Biochemistry the molecular basis of life pdf free download

Biochemistry the molecular basis of life pdf free download

Biochemistry. the molecular basis of life,Account Options

Download Biochemistry: The Molecular Basis Of Life [PDF] Type: PDF. Size: MB. Download as PDF. Download Original PDF. This document was uploaded by user and they Download PDF - Biochemistry: The Molecular Basis Of Life [PDF] [7i70h4lb0]. Biochemistry: The Molecular Basis of Life is an intermediate, one-semester text written for 29/09/ · Biochemistry: The Molecular Basis of Life is a one-semester text focusing on the essential biochemical principles that underpin the modern life sciences. The sixth edition 23/07/ · Biochemistry the Molecular Basis of Life T. Mckee, J. R. McKee Published 23 July Education Biochemistry: The Molecular Basis of Life, Fourth Edition, is the ideal 24/01/ · Book Description Biochemistry: The Molecular Basis of Life is a one-semester text focusing on the essential biochemical principles that underpin the modern life sciences. ... read more

Books Video icon An illustration of two cells of a film strip. Video Audio icon An illustration of an audio speaker. Audio Software icon An illustration of a 3. Software Images icon An illustration of two photographs. Images Donate icon An illustration of a heart shape Donate Ellipses icon An illustration of text ellipses. Search Metadata Search text contents Search TV news captions Search archived websites Advanced Search. Biochemistry : the molecular basis of life Item Preview. remove-circle Share or Embed This Item. EMBED for wordpress. com hosted blogs and archive. Want more? Advanced embedding details, examples, and help! xxi, pages : 29 cm Includes bibliographical references and index Biochemistry : An introduction -- Living cells -- Water : The medium of life -- Energy -- Amino acids, peptides, and proteins -- Enzymes -- Carbohydrates -- Carbohydrate metabolism -- Aerobic metabolism I : The citric acid cycle -- Aerobic metabolism II : Electron transport and oxidative phosphorylation -- Lipids and membranes -- Lipid metabolism -- Photosynthesis -- Nitrogen metabolism I : Synthesis -- Nitrogen metabolism II : Degradation -- Integration of metabolism -- Nucleic acids -- Genetic information -- Protein synthesis -- App.

Access-restricted-item true Addeddate Associated-names McKee, James R. Full catalog record MARCXML. For example, carbon dioxide CO2 has one carbon atom with four valence electrons and two oxygen atoms, each with six valence electrons. Determine the identity of the central atom in the Lewis structure. This atom will often be the one with the lowest electronegativity. Recall that electronegativity decreases from right to left across and from top to bottom of the periodic table. In the case of CO2, carbon is less electronegative than oxygen, so carbon is the central atom. Arrange the electrons so that each atom donates one electron to a single bond between it and another atom and then count the electrons around each atom. Are the octets complete? For CO2, a first try would yield Note, however, that in this structure each carbon only has six electrons and each oxygen atom has seven electrons.

Because the octets are incomplete, more electrons must be shared, an indication that there are double or triple bonds in the molecule. In the case of CO2, rearranging the electrons results in the following Lewis dot structure in which there are two double bonds and all three atoms have a full octet of electrons. In molecules such as ammonia NH3 , it is obvious that although the nitrogen atom is more electronegative than the hydrogens, it is the only atom in the molecule that can form multiple bonds. Hence, the nitrogen atom is the central atom in ammonia molecules. Its Lewis dot structure is For a large number of molecules, there is more than one valid Lewis dot structure. Inasmuch as nitrogen has five valence electrons and each oxygen atom has six valence electrons, the following Lewis dot structure of the nitrate ion satisfies the octet rule: However, there is no reason why the double bond should appear where it does in this formula.

It could easily appear in either of the two other locations around the nitrogen atom. Therefore, there are three valid Lewis dot structures for the nitrate ion. When this situation occurs, the ion or molecule is said to be a resonance hybrid. The doubleheaded arrows are used in the representation of resonance structures. In the case of the nitrate ion, it is considered to have a structure that is the average of these three states. MOLECULAR STRUCTURE Molecules are three-dimensional arrangements of atoms. Understanding molecular structure, also referred to as molecular shape, is important because structure provides insight into the physical and chemical properties of molecules. Physical properties that are affected by molecular shape include boiling point, melting point, and water solubility.

The shape of molecules also powerfully affects chemical reactivity. According to the valence shell electron pair repulsion VSEPR theory, repulsive forces between valence shell bonding and nonbonding electrons lone pairs determine molecular geometry molecular shape. In other words, the valence electron pairs on the central atom in a molecule orient themselves in space so that repulsion is minimized i. Lone pairs of electrons have a greater repulsive effect than bonding pairs. A lone pair is a valence electron pair on a central atom that is not involved in bonding. The term electron group is used in discussions of VSEPR theory. An electron group is defined as a set of valence electrons in a region around a central atom that exerts repulsion on other valence electrons.

Electron groups include bonding and nonbonding electron pairs or the pairs of electrons in double or triple bonds. WORKED PROBLEM 2 What is the Lewis electron dot formula for formaldehyde H2C O? SOLUTION 1. Calculate the number of valence electrons for each atom in the molecule. The valence electrons for hydrogen, carbon, and oxygen are 2 one for each atom , 4, and 6, respectively, for a total of 12 electrons. Single bonds between the elements account for 6 electrons, leaving 6 electrons unaccounted for. Group the remaining 6 electrons around the most electronegative atom oxygen until a total of 8 electrons bonding and nonbonding is reached. Using one pair of these electrons to form a double bond between carbon and oxygen completes the carbon octet. The final Lewis structure is given below.

