Transcript III. Periodic Trends

Ch. 6 - The Periodic Table & Periodic Law
I. Development
of the Modern
Periodic Table
(p. 174 - 181)
I
II
III
A. Mendeleev
Dmitri Mendeleev (1869, Russian)
Organized elements
by increasing
atomic mass
Elements with
similar properties
were grouped
together
There were some
discrepancies
A. Mendeleev
Dmitri Mendeleev (1869, Russian)
Predicted properties of undiscovered
elements
B. Moseley
Henry Moseley (1913, British)
Organized elements by increasing
atomic number
Resolved discrepancies in Mendeleev’s
arrangement
This is the way the periodic table is
arranged today!
C. Modern Periodic Table
Group (Family)
Period
1
2
3
4
5
6
7
1. Groups/Families
Vertical columns of periodic table
Numbered 1 to 18 from left to right
Each group contains elements with similar
chemical properties
2. Periods
Horizontal rows of periodic table
Periods are numbered top to bottom from
1 to 7
Elements in same period have similarities
in energy levels, but not properties
3. Blocks
Main Group Elements
Transition Metals
Inner Transition
Metals
3. Blocks
1
2
3
4
5
6
7
Overall Configuration
Lanthanides - part of period 6
Actinides - part of period 7
Ch. 6 - The Periodic Table
II. Classification of the
Elements
(pages 182-186)
I
II
III
1. Metals
Good conductors of heat and electricity
Found in Groups 1 & 2, middle of table in
3-12 and some on right side of table
Have luster, are ductile and malleable
a. Alkali Metals
Group 1
1 Valence electron
Very reactive
Electron configuration
ns1
Form 1+ ions
Cations
Examples: Li, Na, K
b. Alkaline Earth Metals
Group 2
Reactive (not as reactive as alkali metals)
Electron Configuration
ns2
Form 2+ ions
Cations
Examples: Be, Mg, Ca, etc
c. Transition Metals
Groups 3 - 12
Reactive (not as reactive as Groups 1 or
2), can be free elements
Electron Configuration
ns2(n-1)dx where x is column in d-block
Form variable valence state ions
Cations
Examples: Co, Fe, Pt, etc
2. Nonmetals
Not good conductors
Found on right side of periodic table –
AND hydrogen
Usually brittle solids or gases
a. Halogens
Group 17 (7A)
Very reactive
Electron configuration
ns2np5
Form 1- ions – 1 electron short
of noble gas configuration
Anions
Examples: F, Cl, Br, etc
b. Noble Gases
Group 18
Unreactive, inert, “noble”, stable
Electron configuration
ns2np6 full energy level
Have a 0 charge, no ions
Examples: He, Ne, Ar, Kr, etc
3. Metalloids
Sometimes called semiconductors
Form the “stairstep” between metals and
nonmetals
Have properties of both metals and
nonmetals
Examples: B, Si, Sb, Te, As, Ge, Po, At
C. Valence Electrons
outermost s & p orbital electrons
Stable octet - filled s & p orbitals (8 e-) in one
energy level
Group #A = # of valence e- (except He)
1A
1
2
3
4
5
6
7
8A
2A
3A 4A 5A 6A 7A
C. Valence Electrons
Valence electrons =
electrons in outermost orbitals (highest principle
energy level)
You can use the Periodic Table to determine the
number of valence electrons
Each group has the same number of valence electrons
1A
1
2
3
4
5
8A
2A
3A 4A 5A 6A 7A
Ch. 6 - The Periodic Table
Atomic Radius (pm)
250
III. Periodic
Trends
(p. 187-194)
200
150
100
50
0
0
5
10
Atomic Number
15
20
I
II
III
A. Periodic Law
When elements are arranged in order of
increasing atomic #, elements with similar
chemical and physical properties appear
at regular intervals.
