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2 Chemical Level of Organization PowerPoint® Lecture Presentations prepared by Alexander G. Cheroske Mesa Community College at Red Mountain © 2011 Pearson Education, Inc. Section 1: Atoms and Molecules • Learning Outcomes • 2.1 Describe an atom and how atomic structure affects the mass number and atomic weight of the various chemical elements. • 2.2 Explain the relationship between electrons and energy levels. • 2.3 Compare the ways in which atoms combine to form molecules and compounds. • 2.4 Describe the three states of matter and the importance of hydrogen bonds in liquid water. © 2011 Pearson Education, Inc. Section 1: Atoms and Molecules • Studying the human body begins at the chemical level of organization • Chemistry studies the structure of matter • Matter (anything that takes up space and has mass) • Mass (the amount of material in matter) • On Earth, mass is the same as weight © 2011 Pearson Education, Inc. Section 1: Atoms and Molecules • Atoms (smallest stable units of matter) • Composed of subatomic particles • Protons (p+, have a positive electrical charge) • Neutrons (n or n0, are electrically neutral or uncharged) • Electrons (e–, much smaller and have a negative electrical charge) © 2011 Pearson Education, Inc. An atom and its important subatomic particles Protons (p+) have a positive electrical charge. Neutrons (n or n0) are electrically neutral, which means they are uncharged. Electrons (e-) are much smaller and about 1/1800th the mass of either protons or neutrons. They bear a negative electrical charge. Figure 2 Section 1 © 2011 Pearson Education, Inc. 1 Section 1: Atoms and Molecules • Atoms (continued) • Can be subdivided into: • Nucleus • Center of atom • Contains protons and neutrons • Electron cloud • Electrons whirl around the nucleus creating a cloud © 2011 Pearson Education, Inc. The nucleus and electron cloud of an atom The nucleus of an atom lies at its center. The nucleus contains one or more protons and it may contain neutrons as well. The mass of the atom is primarily determined by the numbers of protons and neutrons in the nucleus. The electrons in the atom whirl around the nucleus, creating an electron cloud. Figure 2 Section 1 © 2011 Pearson Education, Inc. 2 Section 1: Atoms and Molecules • Molecule (atoms interact to produce larger, more complex structures) • Everything is made up of varying combinations of atoms • Unique characteristics of an object result from the particular atoms and how they interact • Mass of any object represents the sum of masses of its atoms © 2011 Pearson Education, Inc. A molecule, which consists of two or more atoms Figure 2 Section 1 © 2011 Pearson Education, Inc. 3 Module 2.1: Atoms • Atoms • Normally contain equal numbers of protons and electrons • Atomic number (number of protons) • Mass number (number of protons + number of neutrons) • Element (pure substance of only atoms with the same atomic number) © 2011 Pearson Education, Inc. Module 2.1: Atoms • Electron cloud • Electrons whirl around the nucleus • Are attracted to positively charged protons in nucleus • Example of an electrical force • Electron shell • Two-dimensional representation of the electron cloud • Outermost shell represents the surface of the electron cloud © 2011 Pearson Education, Inc. Module 2.1: Atoms • Isotopes • Atoms of a single element but with different numbers of neutrons • Distinguishable based on mass © 2011 Pearson Education, Inc. Electron shell models of three isotopes of hydrogen Electron shell Hydrogen-1 mass number: 1 Hydrogen-2, deuterium mass number: 2 Hydrogen-3, tritium mass number: 3 Figure 2.1 © 2011 Pearson Education, Inc. 2 Module 2.1: Atoms • Atomic weight (actual mass of an atom) • Expressed in daltons (also known as the atomic mass unit or amu) • One dalton is very close to the weight of one proton or one neutron • Is an average mass number reflecting the proportion of different isotopes • Example: hydrogen • Atomic number = 1 • Atomic weight = 1.0079 © 2011 Pearson Education, Inc. A representation of an element’s atomic weight, which reflects the proportions of the various isotopes of that element Atomic weight of a mixture of Atomic weight of hydrogen-1 = 1 hydrogen isotopes = 1.0079 Figure 2.1 © 2011 Pearson Education, Inc. 3 Module 2.1: Atoms • Principal elements • Thirteen most abundant elements to total body weight • Trace elements • Fourteen other elements in the body in very small amounts • Each element has a chemical symbol • Most are connected with English names • Example: O for oxygen • A few are from other languages • Example: Na for sodium, from Latin natrium © 2011 Pearson Education, Inc. Figure 2.1 © 2011 Pearson Education, Inc. 4 Figure 2.1 © 2011 Pearson Education, Inc. 4 Figure 2.1 © 2011 Pearson Education, Inc. 4 Figure 2.1 © 2011 Pearson Education, Inc. 4 Module 2.1 Review a. Define atom. b. Describe trace elements. c. How is it possible for two samples of hydrogen to contain the same number of atoms yet have different weights? © 2011 Pearson Education, Inc. Module 2.2: Electrons • Electrons occupy a series of concentric energy levels or electron shells • Reactive elements have unfilled outer energy levels • Will react with other elements to fill outer shell • Once outer shell is full, the atom is stable and unreactive • Examples: hydrogen and helium • Inert elements have full outermost energy levels • Do not participate in chemical processes • Examples: helium and neon © 2011 Pearson Education, Inc. Module 2.2: Electrons • Energy levels (electron shells) • First level: holds 2 electrons • Second level: holds 8 electrons • Third level: holds 8 electrons © 2011 Pearson Education, Inc. Reactive elements Two examples of reactive elements, which have unfilled outermost energy levels Inert elements Two examples of inert elements, which do not react in chemical processes because their outermost energy levels are filled The first energy level can hold a maximum of two electrons. Hydrogen has one electron in the first energy level. Helium has two electrons in the first energy level. Hydrogen, H Atomic number: 1 Mass number: 1 1 electron Lithium has one electron in the second energy level; it is extremely reactive. The second and third energy levels can each contain up to eight electrons. Lithium, Li Atomic number: 3 Mass number: 6 (3 protons + 3 neutrons) 3 electrons © 2011 Pearson Education, Inc. Helium, He Atomic number: 2 Mass number: 4 (2 protons + 2 neutrons) 2 electrons