Chapter 2: Atoms, Ions, and Molecules Chapter Notes PDF

CHAPTER 2: Atoms, Ions, and Molecules CHAPTER NOTES 2.1 Atomic Structure: The simplest level of organization is composed of atoms, ions, and molecules. (pp. 32–35) A. Matter, Atoms, Elements, and the Periodic Table (p. 33) 1. The human body is composed of matter, defined as a substance th...

CHAPTER 2: Atoms, Ions, and Molecules CHAPTER NOTES 2.1 Atomic Structure: The simplest level of organization is composed of atoms, ions, and molecules. (pp. 32–35) A. Matter, Atoms, Elements, and the Periodic Table (p. 33) 1. The human body is composed of matter, defined as a substance that has mass and occupies space. 2. Matter is present in the human body in three forms: solid, liquid, and gas. 3. All matter is composed of atoms. 4. An atom is the smallest particle that exhibits the chemical properties of an element. 5. There are 92 naturally occurring elements, with hydrogen being the smallest and uranium the largest and heaviest. 6. Technical advances in chemistry and physics have resulted in the ability to produce “ultraheavy” elements that are larger than uranium. 7. All elements are organized into a chart form in the periodic table of elements. (Figure 2.1a, p. 33) 8. Elements are grouped into major, lesser, and trace elements based on the percentage each composes by weight in the human body. 9. Major elements comprise almost 99% and minor elements less than 1%. 10. Only 12 elements occur in living organisms in greater than trace amounts: oxygen, carbon, hydrogen, nitrogen, calcium, phosphorus, sulfur, potassium, sodium, chlorine, magnesium, and iron. (Figure 2.1b, p. 33 11. Atoms are composed of three subatomic particles: protons, neutrons, and electrons. (Figure 2.2, p. 33) 12. Two major criteria differentiate subatomic particles—namely, mass and charge. 13. The mass of an atom is expressed as the atomic mass unit (amu) of dalton. 14. Neutrons are uncharged, having a mass of 1 amu. 15. Protons have a positive charge, also having a mass of 1 amu. 16. Neutrons and protons are located in the atomic nucleus. 17. The electrons have a negative charge, having a mass of 1/800th the mass of a proton or neutron. 18. Electrons are located at varying distances from the nucleus in regions called orbitals or shells, often depicted as either an electron cloud or discrete energy levels (shell model). 19. Elements differ in subatomic particles, and the periodic table can be used to determine the number of these subatomic particles. 20. The periodic table shows an element’s symbol, atomic number, and average atomic mass. 21. A unique chemical symbol is assigned to each element. 22. The atomic number of an element indicates the number of protons in an atom. 23. The average atomic mass indicates the mass of both protons and neutrons in the atomic nucleus, and it reflects the “heaviness” of an element’s atoms relative to atoms of other elements. 1  24. The number of protons in an atom is the atomic number. 25. The number of neutrons can be determined by subtracting the atomic number from the atomic mass. 26. The number of electrons is the same as the atomic number because all atoms are neutral. 27. Shells of electrons surround the atomic nucleus, and each shell has a given energy level. 28. Each shell can only hold a limited number of electrons, with the innermost shell holding up to two electrons, and the second shell holding up to eight electrons. 29. All subsequent shells also house eight electrons, but some subsequent shells hold more than eight. 30. Electrons fill the shells from the shells closest to the nucleus first and then move outward, shell by shell. B. Isotopes (p. 34); Figure 2.3 (p. 34) 1. Isotopes are different atoms of the same element that have the same number of protons and electrons but differ in the number of neutrons. 2. Isotopes exhibit essentially identical chemical characteristics but have a different atomic mass. 3. Carbon exists in three isotopes: carbon-12, carbon-13, and carbon-14. 4. All the carbon isotopes have six protons in their nuclei; however, carbon-12 has six neutrons, carbon-13 has seven neutrons, and carbon-14 has eight neutrons. 5. Some isotopes are referred to as radioisotopes, in that they are unstable due to having an excess number of neutrons. 6. Radioisotopes usually lose nuclear components in the form of high-energy radiation that includes alpha particles, beta particles, or gamma rays as they decay or break down to a more stable isotope. 