Carbon-carbon Covalent Bonds

Chapter 2Representative Carbon Compounds:Functional Groups, Intermolecular

Forces and Infrared (IR) Spectroscopy

The ability of carbon to form covalent bonds to other carbon atoms, as well as atoms such as hydrogen, oxygen, nitrogen and others, leads to an enormous variety of structures. Many of these are found in biological systems.

Carbon-carbon Covalent Bonds• Carbon forms strong covalent bonds to other

carbons and to other elements such as hydrogen, oxygen, nitrogen and sulfur• This accounts for the vast variety of organic

compounds possible• Organic compounds are grouped into

functional group families• A functional group is a specific grouping of atoms

(e.g. carbon- carbon double bonds are in the family of alkenes)• An instrumental technique called infrared (IR)

spectroscopy is used to determine the presence of specific functional groups

2

Hydrocarbons: Representative Alkanes, Alkenes Alkynes, and Aromatic Compounds• Hydrocarbons contain only carbon and hydrogen

atoms• Subgroups of Hydrocarbons:

• Alkanes contain only carbon-carbon single bonds• Alkenes contain one or more carbon-carbon double

bonds• Alkynes contain one or more carbon-carbon triple bonds• Aromatic hydrocarbons contain benzene-like stable

structures (discussed later)• Saturated hydrocarbons: contain only carbon-

carbon single bonds e.g. alkanes• Unsaturated hydrocarbons: contain double or triple

carbon-carbon bonds e.g. alkene, alkynes, aromatics• Contain fewer than maximum number of hydrogens per

carbon• Capable of reacting with H2 to become saturated

3

Methane

This first member of the alkane family is the major component in natural gas (60-80%) found in geological deposits. A large industry is based on the recovery and distribution of natural gas that is widely used as a fuel for homes and industry. It is also used as a primary energy source in power plants to produce electricity. About 10% of the electricity produced in the United States comes from natural gas-fired power plants. It is the cleanest of the fossil fuels (coal, petroleum, natural gas).

Methane is produced by bacteria called methanogens under anaerobic (oxygen-free) conditions. These microorganisms are found in ocean trenches, mud, sewage and in the stomach of cows. Very large deposits of methane hydrates (methane trapped in ice crystals) have been discovered in deep ocean trenches.

109.5oThe four equivalent hydrogens in methane are located in the corners of a tetrahedron.

CH

HH

H

Representative Hydrocarbons• Alkanes• Principle sources of alkanes are natural gas and

petroleum– Smaller alkanes (C1 to C4) are gases at room temperature

• Methane is – A component of the atmosphere of many planets– Major component of natural gas – Produced by primitive organisms called methanogens found in

mud, sewage and cows’ stomachs

5

Alkenes

Hydrocarbons that contain a carbon-carbon double bond are called alkenes. The two simplest alkenes, ethene (C2H4) and propene (C3H6), are important industrial chemicals used to prepare polymers.

ethenepolyethylene

(30 billion pounds per year in the United States)

CH2=CH2

propenepolypropene

(15 billion pounds per year in the United States)

CH3CH=CH2

• Alkenes• Ethene (ethylene) is a major industrial feedstock

– Used in the production of ethanol, ethylene oxide and the polymer polyethylene

• Propene (propylene) is also very important in industry– Molecular formula C3H6

– Used to make the polymer polypropylene and is the starting material for acetone

• Many alkenes occur naturally

7

Alkynes

Hydrocarbons that contain the carbon-carbon triple bond are called alkynes. The simplest alkyne, ethyne (C2H2), is an important industrial chemical used as a chemical intermediate (to make other important chemicals) and as a fuel in the oxyacetylene torch for welding.

other organicsethyne

HC CH

• Alkynes• Many alkynes are of biological interest

– Capillin is an antifungal agent found naturally– Dactylyne is a marine natural product– Ethinyl estradiol is a synthetic estrogen used in oral

contraceptives

9

• Benzene: A Representative Hydrocarbon• Benzene is the prototypical aromatic compound

– The Kekulé structure (named after August Kekulé who formulated it) is a six-membered ring with alternating double and single

bonds

10

• Benzene: A Representative Hydrocarbon• Benzene does not actually have discreet single and double

carbon-carbon bonds– All carbon-carbon bonds are exactly equal in length (1.38 Å) – This is between the length of a carbon-carbon single bond and a

carbon-carbon double bond • Resonance theory explains this by suggesting there are two

resonance hybrids that contribute equally to the real structure– The real structure is often depicted as a hexagon with a circle in

the middle

11

• Molecular orbital theory explains the equal bond lengths of benzene by suggesting there in a continuous overlap of p orbitals over the entire ring• All carbons in benzene are sp2 hybridized

– Each carbon also has an empty p orbital • Each p orbital does not just overlap with one adjacent p but overlaps

with p orbitals on either side to give a continuous bonding molecular orbital that encompasses all 6 carbons

• All 6 p electrons are therefore delocalized over the entire ring and this results in the equivalence of all of the carbon-carbon bonds

12

Polar Covalent Bonds• Polar covalent bonds occur when a covalent

bond is formed between two atoms of differing electronegativities• The more electronegative atom draws electron

density closer to itself• The more electronegative atom develops a partial

negative charge (d-) and the less electronegative atom develops a partial positive charge (d+)• A bond which is polarized is a dipole and has a

dipole moment• The direction of the dipole can be indicated by a

dipole arrow – The arrow head is the negative end of a dipole, the

crossed end is the positive end 13

The polarization of electrons in a covalent bond leads to a permanent separation of charge:

Bond Dipoles

B is more electronegativeBAdd

A permanent bond dipole (two centers of charge) is formed where d (delta) means partial. This electrical imbalance leads to electropositive and electronegative regions within the molecule.

