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Chapter 8:
Bonding:Genera
l Concept
s
Cartoon courtesy of NearingZero.net
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Chemical Bonding Forces that hold groups of atoms together and make them function as a unit.
Bond EnergyEnergy required to break a bond
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Bond Polarity and Dipole Moments
Dipolar Molecules1. Molecules with a somewhat negative end and a somewhat positive end (a dipole moment)2. Molecules with preferential orientation in an electric field
+ + + + + + + +
- - - - - - - -3. All diatomic molecules with a polar covalent bond are dipolar
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Bond Polarity and Dipole Moments
Molecules with Polar Bonds but no Dipole Moment1. Linear, radial or tetrahedral symmetry of charge distribution
a. CO2 - linearb. CCl4 – tetrahedral
See table 8.2
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Ionic Bonding
Ionic bond: the electrostatic force that holds ions together in an ionic compound.
Examples of Ionic Compounds (aka Salts):NaClBaCl2
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Ionic Bonding Electrons are transferred
Electronegativity differences are
generally greater than 1.7 The formation of ionic bonds is
always exothermic!
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Determination of Ionic
Character
Compounds are ionic if they conduct electricity in their molten state
Electronegativity difference is not the final determination of ionic character
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Properties of Ionic Compounds
Structure: Crystalline solids
Melting point:
Generally high
Boiling Point:
Generally high
Electrical Conductivity:
Excellent conductors, molten and aqueous
Solubility in water:
Generally soluble
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Coulomb’s Law
r
QQnmJxE 2119 )1031.2(
“The energy of interaction between a pair of ions
• E is in Joules• r is the distance between the center of the ions
• Q1 and Q2 are the charges of the ions• A negative quantity indicates attraction • A positive quantity indicates repulsion
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Coulomb’s Law
• Example: In solid NaCl the distance between the centers of the ions is 2.76 Å (0.276 nm) Calculate the ionic energy per pair of ions:
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Sodium Chloride Crystal Lattice
Ionic compounds form solid crystals at ordinary temperatures.
Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.
All salts are ionic compounds and form crystals.
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12
Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.
Electronegativity - F is highest
X (g) + e- X-
(g)
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Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
Increasing difference in electronegativity
Classification of bonds by difference in electronegativity
Difference Bond Type
0 Covalent
2 Ionic
0 < and <2 Polar Covalent
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14
Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2.
Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3Ionic
H – 2.1 S – 2.5 2.5 – 2.1 = 0.4Polar Covalent
N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent
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Formation of Ionic compounds
• Stable compounds are formed when nonmetallic elements take electrons from metals.
• Atoms usually have a noble gas configuration
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Formation of Ionic compounds
In general: • When a binary ionic compound is
formed – the nonmetal has noble gas
configuration– The valence orbitals of the
representative metal is emptied
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• The term ionic compounds refers to the solid state of the compound
• A collection of positive and negative ions arranged to minimize repulsions and maximize attractions
Predicting formulas of ionic compounds
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Predicting formulas of ionic compounds
• Large electronegativity differences between atoms mean electrons will be transferred
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Predicting formulas of ionic compounds
• Hydrogen typically behaves as a nonmetal
• The number of electrons transferred depends on how many each atom needs to gain or lose to achieve noble gas notation
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Predicting formulas of ionic compounds• EXCEPTIONS:
– Tin forms Sn2+ and Sn4+
– Lead forms Pb2+ and Pb4+
– Bismuth forms Bi3+ and Bi5+
– Thallium forms Tl+ and Tl3+
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Energy and Binary Ionic Compounds
• Factors that influence stability and structure
• Ionic compounds form because together they have lower energy than the original elements
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Lattice Energy
• The energy released when an ionic solid is formed from its ions
• LE is negative (exothermic)• Used as a step to calculate
energy of formation
MX(s)(g)X(g)M
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Lattice energy increases as Q increases and/or as r decreases.
CompoundLattice Energy (kJ/mol)
MgF2
MgO
LiFLiCl
29573938
1036853
Q: +2,-1Q: +2,-2
r F- < r Cl-
Q+ and Q- is the charge on the cation and anionr is the distance between the ions
E is the potential energy
k is a constant based on the compound
r
QQkE
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Calculating Energy of formation Hf
• If we know the steps in the process then we can apply Hess’s law
• Because energy is a state function
• Break the reaction up into steps• Add them up
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Estimate Hf for Sodium Chloride
Na(s) + ½ Cl2(g) NaCl(s)Lattice Energy -786 kJ/mol
Ionization Energy for Na 495 kJ/mol
Electron Affinity for Cl -349 kJ/mol
Bond energy of Cl2 239 kJ/mol
Enthalpy of sublimation for Na
109 kJ/mol
Na(s) Na(g) + 109 kJNa(g) Na+(g) + e- + 495 kJ
½ Cl2(g) Cl(g) + ½(239 kJ)Cl(g) + e- Cl-(g) - 349 kJ
Na+(g) + Cl-(g) NaCl(s) -786 kJ Na(s) + ½ Cl2(g) NaCl(s) -412
kJ/mol
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The energy diagram for the formation of MgO and NaCl.
The Lattice energy to combine Mg2+ and O2- is much more negative than the energy needed for the process that produces Mg2+ and O2- ions.