AP CHEMISTRY CHAPTER 8 BONDING

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AP CHEMISTRY CHAPTER 8BONDING
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bond energy- energy required to break a chemical
bond-We can measure bond energy to determine
strength of interaction
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ionic compound- a metal reacts with a nonmetal

• Ionic bonds form when an atom that loses
  electrons easily reacts with an atom that has a
  high affinity for electrons. The charged ions
  are held together by their mutual attraction.
• Ionic bonds form because the ion pair has lower
  energy than the separated ions. All bonds form
  in order to reach a lower energy level.

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Bond length- the distance where the energy is at
a minimum. We have a balance among proton-proton
repulsion, electron-electron repulsion, and
proton-electron attraction.In H2, the two e-
will usually be found between the two H atoms
because they are spontaneously attracted to both
protons. Therefore, electrons are shared by both
nuclei. This is called covalent bonding.
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Polar covalent bonds occur when electrons are not
shared equally. One end of the molecule may have
a partial charge. This is called a dipole.

• ?
  ?
• H F H H
  ? ?-
  
  O
  
  ? -

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Electronegativity- the ability of an atom in a
molecule to attract shared electrons to
itself.-determined by comparing the measured
bond energy and the expected bond energy.
Expected HX HH bond energy XX bond
energybond energy
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Electronegativity difference (Actual HX
bond energy) (expected HX bond energy) If X
has a greater electronegativity than H, the e-s
are closer to X and the molecule is polar. If
the electronegativities are the same, the
molecule is nonpolar.
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Periodic Trends-Electronegativity generally
increases across a period and decreases down a
group. It ranges from 0.79 for cesium to 4.0 for
fluorine.
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? ?-
HF polar
HH nonpolar
has dipole moment

O?
? H H
O S

O bent, polar
O ?-
has dipole moment planar

no dipole momentCH4 tetrahedral
NH3 trigonal pyramidal
no dipole moment
has dipole moment
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Electron Configurations

• Stable compounds usually have atoms with noble
  gas electron configurations.
• Two nonmetals react to form a covalent bond by
  sharing electrons to gain valence electron
  configurations.

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• When a nonmetal and a group A metal react to form
  a binary ionic compound, the ions form so that
  the valence electron configuration of the
  nonmetal is completed and the valence orbitals of
  the metal are emptied to give both noble gas
  configurations.

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• Ions form to get noble gas configurations.
• -exceptions in Group A metals
• Sn2 Sn4
• Pb2 Pb4
• Bi3 Bi5
• Tl Tl 3
• Metals with d electrons will lose their highest
  numerical energy level electrons before losing
  their inner d electrons.

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Size of Ions

• Positive ions (cations) are smaller than their
  parent atoms since they are losing electrons.
• Negative ions (anions) are larger than their
  parent atoms since they are gaining electrons.
• Ion size increases going down a group.

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Isoelectronic ions ions containing the same
number of electrons O2-, F-, Na, Mg2, Al3 all
have the Ne configuration. They are
isoelectronic. For an isoelectronic series,
size decreases as Z increases.
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Lattice energy- the change in energy that takes
place when separated gaseous ions are packed
together to form an ionic solid. Na(g)
Cl-(g) ? NaCl(s)

• If exothermic, the sign will be negative and the
  ionic solid will be the stable form.
• We can use a variety of steps to determine the
  heat of formation of an ionic solid from its
  elements. This is called the Born-Haber cycle.
  See example on page 355.

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• Lattice energy can be calculated using the
  following
• where k is a proportionality constant that
  depends on the structure of the solid and the
  electron configuration of the ions. Q1 Q2 are
  the charges on the ions. r is the distance
  between the center of the cation and the anion.
• Since the ions will have opposite charges,
  lattice energy will be negative (exothermic).
• The attractive force between a pair of oppositely
  charged ions increases with increased charge on
  the ions or with decreased ionic sizes.

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The Structure of Lithium Fluoride
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There are probably no totally ionic bonds.
Percent ionic character in binary compounds can
be calculated. Percent ionic character increases
with electronegativity difference. See Figure
8.12, pg. 360.

• Compounds with more than 50 ionic character are
  considered to be ionic (electronegativity diff.
  of about 1.7).

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The Relationship Between the Ionic Character of a
Covalent Bond and the Electronegativity
Difference of the Bonded Atoms
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Three Possible Types of Bonds
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• Polyatomic ions are held together by covalent
  bonds. We call Na2SO4 ionic even though it has 4
  covalent bonds and 2 ionic bonds.
• Ionic compound- any solid that conducts an
  electrical current when melted or dissolved in
  water
• Salt- an ionic compound

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• A chemical bond is a model invented by
  scientists to explain stability of compounds. A
  bond really represents an amount of energy. The
  bonding model helps us understand and describe
  molecular structure. It is supported by much
  research data. However, some data suggests that
  electrons are delocalized. That is, they are not
  associated with a particular atom in a molecule.

