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CHAPTER
2
Chemistry Comes Alive
Objectives
Part 1: Basic Chemistry
Definition of Concepts: Matter and Energy
2. Describe the major energy forms.
Composition of Matter: Atoms and Elements
3. Define chemical element and list the four elements that form the bulk of body matter.
How Matter Is Combined: Molecules and Mixtures
6. Define molecule, and distinguish between a compound and a mixture.
Chemical Bonds
8. Explain the role of electrons in chemical bonding and in relation to the octet rule.
Chemical Reactions
11. Define the three major types of chemical reactions: synthesis, decomposition, and
exchange. Comment on the nature of oxidation-reduction reactions and their importance.
Part 2: Biochemistry
Inorganic Compounds
14. Explain the importance of water and salts to body homeostasis.
Organic Compounds
16. Describe and compare the building blocks, general structures, and biological functions of
carbohydrates and lipids.
17. Explain the role of dehydration synthesis and hydrolysis in forming and breaking down
organic molecules.
Suggested Lecture Outline
Part 1: Basic Chemistry
I. Definition of Concepts: Matter and Energy (pp. 23–25)
A. Matter is anything that occupies space and has mass (p. 24).
1. Mass is equal to the amount of matter in the object.
B. States of Matter (p. 24)
1. Matter exists in one of three states: solid, liquid, or gas.
C. Energy (pp. 24–25)
1. Energy is the capacity to do work, and it exists in two forms.
2. Forms of Energy
a. Chemical energy is energy stored in chemical bonds.
II. Composition of Matter: Atoms and Elements (pp. 25–28; Figs. 2.1–2.3;
Table 2.1)
A. Basic Terms (p. 25; Table 2.1)
1. Elements are unique substances that cannot be broken down into simpler substances
by ordinary chemical means.
B. Atomic Structure (pp. 25–27; Figs. 2.1–2.2)
1. Each atom has a central nucleus with tightly packed protons and neutrons.
a. Protons have a positive charge and weigh 1 atomic mass unit (amu).
b. Neutrons do not have a charge and weigh 1amu.
C. Identifying Elements (pp. 27–28; Fig. 2.3)
1. Elements are identified based on their number of protons, neutrons, and electrons.
2. The atomic number of an element is equal to the number of protons of an element.
a. Because the number of protons is equal to the number of electrons, the atomic
number indirectly tells us the number of electrons.
III. How Matter Is Combined: Molecules and Mixtures (pp. 28–30; Fig. 2.4)
A. Molecules and Compounds (pp. 28–29)
1. A combination of two or more atoms is called a molecule.
2. If two or more atoms of the same element combine it is called a molecule of that
element.
B. Mixtures (pp. 29–30; Fig. 2.4)
1. Mixtures are substances made of two or more components mixed physically.
2. Solutions are homogeneous mixtures of compounds that may be gases, liquids, or
solids.
a. The substance present in the greatest amount is called the solvent.
b. Substances present in smaller amounts are called solutes.
4. Suspensions are heterogeneous mixtures with large, often visible solutes that tend to
settle out.
C. Distinguishing Mixtures from Compounds (p. 30)
1. The main difference between mixtures and compounds is that no chemical bonding
occurs between molecules of a mixture.
IV. Chemical Bonds (pp. 30–35; Figs. 2.5–2.10)
A. A chemical bond is an energy relationship between the electrons of the reacting atoms
(p. 30; Fig. 2.5).
1. The Role of Electrons in Chemical Bonding (p. 31)
a. Electrons occupy regions of space called electron shells that surround the nucleus in
layers.
B. Types of Chemical Bonds (pp. 31–35; Figs. 2.6–2.10)
1. Ionic bonds are chemical bonds that form between two atoms that transfer one or more
electrons from one atom to the other.
2. Covalent bonds form when electrons are shared between two atoms.
a. Some atoms are capable of sharing two or three electrons between them, resulting
in double covalent or triple covalent bonds.
b. Nonpolar molecules share their electrons evenly between two atoms.
3. Hydrogen bonds are weak attractions that form between partially charged atoms found
in polar molecules.
a. Surface tension is due to hydrogen bonds between water molecules.
V. Chemical Reactions (pp. 35–38; Fig. 2.11)
A. Chemical Equations (pp. 35–36)
1. Chemical reactions occur whenever bonds are formed, rearranged, or broken.
2. Chemical Equations
B. Patterns of Chemical Reactions (pp. 36–37; Fig. 2.11)
1. In a synthesis (combination) reaction, larger molecules are formed from smaller
molecules.
C. Energy Flow in Chemical Reactions (p. 37)
1. Exergonic reactions release energy as a product, while endergonic reactions absorb
energy.
D. Reversibility of Chemical Reactions (p. 37)
1. All chemical reactions are theoretically reversible.
E. Factors Influencing the Rate of Chemical Reactions (pp. 37–38)
1. Chemicals react when they collide with enough force to overcome the repulsion by
their electrons.
