Chemical Bonding and Molecular Structure
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Bonds are attractive forces that hold groups of atoms together and make them function as a unit.
o Being bound requires less energy than existing in the elemental form. Energy is released
when a bond is formed; energy is required to break a bond. The energy required to break a
bond is called the bond energy.
o Ionic bonds
o Covalent bonds
Coulomb’s Law is used to calculate the energy of an ionic bond (see
equation on the right; k = 2.31 x 10-19 Jnm, Q = charge on each ion)
o The energy of an ionic bond will be negative; it indicates an
attractive force so that the ion pair has lower energy than the
separated ions.
Chemical bond formation
o When two hydrogen atoms approach each other, two repulsions & one attraction occur
 Electron/electron repulsion
 Proton/proton repulsion
 Proton/electron attraction
o When the attractive forces offset the
repulsive forces, the energy of the two
atoms decreases and a bond is formed.
Bond length is the distance between two
nuclei where the energy is at a minimum.
o The electrons and nucleus of one atom
strongly perturb or change the spatial distribution of the other atom’s valence electrons. A
new orbital (wave function) is needed to describe the distribution of the bonding electrons
bond orbital. The energy of the electrons in a bond orbital is lower that their energy in
valence electron orbitals when they are in isolated atoms.
 In an ionic bond, the bonding orbital is strongly displaced toward one nucleus.
 In a covalent bond, the bond orbital is more or less evenly distributed and the
electrons are shared by two nuclei.
 Most chemical bonds are somewhere between purely ionic and purely covalent.
Electronegativity is the ability of an atom in a molecule to attract shared electrons to itself.
o Fluorine is the most electronegative (4.0) due to highest effective nuclear charge (Zeff) and
smallest radius—so that the nucleus is closest to the “action”
o Francium is the least (0.7) due to lowest Zeff and largest radius so that the nucleus is farthest
from the “action”
o This atomic trend is only used when atoms form molecules.
 Ionic: ΔEN >1.67
 Covalent: ΔEN < 1.67
 Nonpolar covalent: ΔEN < 0.4
Covalent bonding
o Most compounds are covalently bonded, especially carbon compounds.
o Localized electron (LE) bonding model: assumes that a molecule is composed of atoms that
are bound together by sharing pairs of electrons using the atomic orbitals of the bound
atoms. Electron pairs are assumed to be localized on a particular atom [lone pairs] or in the
space between two atoms [bonding pairs].
 Lewis structures describe the valence electron arrangement
 Geometry of the molecule is predicted with VSEPR
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Number of bond pairs/Octet rule
Single and multiple bonds
 Single bond: one pair of electrons shared (sigma σ bond)
 Double bond: two pairs of electrons shared (one sigma bond, one pi bond)
 Triple bond: three pairs of electrons shared (one sigma bond, two pi bonds)
 Pi bonds are weaker than sigma bonds but never exist alone. Triple bonds are
stronger than double bonds are stronger than single bonds.
 Multiple bonds are most often formed by C, N, O, and P.
 Multiple bonds increase the electron density
between two nuclei and therefore decrease the
nuclear repulsions while enhancing the attraction
between nucleus/electrons; either way, the nuclei
move closer together and the bond length is shorter
for a double bond than a single bond.
o Exceptions to the Octet rule
 Fewer than eight: H, Be, B
 Expanded valence: can only happen if the central atom has d-orbitals and can thus
be surround by more than four valence pairs in certain compounds.
 Odd-electron compounds: ex. NO, NO2, ClO2
Exercise 1
Relative Bond Polarities
Order the following bonds according to polarity: H-H, O-H, Cl-H, S-H, and F-H
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Bond polarity and electronegativity: electronegativity determines polarity since it measures a
nucleus’s attraction or “pull” on the bonded electron pair. When two nuclei are the same, the
sharing is equal and the bond is described as nonpolar. When two nuclei are different, the
electrons are not shared equally, setting up slight +/- poles, and the bond is described as
polar. When the electrons are shared very unequally, the bond is described as ionic.
 Ionic bonding
o The final result of ionic bonding is a solid, regular array of cations and anions called a crystal
lattice.
o Enthalpy of dissociation: energy required to decompose an ion pair (from a lattice) into ions;
a measure of the strength of the ionic bond (related to Coulomb’s law)
 The energy of attraction depends directly on the magnitude of the charges and
inversely on the distance between them (related to the size of the ion).
