Chemical Bonding
Objectives:
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Understand and be able to describe the difference between ionic, covalent and metallic
bonding. Understand how each model is consistent with observed properties.
Know the definition of Lattice Energy and be able to calculate Lattice Energy from appropriate
thermodynamic quantities (Born-Haber Cycle: Hess’s Law).
Understand and be able to predict trends in Lattice Energies based on ion size and charge;
know how these trends are related to melting points of ionic compounds.
Understand covalent bond formation and properties; bond enthalpy (strength) and bond length.
Be able to predict relative bond enthalpies and lengths based on atomic size and bond order.
Understand and be able to describe electronegativity; use electronegativity to predict polarity of
bonds and diatomic molecules.
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Three Types of Chemical Bonding
(Ionic, Covalent and Metallic)
Octet Rule: Transfer/Sharing of electrons to obtain a complete valence electron shell of 8 electrons (except H has a
valence shell of 2 electrons, or “Duet Rule”. WHY IS THIS?). Applies to ionic and covalent bonding.
Ionic Bonding: Transfer of electrons through redox to obtain 8 valance shell electrons. (Octet Rule applies mainly
to the Main Group Elements.)
Covalent Bonding: Sharing of electrons to obtain 8 valance shell electrons.
(Exceptions to the octet rule exists for covalent compounds.)
Metallic Bonding: “Electron-Sea Model” where atoms in a metal (or metals) pool their valence electrons into a “sea”
of electrons that flow around each metal-ion core (the nucleus plus core electrons).
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Gradations in bond type among Period 3
(black type) and Group 4A (red type)
elements.
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Lewis Symbols of Elements
In writing the Lewis Symbol for an atom, the chemical symbol for the element is
used to represent the core (the inner, non-valence electrons plus the nucleus) and
the valence electrons are represented by dots arranged around the core. The dots
are placed arbitrarily, one at a time, around the symbol. If there are more than four
valence electrons, we continue around the symbol, pairing the dots until all the
valence electrons are represented by dots.
Example: Draw the Lewis Symbols of H and Cl and C.
Can you use the Lewis Symbols of H and Cl to predict that hydrogen and chlorine
should be diatomic elements? (Think “duet rule” for H and “octet rule” for Cl.)
Can you use the Lewis Symbols of C and Cl to predict the formula of the simplest
compound formed between these two elements?
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Chemical Bond Formation: Covalent - Sharing of Electrons.
Covalent bonding: sharing of electrons
by orbital overlap. Each atom (except H)
obtains an octet of electrons.
Consider the formation of H2 and Cl2:
Covalent bond formation in H2.
Covalent bonding results in “molecules”,
individual units of H2 and Cl2. Covalent
molecules exist as discrete molecules.
We will show the bonding in covalent
molecules using Lewis Structures. Lewis
Structures are covered in detail in lab.
The shared electrons are called a shared
pair or bonding pair. An outer-level
electron pair that is not involved in
bonding is called a lone pair, or
unshared pair.
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Attractive
Forces Between
Shared Valence
Electrons and
Nuclei.
Repulsive Forces:
Electron-Electron and
Nucleus-Nucleus
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Chemical Bond Formation: Covalent - Sharing of Electrons.
Covalent bonding: Methane
The formation of methane:
C(s) + 2 H2(g) —> CH4(g)
(8 electrons now surround C)
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Covalent Bonds: Bond Order and Bond Length
Bond order: A single bond is described as having a bond order of 1, a double bond as
having a bond order of 2 and a triple bond as having a bond order of three.
Single bond: A covalent bond involving one shared electron pair.
Triple bond: A covalent bond involving three shared electron pairs.
Double bond: A covalent bond involving two shared electron pairs.
Bond length: The distance between the nuclei of
the atoms involved in a bond. Bond length
depends on the size of the bonded atoms and on
the bond order. Generally, bond length decreases
with a decrease in atomic radii of the bonded
atoms and with an increase in bond order.
Bond Order:
Except for
N≡N, these
bond lengths
are averages.
1
2
C–C
C=C
154 pm 134 pm
3
C≡C
121 pm
N–N
N=N
146 pm 122 pm
N≡N
110 pm
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Covalent Bonds: Bond Strength
Bond enthalpy (energy, strength): The bond enthalpy, BE, is the enthalpy change (∆H) per
mole required to break a bond in the gas phase. Bond enthalpy depends on the two atoms
bonding. Bond enthalpy generally increases with decreasing bond length, shorter bonds
tend to be stronger. A double bond is stronger than a single bond (between the same two
atoms) and a triple bond is stronger than a double.
H2(g) —> 2 H (g)
∆H = bond enthalpy = 432 kJ/mol
Bond breaking is always an endothermic
process, so bond enthalpies are (+) values.
Bond Order:
1
C–C
Except for
347 kJ/mol
N≡N, these
bond energies
N–N
are averages.
2
3
C=C
C≡C
614 kJ/mol 839 kJ/mol
N=N
160 kJ/mol 418 kJ/mol
N≡N
945 kJ/mol
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Chemical Bond Formation: Potential Energy Diagrams
Covalent bond formation in H2.
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Comparison of covalent bond formation of H2
and the halogens.
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Table 9.2 Average Bond Energies (kJ/mol) and Bond Lengths (pm)
Let’s discuss the trend in bond energy and bond length for HF, HCl, HBr, and HI.
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Using Bond Enthalpies (Energies) to Estimate ∆Hrxn
The heat released or absorbed during a chemical change is due to differences between
the bond energies of reactants and products.Therefore, the heat of reaction between
covalent molecules can be estimated from bond energies, BE.
∆Hrxn = ∑ (bond energies broken) - ∑ (bond energies formed)
∆Hrxn= ∑ BEreactants) - ∑ (BEproducts)
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2
= [BE(C-H) + BE(Cl-Cl)]-[BE(C-Cl)+BE(H-Cl)]
= [(413+243)-[339+427] kJ = –110 kJ
The actual value is -101.1 kJ, so our estimate
has a 9% error.
Why is this an estimate?
CH4 (g) + Cl2 (g) —> CH3Cl (g) + HCl (g)
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Using Bond Energies to Estimate ∆Hrxn
Ammonia is produced directly from nitrogen and hydrogen by using the Haber process. The
chemical reaction is
3 H2(g) + N2(g) —> 2 NH3(g)
(a) Use bond energies to estimate the enthalpy change for the reaction. Is it exothermic
or endothermic.
(b) Compare the enthalpy change you calculated in (a) to the true enthalpy change as
obtained using ∆H°f values.
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Using Bond Enthalpies to Calculate Heat of Reaction
An important reaction for the conversion of natural gas to other useful hydrocarbons is
the conversion of methane to ethane.
In practice, this reaction is carried out in the presence of oxygen, which converts the
hydrogen produced to water.
Use bond energies to estimate ΔH for these two reactions. Why is the conversion of
methane to ethane more favorable when oxygen is used?
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