Examples

dd dd

dO

H

HH F

• Example: the molecule HCl• The more electronegative chlorine draws electron density away

from the hydrogen – Chlorine develops a partial negative charge

• The dipole moment of a molecule can be measured experimentally• It is the product of the magnitude of the charges (in electrostatic

units: esu) and the distance between the charges (in cm)• The actual unit of measurement is a Debye (D) which is equivalent

to 1 x 10-18 esu cm

15

• A map of electrostatic potential (MEP) is a way to visualize distribution of charge in a molecule• Parts of the molecule which are red have relatively more

electron density or are negative – These region would tend to attract positively charged species

• Parts of the molecule which are blue have relatively less electron density or are positive– These region would tend to attract negatively charged species

• The MEP is plotted at the van Der Waals surface of a molecule– This is the farthest extent of a molecule’s electron cloud and therefore

indicates the shape of the molecule• The MEP of hydrogen chlorine clearly indicates that the

negative charge is concentrated near chlorine– The overall shape of the molecule is also represented

16

Molecular Dipole• In diatomic molecules a dipole exists if the two

atoms are of different electronegativity• In more complicated molecules the molecular dipole

is the sum of the bond dipoles• Some molecules with very polar bonds will have no

net molecular dipole because the bond dipoles cancel out– The center of positive charge and negative charge coincide

in these molecules

17

• Examples• In carbon tetrachloride the bond dipoles cancel and the

overall molecular dipole is 0 Debye

• In chloromethane the C-H bonds have only small dipoles but the C-Cl bond has a large dipole and the molecule is quite polar

18

19

An unshared pair of electrons on atoms such as oxygen and nitrogen contribute a great deal to a dipole Water and ammonia have very large net dipoles

• Some cis-trans isomers differ markedly in their dipole moment• In trans 1,2-dichloroethene the two carbon-chlorine

dipoles cancel out and the molecular dipole is 0 Debye• In the cis isomer the carbon-chlorine dipoles

reinforce and there is a large molecular dipole

20

Molecular Dipole Moments in Polyatomic Molecules

.

In polyatomic molecules, the molecular dipole moment is a vector sum of the bond dipole moments. From measurements of molecular dipole moments and structural studies of numerous polyatomic molecules, approximate bond dipole moments have been calculated

Some Approximate Bond Dipole Moments (D)

C-C 0.0H-C 0.30H-N 1.31H-O 1.53

H-F 1.98H-Cl 1.03H-Br 0.78H-I 0.38

C-F 1.51C-Cl 1.56C-Br 1.48C-I 1.29

C-O 0.86C=O 2.40C-N 0.40C=N 0.90

from L. N. Ferguson, The Modern Structural Theory of Organic Chemistry

The Importance of Molecular Symmetry

Carbon Tetrachloride, CCl4

.

Each C-Cl bond is polar covalent. But carbon tetrachloride does not have a molecular dipole moment. Because of the tetrahedral symmetry, there is no net molecular dipole moment and CCl4 is a nonpolar molecule

tetrahedral geometry

CCl

ClCl

Cl

1.56 D

1.56 D

1.56 D

1.56 D

The molecular dipole momentis the vector sum of all the bond dipoles.

Evaluate the bonddipole vectors in the -z direction.

1.56 D

+z

-z

109.5o

70.5o

C CCl

Cl

The vector contribution of the bond dipole in the -z direction is1.56 D x COS 70.5 = 1.56 D x .3333

ooverall =1.56 D - 3 x 1.56 D x .3333 = 0 D

+z -z

Methyl Chloride, CH3Cl

.The molecular dipole moment of methyl chloride is 1.86 D

The analysis assumes a tetrahedral geometry.

C

Cl

H

HH

1.560.30

0.30

0.30

The net dipole moment in the +Z direction is

1.560.30 3 x (C-H)

1.86 D

(C-Cl)

Note: The direction of the C-H bond moment is determined from this and similar analyses to be H-C.

Nonbonding Electron Pairs

.The contribution of nonbonding electron pairs to the molecular dipole moment is seen from the following comparison

= 0.24 D = 1.46 D

NF3 NH3

The much larger molecular dipole moment for NH3 is, at first, surprising, because of the similar electronegativity differences of the atoms in the covalent bonds, H-F and N-H. Opposite directions would be expected for the molecular dipole moments.

