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Chapter 2 The Chemistry of Life
Chapter Outline
Module 2.1 Atoms and Elements (Figures 2.1–2.2)
A. Define matter:
List the three states in which matter can exist:
, or
,
. Define chemistry:
B. Atoms and Atomic Structure (Figure 2.1):
1. Define atom:
2. Subatomic particles exist in 3 basic forms: protons, neutrons, electrons.
a. Protons (p+) are found in the central core of the atom known as the
atomic nucleus. What charge do protons carry?
b. Neutrons (n0), also found in the atomic nucleus, are slightly larger
than protons. What charge do neutrons carry?
c. Electrons (e-) are found outside the atomic nucleus. What charge do
electrons carry?
d. All atoms are electrically neutral, meaning they have no charge. The
number of protons and electrons are equal, cancelling out each
subatomic particles charge. The number of neutrons does not have to
be the same as the protons.
3. Electron shells, regions surrounding the atomic nucleus where the likelihood
or probability that an electron may exist, can hold a certain number of
electrons.
a. The 1st shell closest to the nucleus can hold
electrons.
b. The 2nd shell can hold
electrons.
c. The 3rd shell can hold
electrons but is satisfied with 8.
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d. Some atoms may have more than 3 shells.
C. Elements in the Periodic Table and the Human Body (Figure 2.2)
1. The number of
that an atom has in its nucleus is its
atomic number.
2. The atomic number defines an element.
a. An element is a substance that cannot be broken down into a simpler
substance by chemical means.
b. Each element is made of atoms with the same number of protons.
3. The periodic table of the elements lists the elements by their increasing
atomic numbers.
a. This organizes elements into groups with certain properties.
b. Each element is represented by a chemical symbol.
4. What are the four major elements that make up the human body?
,
,
, and
. The human body is also made of 7 mineral elements
and 13 trace elements.
D. Isotopes and Radioactivity:
1. The mass number is equal to the sum of all the
and
found in the atomic nucleus.
2. An isotope is an atom with the same
of
number of
, but different
number and same number
number and different
.
3. Radioisotopes are unstable isotopes have high energy or radiation that can be
released by radioactive decay. This allows the isotope to assume a more stable
form.
Module 2.2 Matter Combined: Mixtures and Chemical Bonds (Figures 2.3–2.7;
Tables 2.1-2.2)
A. Matter can be combined physically to form a mixture. What is a mixture?
B. Mixtures (Figure 2.3): List and describe the 3 types of mixtures.
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1.
2.
3.
C. Chemical Bonds: Matter can be combined chemically to form a molecule, which
is when the atoms of two or more elements are combined by forming chemical
bonds. A chemical bond is not a physical structure, but rather an energy
relationship or an attractive force between atoms
1. Define molecule.
Define compound.
2. Very large molecules composed of many atoms are called macromolecules.
3. Molecular formulas are a way to represent molecules symbolically with letters
and numbers to show the kinds and numbers of atoms in a molecule.
D. Chemical bonds are formed when valence electrons in the valence shell, the
outermost electron shell, of atoms interact.
1. Valence electrons determine how an atom interacts with other atoms and
whether it will form bonds with a specific atom.
2. Define the octet rule.
3. Define the duet rule.
E. Ions and Ionic Bonds (Figure 2.4): An ionic bond is formed when electrons are
from a metal atom to a nonmetal atom and results in the
formation of ions: cations and anions. The attraction between opposite charges
holds or bonds the atoms to one another, forming a compound called a salt.
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1. What is a cation?
2. What is an anion?
F. Covalent Bonds (Figures 2.5, 2.6; Table 2.1): Covalent bonds, the strongest
bond, form when two or more nonmetals
electrons between
themselves.
1. Two atoms can share one (single bond), two (double bond), or three (triple
bond) electron pairs.
2. All elements have protons that can attract electrons, a property known as
electronegativity.
a. An element’s electronegativity increases from the bottom left to the
upper right of the periodic table, making fluorine (F) the most
electronegative element.
b. The more electronegative an element the more strongly it attracts
electrons, pulling them away from less electronegative elements.
3. Nonpolar covalent bonds result when two nonmetals in a molecule with
similar or identical electronegativities pull with the same force and share the
electrons equally.
4. Nonpolar molecules occur in 3 situations:
a.