7. The time it takes for 50% of the radioisotope to become stable is its physical half-life. 8. The time it takes for half of the radioactive material (e.g., from a medical test using radioactive contrast material) to be eliminated from the body is the biological half-life. 9. Clinical View—Medical Imaging of the Thyroid Gland (p. 35)—radioisotopes can be used for diagnosis as well as treatment of certain diseases such as thyroid disease. Scintigraphy is the nuclear medicine imaging technique used to assess metabolic activity using radioactive tracers. C. Chemical Stability and the Octet Rule (pp. 34−35); Figure 2.4 (p. 35) 1. The periodic table is organized into horizontally across the table based on atomic number, and it is organized into columns based on electrons in the outer shell, or what is referred to as the valence shell. 2. Each consecutive column, IA–VIIIA, has the same number of electrons in its valence shell, as represented by the numeral at the top of the column. 3. For example, in Column IIA, all the elements in the column have two electrons in their valence shells. 4. Atoms that have a completely filled valence shell of either two or eight electrons are stable or inert and thus do not typically combine with other elements. 2  5. Elements that do not have a filled valence shell tend to lose, gain, or share electrons to obtain a complete outer shell and are chemically active. 6. The tendency to fill the valence shell is termed the octet rule. 2.2 Ions and Ionic Compounds: The body is composed mostly of compounds. (pp. 36–38) A. Ions (pp. 36–37) 1. Ions are atoms or groups of atoms with either a positive charge or a negative charge and are produced from the loss or gain of an electron or electrons. 2. Whether an atom will lose or gain electrons depends on which action will provide it with a complete valence shell. 3. When an atom loses one or more electrons, it becomes positively charged and called a cation. 4. When an atom gains one or more electrons, it becomes negatively charged and called an anion. 5. When two or more atoms complex together and become ions, they are termed polyatomic ion. 6. The periodic table can be used to predict whether an atom will become a cation or an anion and how many electrons will be lost or gained to attain the octet. 7. Usually, elements on the left side of the periodic chart lose electrons, thus becoming cations, and elements on the right side of the chart gain electrons, thus becoming anions. ​ 8. Certain ions have important physiological functions in the body. (Table 2.1, p. 36) ​ 9. Ions, including the common ones found in the body, function as electrolytes, which are substances that conduct an electric current when dissolved in water. ​ 10. Optimal health requires electrolyte balance, which is maintaining homeostatic blood levels of the different ions. ​ 11. An electrolyte imbalance occurs when blood concentrations are either too high or too low. This can lead to disease and even death. B. Ionic Bonds (pp. 37−38) 1. Positively charged cations and negatively charged anions may bind together by electrostatic interactions called ionic bonds. 2. The structures formed are salts. 3. A classic example is when table salt is formed from the bonding of metallic atoms of sodium with nonmetallic atoms of chlorine. (Figure 2.5, p. 37) 4. In this example, the sodium atom donates one outer shell electron to a chlorine atom. 5. After donating the electron, the sodium atom becomes a cation (Na+) and the chlorine atom becomes an anion (Cl−). ​ 6. The positively and negatively charged ions then attract each other in amounts so that the salt is neutral. ​ 7. The chemical formula for an ionic compound represents the atoms present and how many are needed for electrical neutrality. Examples NaCl (1:1 ratio), MgCl2 (1:2 ratio), Ca3(PO4)2 (3:2 ratio). 2.3 Covalent Bonding, Molecules, and Molecular Compounds: Instead of donating or gaining electrons to form ionic compounds, atoms also have the possibility of reaching stability by sharing electrons. (pp. 38–42) A. Chemical Formulas: Molecular and Structural (pp. 38–39) 1. The sharing of electrons between atoms results in a covalently bonded molecule. 3  2. Most molecules are composed of two or more different elements and are called molecular compounds. 3. The chemical constituents of a molecule and their actual numbers are represented using a molecular formula. 4. The structural formula of a molecule is complementary to its molecular formula and exhibits not only the number and types of atoms but also their arrangements within the molecule. 5. Structural formulas provide a means for differentiating isomers, which are molecules composed of the same number and kind of elements but arranged differently in space. 6. Isomers may have very different properties from one another—therefore, structural formulas are an important piece of chemical information. (Figure 2.6, p. 38) B. Covalent Bonds (pp. 39–41) 1. The bond that is formed when atoms share electrons is a covalent bond. 2. Covalent bonding occurs when both atoms require electrons to become stable. 