Neon, Ne Atomic number: 10 Mass number: 20 (10 protons + 10 neutrons) 10 electrons Neon has eight electrons in the Second energy level; it does not React with other atoms. Figure 2.2 1 - 2 Module 2.2: Electrons • Reactive elements achieve stability by gaining, losing, or sharing electrons Cation – Loss of electron, giving atom a positive charge • • Loss of 1 electron gives the ion a +1 charge • Some ions carry charges of +2, +3, or +4 Anion – Gain of electron, giving atom a negative charge • • • Gain of 1 electron gives the ion a –1 charge • Some ions carry charges of –2, –3, or –4 Stabilizing interactions often form chemical bonds © 2011 Pearson Education, Inc. The loss of an electron from a sodium atom, forming a sodium cation Sodium atom, Na (reactive) Sodium ion, Na+ (stable) Figure 2.2 © 2011 Pearson Education, Inc. 3 The gain of an electron by a chlorine atom, forming a chloride anion Chlorine atom, Cl (reactive) Chloride ion, Cl– (stable) Figure 2.2 © 2011 Pearson Education, Inc. 4 Module 2.2 Review a. Indicate the maximum number of electrons that can occupy each of the first three electron shells (energy levels) of an atom. b. Explain why the atoms of inert elements do not react with one another or combine with atoms of other elements. c. Explain how cations and anions form. © 2011 Pearson Education, Inc. Module 2.3: Chemical bonds • Chemical bonding creates • Molecules (substances with two or more atoms) • Compounds (substances with two or more different atoms) © 2011 Pearson Education, Inc. Module 2.3: Chemical bonds • Ionic bonds • Created by electrical attraction between cations and anions • Involves the transfer of electrons from one atom to another to achieve stability © 2011 Pearson Education, Inc. Step 1: Formation of sodium and chloride ions. The sodium atom loses an electron to the chloride atom. This produces two stable ions with filled outer energy levels. Step 2: Formation of an ionic bond. Because these ions form close together, and have opposite charges, they are attracted to one another. This creates NaCl, an ionic compound. Sodium atom Sodium ion (Na+) Sodium chloride (NaCl) Chlorine atom Chloride ion (CI–) The formation of the compound sodium chloride (table salt), involving the formation of ions and ionic bonding Chloride ions (CI–) Sodium ions (Na+) A crystal of sodium chloride © 2011 Pearson Education, Inc. Figure 2.3 1 - 2 Module 2.3: Chemical bonds • Covalent bonds • Sharing of electrons between atoms • Single covalent bond • • One electron contributed by each atom Double covalent bond • Two electrons contributed by each atom © 2011 Pearson Education, Inc. Molecules, which consist of one or more elements bound by covalent bonds resulting from the sharing of electrons Molecule Description Hydrogen (H2) Hydrogen atoms aren’t found as individuals—they exist as molecules, each containing a pair of hydrogen atoms. The two atoms share their electrons to fill their outer energy levels, and the electron pair orbits both nuclei. One electron is contributed by each atom, so this is called a single covalent bond. Oxygen (O2) Carbon dioxide (CO2) An oxygen atom has 6 electrons in its outer energy level. By forming a double covalent bond with another oxygen atom, an oxygen molecule is created with a stable outer energy level. A carbon atom has 4 electrons in its outer energy level, so it needs to gain 4 from other atoms to achieve stability. In a molecule of carbon dioxide, a carbon atom shares a pair of electrons with each of two oxygen atoms and forms two double covalent bonds. Figure 2.3 © 2011 Pearson Education, Inc. 3 Module 2.3: Chemical bonds • Nonpolar molecule • Electrons shared equally between atoms • No electrical charge on the molecule © 2011 Pearson Education, Inc. Space-filling models of two nonpolar molecules Oxygen (O2) Carbon dioxide (CO2) Figure 2.3 © 2011 Pearson Education, Inc. 4 Module 2.3: Chemical bonds • Polar molecule • Unequal sharing of electrons between atoms • Formed by polar covalent bonds • Example: water molecule (H2O) • Oxygen exerts a stronger attraction for the electrons than the hydrogens • Oxygen carries a slightly negative charge • Hydrogens carry a slightly positive charge © 2011 Pearson Education, Inc. A water molecule, which is a polar molecule (has positively and negatively charged portions) that contains polar covalent bonds Hydrogen atom Positive pole Oxygen atom Negative pole Figure 2.3 © 2011 Pearson Education, Inc. 6 Module 2.3 Review a. Name and distinguish between the two most common types of chemical bonds. b. Describe the kind of bonds that hold the atoms in a water molecule together. c. Relate why we can apply the term molecule to the smallest particle of water but not to that of table salt. © 2011 Pearson Education, Inc. Module 2.4: Three states of matter • Matter exists in one of three states: 1. Solid (particles held tightly together) • Maintain their volume and shape at ordinary temperatures and pressures 2. Liquid (particles held less tightly together) • • Has a constant volume Container determines shape 3. Gas (particles independent of each other) • • • Has neither a constant volume nor a fixed shape Can be compressed or expanded Will fill a container of any size © 2011 Pearson Education, Inc. Figure 2.4 © 2011 Pearson Education, Inc. 1 - 3 Module 2.4: Three states of matter • Water • Only substance that exists in all three states at temperatures compatible with life • Solid (ice) • Liquid (water) • • Exists over a broad range of temperatures due to interactions among the polar water molecules Gas (water vapor) © 2011 Pearson Education, Inc. Figure 2.4 © 2011 Pearson Education, Inc. 4 Module 2.4: Three states of matter • Hydrogen bond • Attraction of the small polar charges on hydrogen atoms to negative charges on other polar molecules • Can change shape of molecules or pull molecules together • Produces surface tension • Slows the rate of evaporation at water surface © 2011 Pearson Education, Inc. The attraction of hydrogen bonds, which occur between the small positive and negative charges on different water molecules KEY Hydrogen Oxygen Hydrogen bond Figure 2.4 © 2011 Pearson Education, Inc. 5 Figure 2.4 © 2011 Pearson Education, Inc. 6 Module 2.4: Three states of matter • Water can dissolve a variety of inorganic compounds due to its polar molecules • Seawater contains almost all naturally occurring elements • Our body fluids contain at least 29 elements © 2011 Pearson Education, Inc. Module 2.4 Review a. Describe the