Chapter 2: The Chemical Basis of Life PDF

Chapter 2 The Chemical Basis of Life Collage...

Chapter 2 The Chemical Basis of Life Collagen and Elastic Fibers Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.  Matter The “stuff” of the universe Anything that has mass and takes up space States of matter – Solid: has definite shape and volume – Liquid: has definite volume, changeable shape – Gas: has changeable shape and volume  Composition of Matter Elements: unique substances that cannot be broken down by ordinary chemical means Atoms: more-or-less identical building blocks for each element Atomic symbol: one- or two-letter chemical shorthand for each element  Properties of Elements Each element has unique physical and chemical properties – Physical properties: those detected with our senses – Chemical properties: pertain to the way atoms interact with one another Major Elements of the Human Body About 96% of body weight results from the following elements: Oxygen (O) Hydrogen (H) Carbon (C) Nitrogen (N) Lesser elements make up 3.9% of the body: Phosphorus (P) Potassium (K) Calcium (Ca) Sulfur (S) Sodium (Na) Chlorine (Cl) Magnesium (Mg) Iodine (I) Iron (Fe) Trace elements make up < 0.01% of the body: – Required in minute amounts – Found as part of enzymes  Atomic Structure The nucleus consists of neutrons and protons – Neutrons: No charge Mass= one atomic mass unit (amu) – Protons: Positive charge Mass of 1 amu Electrons are found orbiting the nucleus – Electrons: Negative charge Mass of 1/2000 amu Fig. 2.1  Atomic Structure Atoms are electrically neutral because the number of protons in atoms equals the number of electrons Fig. 2.2  Identification of Elements Atomic number: equal to the number of protons Mass number: equal to the mass of the protons and neutrons Atomic weight: average of the mass numbers of all isotopes Isotope: atoms with same number of protons but a different number of neutrons  Chemical Bonds Electron shells (energy levels) surround the nucleus of an atom – Bonds are formed using the electrons in the outermost electron shell Valence shell: outermost energy level containing chemically active electrons Octet rule: except for the first shell, which is full with two electrons, atoms interact in a manner to have eight electrons in their valence shell  Types of Chemical Bonds Ionic bond: formed when one atom loses an electron and another accepts that electron Covalent bond: the sharing of electrons Hydrogen bond: hydrogen atoms (bound covalently to either N or O atoms) have a small positive charge that is weakly attracted to the small negative charge of other atoms  Ionic Bonds Ionic bonds form between atoms by the transfer of one or more electrons Ions: charged atoms resulting from the gain or loss of electrons – Anions: negatively charged ions due to gaining one or more electrons – Cations: positively charged ions due to losing one or more electrons Ionic compounds form crystals instead of individual molecules Example: NaCl (sodium chloride) Fig. 2.3  Covalent Bonds Covalent bonds are formed by the sharing of two or more electrons Electron sharing produces molecules – single covalent bond: sharing of a pair of electrons (H — H) – double covalent bond: sharing of two pairs of electrons (O ═ C ═ O) Fig. 2.4  Polar and Nonpolar Molecules Electrons shared equally between atoms produce nonpolar molecules Unequal sharing of electrons produces polar molecules Fig. 2.5  Hydrogen Bonds Too weak to bind atoms together Common in dipoles such as water Responsible for surface tension in water Important as intramolecular bonds, giving the molecule a three-dimensional shape Hydrogen Bonds Figure 2.10a  Molecules and Compounds Molecule: two or more atoms held together by chemical bonds to form a structure that behaves as an independent unit Compound: two or more different kinds of atoms chemically combined – covalent compound: a molecule – ionic compound: organized array of ions  Dissociation Separation of ions in an ionic compound by polar water molecules – Dissociated ions are called electrolytes because they can conduct electricity – Molecules that do not dissociate in water are called nonelectrolytes Fig. 2.7  Chemical Reactions Occur when chemical bonds are formed, rearranged, or broken – Reactants: substances that enter a chemical reaction – Products: substances that result from the chemical reaction Written in symbolic form using chemical equations – Chemical equations contain: Number and type of reacting substances Products produced Relative amounts of reactants and products  Synthesis