3. This bonding takes place when the participating atoms that form the chemical bond have four, five, six, or seven electrons in the outer shell. 4. The most common elements of the human body that form covalent bonds are oxygen, carbon, hydrogen, and nitrogen. 5. The number of covalent bonds formed by an atom may be determined by examining the number of electrons needed to complete the outer shell. 6. Atoms of elements that can form more than one covalent bond may do so through combinations of single, double, or triple covalent bonds. 7. A single covalent bond is one pair of electrons shared between two atoms. (Figure 2.7a, p. 39) 8. A double covalent bond is two pairs of electrons shared between two atoms. (Figure 2.7b, p. 39) 9. A triple covalent bond is three pairs of electrons shared between two atoms. (Figure 2.7c, p. 39) 10. Numerous carbon atoms are sometimes bonded together to form a “carbon skeleton.” 11. Three possible arrangements of the carbon skeleton may occur: a straight chain, branched chain, or a ring. (Figure 2.9, p. 40) 12. Atoms share electrons in a covalent bond either equally or unequally between the atoms. 13. How the atoms share the electrons between them is determined by the relative attraction each atom has for electrons, a concept referred to as electronegativity. 14. Different types of atoms have varying degrees of electronegativity, or attraction for electrons, and thus may share the electrons unequally, resulting in what is termed a polar covalent bond. 15. If the atoms are the same or have very close electronegativities, they share the electrons equally, a nonpolar covalent bond is formed. ​ ​ 16. As a general rule, electronegativity increases both from left to right across a row of the periodic table and from bottom to the top of a column. ​ ​ 17. The more protons in the nucleus of an atom, the greater the pull on shared electrons, while more electron shells will decrease the pull. 4 ​ ​ 18. For the four most common elements in the body, the order of electronegativity from least to greatest is hydrogen, carbon, nitrogen, and oxygen. ​ ​ 19. When a bond is nonpolar, the shared electrons spend equal amounts of time around each atom. ​ ​ 20. When a bond is polar, the shared electrons are more strongly attracted to the more electronegative atom and spend more time circulating around that atom. This causes the more electronegative atom to develop a partial negative charge (δ-). ​ ​ 21. The atom that is more electropositive develops a partial positive charge (δ+) because the shared electrons spend less time in orbit around that atom. ​ ​ 22. These polar covalent bonds act like magnets with opposite poles. C. Nonpolar, Polar, and Amphipathic Molecules (pp. 41−42) 1. If the atoms share the electrons equally, a nonpolar covalent bond is formed. This will form a nonpolar molecule. 2. Nonpolar molecules in the body are formed predominantly by covalent bonding between the same elements and between carbon and hydrogen. (Figure 2.10a, p. 42) 3. Polar molecules are formed by polar covalent bonding between different elements. (Figure 2.10b, p. 42) ​ ​ 4. If a molecule has polar covalent bonds but is symmetrically arranged, the resulting molecule will be nonpolar. (Figure 2.8a and b, p. 40) ​ ​ 5. If the polar bonds are not symmetrically arranged, the resulting molecule will be polar. 4. Sometimes a molecule is large enough that it can have one major part that is nonpolar and another part that is polar, resulting in a molecule termed an amphipathic molecule. D. Intermolecular Attractions (p. 41–42) 1. Molecules sometimes have weak chemical attractions to other molecules, called intermolecular attractions. 2. One important intermolecular attraction is termed a hydrogen bond, which occurs between polar molecules. 3. A hydrogen bond is a weak attraction between a partially positive hydrogen atom within a molecule and a partially negative atom (oxygen or nitrogen) within another molecule. (Figure 2.11, p. 42) 4. A second type of force occurs when electrons orbiting the nucleus of an atom of a nonpolar molecule are for a brief instant distributed unequally, causing one portion of the atom to be slightly negative and one end slightly positive. 5. Nonpolar atoms with an unequal charge distribution cause neighboring atoms to perform similarly, leading to opposition of charges that cause weak intermolecular attractions 6. Hydrophobic interactions are a type of intermolecular attraction between nonpolar molecules when placed in water. ​ 7. When a molecule is large, intermolecular type forces can occur between different parts of that molecule and are called intramolecular attractions. ​ 8. Intermolecular and intramolecular attractions are important in establishing the three-dimensional shapes of complex molecules like proteins and DNA, as well as temporary binding between molecules such as a hormone to its protein receptor. 