Biochemistry

An Organic Chemistry Approach

Michael B. Smith

• Verlag: CRC Press
• Jahr: 2020
• ISBN:
  9781351258067
• Auflage: 1
• Thema: Medicina
• Thema: Biochimica in medicina

"There is a continuing demand for up to date organic & bio-organic chemistry undergraduate textbooks. This well planned text builds upon a successful existing work and adds content relevant to biomolecules and biological activity".

-Professor Philip Page, Emeritus Professor, School of Chemistry University of East Anglia, UK

"Introduces the key concepts of organic chemistry in a succinct and clear way".

-Andre Cobb, KCL, UK

Reactions in biochemistry can be explained by an understanding of fundamental organic chemistry principles and reactions. This paradigm is extended to biochemical principles and to myriad biomolecules.

Biochemistry: An Organic Chemistry Approach provides a framework for understanding various topics of biochemistry, including the chemical behavior of biomolecules, enzyme activity, and more. It goes beyond mere memorization. Using several techniques to develop a relational understanding, including homework, this text helps students fully grasp and better correlate the essential organic chemistry concepts with those concepts at the root of biochemistry. The goal is to better understand the fundamental principles of biochemistry.

Features:

• Presents a review chapter of fundamental organic chemistry principles and reactions.
• Presents and explains the fundamental principles of biochemistry using principles and common reactions of organic chemistry.
• Discusses enzymes, proteins, fatty acids, lipids, vitamins, hormones, nucleic acids and other biomolecules by comparing and contrasting them with the organic chemistry reactions that constitute the foundation of these classes of biomolecules.
• Discusses the organic synthesis and reactions of amino acids, carbohydrates, nucleic acids and other biomolecules.

1 Fundamental Principles of Organic Chemistry

It is likely that this book will be used after, or concurrently with, a course in organic chemistry. This chapter is therefore intended as a review of fundamental principles of organic chemistry. It is not intended as a stand-alone treatment. The intent is to bring the reader “up to speed” with important principles of organic chemistry that are important for understanding the extension of those principles to biochemistry.

Arguably, the most fundamental concept in organic chemistry is the nature of the bond between two carbon atoms or between carbon and another atom. Bonding is an important concept in organic chemistry because chemical reactions involve the transfer of electrons with the making and breaking of chemical bonds. Since the molecules associated with biochemistry processes are organic molecules, the bonding will be similar to those organic molecules commonly discussed in a sophomore organic chemistry course. For the most part, the bonds between carbon and another atom are covalent, so the initial focus will be on covalent bonds between two carbon atoms or covalent bonds on a different atom to carbon. The definition of covalent bonds and polarized covalent bonds will be reviewed, as well as the concept of functional groups. The concept of isomers, different connectivity within organic molecules, and rules for naming organic molecules will also be reviewed.

In order to lay the groundwork for understanding bioorganic molecules, this chapter will review relatively simple molecules, hydrocarbons, that have π-bonds to carbon atoms, both carbon–carbon double bonds and carbon–carbon triple bonds. Compounds that have carbon bonded to heteroatoms (e.g., oxygen and nitrogen) may have both σ- and π-bonds. Functional groups with both types of bonding will be reviewed. The rules of nomenclature will be extended to accommodate each new functional group will be briefly reviewed. Relatively simple physical properties associated with polarized covalent bonds and π-bonds, especially hydrogen-bonding, will be reviewed. Finally, chirality, stereochemistry, and stereoisomers will be reviewed.

1.1 Bonding and Orbitals

Elemental atoms are discrete entities that differ from one another by the number of protons, neutrons and electrons that make up each atom. The motion of electrons with respect to the nucleus has some characteristics of a wave, which is expressed as a wave equation and a solution to this equation is called a wavefunction. Each electron may be described by a wavefunction whose magnitude varies from point to point in space. The wavefunction is described by ψ(x,y,z) by using Cartesian coordinates to define a point, which describes the position of the electron in space. Wavefunction solutions are correlated with the space volume pictorial representations, which are charge clouds called orbitals. As stated by the Heisenberg uncertainty principle, the position and momentum of an electron cannot be simultaneously specified so it is only possible to determine the probability that an electron will be found at a particular point relative to the nucleus. Wavefunction solutions can be correlated with the position of an electron relative to the nucleus for an electron, which leads to the familiar s-, p-, and d-orbitals. The s-orbital is spherical and a p-orbital has a “dumbbell” shape and there are three 2p-orbitals that are degenerate (they have the same energy). (Figure 1.1)

FIGURE 1.1 Common representations of the s- and p-orbitals from Figure 3.1, along with two d-orbitals.

