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Organic Chemistry. William H. Brown Christopher S. Foote Brent L. Iverson. Covalent Bonding & Shapes of Molecules. Chapter 1. Organic Chemistry. The study of the compounds of carbon Over 10 million compounds have been identified about 1000 new ones are identified each day!

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organic chemistry

Organic Chemistry

William H. Brown

Christopher S. Foote

Brent L. Iverson

organic chemistry1
Organic Chemistry
  • The study of the compounds of carbon
  • Over 10 million compounds have been identified
    • about 1000 new ones are identified each day!
  • C is a small atom
    • it forms single, double, and triple bonds
    • it is intermediate in electronegativity (2.5)
    • it forms strong bonds with C, H, O, N, and some metals
schematic view of an atom
Schematic View of an Atom
  • a small dense nucleus, diameter 10-14 - 10-15 m, which contains positively charged protons and most of the mass of the atom
  • an extranuclear space, diameter 10-10 m, which contains negatively charged electrons
electron configuration of atoms
Electron Configuration of Atoms
  • Electrons are confined to regions of space called principle energy levels (shells)
    • each shell can hold 2n2 electrons (n = 1,2,3,4......)
electron configuration of atoms1
Electron Configuration of Atoms
  • Shells are divided into subshells called orbitals, which are designated by the letters s, p, d, f,........
    • s (one per shell)
    • p (set of three per shell 2 and higher)
    • d (set of five per shell 3 and higher) .....
electron configuration of atoms2
Electron Configuration of Atoms
  • Aufbau Principle:
    • orbitals fill in order of increasing energy from lowest energy to highest energy
  • Pauli Exclusion Principle:
    • only two electrons can occupy an orbital and their spins must be paired
  • Hund’s Rule:
    • when orbitals of equal energy are available but there are not enough electrons to fill all of them, one electron is added to each orbital before a second electron is added to any one of them
electron configuration of atoms3
Electron Configuration of Atoms
  • The pairing of electron spins
electron configuration of atoms4
Electron Configuration of Atoms
  • Table 1.3 The Ground-State Electron Configuration of Elements 1-18
lewis dot structures
Lewis Dot Structures
  • Gilbert N. Lewis
  • Valence shell:
    • the outermost occupied electron shell of an atom
  • Valence electrons:
    • electrons in the valence shell of an atom; these electrons are used to form chemical bonds and in chemical reactions
  • Lewis dot structure:
    • the symbol of an element represents the nucleus and all inner shell electrons
    • dots represent valence electrons
lewis dot structures1
Lewis Dot Structures
  • Table 1.4 Lewis Dot Structures for Elements 1-18
lewis model of bonding
Lewis Model of Bonding
  • Atoms bond together so that each atom acquires an electron configuration the same as that of the noble gas nearest it in atomic number
    • an atom that gains electrons becomes an anion
    • an atom that loses electrons becomes a cation
    • the attraction of anions and cations leads to the formation of ionic solids
    • an atom may share electrons with one or more atoms to complete its valence shell; a chemical bond formed by sharing electrons is called a covalent bond
    • bonds may be partially ionic or partially covalent; these bonds are called polar covalent bonds
electronegativity
Electronegativity
  • Electronegativity:
    • a measure of an atom’s attraction for the electrons it shares with another atom in a chemical bond
  • Pauling scale
    • generally increases left to right in a row
    • generally increases bottom to top in a column
formation of ions
Formation of Ions
  • A rough guideline:
    • ions will form if the difference in electronegativity between interacting atoms is 1.9 or greater
    • example: sodium (EN 0.9) and fluorine (EN 4.0)
    • we use a single-headed (barbed) curved arrow to show the transfer of one electron from Na to F
    • in forming Na+F-, the single 3s electron from Na is transferred to the partially filled valence shell of F
covalent bonds
Covalent Bonds
  • The simplest covalent bond is that in H2
    • the single electrons from each atom combine to form an electron pair
    • the shared pair functions in two ways simultaneously; it is shared by the two atoms and fills the valence shell of each atom
  • The number of shared pairs
    • one shared pair forms a single bond
    • two shared pairs form a double bond
    • three shared pairs form a triple bond
polar and nonpolar covalent bonds
Polar and Nonpolar Covalent Bonds
  • Although all covalent bonds involve sharing of electrons, they differ widely in the degree of sharing
  • We divide covalent bonds into
    • nonpolar covalent bonds
    • polar covalent bonds
polar and nonpolar covalent bonds1
Polar and Nonpolar Covalent Bonds
  • an example of a polar covalent bond is that of H-Cl
  • the difference in electronegativity between Cl and H is 3.0 - 2.1 = 0.9
  • we show polarity by using the symbols d+ and d-, or by using an arrow with the arrowhead pointing toward the negative end and a plus sign on the tail of the arrow at the positive end
polar covalent bonds
Polar Covalent Bonds
  • Bond dipole moment (m):
    • a measure of the polarity of a covalent bond
    • the product of the charge on either atom of a polar bond times the distance between the nuclei
    • Table 1.7 shows average bond dipole moments of selected covalent bonds
lewis structures
Lewis Structures
  • To write a Lewis structure
    • determine the number of valence electrons
    • determine the arrangement of atoms
    • connect the atoms by single bonds
    • arrange the remaining electrons so that each atom has a complete valence shell
    • show a bonding pair of electrons as a single line
    • show a nonbonding pair of electrons as a pair of dots
    • in a single bond atoms share one pair of electrons, in a double bond they share two pairs of electrons, and in a triple bond they share three pairs of electrons
lewis structures table 1 3
Lewis Structures - Table 1.3
  • In neutral molecules
    • hydrogen has one bond
    • carbon has 4 bonds and no lone pairs
    • nitrogen has 3 bonds and 1 lone pair
    • oxygen has 2 bonds and 2 lone pairs
    • halogens have 1 bond and 3 lone pairs
formal charge
Formal Charge
  • Formal charge: the charge on an atom in a molecule or a polyatomic ion
  • To derive formal charge

