NH3 (ammonia) (CH3)3C+ (the tert-butyl carbocation)

H2CO (formaldehyde) CH3CO- (the acetate ion)

CH3NC (methyl isonitrile)

Apparent Exceptions to the Octet Rule

• Molecules that contain atoms of Group 3A elements, particularly boron and aluminum

Electronegativity

Cl — Cl, H — Cl, Na — Cl

Electronegativity and Chemical Bonds

• Electronegativity: A measure of an atom’s attraction for electrons that it shares with another atom in a chemical bond

• Pauling scale • In the Periodic Table, electronegativity increases from left to right in a

row and decreases from top to bottom in a column

• Electron affinity• Energy added or released when an electron is added to an atom or a

molecule• Becomes more favorable as one moves from left to right in a row of the

Periodic Table

Table 1.5 - Electronegativity Values for Some Atoms (Pauling Scale)

Incr

ease

Increase

Linus Pauling invented the tremendously useful concept of electronegativity to describe the distribution of electron density in covalent compounds. The electronegativity of an atom is a measure of its tendency to attract electron density in covalent bonds.

Table 1.6 - Classification of Chemical Bonds

Bond Difference in ElectronegativityType of

Interaction

(a) O—H 3.5 – 2.1 = 1.4 Polar covalent

(b) N—H 3.0 – 2.1 = 0.9 Polar covalent

(c) K—Br 2.8 – 0.8 = 2.0 Ions formed

(d) C—Mg 2.5 – 1.2 = 1.3 Polar covalent

H — H Cl — Cl

NaCl

Polar Molecules

• Electrostatic potential maps for water and ammonia that display the computed electronic charge density in water and ammonia are given below

• Red represents negative charge, and blue represents positive charge

Polar Bonds in Nonpolar Molecules

• These molecules have polar bonds, but each molecule, because of its geometry, has a zero dipole moment

Valence-Shell Electron Pair Repulsion (VSEPR)

• Method for predicting bond angles based on the idea that electron pairs repel each other and stay as far apart as possible

Valence-Shell Electron Pair Repulsion (VSEPR)

Note the “wedge and dash” notation that chemists use to represent three-dimensional structures on a flat page.

Hybridization and Bonding

Bonding Angles

In forming covalent bonds, atoms of C, N, and O use 2s and 2p atomic orbitals

Valence Bond (VB) Theory• Basic principle - Bonds are created by the overlap of atomic orbitals on

adjacent atom• If atoms used these orbitals to form bonds, one would expect bond

angles of approximately 90° (2px, 2py, 2pz)• Instead, bond angles of approximately 109.5° in molecules with only

single bonds, 120° in molecules with double bonds, and 180° in molecules with triple bonds are found

• Linus Pauling proposed that atomic orbitals for each of these atoms should be thought of as combining to form new atomic orbitals, called hybrid orbitals

• Hybrid orbitals then interact to form bonds by overlapping with orbitals from other atoms

• Hybrid orbitals are formed by combinations of atomic orbitals, or hybridization

• Accomplished mathematically by combining wave functions of the 2s and three 2p orbital wave functions

Hybridization and Multiple Bonds • A dirty secret:

Hybridization is simply a code for electron domain geometry. Let's take a look:

• Consider ethylene (also called ethene): C2H4. Draw a Lewis structure, and use it to determine the geometry and hybridization of each of the carbon atoms. How can we describe a double bond in terms of the overlap of orbitals?

The Master Table of Bonding and Hybridization • Let's fill out the following table of bonding and hybridization:

• Now let's use this table to describe the bonding in acetonitrile (CH3CN):

Resonance

Resonance

• For many molecules and ions, no single Lewis structure provides a truly accurate representation

• Example - There are three Lewis structures for the carbonate ion• The negative charge is distributed equally among the three oxygens

as shown in the electrostatic potential map

Lewis structures for carbonate CO32-

Theory of Resonance

• Developed by Linus Pauling in the1930s• Many molecules and ions are best described by writing two or more Lewis

structures and considering the real molecule or ion to be a composite of these structures

• Individual Lewis structures are called contributing structures or resonance structures

• Individual contributing structures are interconnected by double-headed arrows

• Help denote that the real molecule or ion is a resonance hybrid

Resonance - Rules for Writing Contributing Structures

• All contributing structures:• Must have the same number of valence electrons• Must obey the rules of covalent bonding

• No more than two electrons can be present in the valence shell of H • No more than eight electrons can be present in the valence shell of

a second-period element• Third-period elements may be drawn indicating up to 12 electrons in

their valence shells• Must have the same number of paired and unpaired electrons

• Positions of all nuclei must be the same in all contributing structures

Resonance - Relative Importance of Contributing Structures

• Preference 1 - Filled valence shells• Structures in which all atoms have filled valence shells contribute more

than those with one or more unfilled valence shells

Resonance - Relative Importance of Contributing Structures (continued 1)

• Preference 2 - Maximum number of covalent bonds• Structures with a greater number of covalent bonds contribute more

than those with fewer covalent bonds

Resonance - Relative Importance of Contributing Structures (continued 2)

• Preference 3 - Least separation of unlike charges• Structures that involve separation of unlike charges contribute less than

those that do not involve charge separation • Separation of charges costs energy

Resonance - Relative Importance of Contributing Structures (continued 3)

• Preference 4 - Negative charge on a more electronegative atom • Structures that carry a negative charge on a more electronegative atom

contribute more than those with the negative charge on a less electronegative atom

• Structures that carry a positive charge on a less electronegative atom contribute more than those that carry the positive charge on a more electronegative atom

Hybridization Considerations in Light of Resonance

• Whenever atoms are involved in resonance that has one or more contributing structures involving a π bond to that atom, the atoms must be sp2 hybridized

• On rare occasions, the atoms can be sp hybridized, but never sp3hybridized

Example 1.19 - Contributing Structures and Hybridization

• Draw three contributing structures of the following amide and state the hybridization of the highlighted O, C, and N

• In which orbitals do the three lone pairs drawn reside?

Example 1.19 - Solution• There is a reasonable contributing structure (C), which places a

double bond between the carbon and the nitrogen• Therefore, all three atoms are sp2

• This means that the lone pair on N is in a 2p orbital

Delocalized 𝜋 Bonding• Draw the Lewis structure for the nitrate ion, NO3–. • Using the "Master Table", how can we describe the bonding in nitrate?

Remember: the actual molecule must somehow be an "average" of the resonance structures.

• Nitrate is a classic example of delocalized 𝜋-bonding. How do the pictures below represent localized and delocalized 𝜋-bonding?

• What is the hybridization of each oxygen atom in nitrate?

Resonance and Hybridization: The Wrong View

• Draw the Lewis structure of formamide, HCONH2.

• What is the hybridization of the nitrogen atom?

• The bond angles around the nitrogen atom in formamide are 120°. Is this consistent with the hybridization that you determined?

Resonance and Hybridization: The Right View

• Draw all the reasonable resonance structures of formamide, HCONH2.

• What is the hybridization of the nitrogen atom in each of these resonance structures?

• Why must the hybridization of an atom be the same in all resonance structures?

• How can we resolve this question of hybridization?