5 2.4 Molecular Structure of Water and the Properties of Water: Water is the substance that comprises approximately two third of the human body by weight. (pp. 43–46) A. Molecular Structure of Water (p. 43) 1. Chemist classify molecules into two broad categories: organic molecules and inorganic molecules. 2. Organic molecules are defined as molecules that contain carbon, which are (or have been) components of living organisms, such as glucose, protein, and triglycerides. 3. Inorganic molecules are all the other molecules that are not organic, such as water, salts, and others. 4. Water is a polar molecule composed of one oxygen atom bonded to two hydrogen atoms (H2O). 5. The polar nature of water is due to the oxygen atom being more electronegative than the hydrogen atom, thus pulling the electrons unequally. 6. Every water molecule has the ability to form four hydrogen bonds with adjacent water molecules. (Figure 2.12, p. 43) B. Properties of Water (pp. 43–44) 1. Water is present in three phases, depending upon the temperature: a gas (water vapor), a liquid (water), and a solid (ice). 2. Water has several functions in the human body: transport, lubricate, cushion, and excrete wastes. 3. Cohesion is the attraction between water molecules due to hydrogen bonding between the water molecules. 4. Surface tension is the inward pulling of cohesive forces at the surface of water due to water at the surface only being able to form three hydrogen bonds rather than the four formed by water in the internal liquid. 5. Adhesion is the attraction between water molecules and a substance other than water, as a result of water forming hydrogen bonds with molecules of the other substance. 6. Temperature is a measure of the kinetic energy, or random movement, of atoms or molecules within a substance. 7. The relationship between temperature and kinetic energy is direct—the temperature is higher when the kinetic energy is greater. 8. Specific heat is the amount of energy required to increase the temperature of 1 g of a substance 1°C. 9. Water has a very high specific heat as a result of the need to break hydrogen bonds in water to raise its temperature. 10. Heat of vaporization is the heat required for the release of molecules from a liquid phase into a gaseous phase for 1 g of a substance. 11. Water has a high heat of vaporization as a result of the need to break hydrogen bonds to release water from the liquid into the gaseous phase. 12. Concept Overview—Water’s Roles in the Body, Figure 2.14 (p. 47) C. Water as the Universal Solvent (pp. 44−46); Figure 2.13 (p. 45) 1. Water is the solvent of the body, and substances that dissolve in water are called solutes. 2. Many substances dissolve in water, but not all. 3. The chemical properties of a substance (whether it is polar, charged, nonpolar, or amphipathic) determine how it interacts with water. 6  4. Since water is polar, substances that can dissolve in it are either polar or charged, such as ions. 5. Substances that can dissolve in water are called hydrophilic (water loving). 6. Substances that dissolve in water are surrounded by many water molecules, known as a hydration shell. 7. Some substances dissolve in water but do not remain intact, thus they break apart: dissociate. 8. Dissociation in water occurs with substances that have ionic bonding. 9. Acids and bases dissociate in water. 10. Substances that both dissolve and dissociate in water, such as salts, acids, and bases, can readily conduct an electric current; thus, they are called electrolytes. 11. Substances that dissolve in water but do not dissociate are nonelectrolytes—do not conduct a current. 11. Hydrophobic (water fearing) substances do not dissolve in water because these substances are not charged (nonpolar). 12. The interaction between the molecules of the “excluded” nonpolar excluded substance is termed hydrophobic interaction because it appears these molecules are avoiding water. 14. Amphipathic molecules have both polar and nonpolar regions. 15. Amphipathic molecules partially dissolve in water due to the polar region seeking water and the nonpolar region avoiding water. 16. Certain molecules (phospholipids) in the cell membrane are amphipathic. 2.5 Acidic and Basic Solutions, pH, and Buffers: Acidic and basic solutions occur when an acid or base is added to water. (pp. 46–49) A. Water: A Neutral Solvent (pp. 46–48) 1. Water can spontaneously dissociate a result of the covalent chemical bond between oxygen and either of the two hydrogen atoms in a water molecule spontaneously breaking apart at a low rate, about 10−7 ions per liter. 