Atomic Radius (pm)
250
200
150
100
50
0
0
5
10
Atomic Number
15
20
B. Chemical Reactivity
Families
Similar valence e- within a group result
in similar chemical properties
1
2
3
4
5
6
7
C. Other Properties
Atomic Radius
size of atom
Ionization Energy
© 1998 LOGAL
Energy required to remove an e- from a
neutral atom
Electronegativity
© 1998 LOGAL
Shielding Effect
There is a Nuclear charge experienced by the outer
(valence) electron(s) in a multi-electron atom is due to
the difference between the charge on the nucleus and
the charge of the core electrons (inner electron shells).
-Results in the reduction of attractive force between
the positive nucleus and the outermost electrons due
to “shielding effect” of the inner electron shells(core
electrons).
Periodic Trend,
1. Shielding effect increases down a group.
2. Shielding effect remains constant across a period.
1. Atomic Radius
Atomic Radius = ½ the distance
between two identical bonded atoms
1. Atomic Radius
Atomic Radius
Increases to the LEFT and DOWN
1
2
3
4
5
6
7
1. Atomic Radius
Why larger going down?
Higher energy levels have larger orbitals
Shielding - core e- block the attraction
between the nucleus and the valence e-
Why smaller to the right?
Increased nuclear charge without additional
shielding pulls e- in tighter
2. Ionization Energy
The minimum energy required to remove an electron
from the ground state of an isolated gaseous atom or ion.
The ease with which an atom loses an e-.
First Ionization Energy = Energy required to remove
one e- from a neutral atom.
He
1st Ionization Energy (kJ)
2500
Ne
2000
Ar
1500
1000
500
Na
Li
0
0
5
10
Atomic Numbe r
15
K
20
2. Ionization Energy
First Ionization Energy
Increases UP and to the RIGHT
1
2
3
4
5
6
7
2. Ionization Energy
Why opposite of atomic radius?
In small atoms, e- are close to the nucleus
where the attraction is stronger
Why small jumps within each group?
Stable e- configurations don’t want to lose
e-
2. Ionization Energy
Successive Ionization Energies
Large jump in I.E. occurs when a CORE
e- is removed.
Mg
Core e-
1st I.E.
736 kJ
2nd I.E.
1,445 kJ
3rd I.E.
7,730 kJ
2. Ionization Energy
Successive Ionization Energies
Large jump in I.E. occurs when a
CORE e- is removed.
Al
Core e-
1st I.E.
577 kJ
2nd I.E.
1,815 kJ
3rd I.E.
2,740 kJ
4th I.E.
11,600 kJ
Electron Affinity
Most atoms can attract e- to form negatively charged
ions
The energy change that occurs when an e- is added
to a gaseous atom or ion.
The ease with which an atom gains an e-.
For most atoms, the energy released when an e- is
added. (in kJ/mol)
Periodic Trend
1. Electron affinity slightly decreases down a group.
2. Electron affinity generally tends to increase
across a period.
3. Electronegativity
The measure of the ability of an atom in a chemical
compound to attract electrons
Given a value between 0 and 4, 4 being the highest
Tendency for an atom to attract e- closer to itself when
forming a chemical bond with another atom.
1
2
3
4
5
6
7
3. Electronegativity
Why increase as you move right?
More valence electrons, need less to fill
outer shell
Why increase as you move up?
Smaller electron cloud, more pull by +
nucleus
Ionic Radius
The size atoms become when losing or gaining
electrons.
Positive Ions – Metal - Atoms that lose e- and
form positive ions become smaller.
The lost e- is a valence e- and the atom may
lose a shell.The repulsion between the
remaining e- is lessened and allows the effective
positive nuclear charge to pull the remaining ecloser.
Negative Ions – Nonmetal - Atoms that
gain e- and form negative ions become
larger.
The repulsion between the added e- and
existing e- is increased and the effective
positive nuclear charge cannot hold onto
the e- tightly.
Examples
Which atom has the larger radius?
Be or Ba
Ca or Br
Examples
Which atom has the higher 1st I.E.?
N or Bi
Ba or Ne
Examples
Which element has the higher electronegativity?
Cl or F
Be or Ca
More Practice
Answer questions 16-19 on page 189 and
20-22 on page 194.