Electrons associated with an atom are assumed to reside in atomic orbitals. The valence electrons are used for bonding and in an atom, valence electrons are those found in the outermost orbitals (those furthest away from the nucleus), and they are more weakly bound than electrons in orbitals closer to the nucleus. The number valence electrons for an atom is calculated by subtracting the last digit of the Group number from 8. Since carbon is Group 14, there are four valence electrons (8-4). Similarly, there are three valence electrons for N (8-5), two for oxygen (8-6) and one valence electron for F (8-7). Note that valence for an atom is the number of bonds a molecule may form using the valence electrons. Therefore, the valence of carbon is 4, that of nitrogen is 3, that of oxygen is 2 and that of fluorine is 1.

The so-called covalent chemical bond between two atoms (see Section 1.2) involves the sharing of electrons. The position of electrons in an atomic orbital of an element such as C, N, O or F can be contrasted with the electrons in a bond between two atoms in a molecule, which are assumed to reside in molecular orbitals. It is important to emphasize that the positions of the atomic orbitals relative to the nucleus have a different energy for the electrons found in molecular orbitals.

1.2 Ionic versus Covalent Chemical Bonds

The skeleton of most molecules to be discussed in this book is made up of carbon, oxygen, nitrogen or sulfur atoms, held together by chemical bonds between carbon and carbon as well as bonds of carbon to the other atoms. Two major types of bonds will be considered for these molecules. A covalent bond is formed by the mutual sharing of electrons between two atoms and the bond holds the atoms together. An ionic bond is formed when one atom in a bond has two electrons and takes on a negative charge, and the other is electron-deficient and takes on a positive charge.

An ionic bond holds two atoms or groups together by electrostatic attraction of positive and negatively charged ions. Most ionic bonds in this book will be the salt of monovalent metals such as Na⁺, K⁺, Li⁺, or a divalent metal cation such as Mg²⁺, and the anion is a halide or the salt of a relatively strong acid: carboxylic acids, sulfonic acids, phosphoric acids, or the salt of a weak acid such as an alcohol. However, ionic bonds that involve the conjugate acid of organic bases such as the ammonium salts formed from amines or the phosphonium salts from phosphines are common.

When one carbon shares electrons another carbon, a hydrogen atom, or another atom, it is a covalent bond, also known as a σ-bond. There are two electrons in a σ-bond, which is commonly called a single covalent bond or just a single bond between the two atoms. In other words, a covalent bond is the mutual sharing of two electrons between two atoms. The electron density of each atom is shared with the other in a covalent bond and not localized on an individual atom. Indeed, the greatest concentration of electron density is between the nuclei. The strength of a covalent bond is related to the amount of electron density concentrated between the nuclei. When two identical atoms share electrons in a covalent bond, most (but not all) of the electron density is equally distributed between the two nuclei (in the “space” between the two atoms), which leads to the strongest type of covalent bond.

If one examines the molecule methane, measurements show that all four C—H bonds are identical, and that the bond angles of each H—C—H unit are 109° 28′, the angles of a regular tetrahedron. Specifically, the four hydrogen atoms are distributed around carbon in the shape of a regular tetrahedron. The electrons in the C—H bonds of methane are in molecular orbitals called sp³ hybrid orbitals. Since there are four bonds to carbon, there are four sp³-hybrid orbitals corresponding to the four covalent bonds.

The tetrahedral array of covalent bonds about carbon, as just described, means that a three-dimensional shape is associated with each. The Valence Shell Electron Pair Repulsion (VSEPR) model is a useful place to begin thinking about the three-dimensional nature of organic molecules. To use this model, a tetrahedron is imagined with C, O, or N at the center of that tetrahedron. The atoms are attached at the corner of the tetrahedron, and any unshared electrons are also attached. Examining only the atoms, carbon has a valence of four and molecules will be tetrahedral about each carbon, a nitrogen has a valence of three and will form three covalent bonds with an unshared pair of electrons that gives the molecule a pyramidal shape. Oxygen has a valence of two, with two unshared electrons so the molecule will have an angular or bent shape. Remember, however, that this model does a poor job of accurately predicting bond lengths and angles. Note that this model underestimates the importance of electron pairs and does not take the size of the atoms or groups...