Organic Chemistry Concepts and Applications for Medicinal Chemistry

Joseph E. Rice

• Publisher: Academic Press
• Year: 2014
• Print ISBN: 9780128007396
• eBook ISBN: 9780128008324
• Topic: Physical Sciences
• Subtopic: Organic Chemistry

Organic Chemistry Concepts and Applications for Medicinal Chemistry provides a valuable refresher for understanding the relationship between chemical bonding and those molecular properties that help to determine medicinal activity. This book explores the basic aspects of structural organic chemistry without going into the various classes of reactions. Two medicinal chemistry concepts are also introduced: partition coefficients and the nomenclature of cyclic and polycyclic ring systems that comprise a large number of drug molecules. Given the systematic name of a drug, the reader is guided through the process of drawing an accurate chemical structure. By emphasizing the relationship between structure and properties, this book gives readers the connections to more fully comprehend, retain, apply, and build upon their organic chemistry background in further chemistry study, practice, and exams.- Focused approach to review those organic chemistry concepts that are most important for medicinal chemistry practice and understanding- Accessible content to refresh the reader's knowledge of bonding, structure, functional groups, stereochemistry, and more- Appropriate level of coverage for students in organic chemistry, medicinal chemistry, and related areas; individuals seeking content review for graduate and medical courses and exams; pharmaceutical patent attorneys; and chemists and scientists requiring a review of pertinent material

Chapter 1

Bonding in Organic Compounds

Abstract

This chapter provides a discussion of the concepts related to the formation of the bonds that make up most organic compounds and how they determine the three-dimensional shape and chemical properties of those compounds. It begins with a review of atomic orbitals, showing how for all but the simplest of elements relying on atomic orbitals alone leads to an incorrect picture of chemical bonding. The concept of hybridization is then introduced to provide a more accurate picture of molecular structure. Overlap of hybrid orbitals with other hybrid orbitals or with hydrogen 1s orbitals leads to σ-bond formation. This chapter then goes on to demonstrate how p-orbitals overlap to form π-bonds. Conjugation and resonance are then introduced leading up to a discussion of aromaticity as applied first to carbocyclic and then heterocyclic systems.

Keywords

π-bond; σ-bond; Aromatic; Conjugated system; Electronegativity; Hybrid orbital; Orbitals; Resonance; Valence-shell electron-pair repulsion

By the time most students in the pharmaceutical sciences get to their third or fourth year, they have very likely already taken general chemistry and organic chemistry and have been exposed to the concepts of chemical bonding several times. The reader is likely asking: Why waste time going over this again when there are more interesting topics to explore? After all, it’s pretty simple. Electrons are shared between two atoms to form bonds. There are single bonds, double bonds, and triple bonds. Can we now move on and get to something new?

The truth is that one cannot begin to understand organic chemistry unless he/she has a thorough understanding of the bonds that hold organic compounds together. It is the bonds that determine the three-dimensional shapes of molecules. With many drug molecules, shape is a factor that determines how the drug molecule interacts with its receptor. The bonds in a molecule also determine its chemical reactivity and properties. Again with drug molecules, reactivity relates to its chemical stability and to the types of metabolic transformations the drug may undergo. Bonds affect the acid/base properties of a compound and also affect the solubility. For drugs, the relative solubility in water and lipids affects how the molecule passes from one bodily compartment to another.

Atomic Orbitals

Most drugs are organic compounds and have structures that are largely composed of carbon atoms. Carbon (atomic number 6) has a nucleus containing six positively charged protons. As a counterbalance, six negatively charged electrons orbit the nucleus. Electrons are restricted to orbitals, which are the functions that represent the probability of finding a given electron in space. Any orbital can be occupied by a maximum of two electrons, and these must have opposite spin quantum numbers (represented as ↑ or ↓). The shapes of orbitals are designated with small letters such as s, p, d, or f. Elements that are in the first row of the Periodic Table (hydrogen and helium) have only a single shell of electrons that reside in an s-orbital. In the second row, two shells exist. The second shell being larger than the first can contain more electrons, which are held in one s-orbital and three p-orbitals. Third row elements have three shells of electrons with the third now capable of having five d-orbitals in addition to the one s- and three p-orbitals. Thus the row number in the Periodic Table tells how many shells of electrons a particular element contains (Table 1.1).

Table 1.1

The First Three Rows of the Periodic Table

	I	II											III	IV	V	VI	VII	VIII
1	H																	He
2	Li	Be											B	C	N	O	F	Ne
3	Na	Mg											Al	Si	P	S	Cl	Ar

Electrons in s-orbitals are distributed spherically symmetrical about the nucleus (like a hollow ball with the nucleus at the center). Carbon, being found in the second row of the Periodic Table, has two shells of electrons. Two electrons are contained within the 1s orbital and the remaining four are found in the orbitals of the second shell. Note the fact that carbon is located in Group IV of the Periodic Table and has four electrons in the outer shell or valence shell. Electrons completely populate lower energy orbitals before higher energy orbitals. This is known as the Aufbau principle. Thus the 1s orbital, being closer to the nucleus, is filled before the 2s orbital. Within a given shell, the s-orbital is lower in energy than the p-orbitals. The three p-orbitals are dumbbell shaped and are mutually perpendicular to one another, directed along the x-, y-, and z-axes (p_x, p_y, and p_z). They are also equivalent in energy (degenerate). The remaining valence electrons first fill the 2s orbital and then partially fill two of the degenerate p-orbitals. Hund’s Rule states that electrons must be pla...

Table of contents

• Cover image
• Title page
• Table of Contents
• Copyright
• Dedication
• Preface
• Chapter 1. Bonding in Organic Compounds
• Chapter 2. The Three-Dimensional Structure of Organic Compounds
• Chapter 3. Functional Groups
• Chapter 4. Acids and Bases
• Chapter 5. Partition Coefficients
• Chapter 6. The Nomenclature of Cyclic and Polycyclic Compounds
• Chapter 7. Drug Metabolism
• Index