2. When a hydrogen ion transfers to a water molecule, giving it an extra hydrogen ion (H+), it is represented as H30+ and termed the hydronium ion. 3. When a water molecule loses a hydrogen ion, it is termed a hydroxide ion and represented as OH−. 4. Water is neutral in that it has as many positively charged hydronium ions as it does negatively charged hydroxide ions. B. Acids and Bases (p. 46) 1. An acid is a substance that dissociates in water to produce an H+ and an anion; thus, it is also called a proton donor. ​ 2. Acid strength is determined by the degree of dissociation; a strong acid produces more hydrogen ions, while a weak acid produces fewer hydrogen ions in solution. 3. A base accepts H+ when added to a solution; thus, it is termed a proton acceptor. ​ 4. Strong bases dissociate to a greater extent and can accept more H+, while a weak acid binds to fewer H+. C. pH, Neutralization, and the Action of Buffers (pp. 46–48) 1. The pH of a solution is a measure of the relative amounts of H+ and is expressed as a number between 0 and 14. (Figure 2.15, p. 48) 7  2. Pure water and other solutions that have equal concentrations of H+ and OH− are neutral and have a pH of 7. 3. Solutions with a pH below 7 are acidic, and solutions with a pH above 7 are basic, or alkaline. ​ ​ 4. The higher the hydrogen ion concentration, the lower the pH, and the lower the hydrogen ion concentration, the higher the number. ​ ​ 5. The pH scale represents a 10-fold change in H+ concentration between two adjacent whole number pH values. A pH of 2 is 10 times more acidic than a pH of 3; a pH of 10 is 10 times less acidic than a pH of 9. 6. Neutralization occurs when a solution that is either acidic or basic is returned to neutral (pH 7). 7. A buffer is a single substance, or an associated group of substances, that functions to help prevent pH changes if either excess acid or base is added. 2.6 Water Mixtures: Mixtures are formed from the combining or “mixing” of two or more substances. (pp. 49–51) A. Categories of Water Mixtures (pp. 49–50); Figure 2.16 (p. 49) 1. Water mixtures are placed into three categories based on the relative size of the substance mixed with water and include suspensions, colloids, and solutions. 2. A suspension, like sand in water, is formed when a material that is larger than 100 nm particulate size is mixed with water, resulting in a mixture that does not remain mixed unless it remains in motion. 3. Blood cells within the plasma (the liquid portion) form a suspension. 4. A colloid is a mixture of smaller particles within water, where the size ranges from 1 to 100 nanometers. May appear milky and scatters light. 5. Colloids in the body include semifluid cell cytosol and proteins within plasma. 6. A solution is a homogeneous mixture in which the particulate size of the substance is smaller than 1 nanometer, and it dissolves in water. 7. In a solution, the substance being dissolved is the solute and the substance doing the dissolving is the solvent. 8. In a solution, the solute is not visible, do not scatter light, and do not settle. 9. Blood plasma is an example of a solution in the body. 7. An emulsion is a type of colloid composed specifically of water and a nonpolar (hydrophobic) liquid substance such as vegetable oil in water; breast milk is an example of an emulsion. B. Expressions of Solute Concentration (pp. 50–51), Table 2.2 (p. 50) 1. The amount of solute dissolved in a solution determines the concentration of a solution and may be expressed in several ways: mass/volume, mass/volume percent, molarity, and molality. 2. Mass/volume is grams of solute per volume of solution. 3. Mass/volume percentage is grams of solute per 100 ml of solution. 4. Molarity is a measure of number of moles per liter of solution. 5. Molality is the moles per kilogram of solvent. 6. An osmole (osm) is another means of expressing concentration, which reflects whether a substance either dissolves, or dissolves and dissociates, when placed in a solution (i.e., whether it is a nonelectrolyte or electrolyte). 7. An osmole measures the number of osmotic active particles in a solution. 8  8. Osmoles can be expressed as either osmolarity or osmolality. 9. Osmolarity is the number of particles in 1 l of solution, whereas osmolality is the number of particles in 1 kg of water. 10. A mole is the mass in grams that is equal to either the atomic mass of an element or the molecular mass of a compound. 11. One mole contains 6.02 × 1023 atoms, ions, or molecules. 2.7 Biological Macromolecules: Four classes of organic biological macromolecules (biomolecules) can be distinguished in living organisms: lipids, carbohydrates, nucleic acids, and proteins. (pp. 51–63) A. General Characteristics (pp. 51–53) 1. Biological macromolecules are large organic molecules that are synthesized by the human body. 2. Biological macromolecules always contain carbon and hydrogen, and generally oxygen; some also contain nitrogen, phosphorus, and sulfur. 