CHAPTER II
BONDING: A Short Treatise
Before examining what a molecular orbital is, it is good to refresh our memories with respect to the
types of bonding that can occur between atoms when molecules are formed. Basically there are two
prevalent types of bonding that occur in organic chemistry. (There are several ways that bonding can occur
but most of these ways can be classified into two very broad categories.) The first category is called “ionic
bonding”. The second category is called “covalent bonding”. Let’s examine these two broad categories
separately even though many organic compounds may display both types in the same molecule.
Ionic bonding is when an atom or molecular entity transfers its valence electrons completely to another
atom or molecular entity so that the bonding force that results is purely electrostatic (held together by
opposite charges). The major consequence of an ionic bonding process is that the “Octet Rule” is satisfied
for each atom or molecular entity involved in the transfer. Neutral atoms or molecular entities are
consequently converted into oppositely charged and stable ions. The ions are attracted to one another but
there is no restriction on any distance of separation that must apply.
Looking at ionic bonding in a simple atomic example we see that a hydrogen atom can transfer a single
valence electron to the valence shell of a chlorine atom. A positively charged hydrogen cation (called a
proton) and a negatively charged chloride anion are produced. (See chemical reaction equation below.)
H.
+
.Cl :
H
+
: Cl :
The ions that result are much more stable than their precursor neutral atoms because both ions satisfy the
“Octet Rule”. The ions are bonded together by the electrostatic force of their opposite charges. Either ion
is considered to be a stable entity (able to stand on its own). The bond that forms between the ions is one of
charge attraction leading to overall electrical neutrality. There is no requirement that the stable ions must be
close in space in terms of distance between the oppositely charged ions. Consequently the overall size of
the neutral molecule is dependent upon the environment only.
The hydrogen cation that is formed in the above example is called a “proton” in the vernacular of
organicese. A proton is a bare nucleus that has lost the single electron occupying the 1s atomic orbital from
its precursor hydrogen atom. Protons are very stable and extremely mobile entities. A proton always needs
the association of some counterion that has negative charge. In our simple example a chloride anion acts as
the counterion that has a single negative charge. The chloride anion is formed when the chlorine atom
precursor accepts a single electron from the hydrogen atom. The pair of electrons that are formed then
occupy the pz atomic orbital of the resultant chloride anion. The resultant argon electron configuration of
the chloride anion satisfies the “Octet Rule”. This latter type of reactivity is characteristic of the halogen
family in general i.e. the elemental atom’s of the halogen family readily accept a single electron to form
anions. The halide anions are very stable because of the noble gas electron configuration they exhibit which
is spherical in shape and satisfies the “Octet Rule”.
Looking at ionic bonding in a more complex molecular example we see that ammonia (NH3) can react
with carbonic acid (H2CO3) to produce two singly charged ammonium cations and one doubly charged
carbonate anion.
2 NH3
+
H2CO3

2 NH4+
+
CO3-2
The starting materials in this example have covalent bonds and are thus neutral molecular entities. The
resultant ionic products are held together by two ionic bonds. (Since the carbonate dianion has two negative
charges and each ammonium cation has only one positive charge there are two ionic bonds between the one
carbonate dianion and the two ammonium cations.) Thus the individual ions of ammonium carbonate are
covalently bonded molecular entities that are held together ionically in the product by electrostatic forces.
(We will focus on the covalent bonding of this example later.) In this example the two ionic bonds of the
product are formed by the transference of two protons. Two protons are transferred from the carbonic acid
to the two ammonia base molecules. Thus carbonic acid reacts with two ammonia base molecules to form
ammonium carbonate. While many ionic bonds are formed when one or more electrons are transferred
between atoms or molecular entities a completely analogous process can occur when the transference of
one or more protons occurs. This latter type of reactivity will be a very common occurrence in organic
chemistry.
Covalent bonding is when atoms share valence electrons with other atoms so that all the atoms that
participate in the creation of the molecular entity satisfy the “Octet Rule”. The covalent bonding process
requires that the participant atom members of the molecular entity must be close to one another in space.
Close proximity in space results from the sharing of valence electrons between atoms and therefore a
covalent bond will have a definitive bond length that is normally measured in angstroms [An angstrom (Ả)
is 10-10 meters (m) or 0.1 nanometers (nm) or 10-8 centimeters (cm)]. Covalent bonds are usually neutral in
charge but there can be charged members or even polar (charge separated) components embedded within
the molecular array. What is important about covalent bonding is that all the participant atoms be near one
another in space so that their shared electrons can confer stability upon the associated atoms of the
molecular array by satisfying the “Octet Rule”. Recognition of covalent bonds is best exemplified by
looking at real examples. Let’s examine the reaction of ammonia with carbonic acid to form ammonium
carbonate from the point of view of the covalent and ionic bonds within this entity.