different states of matter in terms of shape and volume. b. By what means are water molecules attracted to each other? c. Explain why small insects can walk on the surface of a pond, and tears protect the surface of the eye from dust particles. © 2011 Pearson Education, Inc. Section 2: Chemical Reactions • Learning Outcomes • 2.5 Use chemical notation to symbolize chemical reactions. • 2.6 Distinguish among the major types of chemical reactions that are important for studying physiology. • 2.7 Describe the crucial role of enzymes in metabolism. © 2011 Pearson Education, Inc. Section 2: Chemical Reactions • Cells remain alive by controlling chemical reactions • New bonds may form between atoms • Existing bonds may be broken • Reactants (reacting substances) form different products • Metabolism • All the cellular reactions of the body at a given moment © 2011 Pearson Education, Inc. Section 2: Chemical Reactions • Each cell is a “chemical factory” where chemical reactions control: • Growth • Maintenance • Repair • Secretion • Contraction • Energy transfer and use © 2011 Pearson Education, Inc. Figure 2 Section 2 © 2011 Pearson Education, Inc. 1 Section 2: Chemical Reactions • Work • Movement of an object or change in physical structure of matter • Examples: • Walking or running • Synthesis of molecules • Conversion of liquid water to vapor (evaporation) © 2011 Pearson Education, Inc. Section 2: Chemical Reactions • Energy (capacity to perform work) • Kinetic energy • Energy of motion • Example: skeletal muscle contracting • Potential energy • Stored energy (potential to do work) • Example: skeletal muscle at rest • Conversion of energy is never 100% efficient • Some energy is released as heat • Example: body temperature rises as muscles contract © 2011 Pearson Education, Inc. Module 2.5: Chemical notation • Chemical notation • A simple “chemical shorthand” to describe chemical structures and events © 2011 Pearson Education, Inc. Module 2.5: Chemical notation • Symbol of element indicates one atom of that element • H = one atom of hydrogen • O = one atom of oxygen • Number preceding symbol indicates more than one atom of that element • 2 H = two atoms of hydrogen • 2 O = two atoms of oxygen © 2011 Pearson Education, Inc. Module 2.5: Chemical notation • Subscript following symbol indicates a molecule with that number of atoms of that element • H2 = hydrogen molecule, composed of two hydrogen atoms • O2 = oxygen molecule, composed of two oxygen atoms • H2O = water molecules, composed of two hydrogen atoms and one oxygen atom © 2011 Pearson Education, Inc. Module 2.5: Chemical notation • Chemical reaction • Reactants • Participants at reaction start • Usually on the left • Products • Generated at end of reaction • Usually on the right • 2 H + O H 2O © 2011 Pearson Education, Inc. Module 2.5: Chemical notation • A superscript plus or minus following symbol indicates an ion • Single plus sign indicates cation with +1 • Atom has lost one electron • Example: Na+ • Single minus sign indicates anion with -1 • Atom has gained one electron • Example: Cl- • Number before sign indicates more than one electron lost or gained • Example Ca2+ © 2011 Pearson Education, Inc. Module 2.5: Chemical notation • Chemical reactions rearrange atoms into new combinations • Numbers of atoms must be the same one both sides of equation (= balanced) • Unbalanced: H2 + O2 H2O • Balanced: 2 H2 + O2 2 H2O © 2011 Pearson Education, Inc. Module 2.5: Chemical notation • Mole (mol) • Quantity (in grams) equal to an element’s atomic weight • One mole of a given element always contains the same number of atoms of a mole of another element • Weight differs • Example: • 1 mole of oxygen • 16 grams • 1 mole of hydrogen • 1 gram © 2011 Pearson Education, Inc. A mole, a quantity of an element with a weight (in grams) equal to that element’s atomic weight 1 mole of oxygen 1 mole of hydrogen Figure 2.5 © 2011 Pearson Education, Inc. 2 Module 2.5: Chemical notation • Molecular weight • Sum of the atomic weights of all atoms comprising a molecule • Can be used to calculate quantities of reactants or products for a chemical reaction © 2011 Pearson Education, Inc. Atomic weight of H = 1 Molecular weight of H2 = 2 Two moles of H2 weigh 4 g Atomic weight of O = 16 Molecular weight of O2 = 32 One mole of O2 weighs 32 g Molecular weight of 2 H2O: 2 x (2 + 16) 2 x 18 = 36 Two moles of H2O weigh 36 g The use of molecular weights to determine the amounts of reactants needed and the amount of product generated Figure 2.5 © 2011 Pearson Education, Inc. 4 Module 2.5 Review a. The chemical shorthand used to describe chemical compounds and reactions effectively is known as __________________. b. Using the rules for chemical notation,write the molecular formula for glucose, a compound composed of 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms. c. Calculate the weight of one mole of glucose. (The atomic weight of carbon = 12 grams.) © 2011 Pearson Education, Inc. Module 2.6: Types of chemical reactions • Types of chemical reactions 1. Decomposition • Breaks molecule into smaller components • • • AB A + B Example: digestion of food for absorption Hydrolysis • Decomposition using the insertion of water • AB + H2O A-H + OH-B • Catabolism (katabole, a throwing down) • • Collective decomposition reactions of the body Releases kinetic energy • CD C + D + ENERGY © 2011 Pearson Education, Inc. Module 2.6: Types of chemical reactions • Types of chemical reactions (continued) 2. Synthesis • Combining of atoms or smaller molecules into larger molecules • • • Always involves new bond formation Dehydration synthesis (condensation) • • • A + B AB Removal of water to form a larger molecule A-H + OH-B A-B + H2O Anabolism (anabole, a throwing upward) • • Collective synthesis of new molecules in the body Requires energy © 2011 Pearson Education, Inc. Module 2.6: Types of chemical reactions • Many important biological reactions are reversible • A + B ↔ AB • Equilibrium • Rates at which the two reactions occur are in balance © 2011 Pearson Education, Inc. Module 2.6: Types of chemical reactions • Types of chemical reactions (continued) 3. Exchange reactions • Reactants are “shuffled” to produce new products • • AB + CD AD + CB Reactants must break apart (decomposition) in order to form products (synthesis) © 2011 Pearson Education, Inc. Module 2.6 Review a. Identify and describe three types of chemical reactions important in human physiology. b. Distinguish the roles of water in