Polar Bonds in Covalent Molecules
• A polar covalent bond is one in which the bonding electron pair(s) are shared
unequally. This results when two bonding atoms have different electronegativities.
• Electronegativity is
a measure of the
ability of an atom,
when bonded to
another atom, to
attract electrons to
itself.
The Pauling electronegativity (EN) scale.
• The greater the
difference in
electronegativity
between the atoms
the greater the bond
polarity.
most
electronegative
atom becomes
partially (-)
least
electronegative
atom becomes
partially (+)
What is the general trend for electronegativity down a group?
What’s the trend across a period?
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Electronegativity and Oxidation Number
We can now understand the rules for assigning oxidation numbers. They are based on
the following:
•
•
•
•
The more electronegative atom is assigned all the shared electrons.
The less electronegative atom is assigned none of the shared electrons.
Each atom in a bond is assigned all of its unshared electrons.
O.N. = # of valence e– – (# of shared e– + # of unshared e–)
Examples:
CH4 CF4
CO LiH
Remember that oxidation numbers are considered charges only for single atom ions.
For covalently bonded atoms, oxidation number simply reflects the direction of
electron density shift toward the more electronegative atom.
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The Importance of Electronegativity Difference (∆EN)
The magnitude of the difference in electronegativity between two bonded atoms
determines the “percent ionic character” of the bond.
∆EN = 0
∆EN = 1.9
Depicting Polar Bonds
∆EN = 3.0
δ+
δ–
Arrange the
following bonds in
order of increasing
polarity: N–F, Be–F
and O–F
Even in highly ionic LiF some
electron sharing occurs between
the ions in the gaseous ion pair.
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Chemical Bond Formation: Ionic
Ionic bonding: transfer of electrons (through redox) to form ions. For main group elements,
the ions form noble gas configurations. The metal is oxidized while a nonmetal is reduced.
Na(s) + (1/2)Cl2(g) —> [Na+ + Cl–] —> NaCl(s);
∆H°f = –410.9 kJ
• All ionic compounds are solids at room temperature!
• The compounds exist as crystal lattices. Uniform solids with a predictable structure.
• There are no “individual” NaCl units, each Na+ ion is attracted to 6 nearest neighbor
Cl– ions, and each Cl– ion to 6 nearest neighbor Na+ ions.
Three ways to depict electron transfer in the formation
of Na+ and Cl–:
1) Electron Configurations
2) Orbital Diagrams
3) Lewis electron-dot symbols
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Energetics of Ionic Bond Formation – Lattice Energy of Ionic Crystals
The lattice energy is the energy needed to
break apart the solid crystal lattice and form
ions in the gas phase. This is always
ENDOTHERMIC!
NaCl(s) —> Na+(g) + Cl–(g)
∆Hlattice = +788 kJ/mol
The lattice energy IS NOT the opposite of ∆H°f.
For NaCl(s) ∆H°f is –410.9 kJ/mol
The lattice energy is an indication of the force
of attraction between the ions. This attractive
force depends on the charge of the ions and
distance between ions in the solid phase
(Coulomb’s Law); charge having the greater
affect.
∆ H Lattice ∝ attractive force =
kQ1Q2
d
k is a constant
What are the trends in lattice energy shown
above in terms of ion size and ion charge?
Lattice energies can be used to predict
relative melting point for ionic substances. In
general, the greater the lattice energy, the
higher the melting point.
Predict the order of melting points for:
KCl, CaO, CsCl, CaCl2 and BaO.
Q1 and Q2 are the ion charges:
Q1 x Q2 = cation charge x anion charge
d is the distance between ions in the crystal:
d = cation radius + anion radius.
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Lattice Energy cannot be easily measured directly. It is calculated
using the Born-Haber cycle and Hess’ Law.
Steps in the cycle:
1. Na(s) —> Na(g)
2. 1/2 Cl2(g) —> Cl(g) 3. Na(g) —> Na+(g) + e– 4. Cl(g) + e– —> Cl–(g) 5. Na+(g) + Cl–(g) —> NaCl(s)
4.
3.
∆H°f Na(g) ∆H°f Cl(g)
IE1(Na)
EA(Cl)
–∆Hlattice
= 108 kJ (sublimation)
= 122 kJ (1/2 bond energy)
= 496 kJ
= -349 kJ
=?
Add these 5 steps together: to get step 6, ∆H°f NaCl(s):
5.
6. Na(s)+ 1/2 Cl2(g) —> NaCl(s)
∆H°f NaCl(s)= -411 kJ
2.
1.
Start at standard states.
6.
Using Hess’ Law and solving for the energy of step 5:
5 = 6 - {1 + 2 + 3 + 4}
So we have:
–∆Hlattice = 5 = -411 kJ - {108 + 122 +496 + -349} kJ = –788 kJ
∆Hlattice = +788 kJ
NaCl(s) —> Na+(g) + Cl–(g)
End at standard state.