Covalent bonding within the simple example of ammonia can be depicted using the hypothetical reaction
of atoms that follows: N + 3H  NH3 We see that nitrogen needs three electrons to satisfy the
“Octet Rule”. Hydrogen can contribute one electron per each atom. Ammonia must have the molecular
formula of NH3 and thus three hydrogen atoms are required to satisfy the nitrogen atom’s need to acquire
an octet of electrons. We might expect that ionic bonds could form between the nitrogen bearing a triple
negative charge and three associated protons bearing a single positive charge each. Instead ammonia is
found to have covalent bonds only. Consequently nitrogen and the attendant hydrogen atoms share
electrons so that every atom in the molecular array of ammonia has a noble gas electron configuration.
Each hydrogen atom shares two electrons with the central nitrogen atom so that each hydrogen atom can
claim a helium electron configuration. The central nitrogen atom can claim a neon electron configuration
because it shares six electrons with the three hydrogen atoms. The two electrons that are not shared by the
nitrogen atom are called non-bonding electrons or lone pair electrons or an unshared pair of electrons. The
covalent bonds in ammonia require that the associated atoms must be close to one another in space. The
covalent bonds in ammonia are 1.02 Ả in length. The N-H bonds are considered to have polar character
because each hydrogen atom bears a partial positive charge while the nitrogen atom bears a partial negative
charge.
Ammonia is a molecular entity that is well characterized. Although we consider it to be a simple
example of the utilization of the covalent bonding process it can answer some very complex questions that
are raised by this description. What does ammonia look like structurally since the nitrogen atom and the
three hydrogen atoms are required to be within covalent bonding distance of one another? Are the atoms
that make up the ammonia molecule hybridized in a particular way? What is the relative energy
requirement associated with the ammonia molecule? How does the polar character of the bonds affect the
properties of ammonia? Are the non-bonding electrons important contributors to the properties of the
molecule? As potential organic chemists we might begin to formulate many of these types of questions
once we know the structures of particular molecular entities. We will eventually discover the answers but
for this description we will confine ourselves to answering those general questions that pertain to the
covalent bonding process and its implications for structure. When we describe molecular orbital theory we
will return to answer some of these questions in greater detail.
Before answering some of these questions, however, let’s examine some of the rules that have been
developed when the ideas behind hybridization and covalent bonding are combined. Hydrogen does not
hybridize when it participates in bonding of any type but carbon, nitrogen and oxygen do. Carbon, nitrogen
and oxygen can hybridize in different covalent bonding situations but the ways are limited to one of three
possible scenarios (sp1, sp2 & sp3). As a result we will find that when carbon, nitrogen or oxygen covalently
bond to form molecular entities these elements will consistently hybridize to fit the structural situation. In
organic chemistry it is rare to find carbon, nitrogen or oxygen in an ionically bonded situation when a
covalent bond is not present as well. Consequently we can state with some confidence that carbon, nitrogen
and oxygen will hybridize whenever organic compounds are formed. The only decision we have to make
then is to decide which hybridization scheme to utilize in specific and particular structural situations. Most
organic chemists agree that the two empirical rules stated below can be applied in many different situations
although exceptions can arise when unusual circumstances arise.
Rule A: Neutral carbon will have four covalent bonds around each individual carbon atom while a
charge neutral nitrogen atom will have three and a charge neutral oxygen atom will have two covalent
bonds.
Rule B: When a neutral carbon, nitrogen or oxygen atom is the central atom in a covalent bonding
situation the number of other atoms that are arrayed around that central atom will dictate the hybridization
scheme displayed by that central atom according to the particular situation.
The variety of hybridization possibilities that can arise in organic chemistry for the three major elements
of central interest are shown in the Table as a function of the number of atoms arrayed around that central
element. As can be seen carbon and nitrogen can exhibit three possibilities while oxygen can display two
possibilities. These various possibilities will be utilized extensively in future descriptions utilizing
organicese.