hydrolysis and dehydration synthesis reactions. c. In cells, glucose, a six-carbon molecule, is converted into two three-carbon molecules by a reaction that releases energy. What is the source of the energy? © 2011 Pearson Education, Inc. Module 2.7: Enzymes • Most chemical reactions require energy to activate reactants • Activation energy • • Amount of energy to start a reaction Enzymes • Special proteins that lower the activation energy for a reaction © 2011 Pearson Education, Inc. Module 2.7: Enzymes • Catalysts (katalysis, dissolution) • Compounds that accelerate reactions without changing themselves • Reactions continue until equilibrium is reached • Enzymes perform as catalysts © 2011 Pearson Education, Inc. The action of protein catalysts called enzymes in promoting a chemical reaction by lowering its activation energy Activation energy Progress of reaction Specific enzymes lower the activation energy requirements so that important reactions will occur. Energy Energy In the external environment, extreme conditions can provide the activation energy. For example, complex sugars can the broken down in a laboratory by boiling them in an acidic solution. Activation energy Progress of reaction Figure 2.7 © 2011 Pearson Education, Inc. 1 Module 2.7: Enzymes • Metabolic pathway • Series of reactions (each involving an enzyme) that support life • May absorb or release energy on completion • • Exergonic (exo-, outside) • Release energy • Common in the body and help to maintain body temperature Endergonic (endo-, inside) • Absorb energy © 2011 Pearson Education, Inc. A metabolic pathway, a series of interlocking steps, each controlled by a specific enzyme Figure 2.7 © 2011 Pearson Education, Inc. 2 Module 2.7: Enzymes • Metabolites (metabole, change) • All molecules that can be synthesized or decomposed in our bodies Processed by enzymatic reactions Nutrients • • • • Essential metabolites normally obtained from our diet Can be classified as: • • Organic • Always contain carbon and hydrogen • Examples: sugars, fats, proteins Inorganic • Generally do not contain carbon and hydrogen • Examples: carbon dioxide, water, salts © 2011 Pearson Education, Inc. Module 2.7 Review a. What is an enzyme? b. Explain the differences between metabolites and nutrients. c. Why are enzymes needed in our cells? © 2011 Pearson Education, Inc. Section 3: The Importance of Water in the Body • Learning Outcomes • 2.8 Explain how the chemical properties of water affect the solubility of inorganic and organic molecules. • 2.9 Discuss the importance of pH and the role of buffers in body fluids. © 2011 Pearson Education, Inc. Section 3: The Importance of Water in the Body • Water • Most important constituent of body • Accounts for up to 2/3 of total body weight • Changes in body water content can affect all systems © 2011 Pearson Education, Inc. Important Properties of Water Lubrication Water is an effective lubricant because there is little friction between water molecules. Thus even a thin layer of water between two opposing surfaces will greatly reduce friction between them; water reduces friction within joints and in body cavities. Reactivity In our bodies, chemical reactions occur in water, and water molecules are also participants in some reactions, including hydrolysis and dehydration synthesis. Dehydration synthesis Hydrolysis High heat capacity Heat capacity is the ability to absorb and retain heat. Water has an unusually high heat capacity, because water molecules in the liquid state are attracted to one another through hydrogen bonding. • The temperature of water must be high before individual molecules have enough energy to break free to become water vapor, a gas. • Water carries a great deal of heat away with it when it finally does change from a liquid to a gas. This feature accounts for the cooling effect of perspiration on the skin. • A large mass of water changes temperature very slowly. This property is called thermal inertia. Solubility A remarkable number of inorganic and organic molecules will dissolve in water. The individual particles become dispersed within the water, and the result is a solution—a uniform mixture of two or more substances. The medium in which other atoms, ions, or molecules are dispersed is called the solvent; the dispersed substances are the solutes. In aqueous solutions, water is the solvent. Figure 2 Section 3 © 2011 Pearson Education, Inc. Module 2.8: Physiological systems depend on water • Many inorganic compounds have ionic bonds • Ionization or dissociation • Ionic bonds broken in water due to the polarity of the water molecule • Anions surrounded by positive poles of water molecules • Cations surrounded by negative poles of water molecules • Hydration sphere • Sheath of water molecules around an ion © 2011 Pearson Education, Inc. Module 2.8: Physiological systems depend on water • Organic molecules may dissolve in water if they contain polar covalent bonds • = Hydrophilic (hydro-, water + philos, loving) © 2011 Pearson Education, Inc. The roles of hydration spheres in the ionization or dissociation of an ionic compound (center) and in solutions of an organic molecule containing polar covalent bonds (right) Sodium chloride crystal Negative pole Hydration spheres H Positive pole Glucose molecule Cl– Na+ NaCl in solution Glucose molecule in solution Figure 2.1 © 2011 Pearson Education, Inc. 1 - 3 Module 2.8: Physiological systems depend on water • Electrolytes • Ions in solution that conduct electric current • Example: NaCl Na+ + Cl– • Movement of ions across the plasma membrane of cells is important in many body cells • Muscle • Nervous • Ion concentrations in bodily fluids is carefully regulated © 2011 Pearson Education, Inc. The movement of sodium and chloride ions, dissociated from the electrolyte sodium chloride, when subjected to an electric field Cl– Na+ Figure 2.8 © 2011 Pearson Education, Inc. 4 Figure 2.8 © 2011 Pearson Education, Inc. 5 Fats and oils Protein Many organic molecules either lack polar covalent bonds or have very few. Such molecules do not have positive and negative poles and are said to be nonpolar. When nonpolar molecules are exposed to water, hydration spheres do not form and the molecules do not dissolve. Molecules that do not readily interact with water are called hydrophobic (hi-drō-FŌ-bik; hydro-, water + phobos, fear). Among the most familiar hydrophobic molecules are fats and oils of all kinds. Body fluids typically contain large and complex organic molecules, such