Chapter 3: Organic Compounds: Alkanes and Cycloalkanes

Functional Groups

Functional groups are atoms or small groups of atoms (two to four) that exhibit a characteristic reactivity when treated with certain reagents. A particular functional group will almost always display its characteristic chemical behavior when it is present in a compound. Because of their importance in understanding organic chemistry, functional groups have characteristic names that often carry over in the naming of individual compounds incorporating specific groups. In the following table the atoms of each functional group are colored red and the characteristic IUPAC nomenclature suffix that denotes some (but not all) functional groups is also colored.

Functional Group Tables

Exclusively Carbon Functional Groups

Group Formula	Class Name	Specific Example	IUPAC Name	Common Name
	Alkene	H2C=CH2	Ethene	Ethylene
	Alkyne	HC@CH	Ethyne	Acetylene
	Arene	C6H6	Benzene

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Functional Groups with Single Bonds to Heteroatoms

Group Formula	Class Name	Specific Example	IUPAC Name	Common Name
	Halide	H3C-I	Iodomethane	Methyl iodide
	Alcohol	CH3CH2OH	Ethanol	Ethyl alcohol
	Ether	CH3CH2OCH2CH3	Diethyl ether	Ether
	Amine	H3C-NH2	Aminomethane	Methylamine
	Nitro Compound	H3C-NO2	Nitromethane
	Thiol	H3C-SH	Methanethiol	Methyl mercaptan
	Sulfide	H3C-S-CH3	Dimethyl sulfide

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Functional Groups with Multiple Bonds to Heteroatoms

Group Formula	Class Name	Specific Example	IUPAC Name	Common Name
	Nitrile	H3C-CN	Ethanenitrile	Acetonitrile
	Aldehyde	H3CCHO	Ethanal	Acetaldehyde
	Ketone	H3CCOCH3	Propanone	Acetone
	Carboxylic Acid	H3CCO2H	Ethanoic Acid	Acetic acid
	Ester	H3CCO2CH2CH3	Ethyl ethanoate	Ethyl acetate
	Acid Halide	H3CCOCl	Ethanoyl chloride	Acetyl chloride
	Amide	H3CCON(CH3)2	N,N-Dimethylethanamide	N,N-Dimethylacetamide
	Acid Anhydride	(H3CCO)2O	Ethanoic anhydride	Acetic anhydride

Alkanes

Methane CH₄

Ethane CH₃CH₃

Propane CH₃CH₂CH₃

Butane CH₃CH₂CH₂CH₃

Pentane CH₃CH₂CH₂CH₂CH₃

Hexane CH₃CH₂CH₂CH₂CH₂CH₃

Heptane CH₃CH₂CH₂CH₂CH₂CH₂CH₃

Octane CH₃CH₂CH₂CH₂CH₂CH₂CH₂CH₃

Nonane CH₃CH₂CH₂CH₂CH₂CH₂CH₂CH₂CH₃

Decane CH₃(CH₂)₈CH₃

Isomers:

Not all compounds with the same formula are neo (straight chain.) Both the above are C₄H₁₀.

Alkane Substituent Groups

Methyl CH₃-

Ethyl CH₃CH₂-

Propyl CH₃CH₂CH₂ -

Butyl CH₃CH₂CH₂CH₂ -

Pentyl CH₃CH₂CH₂CH₂CH₂ -

Hexyl CH₃CH₂CH₂CH₂CH₂CH₂ -

Heptyl CH₃CH₂CH₂CH₂CH₂CH₂CH₂ –

Octyl CH₃CH₂CH₂CH₂CH₂CH₂CH₂CH₂ –

Nonane CH₃CH₂CH₂CH₂CH₂CH₂CH₂CH₂CH₂ –

Decyl CH₃(CH₂)₈CH₂ –

Degree of substitution

Primary carbon - a terminal carbon

Secondary carbon – with two attached carbons

Tertiary carbon – with three attached carbons

Quaternary carbon – with four attached carbons

Nomenclature

1 Find longest continuous carbon chain with greatest number of branch points

2 Number carbons in chain giving lowest number to branch points

3 Identify and number each substituent

4 Write name as single word

5 Name complex substituents as if they were a compound with –yl ending

Name the following:

Properties of Alkanes

For general purposes it is useful to consider temperature to be a measure of the kinetic energy of all the atoms and molecules in a given system. As temperature is increased, there is a corresponding increase in the vigor of translational and rotation motions of all molecules, as well as the vibrations of atoms and groups of atoms within molecules. Experience shows that many compounds exist normally as liquids and solids; and that even low-density gases, such as hydrogen and helium, can be liquefied at sufficiently low temperature and high pressure. A clear conclusion to be drawn from this fact is that intermolecular attractive forces vary considerably, and that the boiling point of a compound is a measure of the strength of these forces. Thus, in order to break the intermolecular attractions that hold the molecules of a compound in the condensed liquid state, it is necessary to increase their kinetic energy by raising the sample temperature to the characteristic boiling point of the compound.