Reaction Combination of reactants to form a new larger product – Dehydration reaction: a synthesis reaction in which water is a product Fig. 2.8  Decomposition Reaction Breakdown of larger reactants into smaller products – Hydrolysis reaction: a decomposition reaction that uses water Fig. 2.8  Reversible Reactions All chemical reactions are theoretically reversible A + B → AB AB → A + B If neither a forward nor reverse reaction is dominant, chemical equilibrium is reached  Energy The capacity to do work (put matter into motion) Types of energy – Potential: stored (inactive) energy that could do work – Kinetic: energy that does work by causing the movement of an object  Energy Can be neither created nor destroyed Easily converted from one form to another – Mechanical: directly involved in moving matter – Chemical: stored in the bonds of chemical substances – Electrical: results from the movement of charged particles – Radiant or electromagnetic: travels in waves (i.e., visible light, ultraviolet light, and X-rays)  Energy Exists in chemical bonds as potential energy Released when the products contain less potential energy than the reactants – Energy can be “lost” as heat, can be used to synthesize molecules, or can do work. Absorbed in reactions when the products contain more potential energy than the reactants Adenosine Triphosphate (ATP) ATP stores and provides energy Source of immediately usable energy for the cell Fig. 2.9 Fig. 2.10 Factors Influencing Rate of Chemical Reactions Concentration: higher reacting particle concentrations produce faster reactions Temperature: chemical reactions proceed quicker at higher temperatures Catalysts: increase the rate of a reaction without being chemically changed – Enzymes are biological catalysts Particle size: the smaller the particle the faster the chemical reaction  Acids and Bases Acids release H+ and are therefore proton donors HCl → H+ + Cl – Bases release OH– and are proton acceptors NaOH → Na+ + OH– Acid-Base concentration is measured using a pH scale  pH Scale Ranges from 0 to 14 Indicates the H+ concentration of a solution – Neutral solutions have an equal number of H+ and OH– and a pH of 7.0 – Acidic solutions have more H+ than OH– and a pH of less than 7.0 – Basic (alkaline) solutions have fewer H+ than OH– and a pH greater than 7.0  pH Scale Neutral: pH 7.00 Acidic: pH 0–6.99 Basic: pH 7.01–14.00 Fig. 2.11  Acids and Bases Salts are formed by the reaction of an acid and a base HCl + NaOH → NaCl + H2O (acid) (base) (salt) (water) Buffers are chemicals that resist changes in pH when acids or bases are added – Example: Carbonic acid-bicarbonate system Carbonic acid dissociates, reversibly releasing bicarbonate ions and protons The chemical equilibrium between carbonic acid and bicarbonate resists pH changes in the blood  Buffers a) Addition of an acid to a nonbuffered solution results in an increase of H+ and a decrease in pH b) In a buffered solution the added H+ is bound by the buffer and the pH change is much smaller Fig. 2.12  Biochemistry Inorganic chemistry – Mostly concerned with non-carbon-containing substances but does include such carbon-containing substances as CO, CO2, and HCO3- Organic chemistry – Substances contain carbon, are covalently bonded, and are often large – Usually have carbon-carbon or carbon-hydrogen bonding  Inorganic Compounds Oxygen (O2) is involved with the extraction of energy from food molecules to make ATP Carbon Dioxide (CO2) is a by-product of the breakdown of food molecules Water (H2O) has many important properties for living organisms and is essential for life  Properties of Water Stabilizes body temperature – The high heat capacity of water allows it to absorb and release large amounts of heat before changing temperature Protection – acts as a lubricant or cushion Chemical reactions – Most of the chemical reactions necessary for life do not take place unless the reacting molecules are dissolved in water – Water also directly participates in many chemical reactions Transport – Polar solvent properties: dissolves ionic substances, forms hydration layers around large charged molecules, and serves as the body’s major transport medium  Organic Compounds Molecules unique to living systems They include: – Carbohydrates – Lipids – Proteins – Nucleic Acids  Carbohydrates Contain carbon, hydrogen, and oxygen – Ratio of 1:2:1 (C:H:O) Their major function is to supply a source of cellular food Examples – Monosaccharides – glucose and fructose – Disaccharides – sucrose and lactose – Polysaccharides – starch and