(Figure 2.6a)
b.
c.
G. Polar covalent bonds form polar molecules when nonmetals with different
electronegativities interact, resulting in an unequal sharing of electrons (Figure
2.6b).
1. The atom with the higher electronegativity becomes partially
(d-) as it pulls and holds the shared electrons close to itself.
2. The atom with the lower electronegativity becomes partially
(d+) as it allows the shared electrons to be pulled away toward the other atom.
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3. Polar molecules with partially positive and partially negative ends are known
as dipoles.
H. Hydrogen bonds are
attractions between the partially positive end
of one dipole and the partially negative end of another dipole (Figure 2.7).
1. Hydrogen bonds are responsible for a key property of water: surface tension.
2. Where air and water meet, the polar water molecules are more strongly
attracted to one another than they are to nonpolar air molecules.
Module 2.3 Chemical Reactions (Figures 2.8–2.10)
A. Chemical Notation: Chemical notation is a series of symbols and abbreviations
that is used to demonstrate what occurs in a reaction. The chemical equation, the
basic form of chemical notation, has two parts:
1.
2.
B. Energy and Chemical Reactions: A chemical reaction has occurred every time a
chemical bond is formed, broken, or rearranged, or when electrons are transferred
between two or more atoms (or molecules).
1. Define energy.
2. Two general forms of energy are:
a.
energy is stored, but can be released to do work at
some later time.
b.
energy is potential energy that has been released or
set in motion to perform work. All atoms have kinetic energy as they
are in constant motion and the faster they move the greater that energy.
3. Energy is found in 3 forms in the human body: 1)
, and 3)
, 2)
, each of which may be
potential or kinetic depending on the location or process. Describe each type
of energy.
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a. Chemical energy
b. Electrical energy
c. Mechanical energy
4. Energy, inherent in all chemical bonds, must be invested any time a chemical
reaction occurs. Describe the following reactions:
a. Endergonic reactions
b. Exergonic reactions
C. Homeostasis and Types of Chemical Reactions: Three fundamental processes
that occur in the body to maintain homeostasis, breaking down molecules,
converting the energy in food to a usable form, and building new molecules, are
carried out by one of three basic types of chemical reactions:
1. Catabolic reactions (decomposition reactions) occur when
a. The general chemical notation for this reaction is AB A+B.
b. These are usually exergonic because chemical bonds are broken.
2. Exchange reactions occur when
a. The general chemical notation for this reaction is AB + CD AD +
BC.
b. Oxidation-reduction reactions (redox reactions), a special kind of
exchange reaction, occur when electrons and energy are exchanged
instead of atoms. The reactant that loses electrons is oxidized while the
reactant that gains electrons is reduced.
c. Redox reactions are usually exergonic reactions capable of releasing
large amounts of energy
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3. Anabolic reactions (synthesis reactions) occur when
a. The general chemical notation for this reaction is A + B 
AB.
b. These reactions are endergonic, fueled by chemical energy.
D. Reaction Rates and Enzymes (Figure 2.8, 2.9, 2.10): For a reaction to occur,
atoms must collide with enough energy overcome the repulsion of their electrons.
This energy required for all chemical reactions is called the
(Ea). The following factors increase the reaction
rate by either reducing the activation energy or increasing the likelihood of strong
collisions between reactants: concentration, temperature, reactant properties, and
the presence or absence of a catalyst. Describe each of these factors that
influence reaction rate.
1. Concentration:
2. Temperature:
3. Particle size and phase:
a.
b.
4. Catalysts:
5. Most enzymes are macromolecule proteins. Summarize the properties of
enzymes.
a.
b.
c.
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d.
6. Induced-fit mechanism:
(Figure 2.10)
Module 2.4 Inorganic Compounds: Water, Acids, Bases, and Salts Bonds (Figures
2.11–2.13)
A. Biochemistry, grouped into inorganic and organic compounds, is the chemistry of
life.
1. Inorganic compounds generally do not contain
bonded to
hydrogen, and include water, acids, bases, and salts.
2. Organic compounds are defined as those that do contain
bonded to hydrogen.
B. Water (Figure 2.11): Water (H2O) makes up 60-80% of the mass of the human
body and has the several key properties vital to our existence.
1. Summarize the properties of water.
a.
b.
c.
d.