3. Organic biological macromolecules have several distinct features: carbon and carbon skeletons, hydrocarbons, and functional groups. 4. Carbon is the central element of biological macromolecules, existing independently, or covalently bonded into a carbon skeleton. 5. Hydrocarbons contain both carbon and hydrogen, covalently bonded together (nonpolar). 6. Functional groups are two or more atoms bonded together, attached to a carbon skeleton, displaying the same specific chemical characteristics, no matter what carbon skeleton they are attached to. Functional groups are polar. (Table 2.3, p. 52) 7. Polymers, present in biological macromolecules, are composed of repeating subunits called monomers. 8. Polymer forms are seen in carbohydrates, proteins, and nucleic acids, but not in lipids. 9. Dehydration synthesis (condensation) is a process that combines monomers to form polymers. 10. Dehydration synthesis involves the removal of a H from one monomer and an OH from another monomer, thus allowing the formation of a covalent bond between the two monomers and resulting in the formation of a water molecule when the H bonds to the OH (H2O). (Figure 2.17, p. 53) 11. Hydrolysis is a process that splits polymers into monomers by inserting a split water molecule into the polymer, thus breaking the covalent bonding between the two monomers. (Figure 2.17, p. 53) 12. In the process of hydrolysis, water is split into a –H and –OH, the –H is added to one monomeric subunit and the OH is added to the other subunit, resulting in a split covalent bond between the two monomers. B. Lipids (pp. 53–56) 1. Lipids are a diverse group of macromolecules that do not form polymers and are totally (nonpolar) or partially (amphipathic) insoluble in water. 2. Lipids function as stored energy, components of cell membranes, and hormones. 3. The four primary classes of lipids are triglycerides (neutral fats), phospholipids, steroids, and eicosanoids. (Table 2.4, p. 54) 9  4. Triglycerides (triacylglycerols) are formed from a glycerol molecule and three fatty acids bonded together. 5. Triglycerides are the most common form of lipids in living things and function in long-term energy storage (adipose tissue) and provide structural support, cushioning, and insulation of the body. 6. A fatty acid is a lipid molecule with generally 14–20 carbons bonded together in a chain, and terminating in an acidic carboxyl functional group. (Figure 2.18, p. 55) 7. If the fatty acid molecule contains a double bond, it is termed unsaturated, and if it does not contain a double bond, it is termed saturated. 8. If the fatty acid has one double bond, it is unsaturated, and if it has more than one, it is polyunsaturated. 9. Triglycerides are a storage form of energy located in fat (adipose) cells. 10. When cells form triglycerides for energy storage, it is termed lipogenesis, and when cells break down triglycerides to release energy, it is termed lipolysis. 11. A phospholipid is amphipathic and has a chemical structure similar to a triglyceride, except the phospholipid molecule contains a polar phosphate functional group. 12. Phospholipids are amphipathic in that they contain a polar head, where the polar phosphate group is positioned, along with two nonpolar tails, where the fatty acids are positioned. 13. Phospholipids are a main component of cell membranes. 14. Steroids are composed predominantly of hydrocarbons arranged in a distinct four multi-ringed structure, three of the rings having six carbons and one ring having five carbons. 15. Steroids differ according to the side chains extending from the rings. 16. Steroids include cholesterol, steroid hormones (e.g., testosterone, estrogen, and progesterone), and bile salts. 17. Eicosanoids are modified 20-carbon fatty acids that are synthesized from phospholipids of plasma membranes. 18. Four classes of eicosanoids are produced and include prostaglandins, thromboxanes, leukotrienes, and prostacyclins. ​ 19. The primary functions of eiconsanoids are in the inflammatory response of the immune system and communication within the nervous system. 20. Glycolipids are lipid molecules with a carbohydrate covalently attached, having functions associated with cell membranes, such as cell binding. 21. Vitamins A, E, and K are lipids known as the fat soluble vitamins. C. Carbohydrates (pp. 56–57) 1. The term carbohydrate is derived from the fact that a monomer of carbohydrate, termed a monosaccharide, has virtually every carbon hydrated with the equivalent of a water molecule, with a –H and –OH usually attached to each carbon. 2. One monomer of carbohydrate is termed a monosaccharide, two covalently bonded together into one molecule is termed a disaccharide, and several bonded together is a polysaccharide. 