an electron
and becomes Li+. The ionic radius becomes
smaller because it has less electrons.
 For a nonmetal, it gains electrons and
become bigger. They become negative and
are called anions. For example, F gains an
electron and become F-. The ionic radius
becomes larger because it has more
electrons.
REMEMBER
Atomic mass, Atomic number, Oxidation
states, and Electron Configurations are
all listed in the Reference Table On
The Periodic Table.
 Electronegativity, Atomic Radius,
Ionization Energy, and Ionic Radius are
all listed on Reference Table S.

Sample Questions
1.
A diatomic element with a high first
ionization energy would most likely be
a
(1) nonmetal with high
electronegativity
(2) nonmetal with low
electronegativity
(3) metal with high electronegativity
(4) metal with low electronegativity
Sample Questions
2. An atom of which of the following
elements has the smallest atomic
radius?
(1) Li
(2) Be
(3) C
(4) F
3. Which atom has a radius larger than
the radius of its ion?
(1) Cl
(2) Ca
(3) S
(4) Se