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Chapter 6
The Periodic Table
Lavoisier
Dobereiner’s Triads
Newland and Mendeleev
Modern Periodic Table
Section 6.2
Classification of the Elements
Group 1 ↓
Vertical columns are called groups.
Elements are placed in columns by
similar properties.
Labeling and Naming Groups
Three common methods
Traditional – label groups with Roman numerals
and letters A or B

longer columns on right and left are labelled IA –
VIIIA. Shorter columns are labelled IB-VIIIB
Some have switched to 1A-8A for long columns
and 1B-8B for short columns
Label each group 1 -18 (not widely used)
Horizontal rows are called periods
There are 7 periods
Group 1A are the alkali metals
Group 2A are the alkaline earth metals
Group 7A is called the Halogens
Group 8A are the noble gases
1A
2A
The elements in the A groups
8A
0
are called the representative
3A 4A 5A 6A 7A
elements
The group B are called the
transition elements
 These
are called the inner
transition elements and they
belong here
Metals, Non-metals, Metalliods
Section 6.2
Classification of the Elements
Objective: Identify and explain the similarities
between valence electron configurations for
elements in a group of the periodic table.
Electron Configurations and the
Periodic Table
Valence Electrons

the electrons in the HIGHEST energy
level for an element

Determine the properties of an
element

EXAMPLE:
Br: 1s22s22p63s23p64s23d104p5
Highest energy level with electrons = 4
7 electrons in energy level 4 = valence
electrons
Valence Electrons and PERIOD


The energy level of an element’s
valence electrons = the period
number that the element is in
EX: Br is in period 4, so its valence
electrons are in energy level 4
1
2
3
4
5
6
7
Each row (or period) is the energy level
for s and p orbitals.
D orbitals fill up after previous energy level
so first d is 3d even though it’s in row 4.
1
2
3
4
5
6
7
3d
1
2
3
4
5
6
7
4f
f orbitals start filling at 4f
5f
Valence Electrons and COLUMN



Representative elements: group number
corresponds to the number of valence
electrons those elements have (when
using the A/B numbering system)
EX: Group 1A elements all have an
electron configuration that ends in s1, so
these elements all have ONE valence
electron.
Group 6A elements all have an electron
configuration that ends in s2p4. These
elements all have SIX valence electrons
H
Li
1
3
Na
11
K
19
Rb
37
Cs
55
Fr
87
1s1
Alkali Metals
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10
5p66s1
1s22s22p63s23p64s23d104p65s24d105p6
6s24f145d106p67s1
s-block elements



Elements in Groups 1A all have
electron configurations ending in s1
and elements in group 2A on the
periodic table have electron
configurations that end in s2
Groups 1A-2A are s-block elements
s-block is only two groups long
because s can only hold TWO
electrons
S- block
s1
s2
Alkali metals all end in s1
Alkaline earth metals all end in s2
really have to include He but it fits
better later.
He has the properties of the noble
gases.
p-block elements



Elements in group 3A have electron
configurations ending in p1, elements
in 4A have electron configurations
ending in p2, etc.
Continues through group 8A (Group 3A
– 8A or 13- 18 compose the p-block)
Because p orbitals can hold 6
electrons, the p-block spans six groups
The P-block
p1 p2
p3
p4
p5
p6
Noble
Gases
1s2 He 2
Ne
2
2
6
1s 2s 2p
10
1s22s22p63s23p6 Ar18
1s22s22p63s23p64s23d104p6 Kr
36
1s22s22p63s23p64s23d104p65s24d105p6 Xe
54
1s22s22p63s23p64s23d104p65s24d10 Rn
5p66s24f145d106p6 86
d-block elements




Transition metals
First row of d block is actually 3d
(looks like 4d)
B Groups or groups 3 – 12
All the elements in group three have
an electron configuration that ends
in d1, all the elements in group four
have a configuration that ends in d2,
etc.
Transition Metals -d block
d1 d2 d3
s1
d5
s1
d5 d6 d7 d8 d10 d10
f-block elements