The following table illustrates some of the factors that influence the strength of intermolecular attractions. The formula of each entry is followed by its formula weight in parentheses and the boiling point in degrees Celsius. First there is molecular size. Large molecules have more electrons and nuclei that create van der Waals attractive forces, so their compounds usually have higher boiling points than similar compounds made up of smaller molecules. It is very important to apply this rule only to like compounds. The examples given in the first two rows are similar in that the molecules or atoms are spherical in shape and do not have permanent dipoles. Molecular shape is also important, as the second group of compounds illustrate. The upper row consists of roughly spherical molecules, whereas the isomers in the lower row have cylindrical or linear shaped molecules. The attractive forces between the latter group are generally greater. Finally, permanent molecular dipoles generated by polar covalent bonds result in even greater attractive forces between molecules, provided they have the mobility to line up in appropriate orientations. The last entries in the table compare non-polar hydrocarbons with equal-sized compounds having polar bonds to oxygen and nitrogen. Halogens also form polar bonds to carbon, but they also increase the molecular mass, making it difficult to distinguish among these factors.

Boiling Points (°C) of Selected Elements and Compounds
Increasing Molecular Size	Ar (40) -186	Xe (131) -109
	CH4 (16) -161	(CH3)4C (72) 9.5	(CH3)4Si (88) 27	CCl4 (154) 77
Molecular Shape
Spherical:	(CH3)4C (72) 9.5	(CH3)2CCl2 (113) 69	(CH3)3CC(CH3)3 (114) 106
Linear:	CH3(CH2)3CH3 (72) 36	Cl(CH2)3Cl (113) 121	CH3(CH2)6CH3 (114) 126
Molecular Polarity
Non-polar:	H2C=CH2 (28) -104	CH3CH=CH2 (42) -48	(CH3)3CH (58) -10
Polar:	H2C=O (30) -21	CH3CH=O (44) 20	(CH3)3N (59) 3.5	(CH3)2C=O (58) 56
	HC@N (27) 26	CH3C@N (41) 82	(CH2)3O (58) 50	CH3NO2 (61) 101

The melting points of crystalline solids cannot be categorized in as simple a fashion as boiling points. The distance between molecules in a crystal lattice is small and regular, with intermolecular forces serving to constrain the motion of the molecules more severely than in the liquid state. Molecular size is important, but shape is also critical, since individual molecules need to fit together cooperatively for the attractive lattice forces to be large. Spherically shaped molecules generally have relatively high melting points, which in some cases approach the boiling point. This reflects the fact that spheres can pack together more closely than other shapes. This structure or shape sensitivity is one of the reasons that melting points are widely used to identify specific compounds. The data in the following table serves to illustrate these points.

Compound	Formula	Boiling Point	Melting Point
pentane	CH3(CH2)3CH3	36 °C	-130 °C
hexane	CH3(CH2)4CH3	69 °C	-95 °C
heptane	CH3(CH2)5CH3	98 °C	-91 °C
octane	CH3(CH2)6CH3	126 °C	-57 °C
nonane	CH3(CH2)7CH3	151 °C	-54 °C
decane	CH3(CH2)8CH3	174 °C	-30 °C
tetramethylbutane	(CH3)3C-C(CH3)3	106 °C	+100 °C

Notice that the boiling points of the unbranched alkanes (pentane through decane) increase rather smoothly with molecular weight, but the melting points of the even-carbon chains increase more than those of the odd-carbon chains. Even-membered chains pack together in a uniform fashion more compactly than do odd-membered chains. The last compound, an isomer of octane, is nearly spherical and has an exceptionally high melting point (only 6 ° below the boiling point).

Hydrogen Bonding

Hydrogen Bonding

The most powerful intermolecular force influencing neutral (uncharged) molecules is the hydrogen bond. If we compare the boiling points of methane (CH₄) -161 °C, ammonia (NH₃) -33 °C, water (H₂O) 100 °C and hydrogen fluoride (HF) 19 °C, we see a greater variation for these similar sized molecules than expected from the data presented above for polar compounds. This is shown graphically in the following chart. Most of the simple hydrides of group IV, V, VI & VII elements display the expected rise in boiling point with molecular mass, but the hydrides of the most electronegative elements (nitrogen, oxygen and fluorine) have abnormally high boiling points for their mass.

The exceptionally strong dipole-dipole attractions that cause this behavior are called the hydrogen bond. Hydrogen forms polar covalent bonds to more electronegative atoms such as oxygen, and because a hydrogen atom is quite small, the positive end of the bond dipole (the hydrogen) can approach neighboring nucleophilic or basic sites more closely than can other polar bonds. Coulombic forces are inversely proportional to the sixth power of the distance between dipoles, making these interactions relatively strong, although they are still weak (ca. 4 to 5 kcal per mole) compared with most covalent bonds. The unique properties of water are largely due to the strong hydrogen bonding that occurs between its molecules. In the following diagram the hydrogen bonds are depicted as magenta dashed lines.

