Chapter 7
Periodic Properties of the Elements
Learning Outcomes:
Explain the meaning of effective nuclear charge, Zeff, and how Zeff depends on nuclear charge and
electron configuration.
Predict the trends in atomic radii, ionic radii, ionization energy, and electron affinity by using the
periodic table.
Explain how the radius of an atom changes upon losing electrons to form a cation or gaining electrons
to form an anion.
Write the electron configurations of ions.
Explain how the ionization energy changes as we remove successive electrons, and the jump in
ionization energy that occurs when the ionization corresponds to removing a core electron.
Explain how irregularities in the periodic trends for electron affinity can be related to electron
configuration.
Explain the differences in chemical and physical properties of metals and nonmetals, including the
basicity of metal oxides and the acidity of nonmetal oxides.
Correlate atomic properties, such as ionization energy, with electron configuration, and explain how
these relate to the chemical reactivity and physical properties of the alkali and alkaline earth metals
(groups 1A and 2A).
Write balanced equations for the reactions of the group 1A and 2A metals with water, oxygen,
hydrogen, and the halogens.
List and explain the unique characteristics of hydrogen.
Correlate the atomic properties (such as ionization energy, electron configuration, and electron affinity)
of group 6A, 7A, and 8A elements with their chemical reactivity and physical properties.
Development of Periodic Table
•Dmitri Mendeleev and
Lothar Meyer (~1869)
independently came to
the same conclusion
about how elements
should be grouped in
the periodic table.
•Henry Moseley (1913)
developed the concept
of atomic numbers (the
number of protons in
the nucleus of an atom)
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Predictions and the Periodic Table
Mendeleev, for instance, predicted the discovery of germanium
(which he called eka-silicon) as an element with an atomic
weight between that of zinc and arsenic, but with chemical
properties similar to those of silicon.
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.
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Effective Nuclear Charge
The effective nuclear
charge, Zeff, is:
Zeff = Z − S
where Z is the atomic
number and S is a screening
constant, usually close to
the number of inner (core)
electrons.
The valence electrons do not
screen as efficiently as core
electrons.
Zeff is a representation of the
average electrical field
experienced by a single
electron.
Effective nuclear charge is a periodic
property:
• It increases across a period.
• It increases slightly down
a group.
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Sizes of Atoms
• The nonbonding atomic
radius, or van der
Waals radius, is half of
the shortest distance
separating two nuclei
during a collision of
atoms.
• The bonding atomic
radius is defined as one-
half of the distance
between covalently
bonded nuclei.
Periodic Trends in Atomic Radii
Bonding atomic radius tends to…
…decrease from left to right across a row
due to increasing Zeff.
…increase from top to bottom of a column
due to increasing value of n
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Sizes of Ions
Ionic size depends upon:
• Nuclear charge,
number of electrons,
orbitals in which
electrons reside.
• Cations are smaller
than their parent
atoms.
– The outermost
electron is removed
and repulsions are
reduced.
• Anions are larger than
their parent atoms.
– Electrons are added
and repulsions are
increased.
Sizes of Ions
• In an isoelectronic series, ions have the
same number of electrons.
• Ionic size decreases with an increasing
nuclear charge.
• Increasing nuclear charge with decreasing
ionic radius as atomic number increases.
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Ionization Energy
Amount of energy required to remove an electron
from the ground state of a gaseous atom or ion.
– First ionization energy is that energy required
to remove first electron.
– Second ionization energy is that energy
required to remove second electron, etc.
Trends in First Ionization Energies
• I1 generally decreases down a group. For atoms in the same
group, Zeff is essentially the same, but the valence electrons are
farther from the nucleus.
• I1 generally increases across a period.
• The s- and p-block elements show a larger range of values for I1.
• The d-block generally increases slightly across the period; the f-
block elements show only small changes.
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Irregularities in First Ionization Energies
• Discontinuity occurs between
Groups IIA and IIIA.
• Electron removed from p-
orbital rather than s-orbital
– Electron farther from nucleus
– Small amount of repulsion by
s electrons.
• The second occurs between
Groups VA and VIA.
– Electron removed comes
from doubly occupied orbital.
– Repulsion from other
electron in orbital helps in its
removal.