Inner transition metals
Spans 14 columns of the periodic
table
First group of f-block elements have
electron configurations that end in f1
F - block
inner transition elements
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
Section 6.3
Periodic Trends
]
Objective: Be able to identify the major trends
among elements in a group or period on the periodic
table and explain the reason for the trend.
What is a trend?
predictable behavior
on the periodic table, the properties of an
element change in a predictable way
Atomic Radius
I. Atomic Radius: the distance from its nucleus
to its outermost electron (difficult to measure)
Atomic Radius
A. Atomic size – determined by the distance
between the outermost electrons and the nucleus
B. Atomic radius –
 The outermost electrons don’t exist in a
definite boundary (electron cloud)
 Atomic radius = half the distance between
nuclei of two of the same element bound
together in a compound
Atomic Radius
EXAMPLE:
If two atoms of chlorine are bound together like
this: Cl—Cl, the atomic radius of chlorine is
approximately half the distance between their
nuclei
Trend within a Period
Atomic radius DECREASES
 In a period, energy level stays the same and
the nucleus gets more protons as we move
from left to right
 MORE positive nucleus PULLS in electrons
closer, so it gets SMALLER
Atomic Radius DECREASES
across a period
EXAMPLE:
To describe the trend going ACROSS a period,
lets look at the atomic structures of C and N
(RIGHT next to each other on the periodic table):
Carbon (atomic # 6)
Nitrogen (atomic # 7)
Carbon: 6 protons
e- configuration: 1s22s22p2
Nitrogen: 7 protons
e- configuration: 1s22s22p3
Trend within a Group
Atomic size increases as you move down a group
Electrons are in higher energy levels farther from
the nucleus as you go DOWN
EXAMPLE
Li: 1s22s1
Na: 1s22s22p63s1
Na has more electrons, these extra electrons occupy energy levels that are
farther away from the nucleus, so the electron cloud of Na is BIGGER
Atomic Radii Trend
Atomic Size…
What is the smallest atom on the periodic
table?
What is the largest atom on the periodic
table?
Practice
DECIDE WHICH ELEMENT IS BIGGER:
1. Li, Ar
2. Sr, Te
3. Si, F
4. Kr, O
5. Rank from smallest to largest: C, F, Be, Li
6. Rank from smallest to largest: Mg, Si, S, Na
ION SIZE TREND
ION
When an atom loses or gains electrons, it
becomes an ION
EX:
Chlorine has 17 protons (+) and 17 electrons (-).
If Cl GAINS an electron, it becomes Clbecause it has 18 electrons and a charge of -1.
Magnesium has 12 protons (+) and 12 electrons
(-). If Mg looses two electrons it becomes
Mg+2 because it has 10 electrons and has a
charge of + 2.
ION SIZE TREND
POSITIVE IONS:
Positive ions are smaller than the neutral atoms
because they LOOSE electrons
they LOOSE electrons
ION SIZE TREND
NEGATIVE IONS:
Negative ions are bigger than the neutral atom
because they GAIN electrons
Repulsion between electrons increases
(electrons don’t like each other)
TREND: (SAME AS SIZE)


Size of ions INCREASES down a group,
DECREASES across a period
Metals tend to form positive ions (cations) and
nonmetals tend to form negative ions (anions)
Ionic Radii Trend
Practice
Which will be larger, the ion formed by sodium or
the neutral sodium atom?
Which will be larger, the ion formed by bromine or
the bromine atom?
Which will be larger, the ion formed by oxygen or
the ion formed by fluorine?
Which will be larger, the ion formed by calcium or
the ion formed by potassium?
IONIZATION ENERGY
Ionization Energy –
Energy required to remove ONE electron from a
gaseous atom
Measures how strongly the atom holds onto its
outermost electrons
If the atom has a strong hold on the
electrons, will it have a BIG or SMALL
ionization energy?


If it has a strong hold on its electrons, it will have a
LARGE ionization energy
It will take A LARGE AMOUNT of energy to steal the
electron
Which atoms have a strong
hold on their electrons??
Which atoms have a strong hold on
their electrons??
Octet Rule:
atoms want to have a noble gas electron
configuration (want 6 electrons in their p orbitals) for
their outermost electrons
Noble gases are STABLE (every element wants to be stable):
Atoms with electrons in their p orbitals ware closer to
noble gas configuration (only have to gain a few
electrons). These atoms don’t want to lose electrons,
want to GAIN!
Atoms in the s block want to LOSE electrons
because it is easier to lose electrons to get to noble
gas configuration. These atoms want to lose
electrons.
IONIZATION ENERGY TREND
TREND:
Ionization energy gets BIGGER as you move from
left to right across a period
Ionization energy gets SMALLER as you move down
a group because the electrons are FAR from the
nucleus, aren’t held as tightly
First Ionization Energy:
energy required to remove one electron from an
atom
Second Ionization Energy:
Energy required to remove two electrons from an
atom
Third Ionization Energy:
Energy required to remove three electrons from an
atom
Which value, the FIRST ionization energy,
the SECOND ionization energy, or the
THIRD ionization energy do you think has
the LARGEST value? Why?
The third energy has the highest value
because it is harder to take three electrons
away from the positive nucleus
Practice
1. State whether the following elements are more
likely to GAIN or LOSE an electron?
a. Cl
b. O
c. S
d. As
e. Al**
f. C**
g. Sr
h. Ca
i. Mg
j. Br
Practice
2. Order the following elements from LEAST to
GREATEST ionization energy:
a. Br, Cl, F
b. N, Cl, S, Mg, Na
c. K, Rb, S, O
d. Br, As, Se
ELECTRONEGATIVITY
Electronegativity– In a bond between
elements SHARE electrons in a bond
Electronegativity measures how much an
atom HOGS electrons in a bond
Elements that want to GAIN electrons are
going to hog electrons in a bond
ELECTRONEGATIVITY
Which elements on the periodic table have a HIGH
electronegativity? Which have a LOW electronegativity?
 Nonmetals
 Metals have a LOW electronegativity because they
want to LOSE electrons to be like the noble gases
As you move ACROSS the periodic table,
electronegativity INCREASES
As you move DOWN the periodic table, electronegativity
DECREASES
 Elements with electrons far from the nucleus don’t
want more electrons as much, they don’t feel
attracted to the positive nucleus as much
ELECTRONEGATIVITY
http://www.youtube.com/watch?v=YZ8izzEq6zI&feature=related
Practice
Rank the following elements in order of
INCREASING electronegativity:
1. P, Si, Al, Cl
2. K, Rb, Cs
3. Pb, Ba, Br, Cl
4. F, N, O