1. write a correct Lewis structure for the molecule or ion

2. assign each atom all its unshared (nonbonding) electrons and one-half its shared (bonding) electrons

3. compare this number with the number of valence electrons in the neutral, unbonded atom

formal charge1
Formal Charge
  • Example:Draw Lewis structures, and show which atom in each bears the formal charge
exceptions to the octet rule

6 electrons in the

:

:

:

:

:

:

valence shells of boron

F

C

l

and aluminum

:

:

:

F

B

:

C

l

A

l

:

:

:

:

F

:

:

C

l

:

:

Boron trifluoride

Exceptions to the Octet Rule
  • Molecules containing atoms of Group 3A elements, particularly boron and aluminum

Aluminum chloride

exceptions to the octet rule1
Exceptions to the Octet Rule
  • Atoms of third-period elements have 3d orbitals and may expand their valence shells to contain more than 8 electrons
    • phosphorus may have up to 10
exceptions to the octet rule2
Exceptions to the Octet Rule
  • sulfur, another third-period element, forms compounds in which its valence shell contains 8, 10, or 12 electrons
functional groups
Functional Groups
  • Functional group: an atom or group of atoms within a molecule that shows a characteristic set of physical and chemical properties
  • Functional groups are important for three reason; they are

1. the units by which we divide organic compounds into classes

2. the sites of characteristic chemical reactions

3. the basis for naming organic compounds

alcohols
Alcohols
  • contain an -OH (hydroxyl) group
  • Ethanol may also be written as a condensed structural formula
alcohols1
Alcohols
  • alcohols are classified as primary (1°), secondary (2°), or tertiary (3°) depending on the number of carbon atoms bonded to the carbon bearing the -OH group
alcohols2
Alcohols
  • there are two alcohols with molecular formula C3H8O
amines

:

:

:

C

H

N

H

C

H

N

H

C

H

N

C

H

3

3

3

3

H

C

H

C

H

3

3

Methylamine

Dimethylamine

Trimethylamine

(a 1° amine)

(a 2° amine)

(a 3° amine)

