Large Number
Complex Structures
Families
  Structure-Properties

Determining Structure

Determining Reaction Mechanisms

Carbon Geometry

Chains
Multiple bonds
Rings
Branching
Other Atoms

Arrangements determine physical and chemical properties.

{T#121 Cellulose}

{T#125 short segment of the RNA polymer}

{T#126 Double helix of DNA}

{T#135 Alpha-helical arrangement of a protein and pleated sheet arrangement of a protein}

Structural Theory

Systematized

  Arrangement and Structural Formulas

  Electronic Properties

Cyclohexane {structural formula, ball and stick, space filling}
 

Bonding
{NT#2 Table Electronic configurations relative orbital energies Partial periodic table}

Tendency to attain stable electron configuration in the outer shell.

s = 2
p = 8
d = 18

Shells not Orbitals

These are generalized shells and originally were considered to be spherical.  They are in fact composed of orbitals of different energy levels which have different sgeometries.  these geometries are important to chemical structure and properties.

Ionic and Covalent Bonds

Ionic bonds
These are promarily electrostatic attractions between ions.  Ions are species which have transferred electrons and have full positive or negative electric charges.  The bonds are not directed bonds, but in a lattice structure they can assume a specific geometry.  Once in solution, this geometry is lost.
 

{li  f  } {structural formula, lattice model}
 
 

  Electropositive and Electronegative Atoms.

Electropositive atoms tend to give up electrons to electronegative atoms which tend to accept electrons.

Electronegativity and elctropositivity are relative properties.  when considering these properties and their effect on two atoms, their relative electronegativity is the property which determines the polarity of the system.

Covalent Bonds

Two atoms which make up their outer stable electron shell by sharing rather than transferring electrons are said to have a covalent bond.
 

Lewis Structures

{Lewis structures needed}
h2  f2  hf  h2o  NH  ch4
 

Electrons shield positively charged nuclei from each other.
This minimizes the repulsive forces.
Simultaneously, the nuclei and electrons are mutually attracted.
 

This results in the lowest energy of the system which means it is the most stable configuration.
This is what we mean when we say that a bond is formed.

Problems 1.1, 1.2

Quantum Mechanical Description.

Wave Equations

Electron motion as a function of energy

Solutions are wave functions

Wave functions then define the atomic and molecular orbitals.
Each orbital is at a different energy level.

{NT#2 Table Electronic configurations relative orbital energies Partial periodic table}
 

Atomic Orbitals

Wave functions can be used to determine the probability of finding an electron in a given region of space.
The orbital is a region of space for which the probability of finding the electron is some arbitrary value, usually .95 (95%).

The shapes and relative orientations of these favored regions of space (orbitals) are determined by their respective energies.

The electron density (probability) varies in space and is greater in some regions than in others.

Shapes of Orbitals

{ 1s  2s  2p  px py pz}
{NT#1 The atom The 1s orbital}

{T#1 The 2s orbital the 2p orbital}

Electron Configurations.

Pauli Exclusion Principle
Only 2 paired electrons per orbital

Unpaired electrons try to be as far apart as possible.

The balance of these factors of geometry and electrostatic attraction and repulsion, determines the chemical and physical properties of molecules.

{periodic table with configurations,}
See table 1.1.

Problem 1.3

Molecular Orbitals

Atomic orbitals are considered to be centered at one nucleus.

Molecular orbitals are defined over two or more nuclei.

The resulting distribution of nuclei and electrons is the one which results in the lowest energy (most stable) configuration.
This is generally what is referred to as the “structure” of the molecule.

The resulting electron distribution (molecular orbitals) are, to some degree, localized, i.e. they occupy a specific region of space relative to the nuclei, just as atomic orbitals do.
The molecular orbitals also have certain energies and shapes.

Covalent Bonds

{T#2 Bonding overlap of two 1 s orbitals, antibonding overlap of two 1 s orbitals}

When atomic orbitals overlap to form molecular orbitals, energy is released.  This means that the molecule contains less energy than the individual atoms and is therefore more stable than the separate atoms.  If this is not the case, the molecule will not form from the atoms.

When a molecule is formed, energy is released.

This energy is called the Bond Dissociation Energy.
It is also the amount of energy required to break the bond.

When atomic orbitals overlap along a line connecting the two nuclei, the molecular orbital formed is coaxial with the original atomic orbitals.  This head-to-head overlap results in considerable overlap of orbitals and is called a sigma orbital.