CHARGE NEUTRAL
CENTRAL ATOM
# of ATOMS ARRAYED
AROUND CENTRAL ATOM
HYBRIDIZATION SCHEME
DISPLAYED by CENTRAL
ATOM
CARBON
NITROGEN
OXYGEN
4
3
2
sp3
sp3
sp3
CARBON
NITROGEN
OXYGEN
3
2
1
sp2
sp2
sp2
CARBON
NITROGEN
2
1
sp
sp
Two issues to be noted with respect to the above Rules A and B. (1) Rule A implies that in the majority
of organic compounds the number of covalent bonds associated with any of these central atoms is invariant
when the array around that atom is neutral in charge. Consequently for neutral carbon the number of
covalent bonds is equal to four. For neutral nitrogen the number is three and for neutral oxygen the number
is two. (2) Rule B implies that there are only three possible types of covalent bonds to be found in organic
compounds. Those that have two electrons shared between two bonded atoms are called single bonds.
Those that have four electrons shared between two bonded atoms are called double bonds. Those that have
six electrons shared between two bonded atoms are called triple bonds.
Let’s return to the simple example of ammonia briefly. We can speculate about ideas that encompass
some of the properties of ammonia that might result from the structure. The nitrogen in NH3 is probably
hybridized sp3 and thus the ammonia molecule is probably tetrahedral. The hydrogen atoms are probably
attached to the central nitrogen by single covalent bonds utilizing two shared electron pairs each. The
hydrogen atoms probably occupy three of the four corners of the tetrahedral array that surround the central
nitrogen. The remaining corner of this tetrahedral array is most probably occupied by the non-bonding
electrons of the nitrogen atom. The angles made by any two hydrogen atoms with the central nitrogen
should be very close to 109.5 degrees and in fact they are 108 degrees in ammonia. The entire ammonia
molecule is probably a polar entity with some partial charge separation because the plane that contains the
three hydrogen atoms, or proton like entities, has a partial positive charge while a point within the unshared
electrons of the nitrogen atom, or the non-bonding electron pair, has at least a partial negative charge. In
fact the ammonia molecule carries a permanent dipole moment of 1.5 Deybe. We could continue to
speculate about the properties of ammonia, or even about the reactivity of ammonia as a base, but we will
leave such speculation for the future. For now we will move onto the more complex example of carbonic
acid. Carbonic acid has carbon as the central atom. [Note: Polarity in covalent bonds and molecules can
arise for two reasons: (1) When there are electronegativity differences between any two atoms involved in
covalent bonding; or (2) When there are protons or unshared pairs of electrons present in the molecular
entity. Polarity can be measured in a molecule through its dipole moment which is equal mathematically to
the amount of charge separation times the distance between the charge separated entities. It is to be realized
however that the environment surrounding any charge separated entity can have an effect on the magnitude
of a molecule's dipole or its measured dipole moment.]
Carbonic acid is made whenever CO2 dissolves in H2O thus it has a molecular formula of H2CO3. We will
see that in most organic compounds carbon is the dominant atom for determining structure. This means that
carbon will usually dictate covalency in organic chemistry while the hybridization process will control the
number and type of atoms that surround any particular carbon atom. Carbonic acid contains a single carbon
atom that is surrounded in some way by three oxygen atoms and two hydrogen atoms. On the surface of the
earth less energy is required to associate or bond hydrogen to oxygen than is required to bond hydrogen to
carbon. The prime evidence for this dictum is the amount of water that is found on the surface of the earth.
Carbonic acid is also a good example of this dictum and thus we can assume that both hydrogen atoms are
bonded to oxygen by a covalent bond. (This assumption is not far fetched because H2CO3 is formed from
CO2 + H2O.) Consequently the following sequence might apply to the formation of carbonic acid from its
elements.
First 2H + 2O =
Second 2OH + C
Third HO-C-OH + O
2OH
= HO-C-OH
=
HO-C(=O)-OH
There are, of course, other ways to accomplish the task of determining the bonding structure of an organic
compound. Another way is to apply the chemistry we already know is as follows:
First 2O +
C
=
O=C=O
Second H-O-H + O=C=O = H2O-CO2
Third H2O-CO2
= HO-C(=O)-OH
This latter example illustrates a way for rationalizing the bonding that results when organic carbon forms
molecular entities, such as carbonic acid, in the oxygen and water rich environment at the surface of the
earth. Eventually we will have to balance this type of illustration with the realization that most organic
compounds do indeed have numerous single covalent bonds between carbon and hydrogen. (As we will see
later the majority of the organic molecules of life are formed in environments that are depleted in oxygen
and nitrogen and thus carbon to hydrogen covalent bonds are easily formed as long as there is sufficient
energy available to accomplish the task.)