glycogen Figure 2.14a Fig. 2.13  Lipids Dissolve in nonpolar solvents, such as alcohol or acetone, but not in polar solvents, such as water Contain C, H, and O, but the proportion of oxygen in lipids is less than in carbohydrates Examples: – Fats or triglycerides: energy – Phospholipids: structural components of cell membranes – Eicosanoids: regulate physiological processes – Steroids: regulate physiological processes  Examples of Lipids Found in the Body Fats: found in subcutaneous tissue and around organs Phospholipids: chief component of cell membranes Steroids: cholesterol, bile salts, vitamin D, sex hormones, and adrenal cortical hormones Eicosanoids: prostaglandins, leukotrienes, and thromboxanes Fat-soluble vitamins: vitamins A, D, E, and K Lipoproteins: transport fatty acids and cholesterol in the bloodstream  Fats (Triglycerides) Composed of three fatty acids bonded to a glycerol molecule Fig. 2.14  Fatty Acids Saturated: only single covalent bonds between carbons Unsaturated: one or more double covalent bonds between carbons Fig. 2.15  Other Lipids Phospholipids: modified triglycerides with two fatty acid groups and a phosphorus group Fig. 2.16  Other Lipids Eicosanoids: 20-carbon fatty acids found in cell membranes Steroids: flat molecules with four interlocking hydrocarbon rings Fig. 2.17  Proteins Macromolecules Contain C, H, O, N, and some S Composed of 20 basic types of amino acids bound together with peptide bonds – Dipeptide: Two amino acids – Tripeptide: Three amino acids – Polypeptide: Many amino acids Proteins are polypeptides of hundreds of amino acids  Amino Acids (AA) Building blocks of proteins Organic acids containing – amino group (-NH2) – a carboxyl group (COOH) – a hydrogen atom – a side chain designated by the symbol R attached to the same carbon atom as the hydrogen  Structural Levels of Proteins Primary: determined by the number, kind, and arrangement of amino acids Secondary: results from folding or bending of the polypeptide chain caused by the hydrogen bonds between amino acids (helices and pleated sheets) Tertiary: results from the folding of the helices or pleated sheets and the hydrogen bonds formed with water Quaternary: spatial relationships between two or more proteins that associate to form a functional unit  Fig. 2.19ab  Fig. 2.19cd  Proteins Functions – regulate chemical reactions (enzymes) – structural proteins provide the framework for many of the body’s tissues – responsible for muscle contraction – Fibrous proteins Extended and strand-like proteins Examples: keratin, elastin, collagen, and certain contractile fibers – Globular proteins Compact, spherical proteins with tertiary and quaternary structures Examples: antibodies, hormones, and enzymes Denaturation – Disruption of hydrogen bonds, which changes the shape of proteins and makes them nonfunctional  Characteristics of Enzymes Speed up chemical reactions by lowering the activation energy Most are globular proteins that act as biological catalysts Are chemically specific Frequently named for the type of reaction they catalyze Names usually end in -ase Chemical events of the body are regulated primarily by mechanisms that control – concentration of enzymes – activity of enzymes Fig. 2.20  Enzymes Enzymes bind to reactants according to the lock-and-key model – The shape of both the enzyme and reactants are critical to the function of the enzyme – By bringing the two reactants close to each other it reduces the activation energy for the reaction – Each enzyme catalyzes only one type of chemical reaction – After each reaction the enzyme is released and can be used again Fig. 2.21  Nucleic Acids Composed of C, O, H, N, and P The basic unit of nucleic acids is the nucleotide, which is a monosaccharide with an attached phosphate and organic base Five nitrogenous bases contribute to nucleotide structure: – adenine (A) – guanine (G) – cytosine (C) – thymine (T) – uracil (U) Two major classes: DNA and RNA  Deoxyribonucleic Acid (DNA) Double-stranded helical molecule found in the nucleus of the cell Genetic material of the cell Replicates itself before the cell divides, ensuring genetic continuity Provides instructions for protein synthesis Contains the monosaccharide deoxyribose and the organic bases – adenine – thymine – guanine – cytosine Fig. 2.22  Ribonucleic Acid (RNA) Single-stranded molecule found in both the nucleus and the cytoplasm of a cell Composed of the monosaccharide ribose and uses the organic base uracil instead of thymine Three varieties of RNA: – messenger RNA – transfer RNA – ribosomal RNA