Exercise 2
Comparing Lattice Energy
Which compound in each pair will have the higher lattice energy?
NaF or RbF
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MgO or LiCl
Drawing Lewis structures
o H is always a terminal atom
o Atom with lowest EN goes in center
o Find the total number of valence electrons by adding together the valence electrons of every
atom in the compound
 For ions, add for negative charges and subtract for positive charges
o Place one pair of electrons, a sigma bond, between each pair of bonded atoms.
o Complete the octets of all atoms with lone pairs. Leftover pairs are assigned to the central
atom if it can accommodate them. Double/triple bonds may need to be used (pi bonds).
Exercise 3
Writing Lewis Structures
Give the Lewis structure for each of the following.
a) HF
c) NH3
e) CF4
b) N2
d) CH4
f) NO+
Exercise 4
Lewis Structures for Molecules that Violate the Octet Rule
Give the Lewis structure for each of the following.
a) PCl5
c) XeO3
e) BeCl2
b) ClF3
d) RnCl2
f) ICl4-
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Resonance structures
o Ex. ozone has equal bond lengths and
equal bond strengths, implying that there
are an equal number of bond pairs on
each side of the central oxygen atom. The Lewis structure does not agree with this; instead,
we have to use a composite to describe the reality. This composite depicts the blending of
resonance structures for ozone. Instead of truly having a single bond and a double bond,
both of its C-O bonds could be thought of as “a bond and a half”.
o Resonance structures differ only in the assignment of electron pair positions, never atom
positions. They differ in the number of bond pairs between a given pair of atoms.
o Note that the resonance structures and composite are drawn with brackets (required for full
credit on AP exam).
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Exercise 5
Resonance Structures
Draw every resonance structure for the carbonate ion. Also draw the composite structure.
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Bond properties
o Bond order: simply the number of bonding electron
pairs shared by two atoms in a molecule.
 1 = one shared pair; sigma bond between two
atoms
 2 = two shared pairs; sigma bond and pi bond
 3 = three shared pairs; sigma bond and two pi
bonds
 Fractional for resonance structures (3/2 for ozone, 2/3 for carbonate)
Bond order = number of shared pairs linking X and Y
number of X-Y links
o Bond length: the distance between the nuclei of two bonded atoms
 Higher bond order = shorter length
o Bond energy: the greater the number of electron pairs between a pair of atoms, the shorter
the bond. This implies that atoms are held together more tightly when there are multiple
bonds, so there is a relationship between bond order and the energy required to break a
bond.
o Bond dissociation energy (D): the energy supplied to break a chemical bond
 Endothermic; D is positive
o Bonds in reactants are broken while bonds in products are formed. Energy released is
greater than energy absorbed in exothermic reactions. The converse is also true.
ΔH°rxn = ΣmD(bonds broken) - ΣnD(bonds made)
ΔH°rxn = reactants(E cost) - products(E payoff)
 Note that this is “backwards” from thermodynamics. First we must break the bonds of
the reactants (costs energy) then subtract the energy gained by forming new bonds in
the products.
Exercise 6
ΔH from Bond Energies
Using the bond energies from the table on
the right, calculate ΔH° for the reaction of
methane with chlorine and fluorine to give
Freon-12 (CF2Cl2).
CH4 + 2Cl2 + 2 F2  CF2Cl2 + 2 HF + 2 HCl
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Formal charge
o Formal charge is the difference between the number of valence electrons on a free element,
and the number of electrons assigned to the atom once it is in a molecule.
Formal charge = group number – [# of lone electrons – 2(# of bonding electrons)]
o The ideal of formal charge allows us to determine the most favored structure out of a set of
nonequivalent Lewis structures.
o Oxidation states of more than +/- are questionable, while formal charges are more realistic.
o The sum of the formal charges on an ion must equal the ion’s overall charge.
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Use formal charges along with the following to determine resonance structure
 Atoms in molecules (or ions) should have formal charges that are as small (close to
zero) as possible
 A molecule (or ion)is most stable when any negative formal charge resides on the
most electronegative atom.
 Ex. There are three possible structures for the sulfate ion shown below (note that
these are not resonance structures). The third is the most valid of the three; it results
in the fewest (and smallest) formal charges.
Exercise 7
Formal Charges
Give possible Lewis structures for XeO3, an explosive compound of Xenon. Which Lewis structure or
structures are most appropriate according to the formal charges?