Part 2: Biochemistry
VI. Inorganic Compounds (pp. 38–41; Figs. 2.12–2.13)
A. Water (pp. 38–39)
1. Water is the most important inorganic molecule, and makes up 60–80% of the volume
of most living cells.
B. Salts (p. 39; Fig. 2.12)
1. Salts are ionic compounds containing cations other than H+ and anions other than the
hydroxyl (OH–) ion.
C. Acids and Bases (pp. 39–41; Fig. 2.13)
1. Acids are also known as proton donors and dissociate in water to yield hydrogen ions
and anions.
2. Bases are also called proton acceptors and absorb hydrogen ions.
4. Neutralization occurs when an acid and a base are mixed together. They react with
each other in displacement reactions to form a salt and water.
5. Buffers resist large fluctuations in pH that would be damaging to living tissues.
VII. Organic Compounds (pp. 41–56; Figs. 2.14–2.24; Tables 2.2–2.4)
A. Carbohydrates, lipids, proteins, and nucleic acids are molecules unique to living systems,
and all contain carbon, making them organic compounds (pp. 41–43).
B. Carbohydrates (p. 43; Fig. 2.15)
1. Carbohydrates are a group of molecules including sugars and starches.
C. Lipids (pp. 43–47; Fig. 2.16; Table 2.2)
1. Lipids are insoluble in water, but dissolve readily in nonpolar solvents.
2. Triglycerides (neutral fats) are commonly known as fats when solid and oils when
liquid.
D. Proteins (pp. 47–53; Figs. 2.17–2.21; Table 2.3)
1. Proteins compose 10–30% of cell mass.
a. They are the basic structural material of the body.
2. Proteins are long chains of amino acids connected by peptide bonds.
3. Proteins can be described in terms of four structural levels.
a. The linear sequence of amino acids is the primary structure.
4. Fibrous and Globular Proteins
a. Fibrous proteins are extended and strandlike. They are known as structural proteins
and most have only secondary structure.
5. Protein denaturation is a loss of the specific three-dimensional structure of a protein. It
may occur when globular proteins are subjected to a variety of chemical and physical
changes in their environment.
6. Molecular chaperones, or chaperonins, are a type of globular protein that help proteins
achieve their three-dimensional shape.
7. Enzymes and Enzyme Activity
E. Nucleic Acids (DNA and RNA) (pp. 53–55; Fig. 2.22; Table 2.4)
1. Nucleic acids composed of carbon, oxygen, hydrogen, nitrogen, and phosphorus are
the largest molecules in the body.
2. Nucleotides are the structural units of nucleic acids.
3. Each nucleotide consists of three components: a pentose sugar, a phosphate group, and
a nitrogen-containing base.
6. RNA, or Ribonucleic Acid
a. RNA is located outside the nucleus and is used to make proteins using the instruc-
tions provided by the DNA.
F. Adenosine Triphosphate (ATP) (pp. 55–56; Figs. 2.23–2.24)
1. ATP is the energy currency used by the cell.
Cross References
Additional information on topics covered in Chapter 2 can be found in the chapters listed below.
1. Chapter 3: Phospholipids in the composition and construction of membranes; DNA
replication and roles of DNA and RNA in protein synthesis; cellular ions; enzymes and
proteins in cellular structure and function; hydrogen bonding
of proteins, carbohydrates, and lipids
7. Chapter 24: Oxidation-reduction reaction; importance of ions (minerals) in life processes;
metabolism of carbohydrates, lipids, and proteins; basic chemistry of life examples
8. Chapter 25: Renal control of electrolytes
Lecture Hints
1. Introduction to Chemistry for Biology Students, Ninth Edition, by George Sackheim, is
an excellent aid for students who need a quick brushup in chemistry or for those who
2. As an alternative to presenting the chemistry in Chapter 2 as a distinct block of material,
you could provide the absolute minimum coverage of the topics at this time and expand
upon topics later as areas of application are discussed.
3. Students often find the concept of isotopes confusing. A clear distinction between atomic
mass and atomic weight will help clarify the topic.
5. Oxidation-reduction reactions involve the loss and gain of electrons. The reactant oxi-
6. In biological oxidation-reduction reactions the loss and gain of electrons is often associ-
ated with the loss and gain of hydrogen atoms. Electrons are still being transferred since
the hydrogen atom contains an electron.
7. The relationship between the terms catalyst and enzyme can be clarified by asking the
students if all enzymes are catalysts and if all catalysts are enzymes.
9. The notion that ATP is the “energy currency” of the cell should be emphasized. Students
10. The cycling back and forth between ATP and ADP is a simple but important concept
often overlooked by students.
Activities/Demonstrations
1. Audiovisual materials are listed in the Multimedia in the Classroom and Lab section of
this Instructor Guide (p. 387).
3. Bring in materials or objects that are composed of common elements, e.g., a gold chain,
coal, copper pipe, cast iron. Also provide examples of common compounds such as
water, table salt, vinegar, and sodium bicarbonate. Solicit definitions of atom, element,
7. Obtain an electrolyte testing system (lightbulb setup connected to electrodes) and prepare
a series of solutions such as salt, acid, base, glucose, etc. Place the electrodes into the
solutions to illustrate the concept of electrolytes.
