Chemical Bonding and Molecular Architecture
Structure and Shapes of Chemicals
Bonds
Forces that hold groups of atoms together and make them function as a unit.
Bond EnergyIt is the energy required to break or released
in making a bond.
It gives us information about the strength of a bonding interaction.
Ionic bonds—strong attractions between oppositely charged ions
Covalent bonds—attraction between non-metal atoms as both atoms share electrons
Bond Length
The distance where the system energy is a minimum.
08_130
Sufficiently far apartto have no interaction
En
erg
y (k
J/m
ol)
0 Internuclear distance (nm)0.074
-458
0
(H H bond length)
HH
H H
H H
H H
(a) (b)
+
H atom H atom
The atoms begin to interactas they move closer together.
+
H atom H atom
H2molecule
+
Optimum distance to achievelowest overall energy of system
+
+
+
Ionic Bonds
- Formed from electrostatic attractions of closely packed, oppositely charged ions.
- Formed when an atom that easily loses electrons reacts with one that has a high electron affinity.
Ionic Configuration and Size
Ions are formed when electrons are gained or lost from an atom. The gain or loss follows the pattern called the “octet rule”, that an atom forms an ion in which it attains the same electron configuration as the nearest noble gas. Most metals therefore lose electrons, and as a result get smaller. The trend is “greater +, smaller size.”
Likewise, nonmetals gain electrons to form ions, thus increasing in size by the opposite rule to metals.
08_136 Li
(0.60)60
Be
(0.31)31
O
(1.40)140
F
(1.36)136
Na
(0.95)95
Mg
(0.65)65
Al
(0.50)50
S
(1.84)184
Cl
(1.81)181
K
(1.33)133
Ca
(0.99)99
Ga
(0.62)62
Se
(1.98)198
Br
(1.95)195
Rb
(1.48)148
Sr
(1.13)113
In
(0.81)81
Sn
(0.71)71
Sb
(0.62)62
Te
(2.21)221
I
(2.16)216
Isoelectronic Ions
Ions containing the the same number of electrons, due to attaining the configuration of the same noble gas
(O2, F, Na+, Mg2+, Al3+)
All attain to Ne
O2> F > Na+ > Mg2+ > Al3+
largest smallest
Electronegativity
The ability of an atom in a molecule to attract shared electrons to itself.
Periodic trend –increases across the table to the halogen column. Decreases down a group. Least at Cs (0.7), greatest at F (4.0).
08_132
H2.1
Li1.0
Be1.5
Na0.9
Mg1.2
K0.8
Ca1.0
Rb0.8
Sr1.0
Cs0.7
Ba0.9
Fr0.7
Ra0.9
Sc1.3
Y1.2
La-Lu1.0-1.2
Ac1.1
Ti1.5
Zr1.4
Hf1.3
Th1.3
V1.6
Nb1.6
Ta1.5
Pa1.4
Cr1.6
Mo1.8
W1.7
U1.4
Mn1.5
Tc1.9
Re1.9
Np-No1.4-1.3
Fe1.8
Ru2.2
Os2.2
Co1.9
Rh2.2
Ir2.2
Ni1.9
Pd2.2
Pt2.2
Cu1.9
Ag1.9
Au2.4
Zn1.6
Cd1.7
Hg1.9
Ga1.6
In1.7
Tl1.8
Al1.5
B2.0
Ge1.8
Sn1.8
Pb1.9
Si1.8
C2.5
As2.0
Sb1.9
Bi1.9
P2.1
N3.0
Se2.4
Te2.1
Po2.0
S2.5
O3.5
Br2.8
I2.5
At2.2
Cl3.0
F4.0
H2.1
Li1.0
Be1.5
Na0.9
Mg1.2
K0.8
Ca1.0
Rb0.8
Sr1.0
Cs0.7
Ba0.9
Fr0.7
Ra0.9
Sc1.3
Y1.2
La-Lu1.0-1.2
Ac1.1
Ti1.5
Zr1.4
Hf1.3
Th1.3
V1.6
Nb1.6
Ta1.5
Pa1.4
Cr1.6
Mo1.8
W1.7
U1.4
Mn1.5
Tc1.9
Re1.9
Np-No1.4-1.3
Fe1.8
Ru2.2
Os2.2
Co1.9
Rh2.2
Ir2.2
Ni1.9
Pd2.2
Pt2.2
Cu1.9
Ag1.9
Au2.4
Zn1.6
Cd1.7
Hg1.9
Ga1.6
In1.7
Tl1.8
Al1.5
B2.0
Ge1.8
Sn1.8
Pb1.9
Si1.8
C2.5
As2.0
Sb1.9
Bi1.9
P2.1
N3.0
Se2.4
Te2.1
Po2.0
S2.5
O3.5
Br2.8
I2.5
At2.2
Cl3.0
F4.0
Increasing electronegativity
De
crea
sing
ele
ctro
neg
ativ
ity
Increasing electronegativity
De
crea
sing
ele
ctro
neg
ativ
ity
(a)
(b)
Polarity
A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.
