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Chapter 9 MODELS OF CHEMICAL BONDING
HB+ H
AH
BH
A
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A comparison of metals and nonmetals
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Types of Chemical Bonding
Ionic bonding involves the transfer of
electrons and is usually observed when a
metal bonds to a nonmetal.
Covalent bonding involves the
sharing of electrons and is usually
observed when a nonmetal bonds to a
nonmetal.
Metallic bonding involves electron
pooling and occurs when a metal bonds
to another metal.
9.1 Atomic Properties & Chemical Bonds
Chemical bond: A force that holds atoms together
in a molecule or compound
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Three models of chemical bonding
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Gradations in bond type among Period 3
and Group 4A elements
Most bonds are somewhere in between ionic and covalent.
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Lewis Symbols & the Octet Rules
Lewis symbols for atoms show
valence electrons only
Rules of the Game
# of valence electrons of a main group atom = Group number
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Lewis Symbols & The Octet Rule
• The OCTET RULE: Many atoms gain or lose electrons so as to end up with the same # of electrons as the noble gas closest to them on the periodic table
Example: Phosphorus has 5 valence electrons and want to get 3 more electrons in its valence shell
P Group 5A(15) 1s2 2s2 2p6 3s2 3p3
Br Group 7A(17) [Ar] 4s2 3d10 4p5
Core = [Ar] 3d10 , valence = 4s2 4p5
Bromine has 7 valence electrons and want to
get 1 more electron in its valence shell
Why do compounds form? -Elements with unfilled valence shells usually form compounds with other elements to gain stability.
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–Usually made up of a metal ion and one or more nonmetal ions
»Metals have low IE and tend to lose electrons (to form cations) – e.g., Li becomes Li+ ion.
»Nonmetals have large negative values of EA and tend to gain electrons (to form anions) – e.g., F becomes F- ion.
9.2 The Ionic Bonding Model
- Why don’t
compounds such
as Li2F, LiF2
exist?
Li • F ••
• •• •• Li+ + F
••
•• ••
••
-
Li 1s22s1 + F 1s22p5 → Li+ 1s2 + F- 1s22s22p6
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The importance of Lattice Energy DH0lattice
Why do ionic compounds form?
The Born-Haber Cycle
to determine Lattice
Energy
Lattice energy is the energy required to separate 1 mol of
an ionic solid into gaseous ions.
Lattice energy is a measure of the strength of the ionic
bond.
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Periodic Trends in Lattice Energy
Coulomb’s Law:
- As ionic size increases,
lattice energy decreases.
Lattice energy therefore
decreases down a group on
the periodic table.
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Periodic Trends in Lattice Energy
NaCl, Na+ and Cl-,
m.p. 804 oC
MgO, Mg2+ and O2-
m.p. 2800 oC
Coulomb’s Law:
- As ionic
charge increases,
lattice energy
increases.
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Properties of Ionic Compounds
• Ionic compounds tend to be hard, rigid, and brittle, with high melting points.
• Ionic compounds do not conduct electricity in the solid state.
– In the solid state, the ions are fixed in place in the lattice and do not move.
• Ionic compounds conduct electricity when melted or dissolved.
– In the liquid state or in solution, the ions are free to move and carry a current.
ionic
compounds are
easily cracked
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Interionic attractions are so
strong that when an ionic
compound is vaporized, ion
pairs are formed.
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9.3 The Covalent Bonding Model
Covalent bond arises from the mutual attraction of 2
nuclei for the same electrons. Electron sharing
results.
HB+ H
AH
BH
A
Covalent bond is a balance of attractive and repulsive
forces.
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The Covalent Bonding Model
Formation of a
covalent
bond results
in greater
electron
density
between the
nuclei.
Covalent bond
formation in H2.
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Bonding Pairs & Lone Pairs of electrons
• Valence electrons are distributed as
shared or BONDING PAIRS and
unshared or LONE PAIRS.
This is called a LEWIS ELECTRON DOT
structure of a molecule.
G. N. Lewis
1875 - 1946
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Covalent Bond Properties: Order, Energy. & Length
• What is the effect of bonding and structure on molecular properties?
Free rotation
around C–C single
bond
No rotation around
C=C double bond
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Bond Order # of electron pairs being shared between
a given pair of atoms
Double bond Single bond
Triple
bond
Acrylonitrile
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• —measured by the energy required to break a covalent bond. Bond Enthalpy has positive values.
• BOND STRENGTH (kJ/mol)
H—H 436
C—C 346
C=C 610
CC 835
NN 945
The GREATER the bond order the HIGHER the
bond strength and the SHORTER the bond length.
Bond Energy (Bond Strength, Bond Enthalpy)
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Bond Order Length Strength
HO—OH
O=O
1 142 pm 210 kJ/mol
2 121 498 kJ/mol
1.5 128 ?
Bond Strength
O O•••••
••
••
••O
••
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Bond Length
Bond length depends on size of bonded atoms.
H—F
H—Cl
H—I
Bond distances
measured in
Angstrom units
where 1 A = 10-2 pm =
10-8 cm
The distance between the
nuclei of two bonded atoms.
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Bond length also depends on bond order.