The carbon atom has 4 valence electrons and each of the 4 chloride atoms have 7 valence electrons for a total of 32 valence electrons. Determine the number of bonding electrons. Since there are four chlorines bonded to the central carbon atom, there are eight bonding electrons. Carbon is the central atom because it has the lowest electronegativity. Calculate the number of remaining electrons and distribute them about each of the chloride atoms so that each atom in the molecule has an octet. Subtracting 8 bonding electrons from 32 valence electrons yields 24 electrons. The remaining 24 electrons are distributed around the 4 chlorine atoms as lone pairs.

The final Lewis structure is given as follows. If there are two electron pairs, the molecule has a linear shape. Carbon dioxide CO2 , for example, is a linear molecule with two electron groups. Its bond angle is o. Formaldehyde H2C O , with three electron groups, has trigonal planar geometry with bond angles of °. Molecules with a central atom with four pairs of electrons have a tetrahedral shape. Methane CH4 , with its four carbon—hydrogen bonds, has bond angles of If one of the four electron groups in a tetrahedron is a lone pair, the molecular shape is trigonal pyramidal. Because of the strong repulsion of the lone pair, bond angles are less than For example, the lone pair in NH3 forces the NH bonding electron pairs closer together with bond angles of Three-dimensional shape also affects molecular polarity.

In polar covalent bonds, there is an unequal sharing of electrons because the atoms have different electronegativities. This separation of charge is called a dipole. Although a polar molecule always contains polar bonds, some molecules with polar bonds are nonpolar. Molecular polarity requires an asymmetric distribution of polar bonds. For example, CO2 contains two C—O dipole bonds. Carbon dioxide is a nonpolar molecule because of its linear shape i. Water, which also has two polar bonds two O—H bonds , is a polar molecule because of its geometry. Water contains four electron groups: two bonding pairs and two lone pairs.

As a result of the greater repulsions from the lone pairs on the oxygen, however, the bond angle of a water molecule is FIGURE 5 Common Molecular Geometrics These structures illustrate the spatial orientations of electron groups. Note that lone pair electrons are indicated by an enlarged representation of an orbital. WORKED PROBLEM 4 Dimethyl ether has the following formula: CH3—O—CH3. What is the shape of dimethyl ether? Does this molecule have a dipole moment? SOLUTION The oxygen in dimethyl ether has four electron pairs, including two lone electron pairs.

As a result, the oxygen has a tetrahedral shape with an overall bent shape for the dimethyl ether molecule. The concept of orbital hybridization, the result of quantum-mechanical calculations, explains how the mixing of atomic orbitals results in the formation of the more stable hybrid orbitals found in molecules. Each type of hybrid orbital corresponds to a type of electron group arrangement predicted by VSEPR theory. The three most common hybrid orbitals observed in biomolecules are sp3, sp2, and sp. Carbon has an electron configuration of 1s22s22p2, which can also be represented as It appears from this diagram that carbon only has two bonding electrons. The carbon atoms in molecules such as methane, however, are bonded to four hydrogen atoms in a tetrahedral arrangement. During methane formation, as a result of the attraction of each of the hydrogen nuclei i. As they do so, they mix, forming four identical sp3 orbitals Figure 6. FIGURE 6 sp3 Orbitals Hybridization of an s orbital and all three p orbitals gives four identical sp3 orbitals.

In the methane molecule Figure 7 , each of the four sp3 hybrid orbitals overlaps with the 1s orbital of hydrogen to form a sigma bond. A sigma bond σ , which is formed by the overlapping by the outermost orbitals of two atoms, is the strongest type of covalent bond. Each of the two carbon atoms in the molecule ethene H2C CH2 is bonded to three atoms in trigonal planar geometry. Two of the three sp2 orbitals of each carbon atom overlap the orbital of a hydrogen atom, forming a total of four σ bonds. The third sp2 orbital of the two carbon atoms overlaps to form a carbon— carbon σ bond. The p orbitals, one on each carbon, overlap to form a pi π bond Figure 8. A double bond in molecules, such as ethene, consists of a σ bond and a π bond.

Acetylene C2H2 is a molecule with a triple bond, with each carbon bonded to two other atoms in a linear geometry. FIGURE 8 Ethene Structure Each carbon atom in ethene also known as ethylene has three sp2 orbitals with bond angles of °, which have a trigonal planar geometry. b Two of the sp2 orbitals of each carbon green overlap with an s orbital of hydrogen red , forming a total of 4 σ bonds. The remaining two sp2 orbitals, one from each carbon, overlap to form a carbon—carbon σ bond. c Two p orbitals blue , one from each carbon atom, overlap to form a π bond. mixes with one 2p orbital to form 2 sp hybrid orbitals. Each carbon also possesses two unhybridized 2p orbitals. Acetylene has a triple bond consisting of one σ bond and two π bonds. The carbon—carbon σ bond is formed by the overlap of an sp hybrid orbital from each carbon atom. Chemical Reactions In chemical reactions, the atoms in chemical substances are rearranged to form new substances as chemical bonds are broken and formed.

According to the collision theory, the reaction rate in bimolecular reactions depends on the frequency of successful collisions between the chemical species. Successful collisions occur when there is sufficient energy at the moment of impact called the activation energy and the colliding species are oriented during the collision in a manner that favors the rearrangement of the atoms and electrons. Catalysts are substances that increase reaction rates without being affected by the reaction. They do so by lowering the activation energy of the reaction by providing an alternative pathway for the reaction. For example, the metal iron is used as a catalyst in the Haber process, the industrial method of converting nitrogen and hydrogen gases into ammonia NH3. As the N2 and H2 molecules become adsorbed onto the surface of the metal, where they are more likely to be in a favorable orientation for successful collisions, the bonds in both molecules are weakened.

