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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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