3.0 3.0

4.0

4.0

4.0 2.12.1

2.1

N NF FF

H

HH

Contributions from Nonbonding Electrons

.

When dipole contributions from the nonbonding electron pairs are included, the different magnitudes for the molecular dipole moments of NF3 and NH3 are understandable

Partial cancellation of the bond dipoles leads to a small molecular dipolemoment either or .

Bond dipoles are reinforcingleading to a larger moleculardipole moment.

0.24 D 1.46 D

NF FF

NH

HH

Functional Groups• Functional group families are characterized by the

presence of a certain arrangement of atoms called a functional group• A functional group is the site of most chemical

reactivity of a molecule– The functional group is responsible for many of the

physical properties of a molecule• Alkanes do not have a functional groups

– Carbon-carbon single bonds and carbon-hydrogen bonds are generally very unreactive

28

• Alkyl Groups and the Symbol R• Alkyl groups are obtained by removing a hydrogen from an alkane• Often more than one alkyl group can be obtained from an alkane by removal of different kinds of

hydrogens

• R is the symbol to represent a generic alkyl groups– The general formula for an alkane can be abbreviated R-H

29

• A benzene ring with a hydrogen removed is called a phenyl and can be represented in various ways

• Toluene (methylbenzene) with its methyl hydrogen removed is called a benzyl group

30

• Alkyl Halides• In alkyl halides, halogen (F, Cl, Br, I) replaces the

hydrogen of an alkane• They are classified based on the carbon the halogen

is attached to– If the carbon is attached to one other carbon that carbon

is primary (1o) and the alkyl halide is also 1o

– If the carbon is attached to two other carbons, that carbon is secondary (2o) and the alkyl halide is 2o

– If the carbon is attached to three other carbons, the carbon is tertiary (3o) and the alkyl halide is 3o

31

Primary Alkyl Halide

CH

RH

primary(1o)

CH

CH3H

ethyl

CH

CH3H

Cl

ethyl chloridea primary alkyl chloride

Secondary Alkyl Halide

secondary (2o)

CR

RH

isopropyl

CCH3

CH3H

a secondary alkyl chlorideisopropyl chloride

CCH3

CH3H

Cl

Tertiary Alkyl Halide

tertiary (3o)

CR

RR

tertiary butyl

CCH3

CH3CH3

Cl

a tertiary alkyl chloridetertiary butyl chloride

CCH3

CH3CH3

Cl

What kind of a carbon is this?

Alcohols• In alcohols the hydrogen of the alkane is replaced by the

hydroxyl (-OH) group– An alcohol can be viewed as either a hydroxyl derivative of an

alkane or an alkyl derivative of water

34

Alcohols• Alcohols are also classified according to the carbon the

hydroxyl is directly attached to

35

Classification of Alcohols

.

Alcohols are classified as primary, secondary or tertiary according to the structure around the carbon to which the hydroxyl group is attached

ethyl alcohol

CH3CH2OH

a 1o alcohol

CCH3

H

HOH

isopropyl alcohol

(CH3)2CHOH

a 2o alcohol

CCH3

CH3

HOH

tertiary-butyl alcohol

(CH3)3COH

a 3o alcohol

CCH3

CH3OH

CH3

Name the following organic compounds using common alkyl group names.

Quiz Chapter 2 Section 7

CH3Cl CH3CHBrCH3 CH3CH2CH2OH

methyl chloride isopropyl bromide propyl alcohol

• Ethers• Ethers have the general formula R-O-R or R-O-R’ where R’

is different from R– These can be considered organic derivatives of water in which

both hydrogens are replaced by organic groups – The bond angle at oxygen is close to the tetrahedral angle

38

• Amines• Amines are organic derivatives of ammonia

– They are classified according to how many alkyl groups replace the hydrogens of ammonia

– This is a different classification scheme than that used in alcohols

39

Classifications and Names of Amines

.

Simple amines are named using the alkyl group name followed by "amine." When more than one R group is present, the prefixes "di" and "tri" are included. If there are R and R' groups, the alkyl groups are named in alphabetical order

ethylamine(a 1o amine)

CH3CH2NH2

isopropylamine(a 1o amine)

CH3CHCH3NH2

tertiary-butylamine(a 1o amine)

CH3CCH3NH2

CH3

diisopropylamine(a 2o amine)

(CH3)2CHNHCH(CH3)2

piperidine(a 2o amine)

H2C

H2CN

CH2

CH2

H2C

HNH

triethylamine(a 3o amine)

CH3CH2NCH2CH3

CH2CH3

diethylmethylamine(a 3o amine)

(CH3CH2)2NCH3

The Structure and Base Property of Amines

Amines have a trigonal pyramidal structure similar to ammonia.

.

VSEPR theory predicts a tetrahedral geometry

:

107o

ammonia trimethylamine

NH

HH

:

108.7o

NH3C

CH3CH3

Amines are bases.

base

+

acid

:

(CH3)3N HCl

salt

NH

H3CCH3

CH3

Cl-


Identify the following amines as primary, secondary or tertiary types, and give their alkyl group names.