as proteins, that are held in solution by their association with water molecules. A colloid, a solution containing dispersed proteins or other large molecules Figure 2.8 © 2011 Pearson Education, Inc. 6 Module 2.8: Physiological systems depend on water • Many organic molecules lack or have very few polar covalent bonds (nonpolar) • = Hydrophobic (hydro-, water + phobos, fear) • Do not readily interact with water • Examples: fats and oils © 2011 Pearson Education, Inc. Module 2.8: Physiological systems depend on water • Large organic molecules are usually held in solution by water molecules • Example: proteins • Colloid • Large molecules that stay in solution indefinitely • Example: liquid Jell-O • Suspension • Larger molecules that may settle out of solution due to gravity • Example: blood © 2011 Pearson Education, Inc. Module 2.8 Review a. Define electrolytes. b. Distinguish between hydrophilic and hydrophobic molecules. c. Explain how the ionic compound sodium chloride dissolves in water. © 2011 Pearson Education, Inc. Module 2.9: pH • Hydrogen ion (H+) • Hydrogen atom that has lost electron • Extremely reactive in solution • Concentration in body regulated precisely • Hydroxide ion (OH–) • Produced when water decomposes (along with H+) © 2011 Pearson Education, Inc. The dissociation of a water molecule into a hydrogen ion, H+, and a hydroxide ion, OH– Figure 2.9 © 2011 Pearson Education, Inc. 1 Module 2.9: pH • pH • Negative logarithm of the hydrogen ion concentration (mol/L) • Ranges from 0–14 • Acidic: below 7 • Neutral: equal to 7 • Alkaline: above 7 © 2011 Pearson Education, Inc. The pH—the negative logarithm of the hydrogen ion concentration—of a variety of common substances Blood The pH of blood normally ranges form 7.35 to 7.45. Abnormal fluctuations in pH can damage cells and tissues by breaking chemical bonds, changing the shapes of proteins, and altering cellular functions. Acidosis is an abnormal physiological state caused by low blood pH (below 7.35); a pH below 7 can produce coma. Alkalosis results from an abnormally high pH (above 7.45); a blood pH above 7.8 generally causes uncontrollable and sustained skeletal muscle contractions. 1 mol/L hydrochloric acid Urine Beer, vinegar, wine, Tomatoes, pickles grapes Stomach hydrochloric acid Extremely acidic pH 0 [H+] 100 (mol/L)* 1 mol/L sodium hydroxide 1 10-1 Saliva, milk Pure water Increasing concentration of H+ Neutral 2 10-2 3 10-3 4 10-4 A solution with a pH below 7 is acidic (a-SI-dik), meaning that it contains more hydrogen ions than hydroxide ions. 5 10-5 6 10-6 7 10-7 Seawater Eggs Household bleach Household ammonia Increasing concentration of OH– 8 10-8 A solution with a pH of 7 is said to be neutral, because it contains equal numbers of hydrogen and hydroxide ions. 9 10-9 10 10-10 11 10-11 Extremely basic 12 10-12 13 10-13 14 10-14 A pH above 7 is basic, or alkaline (AL-kuh-lin), meaning that it has more hydroxide ions than hydrogen ions. * One liter of pure water contains about 0.0000001 mol of hydrogen ions and an equal number of hydroxide ions. In other words, the concentration of hydrogen ions in a solution of pure water is 0.0000001 mol per liter. This can be written as [H+] = 10–7 mol/L. The brackets around the H+ signify “the concentration of,” another example of chemical notation. © 2011 Pearson Education, Inc. Oven cleaner Figure 2.9 2 Module 2.9: pH • Blood pH • Normally ranges from 7.35–7.45 • Values above or below can alter cell functions • Acidosis: below 7.35 • Alkalosis: above 7.45 © 2011 Pearson Education, Inc. Module 2.9: pH • Acid • Solute that dissociates and releases hydrogen ions • Often referred to as proton donors • After losing an electron, a hydrogen ion consists solely as a proton • Strong acids dissociate completely • Example: hydrochloric acid (HCl) • HCl H+ + Cl– © 2011 Pearson Education, Inc. Hydrochloric acid (HCl), a strong acid because it releases hydrogen ions Figure 2 .9 3 © 2011 Pearson Education, Inc. Module 2.9: pH • Base • Solute that removes hydrogen ions from solution • Proton acceptor • Strong bases dissociate completely • Example: sodium hydroxide (NaOH) • NaOH Na+ + OH– © 2011 Pearson Education, Inc. Sodium hydroxide (NaOH), a strong base because it releases hydroxide ions that remove hydrogen ions from solution Figure 2 .9 4 © 2011 Pearson Education, Inc. Module 2.9: pH • Weak acids and bases fail to dissociate completely • Example: carbonic acid (H2CO3) • H2CO3 ↔ H+ + HCO3– • In body fluids, carbonic acid reversibly dissociates to hydrogen ion and bicarbonate ion © 2011 Pearson Education, Inc. Carbonic acid, a weak acid in body fluids that reversibly dissociates into hydrogen ion and a bicarbonate ion Figure 2.9 © 2011 Pearson Education, Inc. 5 Module 2.9: pH • Salt • Ionic compound with any cation and anion except: • H+ • OH– • Example: NaCl • NaCl Na+ + Cl– © 2011 Pearson Education, Inc. NaCl, a salt—an ionic compound consisting of any cation except a hydrogen ion and any anion except a hydroxide ion. Figure 2.9 © 2011 Pearson Education, Inc. 6 Module 2.9: pH • Buffers • Compounds that stabilize pH by removing or replacing hydrogen ions • Help to maintain normal pH of body fluids • Buffer systems • Usually involve a weak acid and its related salt • Examples: • Carbonic acid (H2CO3) • Sodium bicarbonate (NaHCO3) © 2011 Pearson Education, Inc. Module 2.9 Review a. Define pH. b. Explain the differences among an acid, a base, and a salt. c. What is the significance of pH in physiological systems? © 2011 Pearson Education, Inc. Section 4: Organic Compounds • Learning Outcomes • 2.10 Discuss the structures and functions of carbohydrates. • 2.11 Discuss the structures and functions of lipids. • 2.12 Discuss the structures and diverse functions of eicosanoids, steroids, phospholipids, and glycolipids. • 2.13 Discuss protein structure and the essential functions of proteins within the body. © 2011 Pearson Education, Inc. Section 4: Organic Compounds • Learning Outcomes • 2.14 Explain how enzymes function within the body. • 2.15 Discuss the structure and function of highenergy compounds. • 2.16 Compare and contrast the structures and functions of DNA and RNA. © 2011 Pearson Education, Inc. Section 4: Organic Compounds • Organic compounds • Always contain carbon and hydrogen and generally oxygen • Many are long chains of carbon linked with covalent bonds • Many are soluble in water • Functional groups • Attached groupings of atoms that occur commonly in many organic molecules • Influence