3. Glucose, fructose, and galactose are six-carbon monosaccharide isomers, termed hexose sugars. (Figure 2.20a, p. 57) 10  4. Ribose and deoxyribose are five-carbon monosaccharides, termed pentose sugars, and are found in RNA (ribonucleic acid) and DNA (deoxyribonucleic acid), respectively. 5. Sucrose (table sugar), lactose, and maltose are common disaccharides. (Figure 2.20b, p. 57) 6. Glycogen, the storage carbohydrate in animals, is a polysaccharide, while starch and cellulose (fiber) are plant polysaccharides. ​ 7. Glucose is critical to life processes because it is the primary molecule that provides energy for body cells. The brain uses glucose exclusively for energy. ​ 8. Blood glucose homeostasis is carefully maintained by several hormones and processes that either store or release glucose from storage. ​ 9. Liver and skeletal muscle store excess glucose in the form of glycogen in a process called glycogenesis. (Figure 2.19, p. 56) ​ 10. When blood sugar levels drop, the process of glycogenolysis breaks glycogen back down to glucose. ​ 11. The liver can also make glucose from noncarbohydrate sources through gluconeogenesis. ​ 12. Glycosaminoglycans (GAGs) are attached to proteins to form proteoglycans that are associated with the ground substance of connective tissue. D. Nucleic Acids (pp. 57–59), Figure 2.21 (p. 58), Table 2.5 (p. 59) 1. Nucleic acids are macromolecules that store and transfer genetic information for protein synthesis and were initially discovered in the cell nucleus. 2. Two classes of nucleic acid are deoxyribonucleic acid (DNA) and ribonucleic acid (RNA). 3. A nucleic acid monomer, known as a nucleotide, is composed of a nitrogenous base, pentose sugar, and phosphate group, covalently bonded together. 4. The pentose sugar in DNA is deoxyribose, and the pentose sugar in RNA is ribose. 5. There are five nitrogenous bases classified as either pyrimidines, which are single-ringed, or purines, which are double-ringed. 6. The nitrogenous bases adenine and thymine are pyrimidines, and cytosine, uracil, and guanine are purines. 7. Deoxyribonucleic acid (DNA) is a double-stranded nucleic acid; it can be found both as components of chromosomes within the cell nucleus and as a circular strand of DNA found in the mitochondria. 8. The nitrogenous bases in DNA are adenine, cytosine, guanine, and thymine; DNA contains no uracil. ​ 9. The DNA double strands are held together by hydrogen bonds between adenine and thymine and between cytosine and guanine that formed complementary base pairs. 9. The nitrogenous bases in RNA are adenine, cytosine, guanine, and uracil; RNA contains no thymine; it is replaced by uracil. 10. RNA is a single-stranded nucleic acid located both within the cell nucleus and within the cytoplasm of the cell. 11. Adenosine triphosphate (ATP) is another nucleotide that is composed of adenine, ribose, and three phosphate groups; it provides chemical energy within the cell. (Figure 2.22, p. 59) 11  12. Nicotinamide adenine dinucleotide (NAD+) and flavin adenine dinucleotide (FAD) are nucleotides that participate in the production of ATP in the cell mitochondria. E. Proteins (pp. 59–67) 1. Proteins are polymers composed of one or more linear strands of amino acid monomers that may number in thousands. (Figure 2.23b, p. 61) 2. Proteins have many functions: enzyme, defense, transport, support, movement, regulation, and storage. (Table 2.6, p. 60) 3. Proteins are composed of 20 different amino acids. 4. An amino acid, the monomeric unit of a protein, is a molecule composed of an acidic carboxyl function group and a basic amine group. (Figure 2.23a, p. 61) 5. The differences between amino acids are due to side-chain structures, termed the R groups. 6. Amino acid chemical properties are determined by the chemical nature of the R group; some amino acids are nonpolar, some polar, some charged, and some have special functions. 7. Amino acids are covalently linked together by a peptide bond, which is a type of covalent bond, found only in amino acid polymers and formed by dehydration synthesis. 8. Two amino acids bonded together into one molecule is termed a dipeptide, 3–20 bonded together is termed an oligopeptide, 21–199 bonded together is termed a polypeptide, and 200 or more is termed a protein. 9. A protein covalently bonded to carbohydrate is termed a glycoprotein; the ABO surface marker on a red blood cell is one example of glycoprotein function. ​ 10. The glycocalx of cells is composed of glycoproteins and glycolipids and is used for identification of “self” cells. F. Concept Overview, Figure 2.24 (pp. 62–63) reviews the four primary biological macromolecules. 