(CH3)2NH (CH3)3CN(CH3)2 CH3CH2CH2CH2NH2

(a 1o aminebutylamine

) (a 2o aminedimethylamine

) (a 3o aminetert-butyldimethylamine

)

• Aldehydes and Ketones• Both contain the carbonyl group

• Aldehydes have at least one carbon attached to the carbonyl group

• Ketones have two organic groups attached to the carbonyl group

• The carbonyl carbon is sp2 hybridized – It is trigonal planar and has bond angle about 120o

43

Aldehydes and Ketones

The Carbonyl Functional Group

C O

The carbon-oxygen double bond can be viewed as the combination of sp2 hybridized carbon and oxygen atoms.

1s

2s2p valence

level

atomic C

2s2p

atomic O

1s

Examples of Aldehydes and Ketones

Aldehydes

One of the two groups attached to the carbon atom is H.

general structure R may be alkylor arylOC

R

H

The simplest aldehyde is formaldehyde with R= H.

Ketones

The two groups attached to the carbon are R and R'. No H.

general structure R and R' are alkyl or aryl

OCR

R'

The simplest ketone is acetone with R, R'= CH3.

Aldehydes

Ketones

O

H Hformaldehyde

O

Hacetaldehyde

O

H

benzaldehyde

O

acetone

O

ethyl methyl ketone

O

carvone

• Carboxylic Acids, Esters and Amides• All these groups contain a carbonyl group bonded to

an oxygen or nitrogen• Carboxylic Acids

– Contain the carboxyl (carbonyl + hydroxyl) group

47

Examples of Carboxylic Acids

formic acid(R = H)

O

H OHCO2H HCOOHH

propanoic acid(R = ethyl)

O

OCH3CH2CO2H CH3CH2COOHH

benzoic acid(R = phenyl)

O

OC6H5CO2H C6H5COOHH

Derivatives of Carboxylic Acids

There are several families of organic compounds that are considered to be derivatives of carboxylic acids because their structures can be related to carboxylic acids through reaction with water (hydrolysis).

O

R Xa derivative

O

R OH

carboxylic acid

hydrolysis

H2O

• Esters– A carbonyl group is bonded to an alkoxyl (OR’) group

50

• Amide– A carbonyl group is bonded to a nitrogen derived from

ammonia or an amine

• Nitriles • An alkyl group is attached to a carbon triply bonded to

a nitrogen– This functional group is called a cyano group

51

Summary of Important Families of Organic Compounds

52

Summary (cont.)

53

aldehyde ester ketone

carboxylic acid ketone amide


Identify the functional group in the following carbonyl compounds.

HCOOH

H

O

O

OO

O

NH2

O

Physical Properties and Molecular Structure

Intermolecular forces have a major influence on the physical properties of compounds. These forces arise from permanent and transient (temporary) electrical charges within structures that operate over short distances (several angstroms) and, therefore, are only important in the condensed phases (solids and liquids).

Ion-Ion Forces

Ionic materials generally exist in well-ordered solid state structures of alternating positive and negative charges. The strong electrostatic attractive forces in the solid state are expressed as the Lattice Energy as defined in the example of sodium chloride.

Na(g)+

+ Cl(g)- NaCl(s) Hlattice

energy= -787 kJ/mol

The lattice energy is the energy released when one mole of the ions condense from the gas phase into a solid structure.

.

When ionic materials melt, the well-ordered crystalline solid changes into the more random (less ordered) liquid state with the partial loss of strong (stabilizing) ion-ion attractive forces. To achieve the less stable liquid state, considerable thermal energy must be added, resulting in high melting points for these materials

ionic materials ions mp (oC)

sodium chloride 801

sodium acetate 324

ammonium acetate 114

Na+ Cl-

Na+ CH3CO2-

NH4+ CH3CO2

-

Ion-Ion Forces

The introduction of polyatomic ions into the structure tends to lower the melting point of ionic materials, because these large groups with covalent bonds increase the distance between the ions, decreasing the attractive forces.

stronger ion-ion forces inthe solidstate

.Boiling points of ionic materials are very high and decomposition often occurs before vaporization

Dipole-Dipole Forces

Most organic compounds do not have ionic bonds. They possess covalent bonds. Because of electronegativity differences, many covalent bonds are polar leading to permanent molecular dipole moments, when the internal bond dipole moments do not cancel. A molecule with a permanent dipole moment is polar.

++

+

+ + +

+-

-- -

d

d

d

d d d

d

d

d

d d d - d -

d -

Intermolecular Dipole-Dipole Forces

In the condensed states (liquids and solids), molecules are sufficiently close together (within a few angstroms) for the dipoles of different molecules to interact. The negative pole of one dipole attracts the positive pole of another dipole, while similarly charged poles of different dipoles repel.

Dipole-dipole interactions are stronger in the more ordered solid state compared with the less ordered liquid state. Polar materials tend to have higher melting and boiling points than nonpolar materials of comparable size.

These dipole-dipole interactions introduce order and stabilize the solid and liquid states.