the properties of the overall molecule • Many allow cells to transfer and capture energy as high-energy compounds © 2011 Pearson Education, Inc. Important Functional Groups of Organic Compounds Functional Group Structural Formula* Importance Examples Acts as an acid, releasing H+ to become R–COO– • Fatty acids • Amino acids Amino group, – NH2 Can accept or release H+, depending on pH; can form bonds with other molecules • Amino acids Hydroxyl group, – OH May link molecules through dehydration synthesis; hydrogen bonding between hydroxyl groups and water molecules affect solubility • Carbohydrates • Fatty acids • Amino acids Phosphate group, – PO4 May link other molecules to form larger structures; may store energy • Phospholipids • Nucleic acids • High-energy compounds Carboxylic acid group, – COOH * The term R group is used to denote the rest of the molecule, whatever that might be. The R group is also known as a side chain. Figure 2 Section 4 © 2011 Pearson Education, Inc. Module 2.10: Carbohydrates • Carbohydrates • Contain carbon, hydrogen, and carbon in ratio near 1:2:1 • Examples: sugars and starches • Roughly 1.5% of total body weight • Most important as energy sources © 2011 Pearson Education, Inc. Module 2.10: Carbohydrates • Types of carbohydrates 1. Monosaccharide (mono-, single + sakcharon, sugar) • Simple sugar • Contains 3–7 carbon atoms • Example: glucose (a six-carbon or hexose sugar) © 2011 Pearson Education, Inc. The hexose glucose, a monosaccharide or simple sugar, the simplest form of carbohydrate The atoms in a glucose molecule may form either a straight chain or a ring. In the body, the ring form is more common. Figure 2 .10 © 2011 Pearson Education, Inc. 1 Module 2.10: Carbohydrates • Isomers • Molecules with the same molecular formula but different structures • Can be important in molecular function • Example: glucose and fructose • Both have molecular formula of C6H12O6 • Structures are different © 2011 Pearson Education, Inc. The structures of the monosaccharide isomers glucose and fructose, both with chemical formula C6H12O6 Glucose Fructose The monosaccharides glucose and fructose are isomers. Fructose is a hexose found in many fruits. Although its chemical formula, C6H12O6, is the same as that of glucose, the arrangement of its atoms differs from that of glucose. Figure 2 .10 © 2011 Pearson Education, Inc. 2 Module 2.10: Carbohydrates • Types of carbohydrates (continued) 2. Disaccharide (di-, two) • Two monosaccharides joined • Example: sucrose • Dehydration synthesis creates disaccharides • Hydrolysis breaks them down to monosaccharides © 2011 Pearson Education, Inc. The formation of the disaccharide sucrose by dehydration synthesis DEHYDRATION SYNTHESIS Glucose Fructose Sucrose Figure 2 .10 © 2011 Pearson Education, Inc. 3 Module 2.10: Carbohydrates • Types of carbohydrates (continued) 3. Polysaccharide (poly-, many) • Many disaccharides and/or monosaccharides connected together Example: starches • • • • Major dietary source in potatoes and grains Broken down into monosaccharides by digestive system Glycogen • Animal starch • Can be broken down to glucose to supply energy • Can form from glucose to store energy © 2011 Pearson Education, Inc. The production of two monosaccharides from a disaccharide by hydrolysis, the reverse of dehydration synthesis HYDROLYSIS Glucose Fructose Sucrose Figure 2 .10 © 2011 Pearson Education, Inc. 4 1 The polysaccharide glycogen, produced when multiple dehydration synthesis reactions add additional monosaccharides or disaccharides Glucose molecules The polysaccharide glycogen (GLĪ-kō-jen), or animal starch, has many side branches consisting of chains of glucose molecules. Muscle cells make and store glycogen. When these cells have a high demand for glucose, glycogen molecules are broken down; when the demand is low, they absorb glucose from the bloodstream and rebuild glycogen reserves. Figure 2.10 © 2011 Pearson Education, Inc. 5 Figure 2 .10 © 2011 Pearson Education, Inc. 2 Module 2.10 Review a. List the three structural classes of carbohydrates, and give an example of each. b. A food contains organic molecules with the elements C, H, and O in a ratio of 1:2:1. What class of compounds do these molecules belong to, and what are their major functions in the body? c. Predict the reactants and the type of chemical reaction involved when muscle cells make and store glycogen. © 2011 Pearson Education, Inc. Module 2.11: Lipids • Lipids (lipos, fat) • Carbon to hydrogen ratio is near 1:2 • Much less oxygen compared to carbohydrates with similar number of carbon atoms • Examples: fats, oils, waxes • Most are insoluble in water • Special transport mechanisms for them in the blood • Essential components of all cells • Provide twice as much energy as carbohydrates • Normally account for 12–24% of total body weight • Must obtain some from diet © 2011 Pearson Education, Inc. A triglyceride, a lipid produced by dehydration synthesis between a glycerol molecule and three fatty acids Glycerol Fatty acids Fatty Acid 1 Saturated Fatty Acid 2 Saturated Fatty Acid 3 Unsaturated HYDROLYSIS DEHYDRATION SYNTHESIS Dehydration synthesis can produce a monoglyceride (mon-ō-GLI-ser-īd), consisting of glycerol + one fatty acid. Subsequent reactions can yield a diglyceride (glycerol + two fatty acids) and then a triglyceride (glycerol + three fatty acids). Hydrolysis breaks the glycerides into fatty acids and glycerol. Triglyceride Figure 2 .11 © 2011 Pearson Education, Inc. 4 Module 2.11: Lipids • Fatty acids • Long carbon chains with attached hydrogens • Two ends 1. Head • Has carboxylic group (–COOH) • Hydrophilic 2. Tail • Hydrophobic © 2011 Pearson Education, Inc. A fatty acid, the building block of lipids, which includes a tail and a head composed of a carboxylic acid group: —COOH The carbon chain attached to the carboxylic acid group is known as the hydrocarbon tail of the fatty acid. The hydrocarbon tail is hydrophobic, so fatty acids have very limited solubility in water. In general, the longer the hydrocarbon tail, the lower the solubility of the molecule. Lauric acid (C12H24O2) © 2011 Pearson Education, Inc. Figure 2 .11 1 Module 2.11: Lipids • Fatty acids • Two types 1. Saturated fatty acid • Each carbon has four attached hydrogens 2. Unsaturated fatty acid • Contains double bonds • One double bond = monounsaturated • >1 double bond = polyunsaturated • Has fewer attached hydrogens • Changes metabolism of the molecule © 2011 Pearson Education, Inc. Figure 2 .11 © 2011 Pearson Education, Inc. 2 An unsaturated fatty acid, in which one or more of the single covalent bonds between carbon atoms has been replaced by a double covalent bond Double covalent bond Figure 2 .11 © 2011 Pearson Education, Inc. 3 Module 2.11: Lipids • Glycerides • Fatty acid chains attached to a glycerol molecule • Three types through dehydration synthesis 1. Monoglyceride (glycerol + one fatty acid) 2. Diglyceride (glycerol + two fatty acids) 3. Triglyceride (glycerol + three fatty acids) © 2011 Pearson Education, Inc. Figure 2 .11 © 2011 Pearson Education, Inc. 4 Module 2.11 Review a. Describe lipids. b. Describe the structures of saturated and unsaturated fatty acids. c. In the hydrolysis of a triglyceride, what are the reactants and the products? © 2011 Pearson Education, Inc. Module 2.12: Diverse functions of lipids • Functions of lipids • Chemical messengers • Components of cellular structures (structural lipids) • Maintain cell membrane • Allow separation of different aqueous solutions © 2011 Pearson Education, Inc. Module 2.12: Diverse functions of lipids • Eicosanoids • Derived from arachidonic acid • Examples: • Leukotrienes • • Produced by cells in response to injury Prostaglandins • Released by cells to coordinate local cellular activities © 2011 Pearson Education, Inc. Figure 2 .12 © 2011 Pearson Education, Inc. 1 Module 2.12: Diverse functions of lipids • Steroids • Large molecules with four carbon rings • Differ in attached functional groups • Examples: • Cholesterol (chole-, bile + steros, solid) • • Functions to maintain plasma membranes, growth, and division Hormones • Regulation of sexual and other metabolic functions © 2011 Pearson Education, Inc. Three steroid molecules, which share a distinctive carbon-ring framework Cholesterol Estrogen © 2011 Pearson Education, Inc. Testosterone Figure 2.12 2 Module 2.12: Diverse functions of lipids • Phospholipids and Glycolipids • Phospholipid • • Phosphate linking a diglyceride to a nonlipid group Glycolipid • Carbohydrate attached to a diglyceride © 2011 Pearson Education, Inc. The structures of phospholipids and glycolipids Nonlipid group Carbohydrate Phosphate group Glycerol Glycerol Fatty acids In a phospholipid, a phosphate group links a diglyceride to a nonlipid group. Fatty acids In a glycolipid, a carbohydrate is attached to a diglyceride. Figure 2 .12 © 2011 Pearson Education, Inc. 3 Module 2.12: Diverse functions of lipids • Phospholipids and Glycolipids • Structurally related • Can be synthesized by cells primarily from fatty acids • Like fatty acids • • Tails are hydrophobic • Other end is hydrophilic In water, can form large droplets (micelles) © 2011 Pearson Education, Inc. Hydrophilic heads Water Phospholipid Hydrophobic tails Phospholipid Glycolipid Glycolipid The arrangement of phospholipid and glycoplid molecules in droplets called micelles Micelles Figure 2 .12 © 2011 Pearson Education, Inc. 4 Module 2.12 Review a. Why is cholesterol necessary in the body? b. Describe the basic functions of eicosanoids, steroids, phospholipids, and glycolipids. c. Describe the orientations of phospholipids and glycolipids when they form a micelle. © 2011 Pearson Education, Inc. Module 2.13: Proteins • Proteins • Most abundant organic molecule in the body • In many ways, are most important • Normally account for 20% of total body weight • Contain carbon, hydrogen, oxygen, and nitrogen • • Possibly sulfur and phosphorus as well Consist of long chains of amino acids • 20 amino acids in the body • Typical protein contains 1000 amino acids © 2011 Pearson Education, Inc. Module 2.13: Proteins • Amino acids • All have same structural components • Central carbon • Amino group • Carboxyl group • R group (variable side chain) © 2011 Pearson Education, Inc. Amino group Central carbon Carboxylic acid group The structural elements of amino acids, the basic units of proteins R group (variable side chain of one or more atoms) Figure 2 .13 © 2011 Pearson Education, Inc. 1 Module 2.13: Proteins • Peptides • Amino acids linked through dehydration synthesis • Covalent bond connects the carboxylic acid group of one amino acid to the amino group of another • = Peptide bond • Dipeptide • • Polypeptides • • Two amino acids linked together Three or more amino acids linked together Peptides of over 100 amino acids are called proteins © 2011 Pearson Education, Inc. The linkage of two amino acids by dehydration synthesis, and the breakage of peptide bonds by hyrdolysis Glycine (gly) Alanine (ala) Adjacent amino acids can be linked together by a covalent bond that connects the carboxylic acid group of one amino acid to the amino group of another. Peptide bonds can also be broken through hydrolysis; the hydrolysis of a dipeptide yields a pair of amino acids. The bond between amino acids is known as a peptide bond. Molecules consisting of amino acids held together by peptide bonds are called peptides. This molecule is called a dipeptide because it contains two amino acids. Figure 2 .13 © 2011 Pearson Education, Inc. 2 Module 2.13: Proteins • Protein structure 1. Primary structure • Sequence of amino acids 2. Secondary structure • Molecule shape changes due to bonds at different parts of the polypeptide chain • Example: hydrogen bonds 3. Tertiary structure • • Coiling and folding giving the protein a final 3-D shape Usually from interactions of the protein and surrounding water molecules © 2011 Pearson Education, Inc. Module 2.13: Proteins • Protein structure (continued) 4. Quaternary structure • Interaction between multiple polypeptide chains to form a protein complex • Examples: • Hemoglobin • • Binds oxygen in red blood cells Keratin and collagen • © 2011 Pearson Education, Inc. Fibrous proteins that give strength to tissues The four levels of structural complexity of proteins Primary Structure A polypeptide, a linear chain of amino acids Secondary Structure Hydrogen bond Hydrogen bond or Alpha-helix Pleated sheet Secondary structure results from bonds between atoms at different parts of the polypeptide chain. Hydrogen bonding, for example, may create either a simple spiral, known as an alpha-helix, or a flat pleated sheet. Tertiary Structure Heme unit Quaternary Structure Hemoglobin (globular protein) Tertiary structure results from the complex coiling and folding that gives a protein its final three-dimensional shape. Tertiary structure results primarily from interactions between the polypeptide chain and the surrounding water molecules, and to a lesser