Download - Chapter 7 - Periodic Properties of the Elements

Spring 2016
- Chapter 7 Periodic Properties
of the Elements
Summary
7.1
7.2
7.3
7.4
7.5
7.6
7.7
7.8
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Development of the periodic table
Effective nuclear charge
Size of atoms and ions
Ionization energy
Electron affinities
Metals, Nonmetals, and Metalloids
Group trends for the active metals
Group trends for selected Nonmetals
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Development of Periodic Table


The periodic table is the most significant
tool that chemists use for organizing
elements and recalling chemical facts.
Elements in the same column contain the
same number of outer-shell electrons or
valence electrons and have
consequently similar chemical properties.
Development of Periodic Table


Mendeleev and
Meyer arranged the
elements in order of
increasing atomic
weight.
Certain elements
were missing from
their scheme.
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Development of Periodic Table

in 1871 Mendeleev noted that As properly
belonged underneath P and not Si, which left a
missing element underneath Si. He predicted a
number of properties for this missing element
(which he called eka-silicon or Germanium) with
chemical properties similar to those of silicon.
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Periodic Trends

In this chapter, we will rationalize
observed trends in:
• Sizes of atoms and ions.
• Ionization energy.
• Electron affinity.
Effective Nuclear Charge




In a many-electron atom, electrons are
both attracted to the nucleus and
repelled by other electrons.
The nuclear charge that an electron
experiences depends on both factors.
The electron is attracted to the nucleus,
but repelled by core electrons that shield
or screen it from the full nuclear charge.
This shielding is called the screening
effect.
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Effective Nuclear Charge
The effective nuclear charge, Zeff, is
found this way:
Zeff = Z − S
where Z is the atomic number and S is a
screening constant, usually close to the
number of inner electrons.
Na (Z=11)
 The effective charge Zeff = Z − S
= 11+ -10 + =1+

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Sizes of Atoms




Consider a collection of argon atoms in
the gas phase.
During collisions electron clouds cannot
penetrate each other to a significant
extent.
The apparent radius is determined by
half of the closest distances separating
the nuclei during such collisions.
This radius is the nonbonding radius.
Sizes of atoms in molecules

The bonding
atomic radius is
defined as onehalf of the
distance between
covalently bonded
nuclei which is
shorter than
nonbonding
radius.
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Example

The distance separating the iodine
(I2) nuclei is 2.66 A° , thus the
radius is 1.33 A°.
Sizes of Atoms

Knowing the atomic radii allows the
estimation of the bond lengths between
different elements in molecules. In the
compound CCl4 the measured length of
C-Cl bond is 1.77 A° which is very
close to the sum of (0.77A°+ 0.99 A°)
for C and Cl respectively
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Sizes of Atoms
Bonding atomic radius
tends to decrease from
left to right across a
row
due to increasing ZEFF which
draws the electrons closer
to the nucleus causing the
atom to decrease in size.
Bonding atomic radius
tends to increase
from top to bottom
of a column
due to increasing value of n
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Sizes of Ions


The radii of ions are based on the
distances between ion nuclei in an
ionic compound.
Ionic size depends upon:
• Nuclear charge.
• Number of electrons.
• Orbitals in which electrons reside.
Sizes of Cations

Cations are
smaller than
their parent
atoms.
• The outermost
electron is
removed and
repulsions are
reduced.
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Sizes of Anions

Anions are
larger than their
parent atoms.
• Electrons are
added and
repulsions are
increased.
Sizes of Ions

Ions increase in
size as you go
down a column.
• Due to increasing
value of n.
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Sizes of Ions


In an isoelectronic series, ions have the
same number of electrons.
In an isoelectronic series, ionic size
decreases with an increasing nuclear
charge.
Ionization Energy

Amount of energy required to
remove an electron from the
ground state of a gaseous atom or
ion.
• First ionization energy is the energy
required to remove the first electron.
• Second ionization energy is the
energy required to remove the
second electron, etc.
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Ionization Energy

More energy is needed to remove each
successive electron since the effective
nuclear charges increase (more
attractions).