If the above description of the covalency of carbonic acid is correct, which it is, then we can speculate a
little about the structure and properties of carbonic acid. The carbon atom in carbonic acid is probably sp2
hybridized and thus carbonic acid is trigonal planar around the carbon atom. The three oxygen atoms can
not be differentiated from one another because they occupy the corners of an equilateral triangle around the
carbon atom. Thus the oxygen atoms make 120 degree angles with the carbon atom. At any one time two of
the oxygen atoms are singly bonded to the carbon atom while the third oxygen atom is doubly bonded to
the carbon atom. (The parenthetical oxygen in the formula above represents this oxygen to carbon double
bond.) Consequently the two hydrogen atoms are probably bonded to the singly bonded oxygen atoms. If
the two hydrogen atoms are thought of as protons then these two protons can move with great mobility
around from oxygen to oxygen atom making them equivalent. Since the name carbonic acid implies that
H2CO3 is an acid, that designation must be applied to these two mobile protons. A further implication of
this speculation is that the protons are ionically bonded to a carbonate dianion and therefore carbonic acid
is an example of a molecular entity that has both ionic and covalent bonds present. If ionic and covalent
bonds are present then the entire molecule probably has polar characteristics as well. The properties we
have speculated upon cry out for some additional explanation. We will indeed do that in following chapters
after we have a good understanding of MO theory. Before doing that, however, it might be a good idea to
review some of the rules and concepts that have been presented above and to also add in some additional
ideas that can be drawn from the information that has been presented thus far.
RULES AND CONCEPTS THAT ARE USEFUL IN ORGANICESE
To this juncture we have given a broad description of atoms, electron orbitals and molecular structures
seen through the eyes of chemists with some focus on organicese. We must remember, however, that much
of what we have described is a hypothetical rationalization for what has been determined empirically and
observed by chemists over the years. As a consequence these descriptions are easily manipulated to create
rules that allow chemists to account for and keep track of the valence electrons in atoms and molecules
having known structures. Before we describe a much more realistic and useful theory of molecular electron
distribution and energetics let us list some of the useful rules and concepts of organicese. We will routinely
utilize these rules and concepts in the future and therefore it is helpful to keep them in mind as we go
forward.
The Octet Rule: Atoms that participate in molecule building and bonding will strive to satisfy this rule.
Acquiring a noble gas electron configuration in the electron valence shell is the goal of all participating
atoms. This rule will apply in most, but not all, stable organic compounds whether covalent and/or ionic
bonding is involved.
Rules of Atomic Oxidation States: In most organic compounds the oxidation states for hydrogen and
oxygen are set at +1 and –2 respectively. The oxidation states for participating carbon are calculated taking
into account whether the molecular entity has a charge or is neutral. Nitrogen can have a variety of
oxidation states but the most common is –3 when it is bonded covalently to carbon &/or hydrogen. When
oxygen atoms are covalently bonded to nitrogen, the nitrogen atom will take on a variety of oxidation states
depending upon the number of oxygen atoms that are participants. Halogen atoms normally have an
oxidation state of –1 in covalent organic compounds. Sulfur, phosphorus and silicon mimic their respective
family members O, N and C in most circumstances.
Definition of Oxidation & Reduction States: The oxidation state for an element in the pure atomic form is
assumed to be zero. Thus the definitions of oxidation and reduction require that electrons are subtracted
from or added to the compliment of electrons that identify that particular atomic element. Consequently
oxidation means that electrons have been taken away from an atom, and thus the atomic entity has become
more positive in character, while reduction means that electrons have been given to an atom, and thus the
atomic entity has become more negative in character. Hydrogen normally has an oxidation state of +1
because in most compounds the hydrogen atom will tend to give up one electron to the atom with which it
bonds. Likewise, oxygen normally has an oxidation state of -2 because oxygen is very electronegative and
will tend to take two electrons from the atom with which it bonds. In these instances hydrogen is said to be
oxidized while oxygen is said to be reduced from the elemental state.
