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AP Chemistry--Chapters 8,9,10: Bonding, Orbitals, and Liquids and Solids
I. Bond Formation
A. Bond Energy--Bonds result from the tendency of a system to seek its lowest
possible energy (table pg. 365); bond energy or bond dissociation energy is
the energy required to break a bond
B. Electronegativity is the relative tendency of an atom in a molecule to attract
shared electrons to itself
1. the most active metals have the lowest electronegativities, while fluorine
has the highest electronegativity—what reasons can you give for this
trend? (F2 has a very high bond dissociation energy)
2. notice, pg. 347, electronegativities not given for some the noble gases—
why?
3. Electronegativity developed by Linus Pauling, by comparing expected and
actual bond energies; can use difference in electronegativity to determine
whether a bond is ionic or covalent
4. General rule: large difference = ionic (electrons transferred), in between =
polar covalent (electrons shared but unequally), 0 = covalent (electrons
shared
Practice Problem 8.1
Order the following bonds according to polarity: H–H, O–H, Cl–H, S–H, F–H.
C. Polarity
1. In a polar covalent bond, a partial positive end of the bond is formed,
along with a partial negative end (δ+ and δ–)
2. Polar covalent (or just polar) bonds may produce a polar molecule, a
molecule in which one end is partially negatively charged, and the other
end is partially positively charged (*must take shape into account to
determine if a molecule is polar or not)
3. Polar molecules are called dipoles; they have two oppositely charged
“poles”; dipole moment is a measure of the strength of the dipole and is a
property that results from the asymmetrical charge distribution in a polar
molecule (water most important polar molecule!)
a. a very polar molecule gives a large dipole moment, a large EN
difference makes the molecule more polar
b. the more polar a gas is, the more it deviates from ideal behavior
c. a nonpolar molecule has a dipole moment of 0
Practice Problem 8.2
For each of the following molecules, show the direction of the bond polarities and indicate which
ones have a dipole moment: HCl, Cl2, SO3, CH4, H2S.
AP chem chapter 8,9,10 notes and practice problems
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D. Ionic Bonds
1. An ionic bond is defined as the electrostatic force that holds two ions
together due to their differing charges
2. Oxidation numbers for atoms involved in ionic bonds correspond to how
many electrons are lost or gained to fill an octet when forming compounds
3. Ionic compounds are compounds which contain an ionic bond and are
characterized by high melting points, the ability to conduct electricity in
the molten state, being soluble in water, and crystallizing as sharply
defined particles
4. Sizes of “noble gas configuration” ions, pg. 354, more negative = larger,
more positive = smaller
Practice Problems
2–,
–
+
2+
8.3
Arrange the ions Se Br , Rb , Sr in order of decreasing size.
8.4
Choose the largest ion in each of the following groups of ions.
+
+
+
+
+
2+
+ –
2–
a.
Li , Na , K , Rb , Cs
b. Ba , Cs , I , Te
5. Lattice Energy
a. defined as the energy change that occurs in the conversion of an ionic
solid to widely separated gaseous ions
b. the higher the charge of the ions combining, the higher the lattice
energy (NaF has lower LE than Na2O)
c. higher lattice energies also lead to higher melting points
6. % Ionic Character = (measured dipole moment/calc’d dipole moment) ×
100
E. Covalent Bonds
1. A covalent bond is formed by the shared pair of electrons between two
atoms with the same electronegativity
2. Covalent compounds are made up of covalent bonds and have low melting
points, are unable to conduct electricity, and are brittle
3. When two or more atoms bond covalently, the resulting particle is called a
molecule
4. Some terms:
a. bond dissociation energy—the total energy required to break the bond
between two covalently bonded atoms
b. bond axis—line joining the nuclei of two bonded atoms in a molecule
c. bond angle—the angle between two bond axes
AP chem chapter 8,9,10 notes and practice problems
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d. bond length—the distance between two nuclei along the bond axis (as
bond length decreases, bond strength increases, table, pg. 365)
5. Bonds are not stationary; they can bend, wag, rotate, and vibrate, so values
of the above terms vary, so an average is used
F. Electron Distribution--Electron Dot Diagrams
1. also called Lewis Structures
2. help to show how electrons are arranged in a molecule or polyatomic ion
3. electrons are arranged as dots around the atoms so that each atom ends up
with a full outer level
4. shared pairs of electrons are electrons shared by two different atoms and
are illustrated with a single line
5. unshared pairs of electrons belong to just one atom and are illustrated with
two dots
6. Recall the octet rule (duet rule for hydrogen); also exceptions to the octet
rule exist
a. C, N, O, F always obey octet rule
b. B has 6 e–, not 8 e–, so 3 pairs; Be often fewer than 8 (usually 4)
c. S often extends to 10 e– (5 pairs) or 12 e– (6 pairs); I is another that
sometimes extends its octet
d. any 3rd and higher row element can exceed octet rule (why?)
e. molecules can also “disobey” the octet rule by not pairing all
electrons; paramagnetism is associated with unpaired electrons, while
diamagnetism is associated with paired electrons (ex. O2)
7. Resonance
a. occurs when more than one valid Lewis structure exists
b. real molecule is an average of all possible Lewis structures
8. Formal Charge
a. To calculate formal charge on an atom:
1) Take the sum of the lone pair electrons and ½ the shared electrons.