Once formed, the ammonia molecules desorb from the catalyst. Chemical reactions are described with chemical equations. The substances undergoing the reaction, called reactants, appear on the left-hand side of the equation, and the products of the reaction are on the right-hand side. An arrow between the reactants and products symbolizes the chemical change that occurs as a result of the reaction. For example, the chemical equation for the reaction in which molecule A reacts with molecule B to form molecules C and D is Other symbols that may appear in a chemical equation provide information about the physical state of the reactants and products or a required energy source.

The equation for the decomposition of calcium carbonate, for example, is In this equation, the letter s indicates that the reactant CaCO3 and the product CaO calcium oxide are solids. The letter g indicates that CO2 is a gas. The uppercase Greek letter delta Δ above the arrow indicates that the reaction requires an input of energy in the form of heat. Reactions that require an input of energy in the form of heat are described as endothermic. If light energy is involved in a reaction, hν ν is the lowercase Greek letter nu is placed above the arrow. All chemical equations must obey the law of conservation of matter, which states that during chemical reactions mass is neither created nor destroyed. In other words, the mass of the reactants must be equal to the mass of the products. For example, in the equation for the reaction in which methane CH4 reacts with molecular oxygen O2 to form carbon dioxide CO2 and water, the number of each type of atom on both sides of the arrow must be equal.

In this balanced equation, the same number of carbon, hydrogen, and oxygen atoms is on both sides of the arrow because the number 2 has been placed before the formulas for molecular oxygen and water. WORKED PROBLEM 5 Consider the following reaction equation: Balance the equation and identify the elements that are oxidized, reduced, or unchanged by the reaction. SOLUTION Balancing a chemical reaction equation requires that the number and types of atoms be the same. To satisfy this requirement, the number 4 is placed before the reactant KClO and the product KCl. The reaction becomes Using the oxidation state information given above, oxidation numbers are assigned to each element: Sulfur is the element that is oxidized i. The science of chemical kinetics seeks to answer these and other questions about chemical reaction rates i.

Reaction rate is defined as the change in the concentration of reactant or product per unit time. For the general reaction the rate is equal to k[A]m[B]n, where k is the rate constant and [A] and [B] are the concentrations of the reactants A and B, respectively. For example, if m is equal to 1, the reaction rate doubles when the concentration of reactant A doubles. If m is equal to 2, the rate quadruples when the concentration of reactant A doubles. Refer to pp. The rate constant and order of a reaction can only be determined by experiment. Experiments performed over the course of the past century have revealed that the following factors influence reaction rate: 1. Reactant structure. The nature and the strength of chemical bonds affect reaction rates. For example, salt formation, the exchange of ions, is a fast process compared with the breaking and forming of covalent bonds.

Reactant concentration. The number of molecules of a substance per unit volume affects the likelihood of collisions. Reaction rates increase as the reactant molecules become more crowded. Physical state. Whether the reactants are in the same phase solid, liquid, or gas affects reaction rates because reactants must come into contact with each other. When reactants are in the aqueous phase, for example, thermal motion brings them into contact. When reactants are in different phases, contact only occurs at the interface between the phases. For example, when reactants are in the solid and liquid phases, grinding the solid into small pieces increases its surface area that is in contact with the liquid phase. At higher temperatures, molecules have more thermal energy and are, therefore, more likely to collide with each other.

A catalyst provides a different pathway for the reaction, thereby lowering the activation energy. Reversible reactions are indicated in reaction equations with double arrows. When a reversible reaction begins i. At some point in time, which differs for each reaction, some product molecules are reconverted back into reactant molecules. Eventually, the reaction reaches a dynamic equilibrium state in which both the forward and reverse reactions occur, but there is no net change in the ratio of reactant and product molecules. The extent to which the reaction proceeds to product is measured by an equilibrium constant Keq , which reflects the concentrations of reactants and products under specific conditions of temperature and pressure. For a reaction with the equation Keq is calculated as the ratio of the molar concentrations of product and reactant, each of which is raised to the power of its coefficient. A high Keq value significantly greater than 1 indicates that when a reaction reaches equilibrium, the concentration of the reactant is low i.

If the Keq value is lower than 1, then product concentration is lower than reactant concentration when equilibrium has been reached. When Keq is greater than , the reaction has gone to completion i. In , the French chemist Henri Louis Le Chatelier reported his discovery of a remarkable feature of systems at equilibrium. For a chemical reaction at equilibrium, a change in the conditions of the reaction e. The Haber—Bosch process for making ammonia NH3 from N2 and H2 is a prominent example. All living organisms require a source of usable nitrogen-containing molecules. As a result of the extraordinary difficulty in breaking the stable triple bond of N2, nitrogen fixation the conversion of N2 to NH3, a molecule that can be assimilated into organic molecules such as amino acids is largely limited to a select group of microorganisms.

Note that the synthesis of ammonia is an exothermic reaction i. An iron-based catalyst iron oxide with small amounts of other metal oxides , which increases the rate at which equilibrium is attained, converts a slow reaction to one that is fast enough to be commercially feasible. Ammonia, the product of the reaction, is removed from the reaction vessel. As a result, the system produces more NH3 to reestablish equilibrium. An increase in the pressure within the reaction vessel to atm , obtained by decreasing volume, causes an increase in ammonia synthesis. Note that in this reaction 4 mol of reactant molecules are converted to 2 mol of product. The equilibrium shifts toward ammonia synthesis because there are fewer molecules of this gas.

