Lecture 28: Periodicity; Ionic Bonding
• Readings for next class
– 9.2 Electron Configurations of Ions
– 9.3 Ionic Radii
Periodicity
• The trends in radii, ionization energy and electron
affinity across each period lead to trends in
chemical properties (Periodicity)
• Metallic character decreases across each period,
and increases down each group
• Mainly due to increasing ionization energy
• Group IA (alkali metals) react vigorously with water
to produce hydrogen, while group IIA (alkaline
earth metals) react much more slowly
• Reactivity increases down each group
• Group IIIA metals are even less reactive, and B is
not a metal
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Periodicity
• Oxides become less basic and then more acidic
across each period, and up each group
• Group IA oxides are basic
• Form strong bases in water
Na2O(s) + H2O(l) → 2NaOH(aq)
• Group IIA oxides are basic, but Be(OH)2 is not a
strong base
• B2O3 is acidic, Al2O3 shows both acidic and basic
properties (amphoteric)
• CO2, SiO2, N2O5, P4O10, SO2 are acidic
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Bonding
• Bonding is a strong attractive force between
atoms in a substance
• There are three types of bonding:
• Ionic bonding is due to the electrostatic
attraction between ions of opposite charge
• Covalent bonding results from sharing of
electrons between two atoms
• Metallic bonding occurs in metals and
results from the attraction between metal
cations and their “free” valence electrons
Bonding
• Only electrons in the highest occupied level
are used in bonding - VALENCE
ELECTRONS
• For the main groups (s and p blocks):
number of valence electrons = A-group
number
• LEWIS STRUCTURES (ELECTRON DOT
DIAGRAMS) are use to represent the
valence electrons and track them in bonds
• Draw Lewis structures for elements 1-18
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Ionic Bonding
• The metal loses electrons to reach
a noble gas configuration
Na([Ne ]3s1 ) → Na + ([Ne ]) + e• The non-metal gains electrons to
reach a noble gas configuration
Cl ([Ne ]3s 2 3p 5 ) + e - → Cl − ([ Ne]3s 2 3p 6 )
• The reaction is driven by the large
release of Lattice Energy when
the ions form a crystal
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Figure 9.2: Energetics of Ionic Bonding
Born-Haber Cycle
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Lecture 29: Ionic Bonding; Electronic
Configurations of Ions; Ionic Radii
• Readings for next class
– 9.4 Describing Covalent Bonds
Born-Haber Cycle
electron affinity
Ionization energy
-lattice energy
5
Lewis Structures
• Lewis structures show the valence electrons
in the reactants and products
• Write Lewis structures for the formation of
LiF and Na2S
• Monatomic main group ions in compounds
have 0 or 8 valence electrons (not usually
true for transition metals)
Table 9.2
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Presentation of Lecture Outlines, 9a–
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LATTICE ENERGY
proportional to
(+ve charge) x (-ve charge)
distance2
(-787 kJ/mol for NaCl)
Electron Configurations of Ions
• Common monatomic ions of the main group
elements are:
– Cations of Groups IA to IIIA with noble-gas or
pseudo-noble-gas configurations
– Cations of Groups IIIA to VA which retain
their ns2 valance electrons. E.g. Tl+, Sn2+, Pb2+
– Anions of groups VA toVIIA with noble-gas
configurations
• Write configurations for N3-, Al3+, Ga3+,
Sn2+, Sn4+
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Transition Metal Ions
• Transition metals lose their ns electron(s)
first
– 2+ ions are common
– Cu and Ag form 1+ ions because they only
have one ns electron
• The number of nd electrons lost is variable
– Loss of one d electron to give a 3+ ion is
common
• Write configurations for Fe2+, Fe3+, Ti4+
Figure 9.7: Determining the Iodide Ion Radius in the
Lithium Iodide (Lil) Crystal
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Ionic Radius
• Ionic radii are determined from crystal structures
• Influences properties such as density, melting
point, lattice energy, solubility
• For each element, cations are smaller than the
atom (e.g. Na+ < Na), anions are larger (e.g. F- >
F), because electrons shield each other from the
nucleus. When the nuclear charge is the same, the
species with more electrons is always bigger.
e.g. Fe3+ < Fe2+
because six d electrons shield
each other from the nucleus more than five
Figure 9.8: Comparison of Atomic and Ionic Radii
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Lecture 30: Ionic Radii; Covalent Bonding
• Study for the final exam
The Course Evaluation Questionnaires (CEQ) will be
administered in class on Tuesday 7th April
Table 9.3
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General Trends in Ionic Radii
• Ionic radius increases down each group because
larger levels (shells) are occupied by electrons
• Within an isoelectronic (same number of
electrons) series, ionic radius decreases with
increasing positive charge and increases with
increasing negative charge because the higher
nuclear charge makes all shells contract.
• Put the following ions in order of increasing radius
F-, Mg2+, O2-, Na+, Al3+
Covalent Bonds
• A covalent bond results from sharing of two
electrons by two atoms
– Usually one electron is provided by each atom
• As two atoms come together each nucleus begins
to attract the electrons of the other atom, and the
energy of the system decreases
– See Fig. 9.11
• Eventually the repulsion between the two nuclii
prevents the atoms getting closer, and a stable
bond with a fixed length is formed
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Figure 9.11: Potential Energy Curve for H2
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Covalent Bonds
• Only the valence electrons are used in covalent
bond formation
• Each valence electron can form one bond
• Electrons in bonds are in bonding molecular
orbitals and are called BONDING PAIRS
• Valence electrons that are not used in bonding are
called LONE PAIRS
• Covalent bonds and lone pairs can be represented
by Lewis structures
– Draw Lewis structures for H2, HCl, H2S
Figure 9.10: The Electron Probability Distribution for the H2
Molecule
Bonding molecular orbital for H2
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Figure 9.14: Molecular Model of an HCl Molecule
- the spheres represent the approximate size of each atom’s electron cloud
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reserved.
The Octet Rule
• Each atom in a molecule usually has 8 electrons in
its valence shell
– H only has 2
– Both electrons in a bonding pair are counted for both
atoms in the bond
– Atoms acquire a noble gas configuration, and are
therefore in a stable (low energy) state
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Multiple Bonds
• Atoms can share more than one bonding pair
• Double bonds are formed by two bonding pairs
between the same two atoms
– Draw a Lewis structure for O2
• Triple bonds are formed by three bonding pairs
between the same two atoms
– Draw a Lewis structure for N2
Coordinate Bonds
• When acids and bases bond together, the base
donates both electrons to form a coordinate bond
– Draw Lewis structures to show the
neutralization of OH- by H+
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