Transcript Chapter 4: The Periodic Table

Chapter 4: The Periodic
Table
Section 3: Trends in the Periodic
Table
Key Terms
Ionization energy – removes an
electron from an atom or ion
 Electron shielding – when inner
electrons cancel some of the + charge
of the nucleus & lessen its attraction
of outermost electrons
 Trend – predictable change in a
particular direction

Key Terms 2
Atomic radius – depends on volume
occupied by atom’s electron cloud
 Bond radius – half the distance
between nuclei of atoms that are
bonded together
 Electronegativity – measure of how
much an atom in a chemical compound can attract electrons

Key Terms 3
Electron affinity – change in
energy when a neutral atom gains
an electron
 There are no existing bonds as
with electronegativity.

Things To Know/Answer
What are the periodic trends in
ionization energy, & how are they
affected by atomic structures?
 Answer these questions for atomic
radius, electronegativity, ionic size,
electron affinity, melting points (mp)
& boiling points (bp).

Periodic Trends
All in a group can be explained
by electron configurations.
 Reactivity in alkali metals rise
from top to bottom in group 1.

Periodic Trends 2
When one adds enough energy to
overcome the attraction between
protons & electrons in an atom, it
becomes charged (ion) and an
electron escapes. See page 133
Figure 16.
+
 A + ionization energy → A + e

Ionization Energy
Ionization energy decreases downward in a group because the number
of energy levels increases downward.
 Increasing energy levels have
increasing distance from the nucleus
where the positive charge that attracts
e- in the atom is.

Ionization Energy 2
e
The farther an is from the nucleus, the less attraction protons
there have on the e-.
 Also, the higher an e-’s energy
level is the more full levels of e-’s
there are between it & the nucleus.

Ionization Energy 3
These e-’s “in the middle” reduce the
positive attraction that extends from
the nucleus through the atom
(electron shielding).
 Outermost electrons are less tightly
held for this reason.

Ionization Energy 4
Ionization energy increases as you
move across a period. See Figure
17 on page 134.
 This is because protons increase 1
at a time rightward on a row, but
the energy level stays the same.

Ionization Energy 5 / Atomic Radius
Increasing e-’s get crowded & repel
each other; this counteracts the
positive attrac-tion of the nucleus.
 Here, electrons can only get so close
before increased positive attraction
can-not overcome e- to e- repulsion.
 For this reason, atomic radius stops
decreasing rightward in a row.

Atomic Radius 2
Increasing atomic # across a row produces a much bigger rise in positive
attraction than rise in distance of e-’s
from the nucleus.
 This pattern also causes decreased
atomic radius left to right on a row.

Atomic Radius 3
Rising + attraction also pulls
electrons closer to the nucleus.
 Added inner energy levels are
present downward in groups &
add distance from the nucleus.

Atomic Radius 4
Electron shielding also blocks
rises in + attraction & yields
similar attraction down the group.
 These effects cause increasing
atomic radius downward in a
group.

Electronegativity
Electronegativity is relative attraction
of electrons by nuclei in bonded
atoms where they “play tug of war”
with their shared electrons.
 Linus Pauling, one of America’s most
famous chemists, made a scale of
electronegativity values.

Electronegativity 2
In the scale, he assigned F 4.0 since it
attracts electrons in bonds most then
Pauling calculated values for other
elements relative to this one.
 Electronegativity decreases down a
group mostly because higher energy
levels are farther from the nucleus.

Electronegativity 3
Nuclei cannot attract valence
electrons in these distant energy
levels well.
 For this reason, an element like Cs
has a nucleus w/ more protons but
weak attraction of a valence electron
on its 6th energy level.

Electronegativity 4
However, an element like Li attracts
a valence electron on its 3rd energy
level more strongly.
 This makes Li more electronegative
than Cs.
 Electronegativity increases sharply
rightward across a period.

Electronegativity 5
This trend arises because no change
in electron shielding happens across a
row since no electrons get added to
inner energy levels.
 As atomic # rises quickly, nuclear
charge does also and can attract bond
electrons much more strongly.

Electronegativity 6
Adding inner electrons downward
in a group increases electron
shielding.
 This keeps effective nuclear
charge mostly the same.

Electronegativity 7
Slight drops in electronegativity
result.
 This is b/c distance from the nucleus
is the key factor not nuclear charge.
 Slight rises in distance from the
nucleus downward in a group have
much less effect than boosts in nuclear charge rightward across a row.

Other Periodic Trends
Effective nuclear charge & electron
shielding explain most periodic trends
including ionic size & electron
affinity.
 Ionic size follows trends of atomic
radii for the same reasons.

Other Periodic Trends 2
Metals tend to lose one or more
electrons & become cations (+ions);
whereas, nonmetals tend to gain e-’s
& form anions (-ions).
 Electron affinity follows electronegativity trends (decrease down a group
but increase right across a series) for
the same reasons.

Other Periodic Trends 3
Mp & bp do not generally rise or fall
but reach 2 different peaks as d & p
orbitals fill.
 For example, Cs has low mp & bp
b/c it only has 1 valence electron for
bonding; it is far left in 6th period.

Other Periodic Trends 4
As the electron # increases across a
row, more bonds can form & require
more energy to break.
 This effect peaks near the middle of
d-block elements at W and Re b/c the
d orbitals are half filled.

Other Periodic Trends 5
Further e-’s pair in d orbitals beyond
W & decrease the # of unpaired e-’s
that help strengthen bonds between
atoms by forming multiple bonds.
 More rightward, Hg & Rn have much
lower mp & bp b/c d orbitals are full.

Other Periodic Trends 6
Past Hg, mp & bp rise again as electrons start filling p orbitals until these
are half filled.
 Beyond half filled status, mp & bp
drop again b/c the p orbitals get full
and unable to help strengthen bonds.
 By Rn, p orbitals are full also so mp
& bp are unusually low.