By lowering the temperature of the reaction i. There is a limit to how much the temperature can be lowered, however, because the catalyst requires heat to be efficient. As a result, the reaction vessel operates at °C, a temperature that is hot enough for the catalyst, yet relatively cool for an industrial process. Hydrochloric acid and acetic acid CH3COOH are two well-known acids. Sodium hydroxide NaOH and methylamine CH3NH2 are examples of bases. The strength of an acid or a base is determined by the degree to which it dissociates. HCl is a strong acid because its dissociation in water is complete i. Weak acids and weak bases are so named because they dissociate only to a limited extent and establish a dynamic equilibrium with their ions. The dissociation constants of weak acids and bases are usually expressed as the negative log of the equilibrium constant —log Ka or —log Kb , where the term —log is replaced by the letter p.

The extent to which a weak acid dissociates is referred to as its pKa value. The dissociation constant and pKa for acetic acid at 25°C, for example, are 1. The behavior of weak acids and bases is especially important in biochemistry because many biomolecules possess carboxylate, amino, and other functional groups that can accept or donate hydrogen ions. For example, refer to pp. Water also has a slight capacity to dissociate into ions. The hydrogen ion concentration of pure water at 25°C is 1. Because one hydroxide ion is produced for each hydrogen ion, the hydroxide ion concentration is also 1. The product of these two values i. The concentrations of hydrogen and hydroxide ions change depending on the substances that are dissolved in water, but their product is always 1 × 10— For weak acids and bases, hydrogen ion concentrations in aqueous solution can vary from 1 M to 1 × 10—14 M.

For the sake of convenience, hydrogen ion concentrations are usually converted to pH values. The term pH simply means that the concentration of hydrogen ions in a solution has been converted to its negative log value i. REACTION TYPES There are several basic types of chemical reaction: synthesis reactions, decomposition reactions, displacement reactions, double displacement reactions, acid—base reactions, and redox reactions. Each is discussed briefly. Synthetic reactions also referred to as combination reactions involve two or more substances that combine to form a single new substance. The reaction of sulfur trioxide SO3 with water, for example, yields sulfuric acid H2SO4. WORKED PROBLEM 6 The Ka for acetic acid is 1. Determine the hydrogen ion concentration of a 0.

What is the pH of this solution? SOLUTION The equation for the dissociation of acetic acid is Because acetic acid is a weak acid, it is assumed that the dissociation of acetic acid has no substantive effect on acetic acid concentration. The values of the concentrations of acetate and hydrogen ions are equal to each other and are set at x. The equation for determining the hydrogen ion concentration in a 0. The pH of the 0. For example, ammonium sulfate, NH4 2SO4, decomposes upon heating to yield ammonia NH3 and H2SO4. In displacement or substitution reactions a more reactive element replaces a less active element. For example, if an iron nail is placed in an aqueous solution of copper II sulfate i. The surface of the iron nail turns to reddish brown because of the deposition of metallic copper. Predicting whether a specific metal will displace another is accomplished by referring to the activity series of metals, a list of metals found in general chemistry textbooks that is arranged in order of strength of metal reactivity from highest to lowest.

In double displacement reactions, two compounds exchange their ions to form two new compounds. For example, silver nitrate reacts with potassium bromide in aqueous solution to yield silver bromide and potassium nitrate. The silver bromide product is insoluble in water and precipitates out of solution. Acid—base reactions are a type of double displacement reaction. The Bronsted—Lowry theory defines acids and bases as proton donors and proton acceptors, respectively. Together these two species constitute a conjugate acid—base pair. They also form a conjugate acid—base pair. In another way of explaining acid—base reactions, referred to as the Lewis acid and base theory, acids and bases are defined in terms of atomic structure and bonding.

A Lewis acid is a chemical species that accepts an electron pair and has a vacant low-energy orbital. A Lewis base is defined as a chemical species that donates an electron pair and possesses lone pair electrons. The product of a Lewis acid—base reaction contains a new covalent bond. In the reaction of HCl with ammonia, HCl is polarized, with the slightly positive hydrogen and the chloride slightly negative. Ammonia :NH3 acting as a Lewis base is attracted to the hydrogen atom. As the lone pair on the nitrogen approaches the HCl, the latter becomes more polarized i. Oxidation-reduction reactions, also referred to as redox reactions, involve the exchange of electrons between chemical species. Oxidation occurs when an ion, atom, or molecule loses electrons. In a reduction, there is a gain of electrons. In the reaction of metallic zinc with molecular oxygen to form zinc oxide, for example, the zinc atoms are oxidized i.

Although oxidation and reduction occur simultaneously, for convenience they may be considered two separate half-reactions, one involving oxidation and the other reduction. The oxidation halfreaction is where the two zinc atoms lose two electrons each. In the reduction half-reaction the two atoms of oxygen gain a total of four electrons. The species that accepts the electrons is referred to as the oxidizing agent. In the reaction of zinc with molecular oxygen, zinc is the reducing agent and molecular oxygen serves as the oxidizing agent. It should be noted that any type of reaction in which the oxidation state of the reactants changes could also be classified as a redox reaction. Combustion reactions are a type of redox reaction in which fuel molecules react with an oxidizing agent to release large amounts of energy, usually in the form of heat and light.

Reactions that release energy are described as exothermic. The burning of the hydrocarbon methane natural gas is a typical combustion reaction. The oxidation half-reaction is The reduction half-reaction is Molecular oxygen is the oxidizing agent in the combustion of methane. The eight electrons removed from methane, the reducing agent, are used in combination with four protons to reduce the oxygen atoms to form two water molecules. It should be noted that cellular respiration, the biochemical mechanism whereby aerobic oxygen-utilizing living cells extract energy from fuel molecules such as the sugar glucose, is a slower, controlled combustion reaction process. WORKED PROBLEM 7 Consider the following combustion reaction equation: Identify the elements that are being reduced, oxidized, or unchanged in this reaction, and the reducing or oxidizing agents. The oxidation state of an atom in its elemental state is 0.