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• Single bond- one pair of shared electrons
• Double bond- two pair of shared electrons
• Triple bond- three pair of shared electrons

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• Bond energies and bond lengths are given on page
  374.
• We can use bond energies to calculate heats of
  reaction.
• ?H ?D(bonds broken)- ?D(bonds formed)
• 2H2 O2 ? 2H2O
• Ex. ?H 2(432) 495 4(467) -509 kJ
• 2 H-H OO 4H-O
  exothermic

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• Bonding Models
• Molecular Orbital Model-
• Electrons occupy orbitals in a molecule in much
  the same way as they occupy orbitals in atoms.
• Electrons do not belong to any one atom.
• -very complex model

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Localized electron model-

• molecules are composed if atoms that are bound
  together by sharing pairs of electrons using the
  atomic orbitals of the bound atoms
• traditional model

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lone pair- pair of electrons localized on an atom
(nonbonding)shared pair or bonding pair-
electrons found in the space between atoms
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Lewis structure -shows how the valence
electrons are arranged among the atoms in the
molecule
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The most important requirement for the formation
of a stable compound is that the atoms achieve
noble gas configurationsionic Na
Cl- only valence electrons
are includedmolecular H2O
H O - H
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duet rule- hydrogen forms stable molecules when
it shares two electrons HH-filled
valence shellWhy does He not form bonds?Its
valence orbitals are already filled.octet rule
most elements need 8 electrons to complete
their valence shell Cl-Cl
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Rules for writing Lewis structures1. Add up
the number of valence electrons from all
atoms.2. Use 2 electrons to form a bond
between each pair of bound atoms. A dash
represents a pair of shared electrons.3.
Arrange the remaining electrons to satisfy the
duet rule for H and the octet rule for most
others.
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Ex. H2S of valence electrons 1
1 6 8

• H - S - H

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Ex. CO2 of valence electrons
4 6 6 16

• O C O This uses 20
  electrons!
• O C O

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• NH3 has 8 valence electrons
• H
• N- H
• H

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HCN

• HCN has 10 valence electrons.
• H-CN

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NO

• NO has 5 6 -1 10 electrons
• NO

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CO32-

• Carbonate has 4 18 2 24 valence electrons.
• O 2-
• C O
• O

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ExceptionsBoron and beryllium tend to form
compounds where the B or Be atom have fewer than
8 electrons around them. BF3 24
valence electrons F
B F FCommon AP equation
NH3 BF3 ? H3NBF3
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C, N, O, F always obey the octet rule.
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Some elements in Period 3 and beyond exceed the
octet rule.Ex. SF6 S has 12
electrons around it48 valence electrons
F F
F F S
F
F
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d orbitals are used to accommodate the extra
electrons.Elements in the 1st or 2nd period of
the table cant exceed the octet rule because
there is no d sublevel.If the octet rule can be
exceeded, the extra electrons are placed on the
central atom.
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See examples of exceptions on pg 375.Ex. I3-,
ClF3, RnCl2

• I - I - I F
• F - Cl - F
• Cl - Rn - Cl

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Resonance- -occurs when more than one valid
Lewis structure can be written for a particular
moleculeactual structure is an average of all
resonance structures-this concept is needed to
fit the localized electron model (electrons are
really delocalized)
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Ex. Benzene, C6H6All bond lengths and angles
are the same.
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Ex. SO3
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Formal Charge-used to determine the most
accurate Lewis structure-is the difference
between the of valence electrons on the free
atom and the of valence electrons assigned to
the atom in the molecule
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-atoms try to achieve formal charges as close to
zero as possible-any negative formal charges are
expected to reside on the most electronegative
atoms-Sum of the formal charges must equal the
overall charge on the molecule (zero) or ion.
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Ex. SO42-

• O 2- O
  2-
• O S O O S O
• O O

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VSEPR-Valence Shell Electron Pair
Repulsion-allows us to use electron dot
structures to determine molecular shapes-the
structure around a given atom is determined
primarily by minimizing electron
repulsions-bonding and nonbonding pairs of
electrons around an atom position themselves as
far apart as possible
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Steps1. Draw Lewis structure2. Count
effective electron pairs on central atom (double
and triple bonds count as one)3. Arrange the
electron pairs as far apart as possible
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ShapesAX2 (A represents central atom, X
represents attached atom, E represents unshared
electron pair) X A X linear
180o bond angle OCO
Cl Be Cl
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AX3 Shape is trigonal planar
X X A
120o bond angle
F F
X BF3
B
F Any
resonance SO3 structure can be
used to O- S O determine
shape. O
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AX2E Shape is bentBond angle is lt 120o
X X A


EEx. SnCl2 Cl Cl
Sn
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AX4 Shape is tetrahedral
Bond angle is 109.5o X
Ex. CH4 H
X A X
H C H X
H
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Figure 8.14The Molecular Structure of Methane
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AX3E Shape is trigonal pyramidal
Bond angle is lt 109.5oEx. NH3
H - N- H H
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Figure 8.15The Molecular Structure of NH3
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AX2E2 Shape is bent Bond angle is lt
109.5oUnshared electron pairs repel more than
shared pair. Lone pairs require more space than
share pairs. E
Ex. H2O X A X
E
H O - H
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Figure 8.16The Molecular Structure of H2O
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Figure 8.17The Bond Angles in the CH4, NH3, and
H2O Molecules
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AX5 Shape is trigonal bipyramidal Bond
angles are 120o(equatorial) and 90o(axial)
X X A X X
XEx. PCl5 Cl Cl
P Cl Cl Cl
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AX4E Shape is see-sawBond angles are lt90o and
lt120o X E A X
X XEx. SF4 34 electrons
F S F F
F
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Figure 8.20Three Possible Arrangements of the
Electron Pairs in the I3- Ion
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AX3E2 Shape is T-shaped
Bond angle is lt90o X E
A X E XEx.
ClF3 F Cl F F
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AX2E3 shape is linear
bond angle is 180o X E
A E E XEx. XeF2
F Xe F
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Figure 8.19Possible Electron Pair Arrangements
for XeF4
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AX6 shape is octahedral bond angle
is 90o X X
X AX X
XEx. SF6 F F
F SF F
F
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AX5E Shape is square pyramidal
Bond angle is lt90o X
X X AX
X EEx. BrF5 F
F F BrF
F
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AX4E2 Shape is square planar.
Bond angle is 90o. E X
X AX X
E
animated vsepr table