{H2  F2} {structural formula, ball and stick, space filling}

Hybrid Orbitals
{T#3 sp hybridization, BeH2 is bonding using sp-hybrid orbitals}

best hybrid is directed more than s or p

more overlap

equivalent

far apart as possible

bond angle determined by maximum overlap

{T#4 Sp2 hybridization Sp3 hybridization}

{sp2 bf 3  } {structural formula, ball and stick, space filling}
 

{sp3 ch4} {structural formula, ball and stick, space filling}
 

Hybrid bonds have characteristic bond lengths, and bond angles.  This results in specific geometrical arrangements of atoms in molecules.

Directed bonds.

These directed bonds are responsible for the variety of organic molecular geometries and structures.
 

Unshared Electron Pairs
Two electrons can occupy a molecular orbital without being shared between two atoms.  This orbital has a characteristic shape and extends into space away from the nucleus.

The unshared electron pairs occupy more space than a hydrogen atom since they are not held as tightly by one nucleus.
Their presence can effect the structure and properties of the molecule.

Since they are not held very tightly, they are more available for chemical reactions.

{NH3  sp3 } {structural formula, ball and stick, space filling}
 

{H2O  sp3 } {structural formula, ball and stick, space filling}
 

Problem 1.4

Intramolecular Forces

Within a given molecule, several forces are in effect.  The balance of these forces determine the most stable (lowest energy) structure assumed by the molecule.

Repulsive Forces

The negatively charged electrons tend to stay as far apart from each other as possible.
Electrons of opposite spin also tend to stay as far apart as possible.

The positively charged nuclei also repel each other and stay as far apart as possible.
 

Attractive Forces

The positively charged nuclei attract the negatively charged electrons.
The most stable (lowest energy) orientation for these forces is to have the electrons between the nuclei.
In this case, each nucleus is attracted to the electron pair while the electrons shield the positively charged nuclei from each other to some extent.
 

Bond Dissociation Energy

This is the energy required to break a bond as well as the energy released when the bond is formed.

Bond Breaking

We will mostly consider two types of bond breaking, homolysis and heterolysis.

Homolysis

Recall that the single bond formed when two atomic orbitals overlap to form a molecular orbital, consists of an electron pair.  One way to break this bond is to allow one of the electrons to go with one of the two atoms while the other electron of the bond goes with the other atom.  Thus both atoms (or larger fragments) are separated, each possessing an unpaired electron.  If the original molecule was neutral, both fragments will be neutral.  This is homolytic bond breaking or homolytic bond cleavage.
Heterolytic

Another way to break a bond is to allow both electrons of the molecular orbital to go with either one fragment or the other.  If the original molecule was neutral, then the fragment which now has the two electrons will be negatively charged, while the other fragment will be positively charged.  This is called heterolytic bond breaking or heterolytic bond cleavage.

Study the tables of homolytic and heterolytic bond dissociation energies in the textbook and become familiar with the trends.
 

The term “Bond Energy” is used to describe the average energy of all the bonds in a molecule.  It is not the same as bond dissociation energy.

For example, starting with the molecule methane, a carbon atom with four hydrogens bonded to it, each hydrogen atom can be removed sequentially by breaking each carbon-hydrogen bond.  The four bonds are equivalent, so it does not matter which is broken first.
The bond dissociation energies for the consecutive
breaking of the four bonds are:
 

{CH3-H} 104 KCAL/Mole

{.CH2-H} 106 KCAL/Mole

{..CH-H} 106 KCAL/Mole

{...C-H} 81 KCAL/Mole

Whichever C-H bond is broken first, requires 104 KCAL/Mole.  This is characteristic of the methane molecule.  The second C-H bond requires 106 KCAL/Mole to break.  Of course, this second C-H bond is not in a methane molecule now, but in another species called a methyl radical (more on this in Chapter 2) and so the Bond Dissociation energy can be slightly different.  This is true for the breaking of the third and fourth C-H bonds.
Look at the table in the textbook and note the bond dissociation energy trends of similar bonds in different molecules.

The Bond Energy for the breaking of all four C-H bonds of methane is simply the average of the bond dissociation energies.

397 KCAL/Mole / 4 bonds = 99 KCAL/Mole

This term is useful when estimating the amount of energy available from combustion processes and other types of reactions.  We will use the bond dissociation energy for most of the work in this course.