Bond distances measured in
Angstrom units where 1 A = 10-2 pm.
Bond Length
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Physical Properties of Covalent Substances
1) Poor electrical conductivity
2) Physical properties of
- Molecular covalent substances: gas, liquid, low-
melting point solid due to Strong and localized intramolecular (bonding) forces
Weak intermolecular forces
- Network covalent solids:
Covalent bonds
throughout the sample
Quartz (SiO2) mp 1550 0C
Diamond mp 3550 0C
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9.4 Bond Energy & Chemical Change
∆H0rxn = (Sum of energy required to break bonds)
+ (Sum of energy generated by newly formed bonds)
Using Bond Energies to Calculate DH0rxn
Guide: 1) Break ALL the reactant bonds to obtain individual atoms
2) Use the atoms to form ALL the product bonds
3) Add the bond energies with appropriate signs to obtain DH0rxn
∆H0rxn > 0 for broken bonds
∆H0rxn < 0 for formed bonds
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Problem: Using bond energies to calculate DH°rxn for HF
formation. H2 (g) + F2 (g) 2 HF (g)
∆H0rxn = (BEH2
+ BEF2) + 2 BEHF
∆H0rxn = (432 kJ/mol + 159 kJ/mol) + (– 2x565 kJ/mol)
∆H0rxn = -539 kJ/mol
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Using bond energies to calculate DH°rxn for the combustion of
methane. CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
∆H0rxn = [4 BECH4
+ 2 BEO=O] + [2 BEC=O + 4 BEO-H]
∆H0rxn = [(4 x 413) + (2 x 498)] + [(2 x -799) + (4 x -467)]
∆H0rxn = -818 kJ/mol
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Relative bond
strength and
energy from
fuels.
Bond Strengths &
Heat Released from Fuels & Foods
Weaker bonds
such as C-H
bonds (less
stable, more
reactive) are
easier to break
than stronger
bonds such as
C-O bonds
(more stable,
less reactive)
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Bond Strengths &
Heat Released from Fuels & Foods
Fuels with
fewer bonds
to O release
more energy.
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9.5 Electronegativity and Bond Polarity
A covalent bond in which the shared
electron pair is not shared equally, but
remains closer to one atom than the other,
is a polar covalent bond.
Unequal sharing of electrons causes the
more electronegative atom of the bond
to be partially negative and the less
electronegative atom to be partially
positive.
The ability of an atom in a covalent bond to
attract the shared electron pair is called its
electronegativity.
H Cl••
••
+ -
••
Due to the bond polarity, the H—Cl bond energy is GREATER than expected for a “pure” covalent bond.
BOND ENERGY
“pure” bond BEHCl = 339 kJ/mol
real bond BEHCl = 432 kJ/mol
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A measure of the ability of a bonded atom to attract shared electrons
Relative values of EN determine BOND POLARITY
Electronegativity (EN)
F has
highest
EN
value
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© 2006 Brooks/Cole - Thomson Figure 9.22
Electronegativity and atomic size.
Trends in Electronegativity
EN increases
as atomic size
decreases.
Nonmetals
have higher
EN than
metals.
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Electronegativity and Oxidation Number
• 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.
Electronegativities can be used to assign oxidation numbers (ON):
O.N. = # of valence e- - (# of shared e- + # of unshared e-)
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Example:
Cl is more electronegative than H, so for Cl:
valence e- = 7
shared e- = 2
unshared e- = 6
O.N. = 7 – (2 + 6) = -1
H is less electronegative than Cl, so for H:
valence e- = 1
shared e- = 0 (all shared e- assigned to Cl)
unshared e- = 0
O.N. = 1 – (0 + 0) = +1
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Depicting Polar Bonds
The unequal sharing of electrons can be depicted by a polar
arrow. The head of the arrow points to the more electronegative
element.
A polar bond can also be marked using δ+ and δ- symbols.
Bond Polarity & Partial Ionic Character
H Cl••
••
+ -
••
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The Importance of DEN
- Percent ionic
character of a
bond increases
with DEN .
- No bonds are
purely ionic or
covalent .
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Determine Bond Polarity from DEN
Which bond is more polar?
O—H O—F
DEN 3.5 - 2.1 3.5 - 4.0
DEN 1.4 0.5
OH bond is more polar than OF bond, and direction of polarity is opposite in these cases.
O H
+-
O F+ -
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Properties
of the
Period 3
chlorides.
As DEN decreases, melting point and electrical conductivity
decrease because the bond type changes from ionic to polar
covalent to nonpolar covalent.
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
The Graduation in Bonding across a Period
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9.6 Introduction to 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.
• The metal is held together by the attraction between the metal “cations” and the “sea” of valence electrons.
Group
2A/2
Group
1A/1
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Properties of Metals • Metals are generally solids with moderate to high melting
points and much higher boiling points.
–Melting points decrease down a group and increase across a period.
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Properties of Metals - Mechanical properties: Metals can be shaped without breaking.
– The electron sea allows the metal ions to slide past each other.
Metals dent or
bend rather than
crack
- Electric Conductivity: Metals are good conductors of
electricity in both the solid and liquid states. The electron sea is mobile in both phases.
- Heat Conductivity: Metals are good conductors of heat.