2. Water serves as the body’s primary solvent and is often called the universal
solvent because so many solutes will dissolve in it entirely, or to some degree
(Figure 2.11).
a. Water is a polar covalent molecule where the oxygen pole is partially
negative (d-) and the hydrogen pole is partially positive (d+), which
allows the molecule to interact with certain solutes, surround them,
and keep them apart.
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b. Water is only able to dissolve solutes that are
or those with fully or partially charged ends. “Like dissolves like”:
water generally dissolves ionic and polar covalent solutes.
c. Solutes that do not have full or partially charged ends are known as
and do not dissolve in water. This group includes
uncharged nonpolar covalent molecules such as oils and fats.
C. Acids and Bases (Figures 2.12, 2.13): The study of acids and bases is really the
study of the hydrogen ion (H+) A water molecule in a solution may break apart
or dissociate into a positively charged hydrogen ion and a negatively charged
hydroxide ion (OH-). Acids and bases are defined in the following way
according to their behavior with respect to hydrogen ions:
1. Define acid.
(Figure 2.12b)
2. Define base.
(Figure 2.12c)
3. The pH scale, ranging from
-
, is a simple way of representing the
hydrogen ion concentration of a solution and literally stands for negative
logarithm of the hydrogen ion concentration, or -log [H+] (Figure 2.13).
a. When the pH = 7 the solution is
, where the number of
hydrogen ions and base ions are equal.
b. A solution with a pH less than 7 is
, where hydrogen
ions outnumber base ions.
c. A solution with a pH greater than 7 is
or
, where base ions outnumber hydrogen ions.
d. Most body fluids are slightly basic pH: blood,
-
; inside cells 7.2.
4. What is a buffer?
D. Salts and Electrolytes: A
refers any metal cation and nonmetal anion
held together by ionic bonds. Salts can dissolve in water to form cations and
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anions called
, which are capable of conducting an electrical
current.
Module 2.5 Organic Compounds: Carbohydrates, Lipids, Proteins, and Nucleotides
(Figures 2.14–2.26; Table 2.3)
A. Monomers and Polymers: Each type of organic compound in the body,
carbohydrate, lipid, protein, or nucleic acid, has its own monomer or single
subunit and the corresponding polymer built from those subunits.
1. Describe dehydration synthesis.
2. Describe hydrolysis.
B. Carbohydrates (Figures 2.14, 2.15, 2.16; Table 2.3): Carbohydrates, composed
of carbon, hydrogen, and oxygen, function primarily as fuel in the body with
some limited structural roles.
1. Monosaccharides have from 3 to 7 carbons and are the monomers from
which all carbohydrates are made. List examples of the most abundant
monosaccharides in the body.
,
,
, and
,
(Figure 2.14).
2. Disaccharides are formed by union of two monosaccharides by dehydration
synthesis (Figure 2.15).
3. Polysaccharides consist of many monosaccharides joined to one another by
dehydration synthesis reactions (Figure 2.16).
a. Glycogen is the storage polymer for
found mostly in
skeletal muscle and liver cells.
b. Some polysaccharides are found covalently bound to either proteins or
lipids resulting in the following compounds glycoproteins and
glycolipids, which have various functions in the body.
C. Lipids (Figures 2.17, 2.18, 2.19, 2.20; Table 2.3): The lipids, a group of
nonpolar hydrophobic molecules composed primarily of carbon and hydrogen,
include fats and oils.
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1. Fatty acids are the basic lipid monomers consisting of 4 to 20 carbon atoms
which may have none, one, or more double bonds between the carbons in the
hydrocarbon chain (Figure 2.17).
a. Saturated fatty acids,
at room temperature, have
double bonds between carbon atoms so the carbons are
“saturated” with the maximum number of hydrogen atoms (Figure
2.17a).
b. Monounsaturated fatty acids, generally
temperature, have
at room
double bond between two carbons in the
hydrocarbon chain (Figure 2.17b).
c. Polyunsaturated fatty acids, liquid at room temperature, have
separate double bonds between carbons in the
hydrocarbon chain (Figure 2.17c).
2. Three fatty acids linked by dehydration synthesis to a modified 3-carbon
carbohydrate, glycerol, form a triglyceride, the storage polymer for fatty
acids also called a neutral fat (Figure 2.18).