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Lattice Energy Calculations Using the Born-Haber Cycle
Write the appropriate equations and perform the calculations to determine the lattice energy
for the crystal CaCl2. Data are as follows:
Ca(s) to Ca(g) ∆Hsublimation = 178 kJ/mol
IE1 Ca(g) = 590 kJ/mol
IE2 Ca+(g) = 1145 kJ/mol
Cl2(g) bond energy = 242 kJ/mol
EA Cl(g) = -349 kJ/mol
∆H°f CaCl2(s) = -796 kJ/mol
Answer 2253 kJ
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Conceptual Questions
• lllustrated below are four ions (A1, A2, Z1 and Z2), showing their relative ionic radii. The ions
shown in red carry a 1+ charge, and those shown in blue carry a 1– charge.
(a) Would you expect to find an ionic compound of formula A1A2?
(b) Which combination of ions leads to the ionic compound having
the largest lattice energy?
(c) Which combination of ions leads to the ionic compound having
the smallest lattice energy?
• Energy is required to remove two electrons from Ca to form Ca2+ and is also required to
add two electrons to O to form O2−. Why, then, is CaO stable relative to the free
elements?
• Given that lattice energy increases as ionic charge increases, explain why compounds
such as NaO, where Na is +2, or BaCl, where Cl is –2, do not form.
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Conceptual Questions
• Formic acid (HCOOH) is secreted by certain species of ants when they bite (OUCH!!!).
Using the Periodic Table only, rank the relative lengths and strengths of
(a) the C–O and C=O bonds
(b) the H–C and the O–H bonds
• Use electronegativity values to determine if the bonds in each of the following substances
are mostly ionic, polar covalent or non-polar covalent:
KCl
CH4
SO2
• Using only the Periodic Table for reference, indicate the polarity for a C–Cl bond using
partial charges and a polar arrow
• Using only the Periodic Table, arrange the following in order of increasing ionic character.
SCl2, PCl3, and SiCl4 Larson-Foothill College
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How the Covalent Bonding Model Explains the Properties of
Molecular Compounds
•
•
Covalent bonding forces of attraction hold atoms together (molecules) (relatively
strong). These are intramolecular forces.
Weaker intermolecular forces acting between molecules determine the physical
properties such as melting and boiling point. The stronger these forces, the higher
the melting and boiling points.
Properties
• Fairly soft
• Low to moderate melting
points
• Poor electrical and thermal
conductors
• Solubility in water varies,
depends on polarity.
• Are weak or non-electrolytes
when dissolved in water.
Strong forces within molecules and weak forces between them.
Microscopic view inside a bubble in
boiling water.
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How the Covalent Bonding Model Explains the Properties of Network
Covalent Solids
•
•
Atoms held together in large networks of strong covalent bonds.
Properties reflect the strength of the covalent bonds.
Properties
• Hard
• High melting points
• Poor electrical and thermal
conductors (usually)
• Insoluble in water.
Diamond: mp = 3550°C
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How the Ionic Bonding Model Explains the Properties of Ionic
Compounds
•
•
Ionic forces of attraction (STRONG)
Melting points are HIGH and tend to increase with increasing lattice energy.
Properties
• Hard, but brittle
• High melting points
• Poor electrical conductors as
solid, good as liquid
In the solid state, the ions are
fixed in place in the lattice
and do not move.
• Many are water soluble.
Aqueous solutions conduct
electricity.
In solution, the ions are free
to move and carry a current.
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Why ionic compounds are hard, but
brittle (crack).
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How the Bonding Models Explain the Trend in Properties of
Covalent, Polar Covalent and Ionic Substances
As ΔEN decreases, melting point and electrical conductivity when in the molten
state decreases because the bond type changes from ionic to polar covalent to
nonpolar covalent.
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Metallic Bonding- The “Electron Sea” Model
All metal atoms in the sample contribute their valence electrons to form a
delocalized electron “sea”.
• The metal “ions” (nuclei with core electrons) lie in an orderly array within this mobile
sea.
• All the atoms in the sample share the electrons. In metals, valence electrons are
delocalized throughout the solid.
• Metallic forces of attraction (bonding): Attraction between positive cores and a sea of
valence electrons resulting in bonding that is non-directional in nature.
Electron Sea Model
Metallic Properties that are explained by the model:
• Malleable (can be shaped without breaking/cracking due to nondirectional bonding)
• Ductile (can be drawn into a thin wire)
• Good electrical and heat conductors (electrons are mobile)
• Wide range of melting points, but many have high melting points.
Boiling points are high.
Why metals dent and
bend rather than crack
(non-directional forces
of attration).
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