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
Deduced elements existed, but were
undiscovered elements, their properties
could be predicted
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
1
2
3
4
5
6
7
Group (Family)
Period
1. Groups/Families
Vertical columns of periodic table
Each group contains elements with similar
chemical & physical properties (same amount of
valence electrons in each column)
2 numbering systems exist:
Groups # I through VIII with ea. # followed by A or B
• A groups are Main Group Elements (s&p electrons)
• B groups are Transition Elements (d electrons)
Numbered 1 to 18 from left to right
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
A. Metallic Character
1
2
3
4
5
6
7
Metals
Nonmetals
Metalloids
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
Metallic properties increase as you go
from left to right across a period
a. Alkali Metals
Group 1(IA)
1 Valence electron
Very reactive, form metal oxides
(ex: Li2O)
Electron configuration
ns1
Lowest melting points
Form 1+ ion: Cations
Examples: Li, Na, K
b. Alkaline Earth Metals
Group 2 (IIA)
2 valence electrons
Reactive (not as reactive as alkali metals)
form metal oxides (ex: MgO)
Electron Configuration
ns2
Form 2+ ions
Cations
Examples: Be, Mg, Ca, etc
c. Transition Metals
Groups 3 – 12 (IB – VIIIB)
Reactive (not as reactive as Groups 1 or 2),
can be free elements
Highest melting points
Electron Configuration
ns2(n-1)dx where x is column in d-block
Form variable valence state ions
Always form Cations
Examples: Co, Fe, Pt, 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
2. Nonmetals
Not good conductors
Usually brittle solids or gases (1 liquid Br)
Found on right side of periodic table –
AND hydrogen
Hydrogen is it’s own group, reacts rapidly
with oxygen & other elements (has 1
valence electron)
Nonmetal Groups/Families
Boron Group: IIIA typically 3 valence
electrons, also mix of metalloids and metals
Carbon Group: IVA typically 4 valence
electrons, also has metal and metalloids
Nitrogen Group: VA typically 5 valence
electrons, also has metals & metalloids
Oxygen Group: VIA typically 6 valence
electrons, also contains metalloids
a. Halogens
Group 17 (VIIA)
Very reactive
Electron configuration
ns2np5
Form 1- ions – 1 electron short
of noble gas configuration
Typically form salts (NaCl)
Anions
Examples: F, Cl, Br, etc
b. Noble Gases
Group 18 (VIIIA)
Unreactive, inert, ―noble‖, stable
Electron configuration
ns2np6 full energy level
Have an octet or 8 valence e-
Have a 0 charge, no ions
Helium is stable with 1s2, a duet
Examples: He, Ne, Ar, Kr, etc
B. Chemical Reactivity
Metals
Period - reactivity decreases as you go from left to
right across a period.
Group - reactivity increases as you go down a group
Non-metals
Period - reactivity increases as you go from the left
to the right across a period.
Group - reactivity decreases as you go down the
group.
C. Valence Electrons
Valence Electrons
e- in the outermost s & p energy levels
Stable octet: filled s & p orbitals (8e-) in one
energy level
1A
1
2
3
4
5
6
7
8A
2A
3A 4A 5A 6A 7A
C. Valence Electrons
You can use the Periodic Table to determine
the number of valence electrons
Each group has the same number of valence
electrons
Group #A = # of valence e- (except He)
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
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
Properties of Atoms
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).
As atoms add more protons the nuclear charge increases
Atoms are also adding more e- which are attracted to the p+
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(total charge of
protons in nucleus) 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 (IE1) = Energy required to remove
one e- from a neutral atom.
Na(g) + IE1 (energy) → Na+(g) + e- ; +∆H (positive)
Second Ionization Energy (IE2) = energy needed to remove
a second electron, and so forth
Na+(g) + IE2 (energy) → Na2+ (g) + e- ; +∆H (positive)
2. Ionization Energy
First Ionization Energy
Increases UP and to the RIGHT
1
2
3
4
5
6
7
2. Ionization Energy
Why does it increase up a group?
The closer the e- are to the nucleus the more
difficult it is to remove them
Decreased shielding effect increases the
positive nuclear charge
Why does it increase across a period?
Atomic radius decreases
Positive nuclear charge increases pulling ecloser to the nucleus
2. Ionization Energy
Successive Ionization Energies
Large jump in I.E. occurs when a CORE
e- is removed.
The greater the IE the more difficult it is
to remove an electrons
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)
 Cl(g) + e- → Cl—(g) + EA (kJ/mol) ; -∆H (negative)
Electron Affinity
The greater the attraction between a given atom and
an added e-, the more negative the atom’s EA.
Halogens’ ns2p5 have the most negative EA.
Noble Gases have EA > 0; as do Be, Mg, & N because
e- have to enter previously unoccupied, higher energy
orbitals, an unfavorable energy state.
Periodic Trend
1. Electron affinity slightly increases up a group.
2. Electron affinity generally tends to increase
across a period.
Electron Affinity
Electron affinity increases up a group
decreases the atomic radius taking the electrons
closer to the nucleus’ positive attraction.
decreasing shielding effect increases the effective
positive nuclear charge (+) as additional shells are
added and e- are held on tighter.
Electron affinity increases across a period
atomic radius decreases
effective positive nuclear charge increases steadily
and the e- are drawn closer to the nucleus making it
easier to add e- to unfilled sublevels.
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
Increased nuclear charge
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 eand 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.
Periodic Trend
1. Ionic Radius increases down a group.
2. Ionic radius tends to gradually decrease across
a period for the positive ions, then beginning in
group VA or VIA the much larger negative ions
also gradually decreases
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