Trends in First Ionization Energies

As one goes down
a column, less
energy is required
to remove the first
electron.
• For atoms in the
same group, Zeff is
essentially the
same, but the
valence electrons
are farther from the
nucleus.
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Trends in First Ionization Energies

Generally, as one goes across a
row, it gets harder to remove
an electron.
• As you go from left to right, Zeff
increases.
Trends in First Ionization Energies

Within each group the ionization
energy generally decreases with
increasing atomic number. The alkali
metals have the lowest first
ionization energies
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Irregularities in First Ionization
Energies
There are two apparent irregularities:


The first occurs between Groups IIIA and
IIA. Electrons are more easily removed
from p-orbitals than a s-orbital.
The second occurs between Groups VA and
VIA.
• Electron removed comes from doubly occupied
orbital. Repulsion from other electron in the
same orbital helps in its removal
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Electron Affinity
Energy change accompanying addition
of an electron to a gaseous atom:
Cl (g) + e-  Cl− (g)
Trends in Electron Affinity
In general, electron
affinity becomes
more exothermic as
you go from left to
right across a row.
The greater the
attraction between
atom and added
electron the more
negative is the
electron affinity. A
positive value means
that an electron
cannot be attached
to an atom.
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Trends in Electron Affinity
There are again,
however, two
discontinuities in
this trend.
Trends in Electron Affinity

The first occurs
between Groups
IA and IIA.
• Added electron
must go in porbital, not sorbital (highly
energetic subshell
which is empty in
the neutral atom .
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Trends in Electron Affinity

The second occurs
between Groups
IVA and VA.
• Group VA has no
empty orbitals.
• Extra electron must
go into occupied
orbital, creating
repulsion.
Properties of Metal, Nonmetals,
and Metalloids
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Metals versus Nonmetals


Metals tend to form cations.
Nonmetals tend to form anions.
Metals versus Nonmetals
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Metals


Tend to be malleable and good
conductors of heat and electricity.
Compounds formed between
metals and nonmetals tend to be
ionic.
Metal oxides tend to be basic.
Nonmetals


brittle to hard substances that are
poor conductors of heat and electricity.
Tend to gain electrons in reactions with
metals to acquire noble gas
configuration.
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Nonmetals


Substances containing only nonmetals are
molecular compounds.
Most nonmetal oxides are acidic.
Metalloids


Have some
characteristics of
metals, some of
nonmetals.
For instance,
silicon looks
shiny, but is
brittle and fairly
poor conductor.
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Group Trends
Alkali Metals

Soft, metallic
solids.
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Alkali Metals



Found only as compounds in nature.
Have low densities and melting points.
Also have low ionization energies.
Alkali Metals
Their reactions with water are famously
exothermic.
Used in distillation stills.
Story of the guy burning his hand in Notts.
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Alkaline Earth Metals


Have higher densities and melting
points than alkali metals.
Have low ionization energies, but not as
low as alkali metals.
Alkaline Earth Metals


Beryllium does
not react with
water but others
react readily with
water.
Reactivity tends
to increase as you
go down a group.
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Group 6A



Oxygen, sulfur, and selenium are
nonmetals.
Tellurium is a metalloid.
The radioactive polonium is a metal.
Oxygen

Two allotropes:
• O2
• O3, ozone

Three anions:
• O2−, oxide
• O22−, peroxide
• O2−, superoxide

Tends to take
electrons from other
elements (oxidation)
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Sulfur


Weaker oxidizing
agent than
oxygen.
Most stable
allotrope is S8, a
ringed molecule.
30 allotropes
Group VIIA: Halogens
Name comes from the Greek halos and
gennao: “salt formers”
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Group VIIA: Halogens

Large, negative electron
affinities
• Therefore, tend to oxidize
other elements easily


React directly with metals to
form metal halides
Chlorine added to water
supplies to serve as
disinfectant
Group VIIIA: Noble Gases

Xe forms three
compounds:
• XeF2
• XeF4 (at right)
• XeF6

Kr forms only one
stable compound:
• KrF2

The unstable HArF was
synthesized in 2000.
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