Electronic Configurations of Ions
To form cations from elements remove e-’s
from the subshell with the highest n.
P([Ne]3s23p3) → P3+([Ne]3s2) + 3e-
3p 3p
3s 3s
2p 2p
2s 2s
1s 1s
P P3+
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Electronic Configurations of Ions
For transition metals, remove ns electrons
and then (n - 1)d electrons.
Fe([Ar]4s23d6) → Fe2+([Ar]3d6) + 2e-
Fe Fe2+
4s 4s
3d 3d
To form cations, first
remove electrons with Fe3+
highest n value.
4s
3d
Fe2+([Ar]3d6) → Fe3+([Ar]3d5) + e-
Electron Affinity
Energy change accompanying addition of electron to
gaseous atom:
Cl(g) + e− Cl−(g)
In general, electron affinity becomes more exothermic as you go
from left to right across a row.
Three notable exceptions include the following:
1) Group 2A: s sublevel is full
2) Group 5A: p sublevel is half-full
3) Group 8A: p sublevel is full
Halogens have the most negative electron affinity.
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Metal, Nonmetals, and Metalloids
Metals versus Nonmetals
Metals tend to form cations.
Nonmetals tend to form anions.
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Metals
• Compounds formed between metals and
nonmetals tend to be ionic.
• Metal oxides tend to be basic.
CaO(s) + H2O(l) → Ca(OH)2 (aq)
Nonmetals
• Dull, brittle
substances that are
poor conductors of
heat and electricity.
• Large negative
electron affinity, tend
to gain electrons in
reactions with metals
to acquire noble gas
configuration.
• Substances containing
only nonmetals are
molecular compounds.
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Nonmetals
• Most nonmetal oxides are acidic.
• Nonmetal oxides react with bases to form
salts and water.
Metalloids
• B, Si, Ge, As, Sb, Te
• Have some
characteristics of
metals, some of
nonmetals.
• For instance, silicon
looks shiny (metallic
luster), but is brittle
and fairly poor
conductor
(semiconductor).
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Group Trends
• Elements in a group have similar properties.
• Trends also exist within groups.
• Groups compared:
– Group 1A: the alkali metals
– Group 2A: the alkaline earth metals
– Group 6A: the oxygen group
– Group 7A: the halogens
– Group 8A: the noble gases
– Hydrogen: nonmetal
Alkali Metals
• Soft, metallic solids.
• Name comes from Arabic word for ashes.
• Found only as compounds in nature.
• Have low densities and melting points.
• Also have low ionization energies.
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Alkali Metals
Reactions with water are exothermic
React with oxygen, hydrogen, and halogens.
Alkali Metals
Alkali metals (except Li) react
with oxygen to form
peroxides.
• K, Rb, and Cs also form
superoxides:
K + O2 KO2
Produce bright colors when
placed in flame.
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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
• Be does not react
with water, Mg
reacts only with
steam, but others
react readily with
water.
• Reactivity tends to
increase as go
down group.
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Hydrogen
• 1s1 a metallic electron configuration
like the other ns1 elements
• We do think of acid compounds, like
HCl, as having H+, however they
are really covalent in nature.
• When reacting with metals, hydride
anions (H–) form.
• Forms both ionic (H- with metals)
and molecular compounds (H+ with
nonmetals).
Group 6A - chalcogens
• Oxygen, sulfur, and selenium are
nonmetals.
• Tellurium is a metalloid.
• Radioactive polonium is a metal,
forms a cation.
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Oxygen and Sulfur
• Two allotropes:
– O2, dioxygen
– O3, ozone
• Three anions:
– O2−, oxide
– O22−, peroxide
– O21−, superoxide
• Tends to take electrons from
other elements (oxidation)
• S is weaker oxidizing
agent than oxygen.
• Most stable allotrope
is S8, a ringed
molecule shaped like
a crown.
Group VIIA: Halogens
• Prototypical nonmetals
• Name comes from the Greek halos and
gennao: “salt formers”
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Group VIIA: • Large, negative electron
affinities
Halogens – 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
• Monatomic gases
• Large ionization energies
• Positive electron affinities
• Relatively unreactive
– Only fluorine (F) can remove electrons
to form compounds
– XeF2, XeF4 XeF6, KrF2 are known XeF4
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