{NT#16 Bond dissociation energies}

Bond Plarity

Because of the difference in electronegativity of different atoms, when a covalent bond is formed, the electrons of the bond may not be shared equally between the two bonded atoms.  The electrons will tend to be closer to the more electronegative atom.  This may give the more electronegative atom a greater electron density and a more negative charge than the more electropositive atom of the pair.  As a result the bond is said to be polarized, i.e. one end of the bond is more negative than the other end.  As the relative difference in electronegativities between the two atoms increases, the bond can be more polar.  This is important in determining physical and chemical properties of molecules as will be seen shortly.  It is an important concept which will be used throughout the course.

Become very familiar with the electronegativities of various atoms in the periodic table.  Understanding the structure of molecules and relative electronegativities of their constituent atoms is the most fundamental concept needed to understand a great deal of organic chemistry.

For this course, we will be mostly concerned with the relative electronegativities of some common atoms:

F > O > N, Cl > Br > C, H
{CO and CN bonds with partial charges}

{HF  H2O  NH} {structural formula, ball and stick, space filling}
 

{nt#3 Electronegativities}

Polarity of Molecules

The bond polarity is measured as a vector quantity, the dipole moment, which depends on the two net charges at the ends of the given bond and the distance of separation.

equation mu = e D}

In a molecule with many bonds, these individual dipole moments due to polar bonds can be added vectorially (considering both magnitude and direction) to yield a resultant quantity which is the molecular dipole moment.

Simply, when the center of positive charge does not coincide with the center of negative charge in a molecule, then the molecule has a dipole moment.

{+ ---> -}

Examine table 1.4 in the textbook and note the trends of the dipole moments listed.  Correlate them with what you know about the relative electronegativities of the constituent atoms.

{NH  NF3 copy from book}
Problems:  1.5, 1.6, 1.7

Structure and Physical Properties

The two simple phenomena of geometry due to atomic and molecular orbitals and the relative electronegativities of atoms, can provide a great deal of information which can help predict the physical and chemical properties of molecules.  Here we will discuss several common physical properties and see how they can be related to molecular structure.
 

Melting Point

The process of “melting” is a change of phase from solid to liquid.  In a solid, molecules are held together tightly, while in a liquid the molecules are able to move more freely.  Therefore, to melt a solid substance, some forces holding its molecules together must be overcome.  Since these forces act between molecules, they are called intermolecular forces.  What can be learned about the strength of these forces from the examination of the structure of the molecules?

Ionic Crystals

These are composed of ions in a crystal lattice.  They are held together by strong electrostatic forces.  Each positive ion is attracted by all nearby negative ions and vice versa.

Non-ionic Crystals

These are composed of neutral molecules which are held together in a crystal lattice by electrostatic forces which are weaker than those in ionic crystals.

Intermolecular Forces

The intermolecular forces responsible for the mutual attraction of molecules is electrostatic.  That is, unlike charges attract each other while like charges repel.
In looking at the kind of geometry a molecule assumes, based on a knowledge of the geometry of various types of hybridized bonds, and knowing the relative electronegativities of the constituent atoms, the dipole moment of a molecule can be determined.  Even a brief examination of a small molecule can provide sufficient information to estimate some of its properties.

Dipole-Dipole Interactions

Since a Diple, by definition, has a negatively charged end and a positively charged end, a collection of polar molecules can arrange themselves in such a way that most of the positive ends of each molecule are attracted to negative ends of other molecules.  This is a more stable situation than if similarly charged ends were positioned near each other, maximizing the repulsive forces between them.  The molecules will relax to their most stable, (lowest energy) positions.

The greater the dipole moments of the molecules, the stronger the intermolecular forces between them.

The stronger the attraction between molecules, the more energy is required to break these attractive forces.  Since this energy can be provided by heat, the melting point of a substance can be seen to increase as its dipole moment increases (generally).  See the table in the textbook.

Hydrogen Bonding

An important kind of dipole-dipole interaction is hydrogen bonding or H-bonding.  this is exactly like other dipole-dipole interactions, but one of the atoms making up the polar bond is a hydrogen atom.  The other atom of the bond is usually fluorine, nitrogen or oxygen.  The reason for specifically mentioning this hydrogen bonding interaction is that it can be very strong and make significant differences in physical and chemical properties of molecules.