extent from interactions between the R groups of amino acids in different parts of the molecule. Quaternary structure results from the interaction between individual polypeptide chains to form a protein complex. The protein hemoglobin contains four polypeptide subunits. Hemoglobin is found within red blood cells, where it binds and or transports oxygen. It is an example of a globular protein. In keratin and collagen, three linear subunits intertwine, forming Heme a fibrous protein. The unit three-dimensional shape of a protein plays an essential role in Keratin determining its functional (fibrous protein) properties. Figure 2 .13 © 2011 Pearson Education, Inc. 3 Module 2.13: Proteins • Protein structure • Denaturation • Protein shape and function deteriorate • Occurs under extreme conditions • • Temperature • pH Irreparable damage can occur to tissues and organs © 2011 Pearson Education, Inc. Module 2.13 Review a. Describe proteins. b. What kind of bond forms during the dehydration synthesis of two amino acids? c. Why does boiling a protein affect its structural and functional properties? © 2011 Pearson Education, Inc. Module 2.14: Enzymes are regulatory proteins • Enzymes facilitate most everything that occurs inside the body • Substrates • Reactants in enzymatic reactions • Must bind to a specific region of an enzyme • = Active site • Binding reinforced by hydrogen bonds • Enzyme-substrate complex • Once binding to enzyme occurs Animation: Chemical Reactions: Enzymes © 2011 Pearson Education, Inc. Module 2.14: Enzymes are regulatory proteins • • Structure of interacting molecules is very important in enzymatic reactions • Substrate and enzyme fit in a “lock and key” fashion • Substrate binding results in a temporary, reversible change in shape of enzyme Enzyme is free to catalyze other reactions once product is complete • Saturation limit • Substrate concentration required to have maximum rate of reaction © 2011 Pearson Education, Inc. The events in the actions of enzymes Substrates Active site Enzymesubstrate complex Figure 2 .14 © 2011 Pearson Education, Inc. 1 - 4 Module 2.14 Review a. Define active site. b. What are the reactants in an enzymatic reaction called? c. Relate an enzyme’s structure to its reaction specificity. © 2011 Pearson Education, Inc. Module 2.15: High-energy compounds • High-energy compounds • Donate energy to chemical reactions to form products • Contain high-energy bonds (covalent) • Release energy when broken © 2011 Pearson Education, Inc. Module 2.15: High-energy compounds • Adenosine triphosphate (ATP) • Most common high-energy compound • Provides energy for many vital body functions • Examples: • Contraction of muscles • Synthesis of proteins, carbohydrates, and lipids © 2011 Pearson Education, Inc. Module 2.15: High-energy compounds • Adenosine triphosphate (ATP) (continued) • Formation • • Adenosine monophosphate (AMP) • Adenosine • Ribose sugar • One phosphate Adenosine diphosphate (ADP) • • Adenosine triphosphate (ATP) • • + One more phosphate + One more phosphate Formation of ATP from ADP is reversible © 2011 Pearson Education, Inc. The reversible reaction by which ATP is formed from ADP Figure 2 .15 © 2011 Pearson Education, Inc. 2 Module 2.15 Review a. Where do cells obtain the energy needed for their vital functions? b. Describe ATP. c. Compare AMP with ADP. © 2011 Pearson Education, Inc. Module 2.16: DNA and RNA • Nucleic acids • • Large organic molecules Composed of carbon, hydrogen, oxygen, nitrogen, and phosphorus Primary function is storage and transfer of information • • • • Particularly synthesis of proteins Consist of two long chains formed from dehydration synthesis of subunits (nucleotides) Two classes 1. Deoxyribonucleic acid (DNA) 2. Ribonucleic acid (RNA) © 2011 Pearson Education, Inc. Module 2.16: DNA and RNA • Nucleotide components 1. Nitrogenous base a. Purines • Adenine • Guanine b. Pyrimidines • Cytosine • Thymine (only in DNA) • Uracil (only in RNA) 2. Pentose sugar 3. Phosphate group © 2011 Pearson Education, Inc. The structure of nucleotides, the subunits of a nucleic acid Sugar Phosphate group Nitrogenous base Figure 2 .16 © 2011 Pearson Education, Inc. 1 The structures of the two types of nitrogenous bases of nucleic acids Nitrogenous bases The purines adenine and guanine are found in both DNA and RNA. Adenine Guanine There are three important pyrimidines. DNA and RNA both contain cytosine, thymine is found only in DNA, and uracil is found only in RNA. Cytosine Thymine (DNA only) Uracil (RNA only) Figure 2.16 © 2011 Pearson Education, Inc. 2 Module 2.16: DNA and RNA • Phosphate and sugars of adjacent nucleotides are linked together to form nucleic acids • • Forms the “backbone” of the nucleic acid Sequence of attached nitrogenous bases carries the information © 2011 Pearson Education, Inc. The production of a linear sugar-to-phosphate-to-sugar “backbone” of a nucleic acid by dehydration synthesis DEHYDRATION SYNTHESIS Figure 2.16 © 2011 Pearson Education, Inc. 2 Module 2.16: DNA and RNA • DNA • Pair of nucleotide chains • • • = Complementary strands Twist around each other to form a double helix Hydrogen bonds between opposing nitrogenous bases hold two strands together • = Complementary base pairs • • Adenine – Thymine (A-T) Cytosine – Guanine (C-G) Animation: Protein Synthesis: DNA Molecule Animation: Protein Synthesis: How DNA Information is Used to Make Protein © 2011 Pearson Education, Inc. The structure of a DNA molecule, in which two complementary strands (nucleotide chains) are linked by hydrogen bonds between complementary base pairs Deoxyribose Adenine Thymine Phosphate group Hydrogen bond The two strands of DNA twist around one another in a double helix that resembles a spiral staircase. DNA strand 1 DNA strand 2 Hydrogen bond Cytosine Guanine Figure 2.16 © 2011 Pearson Education, Inc. 3 Module 2.16: DNA and RNA • RNA • Single chain of nucleotides • Three types 1. Messenger RNA (mRNA) 2. Transfer RNA (tRNA) 3. Ribosomal RNA (rRNA) © 2011 Pearson Education, Inc. A molecule of RNA, which consists of a single chain of nucleotides single-strand sugar-phosphate backbone Nitrogenous bases Figure 2.16 © 2011 Pearson Education, Inc. 4 Figure 2 .16 © 2011 Pearson Education, Inc. 3 Module 2.16 Review a. Describe nucleic acids. b. Explain how the complementary strands of DNA are held together. c. A large organic molecule composed of ribose, nitrogenous bases, and phosphate groups is which kind of nucleic acid? © 2011 Pearson Education, Inc.