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