The chemical literature discusses H-bonding for hydrogen bonded to other less electronegative atoms, but in this course we will concentrate on those hydrogen bonds in which H is bonded to F, N or O.
 

{HF   NH  NH H2O} {structural formula, ball and stick, space filling}
 

{NT#8 London dispersion forces}

Van der Waals Forces

A molecule consists of positively charged nuclei surrounded by clouds of electrons.  For simplicity we look at static structures, but in reality the structure is dynamic as we will see in some detail in later chapters.  For now we can consider the relative positions of the nuclei to be constant.  The electrons in the molecular orbitals are, however, subject to very rapid motions around the nuclei.  One result of this motion is that the electron distribution around the nuclei is not constant and for brief instances of time the distribution can result in a temporary (very short lived) dipole moment.  These are very small dipole moments but can have a significant effect on molecular properties.  A dipole moment in one molecule can induce an oppositely oriented dipole in another molecule.  Thus, for a brief period of time, a weak dipole-dipole attractive interaction can exist between two otherwise non-polar molecules.  This is why non-polar molecules such as the hydrogen molecule {H2 } or methane {CH4} can have a non-zero melting point or boiling point.

This interaction can occur anywhere on the surface of the molecule and as a result is stronger as the surface area of the molecule increases.  Thus, as the surface area of a molecule increases, by adding atoms for example, the number of van der Waals interactions can increase and so can the attractive forces between the molecules.  This will be considered in more detail in Chapter 3 concerning the physical properties of alkanes.

{NT#9 Boiling points of alkanes Melting points of alkanes}

This VDW force is attractive and icreases as the distance between molecules decreases.  At some distance, however, the mutual repulsion of the electrons will dominate and the net force acting between the molecules will be repulsive.  the distance of the maximum attractive force is usually defined as the van der Waals radius and can be considered to be the limit of the outer surface of a molecule.  There is much more to it than this, but for our purposes this will be a sufficient working definition.
 

Boiling Point

Ionic compounds are held together by electrostatic forces between positively and negatively charged ions.  These are strong forces and result in high boiling points.

Covalent compounds may be composed of atoms with widely different electronegativities resulting in large dipole moments which can result in significant attractive forces through dipole-dipole interactions.  These are somewhat weaker than ionic forces however and generally result in lower boiling points than for ionic solids.

Hydrogen Bonding and Boiling Point

Molecules which can form hydrogen bonds generally have higher boiling points than molecules of similar molecular weight and dipole moment which cannot form hydrogen bonds.

For example the boiling point of HF is greater than that of the more massive HCl and that of water {H2O} is greater than that of hydrogen sulfide {H2S} and the boiling point of methyl alcohol (methanol) {CH3OH} is much greater than that of methane {CH4}.
 

For organic compounds hydrogen bonding is an important property which should be recognized where it may occur in molecules.  For example, two molecules with the same molecular formula {C2H6O} can have quite different boiling points because of the presence of hydrogen bonding in one molecule and not in the other.
Ethanol (ethyl alcohol) {CH3CH2OH} has a significantly higher boiling point than dimethyl ether {CH3-O-CH3}.

This is a clear example of howthe structure of organic compounds determines their physical and chemical properties.

Example problem 1.8
Solubility

Solubility is important in the study of organic chemistry since most reactions we study occur in solution.  Just as for boiling point and melting point, solubility involves the breaking of attractive intermolecular forces.  In the case of dissolving a solid however, the attractive forces are broken, not by thermal energy, but by solvent molecules.  Consider the very high temperature needed to melt an ionic solid like NaCl.  considerable energy is required to separate the positive and negative ions.  NaCl however can be dissolved easily at room temperature in water.  The process of dissolving also separates positive and negative ions which still requires a significant amount of energy, so where does this energy come from?

When we discussed bond dissociation energy earlier, we saw that when a bond is formed, a certain amount of energy is released and that this same amount of energy is required to break that bond.  Actually, more energy is required to break the bond, but that will be discussed in a later chapter.  It is more correct to say that at least the bond dissociation energy is required to break the bond.

We will see later that in many reactions, much of the energy to break bonds comes from the formation of new bonds.

This is how a solvent can break even very strong intermolecular attractive forces.  If a solvent molecule is polar, it can form an ion-dipole interaction with positive and negative ions.  The positiive ions (cations) attract the negative end of the dipole while the negative ions (anions) attract the positive ends of the dipole.  The formation of these new attractive interactions can be considered the formation of an ion-dipole bond.  This is not a strong bond and is not like a covalent bond in which electrons are shared, but its formation does result in the release of some energy.  When a large number of these interactions occur between an ion and many solvent dipoles, the total energy released may be sufficient to break the ion-ion interactions of the solid and dissolve the compound.