The Concept of Electronegativity: Electronegativity is a concept that allows chemists to easily determine
how electrons “flow” within a covalently bonded molecule. The concept proffers that within a specific
covalent bond electrons will tend to move easily or “flow” towards the most electronegative atom and away
from the most electropositive atom. The most electronegative elements are on the right hand side of the
Periodic Table. The most electropositive elements are on the left-hand side of the Periodic Table. (The
noble gases are considered to be electroneutral.) Fluorine is the most electronegative element of all. Cesium
is the most electropositive element. A fluorine atom is defined as having a numerical value of 4.0
electronegativity units and the greatest ability to cause electrons to “flow” towards it. Fluorine will have an
oxidation state of –1 in the vast majority of those situations where it is part of a covalent compound. Other
atoms of importance to organic chemistry have the following numerical values in terms of electronegativity
units: Oxygen = 3.5; Nitrogen = 3.0; Carbon = 2.2; Hydrogen = 2.0; Sulfur = 3.0; Phosphorus = 2.8;
Chlorine = 3.5; Bromine = 3.0; Iodine = 2.7; Silicon =1.8. Oxygen has a normal oxidation state of –2
mainly because of its high electronegativity value. Hydrogen has a normal oxidation state of +1 in organic
chemistry mainly because it is usually the most electropositive atom in organic compounds. Carbon is
considered to have an intermediate electronegativity value. Consequently carbon can take on a variety of
oxidation states ranging from -4 to +4 depending upon the number and types of atoms to which it is
bonded. Nitrogen can also take on a variety of oxidation states ranging from -3 to +5 but most normal
nitrogen containing organic compounds display an oxidation state of -3 because of its high
electronegativity value.
The Concept of Partial Charges and Polarity: In a covalently bonded situation electronegativity
differences between atomic partners will dictate whether one atomic partner will have more or less electron
density near it than the other atomic partner. In organic chemistry covalent bonds are considered as nonpolar when the bonded atoms exhibit little or no electronegativity differences. Consequently non-polar
bonds have little or no electron density differences between the associated partners and exhibit no partial
charges on either atomic partner. For instance, a carbon to carbon bond will exhibit no partial charge
character on either atom of the bond. In a covalent carbon to carbon bond both atomic partners in that bond
have equal electron density distributions because the electronegativity difference is zero. A carbon to
hydrogen covalent bond has only the smallest of partial charge character and can normally be thought of as
essentially uncharged or non-polar because the electronegativity difference is very small and equal to 0.2.
[Note: We will see that a carbon to hydrogen covalent bond can be polar under very special circumstances
and in these special situations the hydrogen atom will be partially positive and the carbon atom will be
partially negative.] Conversely, the following types of covalent bonds are considered polar in most
situations and thus there is some partial charge separation associated with the covalent bond. A carbon to
oxygen or a carbon to nitrogen covalent bond will always be considered as polar with the most
electronegative atom bearing a partial negative charge and with the partner carbon atom bearing a partial
positive charge. (Partial charges are denoted by the greek symbol delta (δ) followed by a superscript plus or
minus.) The oxygen to hydrogen or the nitrogen to hydrogen covalent bond is also considered to be polar
with the hydrogen bearing a partial positive charge (δ+) while the electronegative heteroatom bears a partial
negative charge (δ-). (Atoms other than carbon and hydrogen are called heteroatoms in organic chemistry.)
The Concept of Formal Charge: When an atom in a molecule has either an excess of electrons, over and
above its atomic number, or a deficiency of electrons, under and below its atomic number, then that atom
will be considered to possess a formal charge. For atoms that participate in an ionic bond it is fairly easy to
determine the formal charge by simply counting electrons and determining the number of electrons an atom
has relative to its atomic number. For atoms that participate in a covalent bond it is more difficult. Most
organic chemists rely on memory and experience to determine the formal charge on a particular atom in an
organic molecule rather than using a mathematical equation. The normal situations in terms of formal
charge that occur for the most common elements of organic chemistry are shown in the Table. We will find
this Table to be useful and it is probably a good idea to remember it so that it is readily available in your
minds eye.