This is the number of valence electrons assigned to the atom in the
molecule.
2) Subtract the number of assigned electrons from the number of
valence electrons on the free, neutral atom to obtain the formal
charge.
b. The sum of the formal charges of all atoms in a given molecule or ion
must equal the overall charge on that species.
c. If nonequivalent Lewis structures exist for a species, those with formal
charges closest to zero and with any negative formal charges on the
most electronegative atoms are considered to best describe the bonding
in the molecule or ion
Example: The sulfate ion
AP chem chapter 8,9,10 notes and practice problems
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Practice Problems
+
8.6 Give the Lewis structure for HF, N2, NH3, CH4, CF4, NO .
8.7 Write the Lewis structure for PCl5.
–
8.8 Write the Lewis structure for ClF3, XeO3, RnCl2, BeCl2, ICl4 .
–
8.9 Write the Lewis structure for NO2 .
8.10
Give the possible Lewis structures for XeO3, an explosive compound of xenon. Which
Lewis structure or structures are most appropriate according to the formal charges?
G. Electron Pair Repulsion
1. Electron pairs spread as far apart as possible to minimize repulsive forces
2. This is sometimes called the VSEPR Theory--valence shell electron pair
repulsion theory
3. Two reasons for VSEPR Theory
a. electrons all have same charge, so they form a charge cloud around an
atom and push away from other charge clouds around other atoms
b. Pauli exclusion principle--electrons with same magnetic spin cannot
occupy the same space (or orbital)
4. The VSEPR theory gives molecules certain shapes and angles between
atoms in the molecule--pg. 386
5. electron pair repulsions in a molecule may not all be equal; unsharedunshared repulsion>unshared-shared repulsion>shared-shared repulsion
6. the above is true because unshared-unshared take up the most space,
shared-shared the least
Practice Problems
8.11
Describe the molecular structure of the water molecule.
AP chem chapter 8,9,10 notes and practice problems
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+
–
8.12
Predict the geometric structures of PCl5, PCl4 , and PCl6 .
8.13
Predict the structure of XeF4 and whether it has a dipole moment.
8.14
Predict the molecular structure of SO2 and whether it has a dipole moment.
II. Orbitals
A. Paramagnetism and Diamagnetism
1. Most substances have no magnetism until they are placed in a magnetic
field.
2. When placed in a magnetic field, two types of magnetism can be induced
(induced here meaning caused; then it goes away when the magnetic field
is removed)
3. Paramagnetism (stronger of the two) causes the substance to be attracted
into the inducing magnetic field; paramagnetism is associated with
unpaired electrons
4. Diamagnetism (much weaker) causes the substance to be repelled from the
inducing magnetic field; diamagnetism is associated with paired electrons.
5. O2 is (oddly) a paramagnetic molecule
B. Hybrid Orbitals
1. The best example of hybridization occurs in carbon, although many other
atoms that form covalent bonds can also hybridize
2. To sp3 hybridize means that one s and three p orbitals merge to form 4
new, identical orbitals
3. The new orbitals are called hybrid orbitals; the bonding of carbon occurs
when orbitals overlap (this is a second way to look at how bonding occurs
and how molecules get their shape)
4. Relationship between shape and hybridization type, pg. 415
a. whenever four effective pairs surround a central atom, that atom is sp3
hybridized
b. whenever three effective pairs surround a central atom, that atom is sp2
hybridized
c. two effective pairs gives sp hybridization
d. five effective pairs give a trigonal bipyramidal shape and sp3d
hybridization; six effective pair give an octahedral shape and sp3d2
hybridization
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C. Sigma and Pi bonds
1. A covalent bond forms when an orbital of one atom overlaps an orbital of
another atom and they share the electron pair in the bond
2. A sigma bond (σ) occurs when two orbitals of any type overlap and form a
bond that lies directly on the bond axis
3. A pi bond (π) occurs when two p orbitals overlap parallel to each other
4. A double bond forms when two pairs of electrons are shared between two
atoms; a double bond always consists of one sigma bond and one pi bond
5. A triple bond forms when three pairs of electrons are shared between two
atoms; a triple bond can contain two pi bonds and one sigma bond
6. Multiple bonds are shorter, stronger and less flexible than single bonds.
Also, pi bonds are broken more easily, so compounds containing multiple
bonds are more reactive
7. An unsaturated hydrocarbon has a carbon chain that contains one or more
multiple bonds; a saturated hydrocarbon has a carbon chain that contains
only single bonds
Notes on Hybridization of Carbon
D. Multiple Bond Molecular Shape
 Since multiple bonds take up more space because of more electrons, they
change expected bond angles in molecules
E. Delocalization of electrons in conjugated systems
1. When a series of unhybridized p orbitals exist in a molecule,
delocalization of electrons occur
2. This means that electrons are shared over a larger region, adding stability
to the molecule
3. Multiple p orbital overlap results in a molecule containing a conjugated
system
AP chem chapter 8,9,10 notes and practice problems
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4. Benzene, C6H6, is a great example of this (pg. 429)
Practice Problems
9.1 Describe the bonding in the ammonia molecule.
9.2 Describe the bonding in the N2 molecule.
–
9.3 Describe the bonding in the triiodide ion (I3 ).
9.4 How is the xenon atom in XeF4 hybridized?
9.5 For each of the following molecules or ions, predict the hybridization of each atom, and
–
describe the molecular structure: CO, BF 4 , XeF2.
