The Chemistry and Physics of Drugs Used in Anesthesia
of certain physical properties of drug molecules. This concept is important for
the understanding of the physical parameters that affect the activity of drugs,
including anesthetics, and their potential interactions with other prescribed
medicines. General knowledge of particular functional groups also makes it
easier to remember certain characteristics of particular drug classes. Later
chapters in the text will describe the functional groups and drug properties
combined into drug molecules that have specific biological activities.
Structure and Bonding
Several types of bonds can occur between 2 atoms. The different bonds vary
in strength and length. Also, not all types of atoms can form all types of bonds.
In this text, our discussion on the types of bonds will cover most bonding forms
between 2 molecules. Chapter 5 describes noncovalent interactions between
molecules, including hydrogen bonds, ionic bonds, and van der Waals forces.
The present chapter focuses on covalent bonds.
Covalent Bonds
A covalent bond is a shared pair of electrons between 2 atoms. These electrons
can originate from 1 atom, or 1 electron can originate from each of the 2 atoms.
The 2 electrons in the bond are attracted to both atomic nuclei and are shared
between the 2 atoms. Two atoms share the electrons because atoms (other
than hydrogen and helium) are most stable when surrounded by 8 electrons
(an octet), which means that an atom with a full octet of electrons has lower
energy (is more stable) than one without a full octet. The electrons surrounding
the nucleus of an atom are contained in orbitals. An orbital is the space around
the nucleus of an atom in which an electron orbits. The shape of an orbital is
determined by the mathematical probability of the location of that electron
around the nucleus. Each orbital contains a maximum of 2 electrons. When 2
electrons are in an orbital, the orbital is considered full, and the electrons are
“paired.” A covalent bond can also be thought of as the overlap of orbitals
surrounding the atomic nuclei.
It is important to understand the following points:
1. Different atoms can support different numbers of covalent bonds.
2. More than 1 covalent bond can be formed between 2 nonhydrogen atoms.
The diversity of atoms and multiplicity of bonds give the tremendous variety
of structures available in organic chemistry.
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The Basics of Organic Chemistry
Chapter 2
Covalent Bonds in Different Atoms
The atomic structure of an individual atom determines the number of covalent
bonds it can form. The number of covalent bonds formed is equal to the
number of valence electrons minus the number of unshared electrons of an
atom. Valence electrons are the electrons that appear in the outermost shell of
electrons surrounding the nucleus of an atom. They include lone pair electrons
and unpaired (bonding) electrons. Lone pair electrons are 2 electrons in the same
orbital. Because an orbital with 2 electrons is full, these electrons participate
in covalent bonding much less than do single electrons. From the number of
covalent bonds formed by an atom, one can tell how many unpaired valence
electrons the atom has in a particular molecule. For example, a halogen (ie,
chlorine, bromine, fluorine, iodine) has 1 unpaired electron. Therefore, a
halogen atom can form a maximum of 1 covalent bond with another atom.
Table 1 lists the common elements in organic chemistry and the maximum
number of covalent bonds that each can form.
Multiple Covalent Bonds
Table 1 shows that several atoms can form more than 1 covalent bond. Because
of the nature of covalent bonding, however, it is possible for an atom to form more
than 1 covalent bond with another atom. Carbon will be used as an example.
When 1 unpaired electron from carbon A and 1 unpaired electron from carbon
B are shared, a covalent bond is formed. This bond is called a single bond and is
indicated by a line drawn between the 2 atoms: CA—CB. Ethane (CH3—CH3) is
an example of a compound having a single bond between 2 carbon atoms. Each
carbon atom in ethane forms 4 bonds: 1 to a carbon atom and 3 to each of the
attached hydrogen atoms. A single covalent bond containing 2 shared electrons
is also called a sigma (σ) bond. The simplest form of a σ bond is the overlap of
two s orbitals.
Table 1. The number of covalent bonds usually formed by common organic atoms.