In the following half-reaction, a decrease in the oxidation number for oxygen indicates that each oxygen atom has gained two electrons; that is, it has been reduced. In a redox reaction, the oxidizing agent is reduced and the reducing agent is oxidized. MEASURING CHEMICAL REACTIONS Chemists use the mole concept as a means of determining the amounts of the reactants and products in chemical reactions. A mole is defined as the amount of a substance that contains as many particles e. This number, which is 6. So there are 6. The molar mass of substances mass per mole of particles is used to determine the amounts of reactants and products in a reaction. For example, in the reaction of methane CH4 with O2 to yield carbon dioxide and water, how much water is produced from the combustion of 8 g of methane?

Solving this problem begins with a balanced equation: According to this reaction equation, the combustion of every mole of methane yields 2 mol of water. The number of moles of methane is calculated by dividing the mass of methane 8 g by the molecular mass of methane, which is 16 g the carbon atom has a mass of 12 g, and each of the four hydrogens is 1 g. Refer to the periodic table for atomic mass numbers. By this calculation, there are 0. Because the ratio of methane to water is 1 to 2, the 0. Because the molecular mass of water is 18 g, the combustion of 8 g of CH4 produces 18 g of H2O.

Moles are also used to express concentrations of substances in solution. Molarity is defined as the number of moles in 1 liter l of solution. For example, what is the molarity of a solution of 5 g of NaCl in 2 l of water? First, the number of moles of NaCl must be determined by dividing the mass of NaCl 5 g by the formula weight of NaCl By this calculation i. Molarity of the solution is determined by dividing the number of moles by the number of liters. The molarity of the solution in this problem is 0. This number is rounded off to 0. Refer to a general chemistry textbook for a discussion of significant figures. WORKED PROBLEM 8 The empirical formula of the sugar glucose is C6H12O6.

a How many moles are there in g of glucose? b Calculate the molarity of a solution of g of glucose dissolved in 2. SOLUTION a The number of moles of glucose is calculated by dividing the molecular mass of glucose by its mass. First, the molecular mass of glucose must be determined by adding the sums of the masses of each atom in glucose. Adding these numbers, the molecular mass m of glucose is determined to be g. The number of moles of glucose in g is calculated by dividing the mass of glucose g by the molecular mass g. There are 1. b The molarity of the glucose solution is calculated by first determining the number of moles of glucose in g. The molarity of the glucose solution is then calculated by dividing the number of moles by the number of liters. The molarity of the glucose solution is 0. An entire field is devoted to molecules composed of carbon because of its astonishing versatility. In addition to its capacity to form stable covalent bonds with other carbon atoms to form long chains, branch chains, and rings, carbon also forms stable covalent bonds with a variety of other elements e.

Carbon can also form carbon—carbon double and triple bonds. As a result of these properties, the possibilities for molecules with different arrangements of carbon and the other elements are virtually limitless. For students embarking on the study of biochemistry, a thorough understanding of the principles of organic chemistry is essential because, as stated previously, with the exceptions of inorganic molecules such as H2O, O2, NH3, and CO2 and several minerals e. The structural and functional properties of proteins, nucleic acids DNA and RNA , fats, and sugars can only be appreciated when students know how carbon-based molecules behave.

This review will focus on the structures and the chemical properties of the major classes of organic molecules: the hydrocarbons molecules containing only carbon and hydrogen and substituted hydrocarbons hydrocarbon molecules in which one or more hydrogens has been replaced with another atom or group of atoms. Hydrocarbons Because hydrocarbon molecules contain only carbon and hydrogen, they are nonpolar. They dissolve in nonpolar solvents such as hexane and chloroform but not in water. The hydrocarbons are classified into four groups: 1 saturated hydrocarbons molecules containing only single bonds , 2 unsaturated hydrocarbons molecules with one or more carbon—carbon double or triple bonds , 3 cyclic hydrocarbons molecules containing one or more carbon rings , and 4 aromatic hydrocarbons molecules that contain one or more aromatic rings, which can be described as cyclic molecules with alternating double and single bonds.

The saturated hydrocarbons, referred to as the alkanes, are either normal straight chains or branched chains. The straight-chain alkanes belong to a homologous series of compounds that differ in the number of carbon atoms they contain. The first six members of this series are methane CH4 , ethane C2H6 , propane C3H8 , butane C4H10 , pentane C5H12 , and hexane C6H Note that the prefix in each of these names indicates the number of carbon atoms e. Hydrocarbon groups that are derived from alkanes are called alkyl groups. For example, a methyl group is a methane molecule with one hydrogen atom removed. As their name suggests, the branched-chain hydrocarbons are carbon chains with branched structures. If one hydrogen atom is removed from carbon-2 of hexane, for example, and a methyl group is attached, the branched product is 2-methylhexane.

Note that branched-chain molecules are named by first identifying the longest chain and that the number of the carbon that is bonded to the sidechain group is the lowest one possible. One of the most remarkable features of the hydrocarbons is the capacity to form isomers, molecules with the same type and number of atoms that are arranged differently. For example, there are three molecules, each with its own set of properties, with the molecular formula C5H12 : pentane, 2-methylbutane, and 2,2-dimethylpropane. The alkanes are unreactive except for combustion p. P and halogenation reactions. In halogenation reactions, alkane molecules react at elevated temperatures or in the presence of light, forming free radicals atoms or molecules with an unpaired electron. For example, when methane reacts with chlorine gas Cl2 , the molecule breaks down to form two chlorine radicals, which then initiate a chain reaction with methane molecules that yields several chlorinated products: CH3Cl methyl chloride , CH2Cl2 methylene chloride , CHCl3 chloroform , and CCl4 carbon tetrachloride.