Amines
  • contain an amino group; an sp3-hybridized nitrogen bonded to one, two, or three carbon atoms
    • an amine may by 1°, 2°, or 3°
aldehydes and ketones
Aldehydes and Ketones
  • contain a carbonyl (C=O) group
carboxylic acids
Carboxylic Acids
  • contain a carboxyl (-COOH) group
carboxylic esters
Carboxylic Esters
  • Ester: a derivative of a carboxylic acid in which the carboxyl hydrogen is replaced by a carbon group
carboxylic amide
Carboxylic Amide
  • Carboxylic amide, commonly referred to as an amide: a derivative of a carboxylic acid in which the -OH of the -COOH group is replaced by an amine
    • the six atoms of the amide functional group lie in a plane with bond angles of approximately 120°
vsepr
VSEPR
  • Based on the twin concepts that
    • atoms are surrounded by regions of electron density
    • regions of electron density repel each other
vsepr model
VSEPR Model
  • Example:predict all bond angles for these molecules and ions
polar and nonpolar molecules
Polar and Nonpolar Molecules
  • To determine if a molecule is polar, we need to determine
    • if the molecule has polar bonds
    • the arrangement of these bonds in space
  • Molecular dipole moment (): the vector sum of the individual bond dipole moments in a molecule
    • reported in debyes (D)
polar and nonpolar molecules1
Polar and Nonpolar Molecules
  • these molecules have polar bonds, but each has a zero dipole moment
polar and nonpolar molecules2
Polar and Nonpolar Molecules
  • these molecules have polar bonds and are polar molecules
polar and nonpolar molecules3
Polar and Nonpolar Molecules
  • formaldehyde has polar bonds and is a polar molecule
resonance
Resonance
  • For many molecules and ions, no single Lewis structure provides a truly accurate representation
resonance1
Resonance
  • Linus Pauling - 1930s
    • many molecules and ions are best described by writing two or more Lewis structures
    • individual Lewis structures are called contributing structures
    • connect individual contributing structures by double-headed (resonance) arrows
    • the molecule or ion is a hybrid of the various contributing structures
resonance2
Resonance
  • Examples:equivalent contributing structures
resonance3
Resonance
  • Curved arrow:a symbol used to show the redistribution of valence electrons
  • In using curved arrows, there are only two allowed types of electron redistribution:
    • from a bond to an adjacent atom
    • from an atom to an adjacent bond
  • Electron pushing is a survival skill in organic chemistry
    • learn it well!
resonance4
Resonance
  • All contributing structures must

1. have the same number of valence electrons

2. obey the rules of covalent bonding

    • no more than 2 electrons in the valence shell of H
    • no more than 8 electrons in the valence shell of a 2nd period element
    • 3rd period elements, such as P and S, may have up to 12 electrons in their valence shells

3. differ only in distribution of valence electrons; the position of all nuclei must be the same

4. have the same number of paired and unpaired electrons

resonance5
Resonance
  • The carbonate ion, for example
    • a hybrid of three equivalent contributing structures
    • the negative charge is distributed equally among the three oxygens
resonance6

+

+

••

C

H

O

C

H

C

H

O

C

H

3

3

••

••

H

H

Greater contribution;

Lesser contribution;

both carbon and oxygen have

carbon has only 6 electrons

complete valence shells

in its valence shell

Resonance
  • 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
resonance7
Resonance
  • Preference 2: maximum number of covalent bonds
    • structures with a greater number of covalent bonds contribute more than those with fewer covalent bonds
resonance8

:

:

-

:

:

:

O

O

C

H

-

C

-

C

H

C

H

-

C

-

C

H

3

3

3

3

Greater contribution

Lesser contribution

(no separation of

(separation of unlike

unlike charges)

charges)