Physical Properties and Molecular Structure• The strength of intermolecular forces (forces

between molecules) determines the physical properties (i.e. melting point, boiling point and solubility) of a compound• Stronger intermolecular forces result in high

melting points and boiling points– More energy must be expended to overcome very strong

forces between molecules• The type of intermolecular forces important for a

molecule are determined by its structure• The physical properties of some representative

compounds are shown on the next slide

60

61

Indentify the intermolecular interaction for each molecule

• Ion-Ion Forces• Ion-ion forces are between positively and negatively charged ions• These are very strong forces that hold a solid compound

consisting of ions together in a crystalline lattice– Melting points are high because a great deal of energy is required to

break apart the crystalline lattice• Boiling points are so high that organic ions often decompose

before they boil• Example: Sodium acetate

62

• Dipole-Dipole Forces• Dipole-dipole forces are between molecules with

permanent dipoles– There is an interaction between d+ and d- areas in each

molecule; these are much weaker than ion-ion forces– Molecules align to maximize attraction of d+ and d-

parts of molecules– Example: acetone

63

• Hydrogen Bonds• Hydrogen bonds result from very strong dipole-

dipole forces• There is an interaction between hydrogens bonded

to strongly electronegative atoms (O, N or F) and nonbonding electron pairs on other strongly electronegative atoms (O, N or F)

64

• Example• Ethanol (CH3CH2OH) has a boiling point of +78.5oC; its isomer

methyl ether (CH3OCH3) has a boiling point of -24.9oC– Ethanol molecules are held together by hydrogen bonds whereas

methyl ether molecules are held together only by weaker dipole-dipole interactions

• A factor in melting points is that symmetrical molecules tend to pack better in the crystalline lattice and have higher melting points

65

The Influence of Hydrogen Bonds on Physical Properties

Alcohols have higher boiling points than other polar compounds of comparable size because they form hydrogen bonds in the liquid state that are lost in the gas state.

=

dipole momentMP

solubility in water

CH3CH2OH CH3CHO

MW

BP

46 44

1.69 D 2.69 D

-115 oC -121 oC

78.5 oC 20 oC

miscible miscible

Hydrogen bonding raises the boiling point because a network of intermolecular bonds (1-9 kcal/mol) must be broken in the transition from the liquid to the gas state. Note this network is not broken in the transition from the solid to the liquid state, as revealed by the similar melting points of acetone and acetaldehyde.


In each pair of structures below, circle the one with the higher boiling point, and indicate the intermolecular attractive force responsible for the difference.

O

H

OOH

OH O

dipole-dipole

hydrogen bonding

hydrogen bonding

Influence of Molecular Shape

In addition to polar influences, the melting point of organic compounds is affected by molecular shape. Molecules that are more symmetrical tend to pack better in the solid state, and have higher melting points than comparable molecules with other shapes.

Alcohol Isomers of C4H10O

tert-butyl alcohol

butyl alcohol isobutyl alcohol sec.-butyl alcohol

MP 25 oC -90 oC -108 oC -114 oC

CH3-C-OHCH3

CH3

CH3CH2CH2CH2OH CH3CHCH2OHCH3

CH3CH2CHOHCH3

Additional Examples

Alkanes and Cycloalkanes

MP -188 oC -127 oC -138 oC -50 oC

-42 oC -32 oC -0.5 oC 12.5 oC

CH3CH2CH3 CH3CH2CH2CH3

BP

.

The more compact cycloalkanes have melting points significantly higher than the comparable alkanes. The boiling points for these small sized alkanes and cycloalkanes are not very different

• van der Waals Forces (London or Dispersion Forces)• Van der Waals forces result when a temporary dipole

in a molecule caused by a momentary shifting of electrons induces an opposite and also temporary dipole in an adjacent molecule– These temporary opposite dipoles cause a weak attraction

between the two molecules– Molecules which rely only on van der Waals forces generally

have low melting points and boiling points

71

Polarizability predicts the magnitude of van der Waals Interactions• Polarizability is the ability of the electrons on an atom to respond to a changing electric field• Atoms with very loosely held electrons are more polarizable• Iodine atoms are more polarizable than fluorine atoms because the outer shell electrons are

more loosely held• Atoms with unshared electrons are more polarizable (a halogen is more polarizable than an

alkyl of similar size)

All things being equal larger and heavier molecules have higher boiling points• Larger molecules need more energy to escape the surface of the liquid• Larger organic molecules tend to have more surface area in contact with each other and so

have stronger van der Waals interactions• Methane (CH4) has a boiling point of -162oC whereas ethane (C2H6) has a boiling point of -

88.2oC72

• Solubilities• Water dissolves ionic solids by forming strong dipole-

ion interactions – These dipole-ion interactions are powerful enough to

overcome lattice energy and interionic interactions in the solid

75

• Generally like dissolves like• Polar solvents tend to dissolve polar solids or polar liquids• Methanol (a water-like molecule) dissolves in water in all

proportions and interacts using hydrogen-bonding to the water

• A large alkyl group can overwhelm the ability of the polar group to solubilize a molecule in water

– Decyl alcohol is only slightly soluble in water– The large alkyl portion is hydrophobic (“water hating”) and

overwhelms the capacity of the hydrophilic (“water loving”) hydroxyl

76

.