2.8 Protein Structure: Protein structure is paramount to its functioning. (pp. 64–67) A. Categories of Amino Acids (pp. 64–66), Figure 2.25 (p. 65) 1. Amino acids are organized into four groups based on the chemical characteristics of their R group: nonpolar amino acids, polar amino acids, charged amino acids, and amino acids with special functions. 2. Nonpolar amino acids contain R groups with either hydrogen (glycine) or hydrocarbons (alanine, valine, leucine, isoleucine, phenylalanine, and tryptophan); they tend to group with other nonpolar amino acids by hydrophobic interactions in water. 3. Polar amino acids contain R groups with elements in addition to carbon and hydrogen (e.g., O, N, or S) (serine, threonine, asparagine, glutamine, and tyrosine); they form interactions with other polar molecules and with water. 4. Charged amino acids can either have a negative charge or a positive charge. 5. Negatively charged amino acids, with a negatively charged R group, include glutamic acid and aspartic acid, and those with a positively charged R group include histidine, lysine, and arginine. 6. Charged and polar amino acids are hydrophilic, and their presence increases the solubility of the protein in water. 7. Three amino acids have unique characteristics: proline, cysteine, and methionine. 12  8. Proline has an R group that attaches to the amino group, forming a ring that bends proteins. 9. Cysteine is an amino acid containing a sulfhydryl group, allowing it to form a disulfide covalent bond (bridge) between the sulfhydryl group of another cysteine amino acid; these disulfide bridges stabilize the construct of a protein. 10. Methionine is always the first amino acid positioned when a protein is synthesized. B. Amino Acid Sequence and Protein Conformation (pp. 66–67) 1. The structuring of a protein is described on a hierarchical level, having four possible structural levels: primary, secondary, tertiary, and quaternary. 2. The primary structure is the linear sequence of amino acids in the protein. (Figure 2.26a, p. 66) 3. The next three structural levels are responsible for the three-dimensional shape of a protein, known as protein conformation. 4. The more complex structural organizations of a protein are dependent upon intramolecular attractions between the amino acids in the linear sequence (primary structure) for proper folding and maintaining of a protein’s conformation. 5. Proteins, known as chaperone proteins, assist in proper protein folding. 6. Four main types of intermolecular and intramolecular interactions contribute to the final conformation of a protein: hydrophobic exclusion, hydrogen bonding, ionic bonding, and disulfide bonds. 7. The primary protein polymer is forced into its initial shape as hydrophobic exclusions “tuck” amino acids with nonpolar R groups into a more central location, limiting their contact with water. 8. Hydrogen bonds form between polar R groups and between amine and carboxylic acid functional groups of closely positioned amino acids. 9. Ionic bonds form between negatively charged and positively charged R groups. 10. Disulfide bonds form between the sulfhydryl groups of two cysteine amino acids. 11. The secondary structure of a protein is a series of repeating patterns within the protein. (Figure 2.26b, p. 66) 12. There are two possible secondary repeating patterns: alpha helix and beta-pleated sheet. 13. An alpha helix pattern is a spiral coiling arrangement of the protein, whereas the beta-pleated sheet is a planar arrangement. 14. The alpha helix gives some elasticity to fibrous proteins that are located in skin or hair. 15. The beta-pleated sheet gives some degree of flexibility to many globular proteins. 16. The tertiary structure of a protein is the final three-dimensional shape exhibited by a completed polypeptide chain. (Figure 2.26c, p. 66) 17. Two categories of proteins, either fibrous or globular, are distinguished by their molecular shape. 18. Globular proteins fold into a compact, often nearly spherical shape such as enzymes, antibodies, and some hormones. 19. Fibrous proteins are extended linear molecules that are constituents of ligaments, tendons, and contractile proteins within muscle cells. 13 20. The quaternary structure of a protein is present only in those proteins with two or more polypeptide chains; hemoglobin is an example of a protein in the quaternary structure. (Figure 2.26, p. 66) 21. The normal function of a protein may also require a prosthetic group which is a nonprotein structure covalently bonded to a protein; heme in hemoglobin is an example of a prosthetic group. 22. The biological activity of a protein is usually disturbed or terminated when its conformation is changed, termed denaturation. 23. Denaturation occurs when a protein is subjected to a nonoptimal chemical environment, such as improper pH and improper temperature. (Figure 2.27, p. 67) 14