There are two types of unsaturated hydrocarbons: the alkenes, which contain one or more double bonds, and the alkynes, which contain one or more triple bonds. The double bond in alkenes is formed from the overlap of two carbon sp2 orbitals a σ bond and the overlap of two unhybridized p orbitals one from each carbon to form a π bond. Ethene H2C CH2 , also known by the older name ethylene, is the first member of the series. For alkenes with more than three carbons, the carbons are numbered in reference to the double bond so that the numbers are the lowest possible. For example, CH2 CH—CH2—CH2— CH2—CH3 is named 1-hexene, not 5-hexene. Alkenes with four or more carbons have structural isomers in which the position of the carbon—carbon double bond is different. For example, 1-butene and 2-butene are referred to as positional isomers.

The rigidity of the carbon—carbon double bond prevents rotation, thereby producing another class of isomers: geometric isomers. Geometric isomers occur when each of the carbons in the double bond has two different groups on it. For example, there are two geometric isomers of 2-butene: cis- 2-butene, in which the methyl groups are on the same side of the double bond, and trans-2butene, in which the methyl groups are on opposite sides of the double bond. Note that 1-butene does not form geometric isomers because one of the double-bonded carbons does not have two different groups. The alkynes such as ethyne or acetylene contain triple bonds composed of 1 σ bond and 2 π bonds. The carbon—carbon triple bond is rare in biomolecules and is not discussed further. The principal reaction of alkenes is the electrophilic addition reaction in which an electrophile an electron-deficient species forms a bond by accepting an electron pair from a nucleophile an electron-rich species.

Electrophiles have positive charges or they may have an incomplete octet. Nucleophiles have negative charges e. Hydrogenation and hydration are two addition reactions that occur frequently in living organisms. In a laboratory or industrial hydrogenation reaction, a metal catalyst e. FIGURE 9 Acid-Catalyzed Hydration of an Alkene In the first step, protonation of the double bond forms a carbocation. Nucleophilic attack by water gives a protonated alcohol. Deprotonation by a water molecule yields the alcohol. Note that the narrow, colored arrows in this and other reaction mechanisms indicate the movement of electrons. Hydration reactions of alkenes are electrophilic addition reactions that yield alcohols.

The reaction Figure 9 requires a small amount of a strong acid catalyst such as sulfuric acid H2SO4 because water is too weak an acid to initiate protonation of the alkene. The newly formed carbocation a molecule containing a positively charged carbon atom is then attacked by the nucleophilic water molecule to yield an oxonium ion a molecule containing an oxygen cation with three bonds. The alcohol product is formed as the oxonium ion transfers a proton to a water molecule. For example, the principal product of propene hydration is 2-propanol and not 1-propanol. As with the alkanes, the cycloalkanes undergo combustion and halogenation reactions. Ring strain, observed in cycloalkane rings with three or four carbons, is caused by unfavorable bond angles that are the result of distortion of tetrahedral carbons.

As a result, the carbon—carbon bonds in these molecules are weak and reactive. There is minimal or no ring strain in cycloalkane rings with five to seven carbons. Cyclohexane has no ring strain because it is puckered so that its bond angles are near the tetrahedral angles. The most stable puckered conformation is the chair form. Because of greater ring strain, cyclopropene and cyclobutene are even less stable than cyclopropane and cyclobutane. Heterocyclic aromatic compounds have two or more different elements in their rings. The simplest aromatic hydrocarbon is benzene.

Cytosine, a pyrimidine base found in DNA and RNA, is an example of a heterocyclic aromatic molecule. Despite the presence of double bonds, benzene and the other aromatic molecules do not undergo reactions typical of the alkenes. In fact, aromatic compounds are remarkably stable. This stability is the result of the unique bonding arrangement of aromatic rings. Each carbon has three sp2 orbitals that form three σ bonds with two other carbon atoms and with one hydrogen atom. The 2p orbital of each of the six carbon atoms overlaps side to side above and below the plane of the ring to form a continuous circular π bonding system.

Instead of two alternate structures of benzene, benzene is a resonance hybrid, which is indicated by the fact that all of the carbon—carbon bonds in benzene are the same length. In alkenes, carbon—carbon double bonds are shorter than carbon— carbon single bonds. Each carbon atom is joined to its neighbors by the equivalent of one and a half bonds. Resonance explains why aromatics do not undergo the addition reactions observed with alkenes: the delocalizing of the π electrons around an aromatic ring confers considerable stability to the molecule. Not all cyclic compounds that contain double bonds are aromatic molecules.

In addition, every atom in an aromatic ring has either a p orbital or an unshared pair of electrons. Despite their resistance to addition reactions, aromatic compounds are not inert. They can undergo substitution reactions. In electrophilic aromatic substitution reactions, an electrophile reacts with an aromatic ring and substitutes for one of the hydrogens. For example, benzene reacts with HNO3 in the presence of H2SO4 to yield nitrobenzene and water Figure In the first step in the reaction, a strong electrophile is generated. In the second step, a pair of π electrons in the benzene attack the electrophile, resulting in the formation of a resonance-stabilized carbocation intermediate. In the final step, the aromatic ring is regenerated when a water molecule abstracts a proton from the carbon atom bonded to the electrophile.

FIGURE 10 Nitration of Benzene The nitronium ion, a powerful electrophile, is created by the protonation of HNO3 by H2SO4. The aromaticity of nitrobenzene is restored with the loss of a proton to a water molecule. Substituted Hydrocarbons Substituted hydrocarbons are produced by replacing one or more hydrogens on hydrocarbon molecules with functional groups. Functional groups also separate the substituted hydrocarbons into families. For example, methanol CH3—OH , a member of the alcohol family of organic molecules, is the product when the functional group —OH is substituted for a hydrogen atom on methane CH4. Three general classes of functional groups are important in biomolecules: oxygen-containing, nitrogen- containing, and sulfur-containing molecules.