Hydrogen bonding increases the intermolecular attraction between solute and solvent, promoting solubility. For example, alcohols with small R groups are miscible (dissolve in all proportions) in water

Hydrogen Bonding Between Solute and Solvent

+ H2O

intermolecularhydrogen bonding

intermolecularhydrogen bonding

ROHdd

d

d

d

d

d

d

d

d

etc.

hydrogen bonding between alcohols and water when R is small

OR

HH

O

O

O

H

H

H

R

H

.

As the size of R increases, the solubility of an alcohol in water decreases. A long hydrocarbon chain is hydrophobic ("water fearing") while the hydroxyl group is hydrophilic ("water loving"). The final solubility is a balance of these two factors

Guidelines for Water Solubility

The solubility limit of a material is usually given in grams of solute per 100-mL of solvent. An organic molecule is considered to be water soluble if at least 3 g of the solute dissolves in 100-mL of water.

.• Organic compounds with one oxygen atom and one to three carbon atoms are water soluble

• Compounds with one oxygen and 4 or 5 carbon atoms have limited solubilities.

• Compounds with one oxygen and 6 carbons or more are usually insoluble in water.

Spectroscopic Methods for Structure Determination

One of the most powerful tools of science today is spectroscopic methods of structure analysis. Many of these methods also are used for chemical analysis to determine the amount and purity of a material.

All of these methods depend on the interaction of a molecule with electromagnetic radiation. Molecules absorb radiant energy in very well defined ways that are characteristic of the molecular structure.

Absorption of Electromagnetic Radiation

M + hn

M+ + e-.

ionization

M* electronic excitation

M* vibrational/rotational excitation

M* nuclear spin excitation

Molecular Changes

The absorption of electromagnetic radiation by a molecule may cause any of several different changes depending on the magnitude of the energy uptake, E. In all cases, the increase in energy is considered to be associated with a transition from a lower to higher energy state, so the uptake in energy is an exact value. Further, only those values of energy are absorbed that match the energy gaps between states in a particular molecule.

ener

gy

E = h

lower state

higherstate

Characterization of Electromagnetic Radiation

.Electromagnetic radiation has the properties of a wave, a repeating mathematical function describing a series of crests and troughs

electric or magnetic field x

The wave describes oscillations in the magnitude of electric and magnetic fields of electromagnetic radiation moving in the x direction

wavelength

one complete cycle one complete cycle

Electromagnetic radiation is characterized either by:(A) the length of one complete cycle of the wave which is the wavelength,d(B) the number of complete cycles passing a point per second which is the frequency, n (nu), called cycles/second, or hertz or s-1 .

cycles per second

moving wave

Some Important Relationships

Wavelength and Frequency

The wavelength and frequency of electromagnetic radiation are related: x = c

where C is the "speed of light"3.0 x 1010 cm/s (vacuum)

.

This relationship means that an electromagnetic wave may be characterized by either the wavelength or frequency of the radiation.Some fields of spectroscopic study, by tradition, identify the electromagnetic radiation by wavelength and others use the frequency

Energy Content of Electromagnetic Radiation

The energy content of electromagnetic radiation, which defines the increase in energy when radiation is absorbed by a molecule, is given by

E = h = hcwhere h is Planck's constant.

The Electromagnetic Spectrum

.

The electromagnetic spectrum is a continuum of radiation energy. Different sources will emit radiation in specific regions of the spectrum, and some sources (such as atoms) may emit only very narrow bands or lines of radiation energy

The Electromagnetic Spectrum

visible near uv

vacuum uv

x-rayscosmic and-rays

near ir

infrared micro-wave

radio

increasing frequency,

Hz 1019 1015 1013 108

increasing wavelength,

0.1 nm 200 nm 400 nm 800 nm 2 mm 50 mm

energyequivalentin kcal/mol

molecularchange

~106 ~102 ~100~10-5

ionization electronicexcitation

vibrational-rotationaltransition

nuclear spin transition

Problem

What is the energy equivalent in kcal/mol of radiation absorbed by a molecule at 300 nm (nanometers)?

Frequency = c/ = (3.00 x 1010 cm/s)/(3.00 x 10-5 cm) = 1.00 x 1015 s-1

E = h = (6.62 x 10-27 erg-sec/molecule) x 1.00 x 1015 s-1

= 6.62 x 10-12 ergs/molecule

E = (6.62 x 10-12 ergs/molecule) x 1.44 x 1016 (conversion factor) = 9.55 x 104 cal/mol

E = 9.55 x 104 x (1 kcal/103 cal) = 95.5 kcal/mol

Absorption of this radiation in the near ultraviolet causes an electronic transition within the molecule.

Wavelength in cm = 300 nm x 10-9 m/nm x 102 cm/m = 3.00 x 10-5 cm

Solution

.