+
FH
Polar bonds shown as arrow with point toward negative pole, + toward the positive pole
Electronegativity and Polarity of Bonds
Subtract lower EN from higherSubtract lower EN from higher
EN DifferenceEN Difference % Ionic Character% Ionic Character Type of BondType of Bond
00 00 Nonpolar CovalentNonpolar Covalent
0.1-0.50.1-0.5 1-5%1-5% Slightly polar covalentSlightly polar covalent
0.6-1.50.6-1.5 6-40%6-40% Polar CovalentPolar Covalent
> 1.5> 1.5 over 40% over 40% IonicIonic
Compounds with over 50% ionic character are considered to be totally ionic solids. These compounds are often called salts.
08_131
F
H
F
H
F
HF
H
F
H
(a)
H F
(b)
H F
H F
H F
H F
08_133
H
O
H
(a)
+
(b)
08_134
HH
N
H
3
(a)
+
(b)
Homework!!
p. 395ff 11, 14, 15, 20
Achieving Noble Gas Electron Configurations (NGEC)
Two nonmetals react: They share electrons to achieve NGEC.
A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC.
Binary Ionic--Lattice Energy
The change in energy when separated gaseous ions are packed together to form an ionic solid.
M+(g) + X(g) MX(s)
Lattice energy is negative (exothermic) from the point of view of the system.
Formation of an Ionic Solid1. Sublimation of the solid metal
M(s) M(g) [endothermic]2. Ionization of the metal atoms
M(g) M+(g) + e [endothermic]3. Dissociation of the nonmetal
1/2X2(g) X(g) [endothermic]4. Formation of X ions in the gas phase:
X(g) + e X(g) [exothermic]5. Formation of the solid MX
M+(g) + X(g) MX(s) [quite exothermic]
08_139Mg2+(g) + O2-(g)
737 Electron affinity
247
2180 Ionization energy
150
-602 -570Overallenergychange
NaF(s)
-923 Latticeenergy
-328 Electronaffinity
-3916 Latticeenergy
109
495Ionizationenergy
77
Mg2+(g) + O(g)
Mg2+(g) + 12 O2(g)
Na(g) + F(g)
Na+(g) + F-(g)
Mg(g) + 12 O2(g)
Mg(s) + 12 O2(g)
Na(g) + 12 F2(g)
Na(s) + 12 F2(g)
Na+(g) + 12 F2(g)
MgO(s)
Covalent Chemical Bonds
Happen when collections of atoms are more stable than the separate atoms. They provide a method for dividing up energy when stable molecules are formed from atoms.
Covalent bonds are due to shared electron pairs. One pair shared is a single bond, two makes a double bond, three make a triple bond.
As bond order increases (single, double, triple), bond length shortens
Bond Energies
Bond breaking requires energy (endothermic).
Bond formation releases energy (exothermic).
H = D(bonds broken) D(bonds formed)
energy required energy released
Localized Electron Model
A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
Two types of electron pairs: bonding pairs and lone pairs. Bonding pairs are linkages between atoms, lone pairs are electrons solely owned by an atom.
Localized Electron Model
Elements of the Model
1. Description of valence electron arrangement (Lewis structure).
2. Prediction of geometry (VSEPR model).
3. Description of atomic orbital types used to share electrons or hold lone pairs.
Lewis StructureShows how valence electrons are arranged
among atoms in a molecule.
Reflects central idea that stability of a compound relates to noble gas electron configuration.
Comments About the Octet Rule
2nd row elements C, N, O, F observe the octet rule.
2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.
3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.
When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.
Rules for Drawing Lewis StructuresAdd up all of the valence electrons for the
atoms involved in the molecule
Select a most likely central atom and arrange other atoms around it. Place pairs of electrons between atoms.
Arrange the remaining electrons around external atoms first. If the central atom is not satisfied, form double or triple bonds to make the molecule work.
Resonance
Occurs when more than one valid Lewis structure can be written for a particular molecule.
These are resonance structures. The actual structure is an average of the resonance structures.
Homework
p. 397ff 31, 36, 39, 42, 50, 57
Molecular Architecture
The structure of a molecule is important in how it reacts and to its physical properties
Once the Lewis structure of a molecule is determined, the shape of the molecule then can be predicted according to the VSEPR model.