ELEMENT
ATOMIC FORM
# of COVALENT BONDS
FORMAL CHARGE
HYDROGEN
Proton
Hydrogen atom
Hydride ion
0
1
0
+1
0
-1
CARBON
Carbocation
Carbon atom
Carbanion
3
4
3
+1
0
-1
NITROGEN
Nitrogen cation
Nitrogen atom
Nitride anion
4
3
2
+1
0
-1
OXYGEN
Oxygen cation
Oxygen atom
Oxide anion
3
2
1
+1
0
-1
The culmination of these rules and concepts is that organicese specifies ways for drawing organic structures
which utilize symbols to represent electrons, atoms, and the movement of each when certain circumstances
arise. Organic chemists are especially interested in representing electron distributions among different types
of bonds within molecular entities. The symbols utilized can be defined in a comprehensible way and we
will attempt to follow these definitions in most circumstances. (1) A dot will represent one electron that is
involved in an ionic bond or non-bonding circumstance. Two dots together will imply a pair of electrons
that obey the Pauli Exclusion Principle. (2) A straight line between two atoms will represent a pair of
electrons that hold two atoms together in a single covalent bond. Two parallel straight lines between two
atoms will represent a double covalent bond and three parallel straight lines between two atoms will
represent a triple covalent bond. (3) When we need to represent a covalently bonded molecular array in
three dimensions (3D) we will use a wedge to represent a two electron single covalent bond that comes out
of the page towards us. We will use a hatched wedge to represent a two electron single covalent bond that
goes behind the page and away from our eye. The normal two electron single covalent bond or straight line
will be as in (2) but in 3D drawings it will represent a bond in the plane of the page and perpendicular to
our line of sight. (4) Formal charges, when necessary to express, will be represented by a circled plus or
minus sign as will ionic charges. A plus or minus sign that is not circled can denote many different things
depending upon circumstances. A partial charge will be represented by a superscript plus or minus sign
preceded by the Greek letter delta (δ). (5)Anticipating what will be represented shortly we can admit that
arrows will play an important role in the symbolic language we have labeled organicese. (a) A straight
horizontal arrow will represent a reaction that begins at the tail of the arrow and proceeds to the head of the
arrow. Short vertical arrows will usually represent electron spin in energy diagrams. (b) A set of arrows that
go in opposite directions between reactants and products will represent a dynamic equilibrium reaction. (c)
A double headed arrow connecting different electronic representations of the same molecular entity will
represent resonance where electron flow within equivalent or nearly equivalent structures occurs and
represents structures having the extreme electron configurations that are theoretically possible within a
single molecular entity. (d) Curved arrows play a special role in organicese in that they represent the
movement of atoms, groups of atoms or electrons from one place in a molecular, or reacting, system to
another place in the system. [Note: Sometimes curved arrows are also utilized to point out a particular
phenomenon or other concept of interest that is illustrative. In these noted situations the curved arrow will
not be part of a molecular entity and will have the word “See” encompassed by the curved arrow.]
SOME EXAMPLES OF THESE RULES AND CONCEPTS
Let us look at some of the reactions that can produce a few of the molecular entities we have discussed
previously. This will allow us to begin to see how organic chemists might represent these structural and
reactive forms using drawings and symbols that have been or will be described in the forthcoming sections
of this manuscript. We will attempt to use some of the concepts and symbolism defined above to help
clarify these reaction examples. The first three example reactions illustrate the concepts where ionic
bonding, covalent bonding and multiple bonding are respectively pictured in the product. The last three
examples illustrate electron or atomic movement during the course of a reaction. The curved arrows in
these latter examples attempt to represent the flow of electron density that occurs during the course of a
reaction. The reactants are always shown at the tail of the reaction arrow while the product is shown at the
head of the arrow. Resonance is illustrated in the latter two examples. [Note: Non-bonding electron pairs
that are not deemed important for a reaction to move forward are rarely displayed in pictorial organicese
although they are understood to be present if not directly involved.] {An additional note: The concepts of
partial bonds and polarity are not illustrated in any of these examples although both phenomena are
probably of some importance in all six of these reactions depending upon environmental conditions.}
Formation of Hydrochloric acid from its elements
+
H2
2 Cl
+
2 H
Cl2
Ionically bonded products do not need to
symbolically show all the valence
electrons
See
Formation of Ammonia from its elements
N2
2
3H2
+
N
H
H
H
See
In covalently bonded molecules the non-bonding
electrons should be symbolically shown
Formation of Carbon Dioxide from its elements
Cn
+
O2
n
O
Multiple covalent bonds are represented by
multiple straight lines as in carbon dioxide
C
See
O
Reaction of HCl with NH3 to form Ammonium Chloride
H
H H
H
N
H
+
Cl
N
H
Cl
H
H
The lone pair of electrons on nitrogen in ammonia attacks the proton of HCl
to create the ionically bonded ammonium chloride.
Reaction of H2O with CO2 to produce Carbonic Acid
O
O
H
Attack
O
C
C
H
H
O
O
O
H
Resonance
O
O
C
H
C
H
O
Shift of Proton
O
H
O
O
H
Note the shift of the proton from one location in this molecular entity.
This is an example of a tautomeric shift which will be defined in a later chapter.
Reaction of NH3 with H2CO3 to produce Ammonium Carbonate
O
H
H
N
+
C
H
H
O
H
+
N
O
H
H
H
O
H
N
H
H
H
C
+
H
O
H
+
N
O
Resonance
O
O
O
C
C
O
H
H
O
O