3. Phospholipids are composed of a glycerol backbone, two fatty acid “tails”,
and one phosphate group “head” in place of the third fatty acid (Figure 2.19).
a. A molecule with a polar group, the phosphate head, and a nonpolar
group, the fatty acid tail, is called amphiphilic.
b. This amphiphilic nature makes phospholipids vital to the structure of
cell membranes.
4. Steroids are nonpolar and share a four-ring hydrocarbon structure called the
steroid nucleus.
is the steroid that forms the basis for all
the other steroids in the body (Figure 2.20).
D. Proteins (Figures 2.21, 2.22, 2.23; Table 2.3): Proteins are macromolecules that
are involved in movement, function as enzymes, play structural roles, function in
the body’s defenses, and can be used as fuel.
1. Twenty different amino acids, the monomers of all proteins, can be linked by
bonds into polypeptides (Figure 2.21).
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2. Peptides are formed when two or more amino acids are linked together by
peptide bonds as a result of dehydration synthesis (Figure 2.22).
a.
consist of two amino acids, tripeptides have
three amino acids, andcontain 10 or more amino acids.
b.
consist of one or more polypeptide chains folded into
distinct structures that must be maintained to be functional.
3. There are two basic types of proteins classified according to their structure:
fibrous and globular (Figure 2.23).
a. Describe fibrous proteins.
b. Describe globular proteins.
4. The complex structure of a complete protein is divided into four levels
(Figure 2.23). Describe each level of protein structure.
a.
(Figure 2.23a)
b.
(Figure 2.23b)
c.
(Figure 2.23c)
d.
(Figure 2.23d)
5. Explain protein denaturation.
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E. Nucleotides and Nucleic Acids (Figures 2.24, 2.25, 2.26; Table 2.3):
Nucleotides are the monomers of nucleic acids, so named because of their
abundance in the nucleus of cells, and make up our genetic material.
1. The nucleotide structure is composed of 3 parts. List the three parts
(Figure 2.24a).
a.
b.
c.
2. There are two types of nitrogenous bases: purines and pyrimidines (Figure
2.24b).
a. Purines, a double-ringed molecule, include
(A) and
(G).
b. Pyrimidines, a single-ringed molecule, include
(U), and
(C),
(T).
3. Adenosine triphosphate (ATP), adenine attached to ribose and three
phosphate groups, is the main source of chemical energy in the body (Figure
2.25a).
a. ATP is synthesized from
and a
using energy from the oxidation of fuels
such as glucose. It is the potential energy in this “high-energy” bond
that can be released to as kinetic energy to do work (Figure 2.25b).
b. The production of large quantities of ATP requires oxygen, which is
the main reason why we breathe air.
4. Deoxyribonucleic acid (DNA) and ribonucleic acid (RNA) are the two main
nucleic acids that together are responsible for the storage and execution of the
genetic code (Figure 2.26).
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5. DNA, an extremely large molecule found in the nucleus of the cell, is
composed of two long chains that twist around each other to form a double
helix (Figure 2.26a). Other structural features of DNA include:
a. DNA contains the pentose sugar
, so called
because it is missing an oxygen-containing group found in ribose.
b. DNA contains the following bases:
,
, and
,
.
c. The two stands of the double helix are held together by hydrogen
bonding between the bases of each strand.
d. DNA exhibits complementary base pairing where the purine
always pairs with the pyrimidine T and the purine G always pairs with
the pyrimidine
.
e. A = T (where = denotes 2 hydrogen bonds) and C ≡ G (where ≡
denotes 3 hydrogen bonds). This arrangement is allowed because each
base faces the inside of the double helix as they run in opposite
directions.
f. DNA contains the genes that provide the recipe or code for protein
synthesis, the process of making every protein in the body.
6. RNA, a single strand of nucleotides, can move between the nucleus of a cell
and its cytosol and is critical to the making of proteins (Figure 2.26b).
a. RNA contains the pentose sugar
.
b. RNA contains the pyrimidine uracil instead of thymine, which still
pairs with adenine (A = U).
c. RNA copies the recipe for a specific protein found in a gene on DNA,
a process called
.
d. RNA is free to exit the nucleus to a location where protein synthesis
occurs then proceeds to direct the making of the protein from the
recipe, a process called
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.
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