10. Use a metal or plastic “coil” toy to demonstrate denaturation of an enzyme. Tie colored
yarn on the coil at two sites that are widely separated, and then twist the coil upon itself
to bring the two pieces of yarn next to each other. Identify the site where the yarn pieces
are as the active site. Then remind students that when the hydrogen bonds holding the
enzyme (or structural protein) in its specific 3-D structure are broken, the active site (or
structural framework) is destroyed. Untwist the coil to illustrate this point.
Critical Thinking/Discussion Topics
1. Discuss how two polysaccharides, starch and cellulose, each having the same subunit
(glucose), have completely different properties. Why can we digest starch but not
cellulose?
4. Describe how weak bonds can hold large macromolecules together.
5. Why can we state that most of the volume of matter, such as the tabletop you are writing
on, is actually empty space?
Library Research Topics
1. Explore the use of radioisotopes in the treatment of cancers.
2. Study the mechanisms by which DNA can repair itself.
3. Locate the studies of Niels Bohr concerning the structure of atoms and the location of
electrons. Determine why his work with hydrogen gas provided the foundation of our
knowledge about matter.
6. What are the problems associated with trans fatty acids in the diet? How has awareness
of these effects changed our food practices?
7. Virtually every time an amino acid chain consisting of all 20 amino acids is formed in the
cell, it twists into an alpha helix, then folds upon itself into a glob. Why?
List of Figures and Tables
All of the figures in the main text are available in JPEG format, PPT, and labeled & unlabeled
format on the Instructor Resource DVD. All of the figures and tables will also be available in
Transparency Acetate format. For more information, go to www.pearsonhighered.com/educator.
Figure 2.1 Two models of the structure of an atom.
Figure 2.2 Atomic structure of the three smallest atoms.
Figure 2.9 Ionic, polar covalent, and nonpolar covalent bonds compared along a
continuum.
Figure 2.10 Hydrogen bonding between polar water molecules.
Figure 2.11 Patterns of chemical reactions.
Figure 2.12 Dissociation of salt in water.
Figure 2.13 The pH scale and pH values of representative substances.
Figure 2.14 Dehydration synthesis and hydrolysis.
Figure 2.21 Mechanism of enzyme action.
Figure 2.22 Structure of DNA.
Figure 2.23 Structure of ATP (adenosine triphosphate).
Answers to End-of-Chapter Questions
Multiple-Choice and Matching Question answers appear in Appendix H of the main text.
Short Answer Essay Questions
23. Energy is defined as the capacity to do work, or to put matter into motion. Energy has no
mass, takes up no space, and can be measured only by its effects on matter. Potential
energy is the energy an object has because of its position in relation to other objects.
Kinetic energy is energy associated with a moving object. (p. 24)
27. a. Add molecular weight of all atoms: 9 x 12 (C) + 8 x 1 (H) + 4 x 16 (O) = 180 g.
b. Total molecular weight equals the number of grams in one mole, in this case 180.
c. Divide the number of grams in the bottle by the number of grams in one mole of
aspirin. This equals the total number of moles in the bottle.
d. Answer = 2.5 moles (p. 30)
28. a. Covalent. b. Covalent c. Ionic (pp. 32–33)
31. Primary structure—linear sequence of amino acids in a polypeptide chain; second struc-
ture—coiling of primary structure into alpha helix or ß-pleated sheet; tertiary structure—
folding of alpha helices or beta-pleated sheets into a ball-like, or globular, molecule.
(pp. 48–50)
32. Dehydration refers to the joining together of two molecules by the removal of water.
Monosaccharides are joined to form disaccharides and amino acids are joined to form
33. Enzymes decrease activation energy and decrease the randomness of reactions by binding
34. Molecular chaperones are proteins that aid the folding of other proteins into their
35. The surface tension of water tends to pull water molecules into a spherical shape, and
since the glass does not completely overcome this attractive force, water can elevate
slightly above the rim of the glass. (pp. 34–35)
Critical Thinking and Clinical Application Questions
1. In a freshwater lake, there are comparatively few electrolytes (salts) to carry a current
away from a swimmer’s body. Hence, the body would be a better conductor of the cur-
rent and the chance of a severe electrical shock if lightning hit the water is real. (p. 39)
2. a. Some antibiotics compete with the substrate at the active site of the enzyme. This
would tend to reduce the effectiveness of the reaction.
3. a. pH is defined as the measurement of the hydrogen ion concentration in a solution. The
4. The blood pH is rising, thus becoming more basic or alkaline. This is due to changes in
the carbonic acid-bicarbonate buffer system in the blood. Hyperventilation will cause an
increase in blood pH by reducing the amount of carbonic acid in the blood. (p. 39)
Suggested Readings
Gorman, Jessica. “Getting Out the Thorn: Biomaterials Become Friendlier to the Body.”
Science News 161 (1) (Jan. 2002): 13–14.