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Valence shell electron pair repulsion theory (VSEPR)
Sigma bonds + Structural Pair or Molecular
o Molecular shape changes with the number of
lone pairs on
Geometry (not the same as
sigma bonds plus lone pairs about the central
central atom
molecular geometry)
atom
2
Linear
3
Trigonal Planar
o Molecular geometry is the arrangement in
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Tetrahedral
space of the atoms bonded to a central atom
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Trigonal Pyramidal
 lone pairs take up more space around
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Octahedral
an atom than bonds
 Each lone pair or bond pair repels all other electron pairs; they try to avoid each other
making as wide an angle as possible.
 Ex. Water: the two lone pairs on oxygen “warp” the normal 109.5 angle
through repulsion, resulting in a bond angle of 104.5.
 To determine molecular geometry:
 Sketch the Lewis dot structure
 Describe the structural pair or electronic geometry (the shape of the
molecule considering both its bonds and lone pairs)
 Focus on the bond locations (ignore lone pairs) and assign a molecular
geometry based on their locations
 Molecular geometry and electronic geometry are only the same in the
absence of lone pairs on the central atom.
o Works well for elements of the s and p-blocks; does not apply to transition element
compounds (exceptions)
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Molecular shapes for central atoms with normal valence
 No more than 4 structural pairs if the atom obeys the Octet rule
 The combination of s and p orbitals provides four bonding sites
Molecular shapes for central atoms with expanded valence
 Only elements with a principal energy level of 3 or higher can expand their valence
and violate the octet rule on the high side. This is because d orbitals are needed for
expansion to a 5th or 6th bonding location.
 Lone pairs “want” to be as far apart as possible.
For example, look at the two possible structures
for XeF4 on the right. The equatorial
configuration is favored (lower energy) because
it allows the lone pairs to be as far from each
other as possible.
 In general, when you need to determine
the most stable shape for a molecule,
you should think about repulsions.
Everything wants to be as far from
everything else as possible (“everything”
being electrons, of course).
 Another example: SF4 has one lone pair and four
S-F bonds. Configuration (a) is favored because
it minimizes total repulsions (two 90° and two
129° lone-bonding repulsions vs. three 90° and one 180° lone-bonding repulsions)—
see saw shape.
Exercise 8
Prediction of Molecular Structure I
Which of the three following arrangements do you predict to be the most stable for I 3-?
Exercise 9
Prediction of Molecular Structure II
Draw and describe the molecular structure of the water molecule.
Exercise 10
Prediction of Molecular Structure III
When phosphorous reacts with excess chlorine gas, the compound phosphorous pentachloride is
formed. In the gaseous and liquid states, this substance consists of PCl5 molecules, but in the solid
state it consists of a 1:1 mixture of PCl4+ and PCl6- ions. Predict the geometric structures of PCl5,
PCl4+, and PCl6-.
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Molecular Polarity
o Bonds can be polar while the entire
molecule isn’t, and vice versa.
o Dipole moment: separation of the
charge in a molecule; product of the
size of the charge and the distance
of separation.
 Molecules align themselves
with an electric field
 Molecules align with each
other in the absence of an electric field
 The direction of the “arrow” indicating the dipole
moment always points to the negative pole with the
cross hatch on the arrow at the positive pole.
 In water, two lone pairs establish a strong negative
pole. Similarly, a negative pole is established by
the one lone pair in ammonia. This negative pole
is flipped if we substitute the hydrogen atoms for
fluorine to make nitrogen trifluoride.
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If the Octet rule is obeyed and all the surrounding bonds are
the same, then the molecule is nonpolar since all the dipole
moments cancel each other out.
 Ex. carbon dioxide is nonpolar since the dipole
moments cancel
 Ex. methane, CH4, and carbon tetrachloride, CCl4, are
both nonpolar since the dipole moments in each
molecule cancel. However, the molecules “between”
these two (shown below) all have net dipoles and are therefore polar.
Exercise 11
Bond Polarity and Dipole Moment
For each of the following molecules, show the direction of the bond polarities and indicate which ones have
a dipole moment: HCl, Cl2, SO3, and H2S.
Exercise 12
Prediction of Molecular Structure IV
Predict the molecular structure of the sulfur dioxide molecule. Is this molecule expected to have a dipole
moment?
6) -1194 kJ/mo
7)
8) Arrangement (c) is the most stable due to 180° separation of I-I bonds and 120° separation of three lone
pairs.
9) Two bonding and two lone pairs  v-shaped/bent