Atom
C (carbon)
N (nitrogen)
O (oxygen)
H (hydrogen)
S (sulfur)
Cl (chlorine)
F (fluorine)
Br (bromine)
I (iodine)
Number of Covalent Bonds
4
3 or 4
2
1
2, 4, or 6
1
1
1
1
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The Chemistry and Physics of Drugs Used in Anesthesia
Certain atoms (generally C, N, O, S) may also form double bonds (Figure 1).
A double bond is 4 electrons shared between 2 atoms. In a double bond, 2 of
the electrons form a σ bond, and the other 2 electrons form a pi (π) bond. The
simplest organic molecule with a double bond is ethylene, CH2=CH2 (also
called ethene). The bond between the 2 carbon atoms in ethylene is a double
bond and is indicated by a double line (=) between the 2 carbon atoms.
Two atoms may also form a triple bond, which is 6 electrons shared between
2 atoms. In a triple bond, 2 of the electrons form a σ bond, and the other 4
electrons form 2 π bonds. The simplest organic molecule with a triple bond is
acetylene, HC≡CH (also called ethyne). The bond between the 2 carbon atoms
in acetylene is a triple bond and is indicated by a triple line (≡) between the
2 carbon atoms.
Figure 1. A covalent bond is the sharing of 2 electrons between 2 atoms. Chlorine atoms
form a σ bond with each other to make Cl2. In ethane, a pi (π) bond is formed from the
overlap of p orbitals. Note that a π bond can be formed only after a σ bond is already
present between 2 atoms. Electrons are free to travel (delocalize) between both atoms of
the π bond.
Double and triple bonds are due to the overlap of p orbitals, which contain
π electrons. Other common double and triple bonds include C=O, C=N, C=S,
C≡N, and S=O. Figure 2 shows the bonding patterns of several common
molecules.
The strength of a chemical bond is measured by the amount of energy
required to break it. Although the exact strength of a bond is dependent on the
particular atoms forming it, typical bond energies are useful to know. Breaking
a carbon-carbon single bond requires about 70 to 90 kcal/mol of energy
(approximately 348 kJ/mol). Breaking a double bond between 2 carbon atoms
requires about 145 to 160 kcal/mol (approximately 612 kJ/mol). Breaking a
carbon-carbon triple bond requires still more energy: 190 to 230 kcal/mol
(approximately 837 kJ/mol).
26
The Basics of Organic Chemistry
Chapter 2
Figure 2. Common molecules with single, double, and triple bonds.
Some useful points to remember are the following:
• A multiple bond is almost always stronger than a single bond.
• A double bond is always stronger than a single bond between the same
2 atoms.
• A triple bond is always stronger than a double bond between the same
2 atoms.
An important detail to note is that the strength of a double bond is less than
twice the strength of a single bond. Based on the average strengths, one would
expect that it would require 696 kJ/mol (2 × 348) to break a double bond;
however, the actual energy required is 612 kJ/mol. This means that the second
bond of the double bond (ie, the π bond) is itself not as strong as a single bond.
Although the double bond is stronger than a single bond overall, the π bond of
the double bond is not as strong as the σ bond of the single bond. Because the π
bond is not as strong as the σ bond, the π bond is usually more reactive than the
σ bond. This is reflected in the metabolism of certain drugs containing a double
bond in which metabolic enzymes functionalize the double bond of the drug but
not the single bond.
Even though 2 electrons are shared in a covalent bond, they may not always
be shared equally between the 2 atoms. When the 2 electrons in a covalent bond
are shared unequally between the 2 atoms, a dipole exists. A dipole is a molecule,
or part of a molecule, with a polarized, covalent bond, such that the electrons
of the bond surround one atomic nucleus more than the other. In this situation,
the atom surrounded by increased electron density has a slight negative charge,
while the atom with less electron density has a slight positive charge. Dipoles
exist when atoms of different electronegativities are bonded to each other. In
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