The structural and chemical properties of each class are briefly discussed. Also refer to Table 1. There are six major families of organic molecules that contain oxygen: alcohols, aldehydes, ketones, carboxylic acids, esters, and ethers. Amines and amides possess nitrogen-containing functional groups. The major type of sulfur- containing functional group is the sulfhydryl group, which occurs in thiols. ALCOHOLS In alcohols, the hydroxyl group —OH is bonded to an sp3 hybridized carbon. The presence of the polar —OH group makes alcohol molecules polar, allowing them to form hydrogen bonds with each other and with other polar molecules. A hydrogen bond is an attractive force between a hydrogen atom attached to an electronegative atom e. For alcohols with up to four carbons methanol, ethanol, propanol, and butanol , the polar OH group allows them to dissolve in water because hydrogen bonds form between the hydrogen of the OH group of the alcohol and the oxygen of a water molecule.

Alcohols with five or more carbons are not water-soluble because the hydrophobic properties of hydrocarbon components of these molecules are dominant. Alcohols can be classified by the number of alkyl groups designated as R groups attached to the carbon adjacent to the —OH group. Ethanol CH3CH2—OH is a primary alcohol. Alcohols are weak acids i. Tertiary alcohols are less acidic than primary alcohols because the alkyl groups inhibit the solvation of the alkoxide ion. The increased electron density on the oxygen atom in these molecules also decreases proton removal.

Alcohols react with carboxylic acids to form esters. They can also be oxidized to give the carbonyl group-containing aldehydes, ketones, or carboxylic acids. The carbonyl group C O in which a carbon is double-bonded to an oxygen atom is a structural feature of the aldehydes, ketones, carboxylic acids, and esters. Amides, which contain both nitrogen atoms and carbonyl groups, are described on p. The carbonyl group is polar because of the difference in electronegativity between oxygen and carbon. The slightly positive carbon is electrophilic and therefore able to react with nucleophiles. ALDEHYDES The functional group of the aldehydes is a carbonyl group bonded to a hydrogen atom [— C O —H].

The simplest aldehyde is formaldehyde also referred to as methanal in which the aldehyde group is bonded to a hydrogen atom. In all other aldehydes the aldehyde group is bonded to an alkyl group. The general formula for aldehydes is abbreviated as R-CHO. Acetaldehyde CH3CHO is the oxidation product of ethanol. The reaction of an aldehyde with an alcohol yields a hemiacetal Figure Note that hemiacetals are unstable and their formation is readily reversible. The reaction of the aldehyde group of aldose sugars with an intramolecular OH group to form the more stable cyclic hemiacetals see p. FIGURE 11 Hemiacetal Formation The reaction begins with the acid catalyst protonating the carbonyl group. The alcohol, acting as a nucleophile, attacks the resonance stabilized carbocation. The hemiacetal product forms with the release of a proton from the positively charged intermediate.

The names of members of the ketone family end in -one. For example, dimethyl ketone is usually referred to by its original name, acetone. Ethylmethyl ketone is also referred to as 2butanone. In ketoses, sugars with a ketone group most notably fructose , the carbonyl group reacts with an OH group on the sugar molecule to form a cyclic hemiketal. These molecules function as weak acids i. Carboxylic acids react with bases to form carboxylate salts. For example, acetic acid reacts with sodium hydroxide to yield sodium acetate and water. The simplest carboxylic acid is formic acid HCOOH , which is found in ant and bee stings. Carboxylic acids with more than two carbon atoms are often named using the hydrocarbon precursor name followed by the ending -oic acid. For example, the carboxylic acid derived from the four-carbon molecule butane is butanoic acid.

In living organisms, the longer-chained carboxylic acids, called fatty acids, are important components of biological membranes and the triacylglycerols, a major energy-storage molecule. An ester is the product of a nucleophilic acyl substitution reaction in which a carboxylic acid reacts with an alcohol. For example, isobutanol reacts with acetic acid to form isobutyl acetate, an ester found in cherries, raspberries, and strawberries. The formation of the ester methyl acetate from acetic acid and methanol is illustrated in Figure Fats and vegetable oils, also called triacylglycerols see p.

Diethyl ether CH3CH2—O— CH2CH3 , the best-known ether, was the first anesthetic used in surgery late nineteenth century and is still used as a solvent. Ethers are relatively inert chemically, but they do convert over time into explosive peroxides e. FIGURE 12 Formation of Methyl Acetate Step 1 Acetic acid is protonated on its carbonyl oxygen to form the conjugate acid of acetic acid. Step 2 A molecule of methanol, acting as a nucleophile, attacks the electrophilic carbon of the protonated acetic acid. Step 3 The oxonium ion the protonated intermediate formed in step 2 loses a proton to form a neutral tetrahedral intermediate. Step 4 The tetrahedral intermediate is protonated on one of its hydroxyl oxygens. Step 5 The hydroxyl protonated intermediate formed in step 4 loses a molecule of water to yield the protonated form of the ester.

Step 6 The loss of a proton from the protonated product of step 5 the conjugate acid of methyl acetate yields methyl acetate. In living organisms, the ether linkage occurs in biomolecules such as the carbohydrates. AMINES Amines are organic molecules that can be considered derivatives of ammonia NH3. Primary amines R—NH2 are molecules in which only one of the hydrogen atoms of ammonia has been replaced by an organic group e. Methylamine CH3NH2 is an example of a primary amine. In secondary amines, such as dimethylamine CH3—NH—CH3 , two hydrogens have been replaced by organic groups.

