Download ap chemistry chapter 8 bonding

AP CHEMISTRY
CHAPTER 8
BONDING
bond energy- energy required to
break a chemical bond
-We can measure bond energy to
determine strength of interaction
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.
Bond length- the distance where the
energy is at a minimum. We have a
balance among proton-proton repulsion,
electron-electron repulsion, and protonelectron 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.
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
-
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 H—X = H—H bond energy + X—X bond energy
bond energy
2
Electronegativity difference =
(Actual H—X bond energy) – (expected H—X 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.
Periodic TrendsElectronegativity generally increases across a
period and decreases down a group. It ranges
from 0.79 for cesium to 4.0 for fluorine.
+ H—F
polar
has dipole moment
+
H
O
-
+
H
bent, polar
has dipole moment
CH4 tetrahedral
no dipole moment
H—H nonpolar
O
O
S
O
planar
no dipole moment
NH3 trigonal pyramidal
has dipole moment
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.
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.
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.
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.
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.
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.
Lattice energy can be calculated using the following:
 Q1Q2 
Lattice energy  k

 r 
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.
The Structure of
Lithium Fluoride
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).
The Relationship Between the Ionic Character of a Covalent Bond
and the Electronegativity Difference of the Bonded Atoms
Three Possible Types of Bonds
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
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.
• Single bond- one pair of shared electrons
• Double bond- two pair of shared electrons
• Triple bond- three pair of shared electrons
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 O=O 4H-O
exothermic
Bonding Models:
Molecular Orbital ModelElectrons 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
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
lone pair- pair of electrons
localized on an atom
(nonbonding)
shared pair or bonding pairelectrons found in the space
between atoms
Lewis structure -shows how the
valence electrons are arranged
among the atoms in the molecule
The most important requirement for the
formation of a stable compound is that the atoms
achieve noble gas configurations
ionic [ Na ]+
[Cl]only valence electrons are included
molecular H2O
H–O-H
duet rule- hydrogen forms stable molecules
when it shares two electrons
H:H
-filled valence shell
Why 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
Rules for writing Lewis structures
1. 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.
Ex. H2S
# of valence electrons: 1 + 1 + 6 = 8
H - S - H
Ex. CO2
# of valence electrons = 4 + 6 + 6 = 16
O–C–O
O=C=O
This uses 20 electrons!
NH3 has 8 valence electrons
H
N- H
H
HCN
HCN has 10 valence electrons.
H-C≡N
NO+
NO+ has 5 + 6 -1 = 10 electrons
N≡O
CO32Carbonate has 4 + 18 + 2 = 24 valence electrons.
O
C O
O
2-
Exceptions:
Boron and beryllium tend to form compounds
where the B or Be atom have fewer than 8
electrons around them.
BF3 = 24 valence electrons
F
BF
F
Common AP equation:
NH3 + BF3  H3NBF3
C, N, O, F always obey the
octet rule.
Some elements in Period 3 and beyond exceed
the octet rule.
Ex. SF6
S has 12 electrons around it
48 valence electrons
F
F
F
S
F
F
F
d orbitals are used to accommodate the
extra electrons.
Elements in the 1st or 2nd period of the
table can’t 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.
See examples of exceptions on pg 375.
Ex. I3-, ClF3, RnCl2
I-I-I
Cl - Rn - Cl
F
F - Cl - F
Resonance-occurs when more than one valid Lewis
structure can be written for a particular molecule
actual structure is an average of all resonance
structures
-this concept is needed to fit the localized
electron model (electrons are really delocalized)
Ex. Benzene, C6H6
All bond lengths and angles are the same.
Ex. SO3
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
-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.
Ex. SO42-
O
O
S
O
2O
O
O S O
O
2-
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
Steps:
1. Draw Lewis structure
2. Count effective electron pairs on central atom
(double and triple bonds count as one)
3. Arrange the electron pairs as far apart as
possible
Shapes
AX2 (A represents central atom, X represents
attached atom, E represents unshared electron
pair)
X–A–X
linear 180o bond angle
O=C=O
Cl – Be – Cl
AX3
Shape is trigonal planar
X
X
A
120o bond angle
F
X
Any resonance
structure can
be used to
determine shape.
BF3
SO3
O- S = O
O
F
B
F
AX2E
Shape is bent
Bond angle is < 120o
X
X
A
E
Ex. SnCl2
Cl
Cl
Sn
AX4
X
X A X
X
Shape is tetrahedral
Bond angle is 109.5o
Ex. CH4
H
H C H
H
Figure
8.14
The
Molecular
Structure
of
Methane
AX3E
Ex. NH3
H - N- H
H
Shape is trigonal pyramidal
Bond angle is < 109.5o
Figure 8.15
The Molecular Structure of NH3
AX2E2 Shape is bent
Bond angle is < 109.5o
Unshared 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
Figure 8.16
The Molecular Structure of H2O
Figure 8.17
The Bond Angles in the CH4, NH3, and
H2O Molecules
AX5 Shape is trigonal bipyramidal
Bond angles are 120o(equatorial) and 90o(axial)
X
X A X
X
X
Ex. PCl5
Cl
Cl P Cl
Cl Cl
AX4E Shape is see-saw
Bond angles are <90o and <120o
X
E A X
X
X
Ex. SF4 34 electrons
F
S F
F
F
Figure 8.20
Three Possible Arrangements of the
Electron Pairs in the I3- Ion
AX3E2
Shape is T-shaped
Bond angle is <90o
X
E A X
E
X
Ex. ClF3
F
Cl F
F
AX2E3
shape is linear
bond angle is 180o
X
A E
E
E
X
Ex. XeF2
F
Xe
F
Figure 8.19
Possible
Electron Pair
Arrangements
for XeF4
AX6 shape is octahedral
bond angle is 90o
X
X
X
A
X
X
X
Ex. SF6
F
F
F
S
F
F
F
AX5E
Shape is square pyramidal
Bond angle is <90o
X
X
X
A
X
E
Ex. BrF5
F
F
Br
F
X
F
F
AX4E2
Shape is square planar.
Bond angle is 90o.
E
X
X
A
X
X
E
animated vsepr table