The molecule providing a polar hydrogen for a hydrogen bond is called a donor. The molecule that provides the electron rich site to which the hydrogen is attracted is called an acceptor. Water and alcohols may serve as both donors and acceptors, whereas ethers, aldehydes, ketones and esters can function only as acceptors. Similarly, primary and secondary amines are both donors and acceptors, but tertiary amines function only as acceptors. Once you are able to recognize compounds that can exhibit intermolecular hydrogen bonding, the relatively high boiling points they exhibit become understandable. The data in the following table serves to illustrate this point.

Compound	Formula	Mol. Wt.	Boiling Point	Melting Point
dimethyl ether	CH3OCH3	46	-24 °C	-138 °C
ethanol	CH3CH2OH	46	78 °C	-130 °C
propanol	CH3(CH2)2OH	60	98 °C	-127 °C
diethyl ether	(CH3CH2)2O	74	34 °C	-116 °C
propyl amine	CH3(CH2)2NH2	59	48 °C	-83 °C
methylaminoethane	CH3CH2NHCH3	59	37 °C
trimethylamine	(CH3)3N	59	3 °C	-117 °C
ethylene glycol	HOCH2CH2OH	62	197 °C	-13 °C
acetic acid	CH3CO2H	60	118 °C	17 °C
ethylene diamine	H2NCH2CH2NH2	60	118 °C	8.5 °C

Alcohols boil considerably higher than comparably sized ethers (first two entries), and isomeric 1°, 2° & 3°-amines, respectively, show decreasing boiling points, with the two hydrogen bonding isomers being substantially higher boiling than the 3°-amine (entries 5 to 7). Also, O-H---O hydrogen bonds are clearly stronger than N-H---N hydrogen bonds, as we see by comparing propanol with the amines. As expected, the presence of two hydrogen bonding functions in a compound raises the boiling point even further. Acetic acid (the ninth entry) is an interesting case. A dimeric species, shown on the right, held together by two hydrogen bonds is a major component of the liquid state. If this is an accurate representation of the composition of this compound then we would expect its boiling point to be equivalent to that of a C₄H₈O₄ compound (formula weight = 120). A suitable approximation of such a compound is found in tetramethoxymethane, (CH₃O)₄C, which is actually a bit larger (formula weight = 136) and has a boiling point of 114 °C. Thus, the dimeric hydrogen bonded structure appears to be a good representation of acetic acid in the condensed state.

A related principle is worth noting at this point. Although the hydrogen bond is relatively weak (ca. 4 to 5 kcal per mole), when several such bonds exist the resulting structure can be quite robust. The hydrogen bonds between cellulose fibers confer great strength to wood and related materials.

Water Solubility

Water Solubility

Water has been referred to as the "universal solvent", and its widespread distribution on this planet and essential role in life make it the benchmark for discussions of solubility. Water dissolves many ionic salts thanks to its high dielectric constant and ability to solvate ions. The former reduces the attraction between oppositely charged ions and the latter stabilizes the ions by binding to them and delocalizing charge density. Many organic compounds, especially alkanes and other hydrocarbons, are nearly insoluble in water. Organic compounds that are water soluble, such as most of those listed in the above table, generally have hydrogen bond acceptor and donor groups. The least soluble of the listed compounds is diethyl ether, which can serve only as a hydrogen bond acceptor and is 75% hydrocarbon in nature. Even so, diethyl ether is about two hundred times more soluble in water than is pentane.

The chief characteristic of water that influences these solubilities is the extensive hydrogen bonded association of its molecules with each other. This hydrogen bonded network is stabilized by the sum of all the hydrogen bond energies, and if nonpolar molecules such as hexane were inserted into the network they would destroy local structure without contributing any hydrogen bonds of their own. Of course, hexane molecules experience significant van der Waals attraction to neighboring molecules, but these attractive forces are much weaker than the hydrogen bond. Consequently, when hexane or other nonpolar compounds are mixed with water, the strong association forces of the water network exclude the nonpolar molecules, which must then exist in a separate phase. This is shown in the following illustration, and since hexane is less dense than water, the hexane phase floats on the water phase.

It is important to remember this tendency of water to exclude nonpolar molecules and groups, since it is a factor in the structure and behavior of many complex molecular systems. A common nomenclature used to describe molecules and regions within molecules is hydrophilic for polar, hydrogen bonding moieties and hydrophobic for nonpolar species.

Cycloalkanes

Cycloalkanes are ring structures of carbon. The nomenclature is similar to the linear alkanes:

1 Parent is the cyclic structure

2 Number carbons in cyclic structure giving lowest number to substituents

3 Identify and number each substituent

4 Write name as single word

Name complex substituents as it they were a compound with –yl ending

Name the following

There is a consequence for the rigidity of these cycloalkanes and that is the possibility of cis-trans isomerism when substituent groups are attached.