Tertiary amines such as triethylamine [ CH3CH2 3N] are molecules in which all three hydrogens have been replaced with organic groups. Amines with small organic groups are water-soluble, although the solubility of tertiary amines is limited because they do not have any hydrogen atoms bonded to the electronegative nitrogen atom. Like ammonia, amines are weak bases because of the lone pair of electrons on the nitrogen atom, which can accept a proton. Protonation of the nitrogen converts the amine into a cation. There are an enormous number of biomolecules that contain amine nitrogens. Examples include the amino acids components of proteins , the nitrogenous bases of the nucleic acids, and the alkaloids complex molecules produced by plants such as caffeine, morphine, and nicotine that have significant physiological effects on humans.

AMIDES Amides are amine derivatives of carboxylic acids with the general formula [R C O — NR2] where the R groups bonded to the nitrogen can be hydrogens or hydrocarbon groups. In contrast to the amines, amides are neutral molecules. As a result, amides are not weak bases i. In living organisms, the amide functional group is the linkage, referred to as the peptide bond, that connects amino acids in polypeptides. Amides are classified according to how many carbon atoms are bonded to the nitrogen atom. Molecules with the molecular formula R C O —NH2 are primary amides. Amides with two alkyl groups attached to the nitrogen are tertiary amides [R C O —NR2].

THIOLS A thiol is a molecule in which an sp3 carbon is bonded to a sulfhydryl group —SH. Although thiols are considered the sulfur analogues of alcohols, the low polarity of the SH bonds limits their capacity to form hydrogen bonds. As a result, thiols are not as soluble in water as their alcohol counterparts. Thiols are stronger acids than their alcohol equivalents, however, in part because of the weakness of the S—H bond. The sulfhydryl group of thiols is easily oxidized to form disulfides RS—SR. For example, two molecules of the amino acid cysteine react to form cystine, which contains a disulfide bond.

Better World Books. Uploaded by station cebu on August 31, Internet Archive logo A line drawing of the Internet Archive headquarters building façade. Search icon An illustration of a magnifying glass. User icon An illustration of a person's head and chest. Sign up Log in. Web icon An illustration of a computer application window Wayback Machine Texts icon An illustration of an open book. Books Video icon An illustration of two cells of a film strip. Video Audio icon An illustration of an audio speaker. Audio Software icon An illustration of a 3. Software Images icon An illustration of two photographs. Images Donate icon An illustration of a heart shape Donate Ellipses icon An illustration of text ellipses. Search Metadata Search text contents Search TV news captions Search archived websites Advanced Search. Biochemistry : the molecular basis of life Item Preview. remove-circle Share or Embed This Item. EMBED for wordpress. com hosted blogs and archive.

Want more? Advanced embedding details, examples, and help! xxi, pages : 29 cm Includes bibliographical references and index Biochemistry : An introduction -- Living cells -- Water : The medium of life -- Energy -- Amino acids, peptides, and proteins -- Enzymes -- Carbohydrates -- Carbohydrate metabolism -- Aerobic metabolism I : The citric acid cycle -- Aerobic metabolism II : Electron transport and oxidative phosphorylation -- Lipids and membranes -- Lipid metabolism -- Photosynthesis -- Nitrogen metabolism I : Synthesis -- Nitrogen metabolism II : Degradation -- Integration of metabolism -- Nucleic acids -- Genetic information -- Protein synthesis -- App. Access-restricted-item true Addeddate Associated-names McKee, James R. Full catalog record MARCXML. plus-circle Add Review. There are no reviews yet. Be the first one to write a review. download 1 file. Books for People with Print Disabilities. Internet Archive Books.

SIMILAR ITEMS based on metadata.

Biochemistry : the molecular basis of life,Item Preview

29/09/ · Biochemistry: The Molecular Basis of Life is a one-semester text focusing on the essential biochemical principles that underpin the modern life sciences. The sixth edition 24/01/ · Book Description Biochemistry: The Molecular Basis of Life is a one-semester text focusing on the essential biochemical principles that underpin the modern life sciences. Biochemistry: The Molecular Basis of Life is the ideal text for students who do not specialize in biochemistry but who require a strong grasp of biochemical principles. The goal of this 21) Metabolic Pathways and Energy Metabolism. Wilson Lab York University. Source: Wilson Lab York University. Read Download. Well, this was our collection of Biochemistry books in PDF Download Biochemistry: The Molecular Basis Of Life [PDF] Type: PDF. Size: MB. Download as PDF. Download Original PDF. This document was uploaded by user and they 23/07/ · Biochemistry the Molecular Basis of Life T. Mckee, J. R. McKee Published 23 July Education Biochemistry: The Molecular Basis of Life, Fourth Edition, is the ideal ... read more

For example, refer to pp. The science of chemical kinetics seeks to answer these and other questions about chemical reaction rates i. As with the alkanes, the cycloalkanes undergo combustion and halogenation reactions. SN2 reactions proceed in one step: as the nucleophile, functioning as a Lewis base an electron pair donor , donates its electron pair to an electrophilic carbon that has been polarized by an electronegative atom. Knowledge to Practical Applications Ewa Zymanczyk-Duda Source: Austin Publishing Group.

A single covalent bond consists of two shared electrons. Reactant structure. For this reason E1 reactions are similar to SN1 reactions because the reaction rates of both depend on one molecule, the precursor of the carbocation. There are two additional reaction classes, aliphatic substitution and elimination reactions, biochemistry the molecular basis of life pdf free download, which students should be familiar with. In the first step of an SN1 reaction, a planar carbocation forms as the leaving group a stable ion or a neutral molecule is displaced. Hydrocarbons Because hydrocarbon molecules contain only carbon and hydrogen, they are nonpolar. Remember me Forgot password?