Almost all organic molecules are "infrared active," meaning they absorb radiation in the infrared region. These absorptions are associated with transitions from lower to higher vibrational levels within molecules. The energy increase due to absorption of an infrared photon is typically 1 - 10 kcal/mol

E = h

lower vibrationalstate

higher vibrationalstate

.

The energy gap between the statesdepends on the atoms present, the bond strength, and the details of the molecular change

Infrared Spectroscopy: An Instrumental Method for Detecting Functional Groups

Infrared Spectroscopy: An Instrumental Method for Detecting Functional Groups• Electromagnetic radiation in the infrared (IR)

frequency range is absorbed by a molecule at certain characteristic frequencies• Energy is absorbed by the bonds in the molecule and they

vibrate faster• The bonds behave like tiny springs connecting the atoms

– The bonds can absorb energy and vibrate faster only when the added energy is of a particular resonant frequency

• The frequencies of absorption are very characteristic of the type of bonds contained in the sample molecule

• The type of bonds present are directly related to the functional groups present

• A plot of these absorbed frequencies is called an IR spectrum

88

Spectral Presentation: A Preview

Infrared spectra are recorded by passing a broad spectral range of the radiation through a sample, and noting the wavelengths that are absorbed. The pattern of wavelengths absorbed provides structural information on the organic compound.

bands indicateabsorption of IRradiation

m (10-4 cm)wavelength

4000 3000 2000 1000

wavenumber, cm-1

2.5 3 4 5 6 7 10 15

tran

smitt

ance

(%) 100

50

0

absorbance

.1

.4

.8

8

In addition to the spectral wavelength presented in micrometer or micron units (10-4 cm or 10-6 m), the spectrum is displayed in wavenumber units (). This unit is linear with frequency with = 1/ (with l in cm) = /c.

A Mechanical Model: Hooke's Law

.

Two atoms connected by a covalent bond behave like two weights on the ends of a spring. When the weights are displaced from their rest position and released, they vibrate only with certain frequencies that depend on the masses of the weights, and a constant related to the stiffness of the spring.

The mathematical relationship between the allowed frequency (n), the masses and the constant (k) is called Hooke's Law..

(s-1) = x c = 2p

1 MAMB MA + MB

where is the force constant in dynes/cm or g-cm s-2/cm.

.

In a similar way, the frequency of the radiant energy absorbed during a vibrational transition depends only on the masses of the atoms and a force constant related to the strength of the covalent bond

Vibrational States in Molecules

.

In molecules, the vibrational states are quantized in a series of allowed states of increasing energies Evib = (n + 1/2)h where n = 0, 1, 2, 3...etc. The allowed transitions are between states where n = + 1, which means only adjacent states. Accordingly, in an allowed transition between vibrational states, the change in energy is E = h, with the frequency expressed by Hooke's Law

lower vibrational state

higher vibrational state

+h

(s-1) = x c = 2p1

MAMB MA + MB

where MA and MB are the masses of the atoms

.

The infrared spectra of the H-X compounds show a single absorption band due to a vibrational transition from state n = 0 to state n = 1 with an increase in internal energy of h

Some Vibrational Transitions in Diatomics

molecule (cm-1) (dynes/cm)(calculated)(measured)

(s-1)

3958 1.187 x 1014 8.8 x 105

2885 8.655 x 1013 4.8 x 105

2559 7.677 x 1013 3.8 x 105

2230 6.690 x 1013 2.9 x 105

H-FH-ClH-BrH-I

.

From the experimentally measured values of , and the masses of the atoms, the force constants may be calculated by Hooke's Law. The decreasing values of from H-F through H-I parallel the decreasing bond strengths of the H-X compounds

Additional Meaningful Trends

.

There are several additional meaningful correlations between observed infrared transitions and covalent bond strengths that follow from Hooke's Law. In each example below, the atoms (and therefore the masses) connected by the covalent bond remain the same. Therefore, the trend in the values of is due to changing values for the force constants (bond strengths)

C-H Bond Stretch

characteristic ranges (cm-1)

alkanes alkenes alkynes

3000-2850 3100-3000 ~3300

increasing force constants(bond strengths)

Carbon-Carbon Bond Stretch

characteristic ranges (cm-1)

C C C C C C

800-1200 1600-1680 2100-2250


Carbon-Nitrogen Bond Stretch

characteristic ranges

(cm-1)C N C N C N

1000-1350 1640-1690 2240-2260


Other Normal Modes of Vibration

The normal modes of vibration refers to the possible motions of atoms with no change in the center of mass of the molecule. The two fundamental types of vibrations are stretching and bending. As shown above, stretching is the back and forth movement of atoms along the bond axis. Bending is the opening and closing of a bond angle. Various names such as twisting, scissoring and rocking are used to describe the bending motions of groups of atoms .

Polyatomic Molecules

The only normal mode in a diatomic molecule is the stretch.

Larger Polyatomics

The number of normal modes in larger molecules depends on the number of atoms.

3n - 6polyatomics linear polyatomics

3n - 5where n is the number of atoms

• There are 2 basic different types of stretching and bending vibrations induced by the absorption of infrared energy

96

.

