Chapter 1Chemical Bonding and Chemical Structure

Organic chemistry

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• The branch of chemistry that deals with carbon based compounds– Organic compounds may contain any number of

other elements, including hydrogen, nitrogen, oxygen, halogens, phosphorus, silicon, and sulfur

Methane Sucrose Morphine

History• Vitalism: Only biological

systems (e.g., plants, animals) could produce organic compounds

• Wohler’s synthesis of urea (1828), began to undermine vitalism

Why Study Organic Chemistry?

• Organic chemistry lies at the heart of the modern chemical industry

• Central to medicine and pharmacy• Interface of physical and biological sciences• Everyday applications: Plastics, textiles,

communications, transportation, food, clothing, cosmetics, etc.

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Review of Chemical Bonding

• Valence Electrons: Outermost electrons• s and p electrons for main group elements• Responsible for chemical properties of

atoms• Participate in chemical reactions

Core Electrons Valence Electron

Octet Rule

• Octet Rule: the tendency for atoms to seek 8 electrons in their outer shells– Natural electron configuration of the Noble Gases– Done by gaining, losing, or sharing electrons– Increases stability– H and He seek a “Duet”

Ionic Bonding• Ions: atoms that have a charge due to gain or loss of

electrons– Anion: (-) charged atom– Cation: (+) charged atom

• Ionic Bond: a bond formed through the transfer of one or more electrons from one atom or group of atoms to another atom or group of atoms

Formula Unit

• Ionic bonds are omni-directional• Can dissociate into free ions

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Covalent Compounds• Covalent Compounds: compounds composed of atoms

bonded to each other through the sharing of electrons• Electrons NOT transferred• No + or – charges on atoms• Non-metal + Non-metal• Also called “molecules”• Examples:– H2O– CO2

– Cl2

– CH4

or H-H

or

Duet

Covalent Bonds

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Electronegativity• The measure of the ability of an atom to

attract electrons to itself– Increases across period (left to right) and– Decreases down group (top to bottom)– fluorine is the most electronegative element– francium is the least electronegative element

Electronegativity Scale

Types of Bonding

1) Non-Polar Covalent Bond:• Difference in electronegativity

values of atoms is 0.0 – 0.4• Electrons in molecule are

equally shared• Examples: Cl2, H2, CH4

ENCl = 3.03.0 - 3.0 = 0

Pure Covalent

2) Polar Covalent Bond:• Difference in

electronegativity values of atoms is 0.4 – 1.7/2.0

• Electrons in the molecule are not equally shared• The atom with the higher

EN value pulls the electron cloud towards itself

• Partial charges• Examples: HCl, ClF, NO

ENCl = 3.0ENH = 2.1

3.0 – 2.1 = 0.9Polar Covalent

Electrostatic Potential Maps

• A graphical depiction of electron distribution

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3) Ionic Bond: • Difference in EN

above 1.7-2.0• Complete transfer

of electron(s)• Whole charges

ENCl = 3.0ENNa = 1.0

3.0 – 0.9 = 2.1Ionic

Dipole Moment ()

• Depends on charge separation and distance• = qr (a vector quantity)

• q = magnitude of charge• r = vector from site of + charge to site of – charge

• Units = Debyes (D)

Molecular Polarity

Lewis Dot Structures1) Count the number of valence electrons present in

the molecule2) Determine the arrangement of atoms. Generally,

the atom that occurs least often is central. Join the terminal atoms to the central atom(s) using shared pairs of electrons (bonds)

3) Place any remaining electrons around the terminal atoms to satisfy the octet rule

• Exception: Hydrogen

4) Place any remaining electrons on the central atom(s) to satisfy the octet rule

5) Check to make sure:• You’ve used the correct number of valence

electrons• Everyone has an octet (or duet)• Everyone is doing what they like to do

6) If the number of electrons around the central atom is less than 3, change the single bonds to multiple bonds

What Things Like To Do1) Halogens

• Like to be terminal• Like to have one bonding pair

(two shared electrons) and 3 lone pairs (non-bonding electrons)

2) Carbon• Likes to have 4 bonding pairs

and no lone pairs• Likes to bond to other carbons• Likes to be central

3) Silicon• Likes to do what carbon does• Notice, it sits under C on the

periodic table

4) Oxygen• Like to have 2 bonding pairs

and 2 lone pairs

5) Sulfur• Likes to do what O does

6) Nitrogen• Likes to have 3 bonding pairs

and 1 lone pair

7) Phosphorous• Likes to do what N does

8) Hydrogen• Likes to be terminal with only 1

bond• Do not put lone pairs on H

9) Boron• Likes to have 3 bonds and no lone

pairs• Likes a sextet instead of an octet

(what everybody else besides Hydrogen likes)

10) *Note: • A double bond = 2 bonding pairs• A triple bond = 3 bonding pairs

Problems

• Draw the Lewis Dot Structures for the following molecules

1) CO2

2) P2H4

3) O3

4) NO3-

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Drawing Resonance Structures

1. Draw first Lewis structure that maximizes octets

2. Assign formal charges3. Move electron pairs from atoms with (-)

formal charge toward atoms with (+) formal charge

-1

-1

Formal Charge• Assigned charge for each atom in a molecule/ion– Electronic bookkeeping – may or may not correspond

to a real charge– Sum of formal charges on each atom must equal the

total charge on the molecule/ion

• FC = Valence e-’s – Lone Pair e-’s – ½ bonding e-’s

Molecular Structures of Covalent Compounds

• Atomic connectivity: How atoms in a molecule are connected

• Molecular geometry: How far apart atoms are and how they are arranged in space– Bond lengths– Bond angles– Dihedral angles

OR

Bond Length

• Distance between nuclei

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• Increases with atoms in higher rows• Decreases toward higher atomic number along a row• Decreases with increasing bond order

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Bond Angles

• Angle between each pair of bonds• Contribute to molecular shape• Determined by Valence-shell electron-pair

repulsion (VSEPR)• Use molecular models!• Line-and-wedge structures

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Drawing LDS With Correct Geometry

Valence Shell Electron Pair Repulsion Theory

• VSEPR theory:– Electrons repel each other– Electrons arrange in a

molecule themselves so as to be as far apart as possible• Minimize repulsion• Determines molecular

geometry

Defining Molecular Shape• Electron pair geometry: the geometrical

arrangement of electron groups around a central atom– Look at all bonding and non-bonding e-’s

• Molecular Geometry: the geometrical arrangement of atoms around a central atom– Ignore lone pair electrons

Problems• Predict the approximate geometry in each of

the following molecules– BF3

– HCN– CO3

2-

• Estimate the bond angles and relative bond lengths in the following molecule

Dihedral Angle

• Also known as the torsional angle • Rotation can occur along single bonds

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Valence Bond Theory

Types of Bonds• A sigma () bond results when the bonding orbitals

point along the axis connecting the two bonding nuclei– either standard atomic orbitals or hybrids

• s-to-s, p-to-p, hybrid-to-hybrid, s-to-hybrid, etc.

• A pi () bond results when the bonding orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei– between unhybridized parallel p orbitals

• the interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore bonds are stronger than bonds

Problems

• Write a hybridization and bonding scheme for acetaldehyde

Molecular Orbital Theory

Bond Order: ½ (# of electrons in bonding MO’s - # of electrons in antibonding MO’s)

Problems

1) Draw an MO diagram to predict the bond order of N2

2) Draw an MO diagram to predict the bond order of CN-

3) Use MO theory to determine the bond order of Ne2