VSEPR Model
The structure around a given atom is determined principally by minimizing electron pair repulsions.
08_06T
Number ofElectron Pairs
Table 8.6 Arrangements of Electron Pairs Around an Atom Yielding Minimum Repulsion
Arrangement of Electron Pairs Example
2 Linear
3 Trigonalplanar
4 Tetrahedral
5 Trigonalbipyramidal
6 Octahedral
A
A
A
A120°
90°
A
Predicting a VSEPR Structure
1. Draw Lewis structure.
2. Count pairs, both bonding and lone pairs around the central atom.
3. Determine positions of atoms from the way electron pairs are shared.
4. Determine the name of molecular structure from the number of bonding and lone pairs and their necessary arrangements. Remember that lone pairs prefer to be at 120º or greater from each other.
Homework!!
p. 399ff 59, 62, 73, 78, 79, 91
Hybridization
The mixing of atomic orbitals to form special orbitals for bonding.
The atoms are responding as needed to give the minimum energy for the molecule.
To determine hybridization, count lone and bonding pairs, but count multiple bonds only once.
09_158
z
y
z
x
y
x
x
y
z
y
x
sp3
sp3
sp3
sp3
Hybridization
gives a tetrahedralarrangement
s
p y
p x
p z
x
y
z
z
y
x
y
z
x
y
z
z
x
09_179 Number ofEffective Pairs
Arrangementof Pairs
HybridizationRequired
2 Linear sp
180°
3 Trigonalplanar
sp2
120°
4 Tetrahedral
109.5°
5 Trigonalbipyramidal
dsp3
90°
120°
90°
90°
6 Octahedral d2sp3
sp3
09_161
sp3
sp3
H1s
H1s H1s
H1s
C
sp3
sp3
09_162
sp3
sp3
H1s
H1s H
1s
lone pair
N
sp3
sp3
09_166
C C
sp 2
sp 2
sp2
sp2
H1s
H1s
H1s
H1s sp2sp2
A sigma () bond centers along the internuclear axis.
A pi () bond occupies the space above and below the internuclear axis.
CCH H
HH
09_167
sigmabond
pi bondC C
p orbital p orbital
09_168
(b)
H
C C
H
H
Hsp 2
sp 2
sp 2
sp 2
H1sH1s C C
2p
sp 2 sp 2
(a)
09_174
O C O
sigma bond(1 pair of electrons) pi bond
(1 pair ofelectrons)
pi bond(1 pair ofelectrons)
(a)
(b)
O C O
09_189
(a)
(b) (c) (d)
B B
The Localized Electron Model
- Draw the Lewis structure(s)
- Determine the arrangement of electron pairs (VSEPR model).
- Specify the necessary hybrid orbitals.
Homework
p. 432ff 5, 8, 11
Molecular Orbitals (MO)
Analagous to atomic orbitals for atoms, MOs are the quantum mechanical solutions to the organization of valence electrons in molecules.
Types of MOs
bonding: lower in energy than the atomic orbitals from which it is composed.
antibonding: higher in energy than the atomic orbitals from which it is composed.
09_556
EMO2
1sA
H2HA HB
1sB
MO1
Energy diagram
(a)
Electron probability distribution
+ +
+ +(b)
09_190
2py 2py
2px 2px
Antibonding
Bonding
2p
*2p
Antibonding
Bonding
*2p
2p
(b)
(a)
Bond Order (BO)
Difference between the number of bonding electrons and number of antibonding electrons divided by two.
BO = # bonding electrons # antibonding electons
2
Paramagnetism
- unpaired electrons
- attracted to induced magnetic field
- much stronger than diamagnetism
Outcomes of MO Model1. As bond order increases, bond energy increases
and bond length decreases.
2. Bond order is not absolutely associated with a particular bond energy.
3. N2 has a triple bond, and a correspondingly high bond energy.
4. O2 is paramagnetic. This is predicted by the MO model, not by the LE model, which predicts diamagnetism.
09_195
E
2p*
2p*
2p
2p
2s*
2s
B 2 C2 N 2 O2 F2
MagnetismPara–
magneticDia–
magneticDia–
magneticPara–
magneticDia–
magnetic
Bond order 1 2 3
2s
2s*
2p
2p
2p*
2p*
2 1
Observedbonddissociationenergy(kJ/mol) 290 620 942 495 154
Observed bondlength(pm) 159 131 110 121 143
Combining LE and MO Models
bonds can be described as being localized.
bonding must be treated as being delocalized.
09_203
H H
H
H
H H
H H
H H
(a) (b)
H
H
Homework
p. 434 ff 17, 22, 25, 37