Resonance
  • Preference 3: least separation of unlike charge
    • structures with separation of unlike charges contribute less than those with no charge separation
resonance9
Resonance
  • Preference 4: negative charge on the more electronegative atom
    • structures that carry a negative charge on the more electronegative atom contribute more than those with the negative charge on the less electronegative atom
quantum or wave mechanics
Quantum or Wave Mechanics
  • Albert Einstein: E = hn (energy is quantized)
    • light has particle properties
  • Louis deBroglie: wave/particle duality
  • Erwin Schrödinger: wave equation
    • wave function, : a solution to a set of equations that depicts the energy of an electron in an atom
    • each wave function is associated with a unique set of quantum numbers
    • each wave function occupies three-dimensional space and is called an orbital
    •  2 is the probability of finding an electron at a given point in space
shapes of 1 s and 2 s orbitals
Shapes of 1s and 2s Orbitals
  • Probability distribution (2) for 1s and 2s orbitals showing an arbitrary boundary surface containing about 95% of the electron density
shapes of a set of 2 p atomic orbitals
Shapes of a Set of 2p Atomic Orbitals
  • Three-dimensional shapes of 2p atomic orbitals
molecular orbital theory
Molecular Orbital Theory
  • Electrons in atoms exist in atomic orbitals
  • Electrons in molecules exist in molecular orbitals (MOs)
  • Using the Schrödinger equation, we can calculate the shapes and energies of MOs
molecular orbital theory1
Molecular Orbital Theory
  • Rules:
    • combination of n atomic orbitals (mathematically adding and subtracting wave functions) gives n MOs (new wave functions)
    • MOs are arranged in order of increasing energy
    • MO filling is governed by the same rules as for atomic orbitals:
      • Aufbau principle: fill beginning with LUMO
      • Pauli exclusion principle: no more than 2e- in a MO
      • Hund’s rule: when two or more MOs of equivalent energy are available, add 1e- to each before filling any one of them with 2e-
molecular orbital theory2
Molecular Orbital Theory
  • Terminology
    • ground state = lowest energy state
    • excited state = NOT lowest energy state
    •  = sigma bonding MO
    • * = sigma antibonding MO
    •  = pi bonding MO
    • * = pi antibonding MO
    • HOMO = highest occupied MO
    • LUMO = lowest unoccupied MO
molecular orbital theory3
Molecular Orbital Theory
  • Sigma 1s bonding and antibonding MOs
molecular orbital theory4
Molecular Orbital Theory
  • MO energy diagram for H2: (a) ground state and (b) lowest excited state
molecular orbitals
Molecular Orbitals
  • computed sigma bonding and antibonding MOs for H2
molecular orbitals1
Molecular Orbitals
  • pi bonding and antibonding MOs
molecular orbitals2
Molecular Orbitals
  • computed pi bonding and antibonding MOs for ethylene
molecular orbitals3
Molecular Orbitals
  • computed pi bonding and antibonding orbitals for formaldehyde
hybrid orbitals
Hybrid Orbitals
  • The Problem:
    • bonding by 2s and 2p atomic orbitals would give bond angles of approximately 90°
    • instead we observe bond angles of approximately 109.5°, 120°, and 180°
  • A Solution
    • hybridization of atomic orbitals
    • 2nd row elements use sp3, sp2, and sp hybrid orbitals for bonding
hybrid orbitals1
Hybrid Orbitals
  • Hybridization of orbitals (L. Pauling)
    • the combination of two or more atomic orbitals forms a new set of atomic orbitals, called hybrid orbitals
  • We deal with three types of hybrid orbitals

sp3(one s orbital + three p orbitals)

sp2 (one s orbital + two p orbitals)

sp (one s orbital + one p orbital)

  • Overlap of hybrid orbitals can form two types of bonds depending on the geometry of overlap

 bondsare formed by “direct” overlap

 bondsare formed by “parallel” overlap

sp 3 hybrid orbitals
sp3 Hybrid Orbitals
  • each sp3 hybrid orbital has two lobes of unequal size
  • the sign of the wave function is positive in one lobe, negative in the other, and zero at the nucleus
  • the four sp3 hybrid orbitals are directed toward the corners of a regular tetrahedron at angles of 109.5°
sp 3 hybrid orbitals1
sp3 Hybrid Orbitals
  • orbital overlap pictures of methane, ammonia, and water
sp 2 hybrid orbitals
sp2 Hybrid Orbitals
  • the axes of the three sp2 hybrid orbitals lie in a plane and are directed toward the corners of an equilateral triangle
  • the unhybridized 2p orbital lies perpendicular to the plane of the three hybrid orbitals
sp hybrid orbitals
sp Hybrid Orbitals
  • two lobes of unequal size at an angle of 180°
  • the unhybridized 2p orbitals are perpendicular to each other and to the line created by the axes of the two sp hybrid orbitals
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