The interaction between the vibrating molecule and the radiation field requires a change in dipole moment during the vibrational transition. This condition must be met for the vibrational mode to be "infrared active "

Selection Rule for Light Absorption

Examples

A

H Cl+h

stretch

B

B > AH Cl

.Because the bond length increases in the higher vibrational state, the dipole moment is larger and this vibrational transition is infrared active

AH H

O +h

symmetrical stretch BB > AH H

O

.The molecular dipole moment increases in the higher vibrational state because of longer bonds, so this vibrational transition is infrared active

N N

= 0 Dstretch

N N

= 0 D

Because there is no changein dipole moment during the vibrational transition, N2 is infrared inactive.

Diagram of a Dual Beam Spectrometer

source of ir radiation

Io

Io

Io

reference cellsample cell

Io'wavelength selectorI

Io'()

()

photometer

I

recorder orcomputer

The intensities of the transmittedradiation Io' and I are comparedat different l. When Io' = I , there is no net absorption by the sample.100% transmission.

99

• The actual relative frequency of vibration can be predicted – Bonds with lighter atoms vibrate faster than those with heavier

atoms

100

• Triple bonds (which are stiffer and stronger) vibrate at higher frequencies than double bonds– Double bonds in turn vibrate at higher frequencies than

single bonds

• The IR spectrum of a molecule usually contains many peaks– These peaks are due to the various types of vibrations

available to each of the different bonds– Additional peaks result from overtone (harmonic) peaks

which are weaker and of lower frequency– The IR is a “fingerprint” of the molecule because of the

unique and large number of peaks seen for a particular molecule

101

102

Interpreting IR Spectra• Generally only certain peaks are interpreted in the IR

– Those peaks that are large and above 1400 cm-1 are most valuable

• Hydrocarbons• The C-H stretching regions from 2800-3300 cm-1 is characteristic

of the type of carbon the hydrogen is attached to• C-H bonds where the carbon has more s character are shorter,

stronger and stiffer and thus vibrate at higher frequency– C-H bonds at sp centers appear at 3000-3100 cm-1

– C-H bonds at sp2 centers appear at about 3080 cm-1

– C-H bonds at sp3 centers appear at about 2800-3000 cm-1

• C-C bond stretching frequencies are only useful for multiple bonds– C-C double bonds give peaks at 1620-1680 cm-1

– C-C triple bonds give peaks at 2100-2260 cm-1

– These peaks are absent in symmetrical double and triple bonds103

• Example: octane

104

• Example: 1- hexyne

105

• Alkenes• The C-H bending vibration peaks located at 600-

1000 cm-1 can be used to determine the substitution pattern of the double bond

106

• Example: 1-hexene

107

• Aromatic Compounds• The C-C bond stretching gives a set of characteristic

sharp peaks between 1450-1600 cm -1

• Example: Methyl benzene

108

octaneCH3CH2CH2CH2CH2CH2CH2CH3

1-hexeneCH3CH2CH2CH2CH=CH2

CH3CH2CH2CH2C CH

1-hexyne

110

CH3CH2CH2CH2C CH

1-hexyne

Other Functional Groups• Carbonyl Functional Groups• Generally the carbonyl group gives a strong peak

which occurs at 1630-1780 cm-1 – The exact location depends on the actual functional

group present

111

Alcohols and Phenols • The O-H stretching

absorption is very characteristic• In very dilute solutions,

hydrogen bonding is absent and there is a very sharp peak at 3590-3650 cm-1

• In more concentrated solutions, the hydroxyl groups hydrogen bond to each other and a very broad and large peak occurs at 3200-3550 cm-1

• A phenol has a hydroxyl group directly bonded to an aromatic ring

112

sec-butyl alcohol

benzyl alcohol

A Comparison of Aliphatic and Aromatic Alcohols

CH3CHCH2CH3

OH

CH2-OH

Some Characteristic Infrared Absorptions

Bond Compound Type Frequency Range, cm-1

alkanes 2850-29601350-1470

alkenes 3020-3080 675-1000

alkynes 3300

alkenes 1640-1680

alkynes 2100-2260

alcohols 3610-3640 (monomeric)3200-3600 (H-bonded)

C=O carbonyl stretchaldehydes 1690-1740ketones 1680-1750esters 1735-1750carboxylic acids 1735-1750amides 1630-1690

amines 3300-3500

R C H

C C H

C C H

C C H

C C H

R O H

R N H

Carboxylic Acids• The carbonyl peak at 1710-1780 cm-1 is very

characteristic• The presence of both carbonyl and O-H stretching

peaks is a good proof of the presence of a carboxylic acid

• Example: propanic acid

115

• Amines• Very dilute solution of 1o and 2o amines give sharp

peaks at 3300-3500 cm-1 for the N-H stretching– 1o amines give two peaks and 2o amines give one peak– 3o have no N-H bonds and do not absorb in this region

• More concentrated solutions of amines have broader peaks• Amides have amine N-H stretching peaks and a

carbonyl peak

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Documents

Carbon-carbon Covalent Bonds