Transcript Atomic Radius

Periodic Table Trends
Unit 3: Electrons and the
Periodic Table
I
II
III
Periodic Law
When elements are arranged in order of
increasing atomic #, elements with similar
properties appear at regular intervals.
Atomic Radius (pm)
250
200
150
100
50
0
0
5
10
Atomic Number
15
20
Atomic Radius
Atomic Radius
size of atom
© 1998 LOGAL
Atomic Radius
Average distance in an atom between
the nucleus and the outermost electron
Atomic Size Trend
Atomic Size increases down a group
Why larger going down?
Adding more energy levels.
Atomic Size decreases across a period
Why smaller across?
 Increased nuclear charge (more
protons) without additional energy levels
pulls e- in closer.
Atomic Radius
Atomic Radius
Increases to the LEFT and DOWN
1
2
3
4
5
6
7
Ionization Energy
Ionization energy is the amount of energy
needed to remove an electron.
M + energy  M+1 + eElectrons that are close to the nucleus
are hard to remove because they are
under a strong force of attraction
Ionization Energy Trend
Ionization Energy Increases across a period
Why? Valence electrons experience a greater
nuclear force because they are closer to the
nucleus.
Smaller atoms have higher Ionization energy.
Ionization Energy Decreases down a group.
Why? Valence electrons removed are farther
from the nucleus because they are in higher
energy levels.
Bigger atoms have lower Ionization energy.
Ionization Energy Trends
Why opposite of atomic radius?
In small atoms, e- are close to the
nucleus where the attraction is stronger
E. Ionization Energy
First Ionization Energy
Increases UP and to the RIGHT
1
2
3
4
5
6
7
Which would have a higher Ionization
energy, Sodium or Chlorine?
Chlorine is higher.
It does not want to lose electrons, it wants
to gain. (non-metals want to gain
electrons!)
Sodium wants to give it away, so it is
easier. (metals want to give away!)
E. Ionization Energy
1st Ionization Energy (kJ)
First Ionization Energy
He
2500
Ne
2000
Ar
1500
1000
500
Li
Na
K
0
0
5
10
Atomic Number
15
20
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
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
Electronegativity
The ability of an atom to attract an
electron to itself for bonding.
The smaller the atom, the more
electronegative it is because of a greater
nuclear force.
Electronegativity Trends
Electronegativity Increases across a period.
Why? Non-metals such as F, O and N want more
electrons to complete their valence shell.
Smaller atoms have greater nuclear charge and
thus, more force to attract electrons.
Exception: Noble gases are not included
because they generally do not want to gain
electrons. They are already stable.
Electronegativity Trends
Electronegativity Decreases Down a Group
Why? Atomic size increases and valence
electrons are farther from the nucleus.
More energy levels increases shielding. So
the pull from the positive nuclear charge is less.
In General:
Non-Metals have high Electronegativities
Metals have low Electronegativities
Ionic Radius
Ionic Radius
Cations (+ ions) the ionic radius is
smaller than the original atom.
Why? There is an increased
attraction for the fewer electrons that
remain.
Na  Na+
Ionic Radius
For Anions (– ions) the ionic radius is
larger than the original atom.
Why? The nuclear attraction is less for
an increased number of electrons.
Extra electrons repel each other and
spread out – larger!)
Cl  Cl-1
© 2002 Prentice-Hall, Inc.
Practice
Which atom is larger H or He?
Which atom has a greater ionization
energy, Ca or Sr?
Which atom is more electronegative, F or
Cl?
Examples
Which atom has the larger radius?
Be or Ba
Ba
Ca or Br
Ca
Examples
Which atom has the higher 1st I.E.?
N or Bi
N
Ba or Ne
Ne
Examples
Which particle has the larger radius?
S or
2S
2S
Al or
3+
Al
Al