III. Liquids and Solids
A. Intermolecular Forces
1. Intramolecular forces are bonds; intermolecular forces are the forces that
attract molecules to each other to create solids and liquids (and gases)
2. Dipole-Dipole Forces—the force that holds polar molecules together, the +/–
interaction from the negative end of one polar molecule to the positive end of
the next molecule
3. Hydrogen Bonding—a unique and much stronger type of dipole-dipole force
in which hydrogen interacts with a highly electronegative atom such as
oxygen, nitrogen, or fluorine
a. hydrogen is very small so it can exist closer to the other atom,
strengthening the +/– interaction between the molecules
b. hydrogen bonding will cause substances to exist in condensed states of
matter, lead to higher boiling points and Hvapº’s
c. water exhibits strong hydrogen bonding; substances with –OH groups
also exhibit H-bonding, the more –OH groups, the more the substance
behaves like water and becomes more soluble in water (like dissolves
like!!!)
4. London Dispersion Forces—very weak intermolecular forces between noble
gas atoms and nonpolar molecules, sometimes referred to as an instantaneous
and/or temporary dipoles
AP chem chapter 8,9,10 notes and practice problems
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a. an accidental arrangement of electrons in one atom (molecule) induces a
dipole on a neighboring atom (molecule)
b. London dispersion forces become stronger and more significant as the
size of the atom, and therefore number of electrons, increases which
would explain why F2 is a gas and I2 is a solid
c. Recall the hydrocarbon trend of gls as the carbon chain length
increases; more electrons present increasing the strength of the London
dispersion forces
B. Liquids
1. Exhibit properties of surface tension (the resistance of a liquid to an increase
in its surface area), capillary action, and viscosity (a measure of a liquid’s
resistance to flow)
2. Polar liquids have high surface tension, a concave meniscus in glass, and
high viscosity, while nonpolar liquids are the opposite
C. Solids
1. Ionic solids have a lattice comprised of positive and negative ions held
together by electrostatic forces.
a. recall that a larger + and – charge on the ions gives a larger lattice energy
and therefore higher melting points
b. when comparing ionic solids of equal +/– charge, smaller ions in the solid
will be pulled closer together resulting in higher melting points
2. Metal solids have a closely packed lattice with delocalized electrons
throughout the crystal.
3. Carbon (pg. 458-459)--The two most common forms of carbon are diamond
and graphite
a. diamond is a network atomic solid in which each carbon atom is
surrounded by a tetrahedral arrangement of other carbon atoms to form a
huge covalently bonded molecule; diamond is the hardest naturally
occurring substance
b. graphite consists of layers of carbon atoms arranged in fused six-member
rings; strong bonding exist within the layers but little bonding between
the layers since all valence electrons are used in forming the rings
4. Elemental silicon has the same structure of diamond, but since it is a larger
atom, a few electrons can pass from one atom to another, making silicon a
semiconductor. This conductivity of silicon is enhanced by “doping” silicon
with arsenic.
Practice Problems
10.1 Classify each of the following substances according to the type of solid it forms: gold, carbon
dioxide, lithium fluoride, krypton.
10.2 Identify the most important types of interparticle forces present in the solids of each of the
following substances.
a.
NH4Cl
b.
Steel
c.
Teflon, CF3(CF2CF2)nCF3
d.
Polyethylene, CH3(CH2CH2)nCH3
AP chem chapter 8,9,10 notes and practice problems
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e.
f.
g.
h.
CHCl3
Ge
NO
BF3
10.3 For which molecule in each of the following pairs would you expect greater interparticle
forces?
a.
CH3CH2CH2NH2 or H2NCH2CH2NH2
b.
B(OH)3 or BH3
c.
CH3OH or H2CO
d.
HF or HI
10.4 In each of the following groups of substances, pick the one that has the given property.
Justify your answer.
a. lowest boiling point: CH3CH3 or CH3CH2CH2CH3
b. highest freezing point: H2O, NaCl, CH4
c.
highest vapor pressure at 25°C: Hg, Ti, H2O, gasoline
d. lowest melting point: N2, CO, KCl, CO2
e. strongest hydrogen bonds: H2O, (CH3)2NH, HF
f.
smallest heat of vaporization: CH4, CH3CH3, CH3CH2CH3
g. largest enthalpy of fusion: H2O, CO2, MgO, Li2O
h. lowest vapor pressure at 25°C: CH3CH2CH3, CH3C=OCH3, CH3CH2CH2OH
AP chem chapter 8,9,10 notes and practice problems