Additionally, the solvent molecules, e.g. water, may have a high dielectric constant which can insulate the electrostatic effects between the ions, reducing the attractive forces between them.

Thus, the more polar the solvent (larger dipole moment), the more effective it can be in dissolving ionic solids.

Similarly, covalent solids with large dipole moments can be dissolved in polar solvents.  There are several factors competing here and they will be discussed in Chapter 7.

Non-polar solids which are held together by van der Waals forces can be dissolved in non-polar solvents.  In this case, the same kind of interactions occur between solute and solvent molecules as occur between solute molecules.

In the case of polar solvents and non-polar solutes or non-polar solvents and polar solutes, the molecules of the polar components are held together by relatively strong electrostatic forces.  The non-polar components cannot form bonds to the polar molecules (either solvent or solute) which release sufficient energy to break the stronger interactions.

{T#6 Polarity effects on solubilities}

Acids and Bases

This is another important concept which will be used in the study of organic chemistry.  Many reactions studied can be reduced to acid-base type reactions.  We consider two definitions of acidity and basicity.

The Lowry-Bronsted Concept of Acids and Bases

A Lowry-Bronsted acid is a molecule which can give up (donate) a hydrogen ion {H+) to a base.  It may be neutral or positively charged.  The fragment of the molecule remaining after the loss of the hydrogen ion is called the conjugate base of the acid.

A Lowry-Bronsted base is a molecule which can accept a hydrogen ion {H+}.  It may be neutral or negatively charged.

{H2SO4 + H2O  --> H3O+ + HSO4-}
The {H+} is removed from the sulfuric acid and goes to the water molecule.  There is in reality no free proton {H+} in the solution.  It must leave a molecule which is less able to accomodate it and go to a molecule which is more able to accomodate it.  In other words it leaves a weaker base for a stronger base.  This is an equilibrium with specific equilibrium constants, so some {H+} is always attached to both bases and constantly exchanging between them.

{NH + HCl --> NH4+ + Cl-}

{H3O+  + NaOH --> H2O + H2O + Na+}
Charge must be conserved in the equilibrium.

{NH4Cl + NaOH --> H2O + NH}

Order of Acidity

Acids and bases can be categorized according to their relative acidity and basicity (acid strenght and base strength).  These are relative scales, so acids are compared to each other to determine which is stronger.

Since a Lowry-Bronsted Acid consists of a hydrogen ion and its conjugate base, if we compare one acid to another, we are really comparing the relative strenghts of their conjugate bases.  A weaker conjugate base gives up the hydrogen ion more easily to a stronger conjugate base.  Thus, {H2SO4} consists of the {H+} and the conjugate base {HSO4-}.  Sulfuric acid can readily lose its {H+} to water {H2O} to form the new acid, hydronium ion, {H3O+}.  This shows that the tendency of the {H+} is to leave the conjugate base {HSO4+-} and go to the neutral conjugate base water, {H2O}.  In this case the {H2SO4} is a stronger acid than {H3O+} and {HSO4-} is therefore a weaker base than {H2O}.

Comparing a series of acids we can order them according to acidity or acid strenght by determining which acid will give up a proton more readily.  For example:

{H2SO4 , HCl > H3O+ > NH4+ > H2O)

Order of Basicity

Similarly, we can determine the strengths of bases, by noting which base will displace a hydrogen ion {H+} from a weaker base.  This is exactly the same as was done for acids, but now the focus is on the conjugate bases.

The order of basicity for a series of conjgate bases is exactly the opposite of the order of acidity of the corresponding acids.
hso4-, cl-  h2o  NH  oh-

If you look at the question of acidity and basicity in terms of competition for hydrogen ions between two conjugate bases, the concept should be clear.
{HSO4-......H+..H2O}

{H2O......H+..NH}

{NH......H+..OH-}

{OH-......H+..NH2-}

From the above chart the order of acidity will be:

{H2SO4 > H3O+ > NH4+ > H2O > NH}

The order of basicity of the corresponding conjugate bases is:

NH2- > OH- > NH > H2O > HSO4-}

Oxonium Ions

In the presence of a stronger acid, certain organic molecules containing oxygen can act as Lowry-Bronsted bases and accept a hydrogen ion at the oxygen atom.  This cation is called an oxonium ion and occurs frequently in many organic reactions.  It is an important intermediate in many mechanisms which will be covered in this course.

{CH3-CH2-OH + H2SO4  <--> CH3-CH2-OH2+ + HSO4-}
 

Lewis Acids and Bases

The Lewis definition of acids and bases is more general and includes the Lowry-Gronsted definition as a subset.  The Lewis definition focuses on electron pairs instead of hydrogen atoms and so can identify molecules as acids or bases whether or not they donate or accept hydrogen ions.

A Lewis acid is a molecule which can accept an electron pair from a Lewis base to form a covalent bond.  The Lewis acid must be electron deficient and can be either neutral or positively charged.

A Lewis base is a molecule which can donate an electron pair to a Lewis acid to form a covalent bond.  The Lewis base is thus electron rich (having unshared electron pairs) and can be either neutral or negatively charged.

{H+  + OH- <--> HOH}

{H3N + BF3 <--> H3N-BF3}

{F3B + O(CH2CH3)2 <--> F3B-O(CH2CH3)}

Relative Strengths of Lowry-Bronsted Acids and Bases

To judge the relative strengths of Lowry-Bronsted acids it is necessary to look at the atoms to which the hydrogen atom is bound.

In any given row of the periodic table, the more electronegative that atom is, the more readily it can accomodate a negative charge and therefore, the more easily it can donate the positively charged hydrogen ion.  In a given row, the size of the atoms are similar.

HF is a stronger acid than {H2O} largely because F is more electronegative than O.

Looking at hydrides of some first row elements we see that the order of acidity is:

{CH4 < NH < H2O < HF}

In the second row:

{H2S < HCl}

The size of the atom to which the hydrogen is bound can also indicate relative acidity.  As the size of the atom to which the hydrogen is bound increases (goes down a column in the periodic table) the molecule becomes more acidic.

HF < HCl < HBr < HI

{H2O < H2S < H2Se}

This latter case contradicts the statement on electronegativity.  One reason for the size effect is that the larger atoms can better accomodate a negative charge after the hydrogen atom is removed.  this is partly because the negative charge would be delocalized over a greater volume of space and a less concentrated charge is more stable than a case in which the same charge exists in a smaller volume of space.

Regarding Lewis acids, electron deficient atoms are stronger acids while electron rich atoms may be stronger bases.

In all cases these are just general rules and may not correctly predict relative acidity or basicity.  For example, {NH} is a stronger base than {H2O} though {H2O} has twice as many unshared electron pairs as {NH}.  This is largely due to the greater electronegativity of oxygen than nitrogen, thus it has less of a tendency to donate its electrons.

Throughout this course, situations will be encountered in which many factors combine to determine the properties of molecules.  There are many general rules which can suggest how these factors manifest themselves, but in many cases it is difficult to make a generalization.

Example problems 1.9, 1.10, 1.11, 1.12

Isomerism

 In organic chemistry, many compounds will be encountered which have identical empirical formulas, i.e. the same number and kinds of atoms.  These compounds are called isomers.  We will encounter this extensively beginning in Chapter 3.  There are many different kinds of isomerism, which means that we can classify isomers in ways which focus on some aspect of their structures.  Now you should be getting used to the different ways in which a given collection of atoms can be connected to form different compounds.  You need to remember how many bonds are possible to each atom type.  Now we are concerned with single bonds, but later will have to deal with double and triple bonds as well.  This is an important concept to understand and to be able to apply since two isomers may have very different chemical and/or physical properties.  We will be considering isomerism of several kinds throughout the course, e.g. stereoisomers, enantiomers, diastereomers, geometric isomers, conformational isomers and tautomers.

For example, two common compounds have the empirical formula {C2H6O}.  considering only single bonds, we know that carbon atoms must have a total of four bonds, hydrogen atoms have one bond and oxygen atoms have two bonds.  Two different molecules can be constructed from this collectoon of atoms:

Ethanol (ethyl alcohol) {CH3-CH2-O-H}
It has a boiling point of {78 degrees C).

Dimethyl ether  {CH3-O-CH3}
It has a boiling point of {24 degrees C}.

Why are the boiling points so different?
 

Problems at the end of Chapter 1.

1, 2, 3, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14