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LEARNING OUTCOMES
2.1
Introduction
1.
2.2
2.3
Give examples of how the study of living materials requires an understanding of
chemistry.
Structure of Matter
2.
Describe how atomic structure determines how atoms interact.
3.
Describe the relationships among matter, atoms, and molecules.
4.
Explain how molecular and structural formulas symbolize the composition of
compounds.
5.
Describe the three types of chemical reactions.
6.
Explain acids, bases, and buffers.
7.
Define pH.
Chemical Constituents of Cells
8.
List the major groups of inorganic chemicals common in cells.
9.
Describe the general functions of the main classes of organic molecules in cells.
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2.1. INTRODUCTION
2.2
A.
Chemistry = the study of matter.
1.
Biochemistry = chemistry of living organisms.
B.
Matter = anything that occupies space and has mass; (i.e. solids, liquids, gases)
STRUCTURE OF MATTER
A.
Elements and Atoms
1.
Atom =the smallest particle of an element.
2.
Element = a basic chemical substance composed of atoms.
3.
Elements are represented by a 1 or 2 letter symbol that are shown in the
Periodic Table of the Elements in the Appendix.
4.
116 elements exist in nature, however only approximately 26 are naturally
occurring in humans.
5.
Learn the elements (and chemical symbols) listed in Table 2.2, page 52.
6.
The most abundant of the naturally occurring elements are carbon (C),
Hydrogen (H), Oxygen (O) and Nitrogen (N) = CHON.
B.
Atomic Structure
3 Subatomic Particles
See Fig 2.1 page 52 & Table 2.1, page 51.
1.
Proton =
a positively charged particle in the nucleus of an atom;
Mass = 1.
2.
Neutron =
an electrically neutral particle in the nucleus of an atom;
Mass = 1.
3.
Electron =
an electrically negative particle that revolves around the
nucleus; Mass = 0.
SUBATOMIC PARTICLE SUMMARY TABLE (Keyed at the end of the outline)
SUBATOMIC
PARTICLE
CHARGE
LOCATION
MASS
(WEIGHT)
4.
*
Atoms are neutral in charge - The number of protons is equal to the
number of electrons.
5.
The Atomic Number (A#) of an atom represents the number of protons in
its nucleus.
a.
A# of H = 1
b.
A# of He = 2
c.
A# of O = 8
6.
The Atomic Weight (AW) of an atom is nearly equal to the number of
protons plus the number of neutrons in its nucleus.
See Table 2.3, page 53, Atomic Structure of Elements 1-12.
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2.2. STRUCTURE OF MATTER
C.
Isotopes = atoms of an element that have the same A#'s, but different AW's (i.e.
same # of protons, different # of neutrons).
1.
The nuclei of some isotopes are stable;
2.
The nuclei of other isotopes are unstable and break apart to become more
stable;
a.
b.
c.
When the nucleus of an atom breaks apart, it releases radioactive
energy;
Radioactive isotopes have many biological uses.
a. See From Science to Technology 2.1, pages 54-55 re:
Radioactive Isotopes Reveal Physiology
See From Science to Technology 2.2, pages 56-57 re: Ionizing
Radiation: From the Cold War to Yucca Mountain, and Table 2A,
page 56 re: sources of ionizing radiation.
D.
Molecules and Compounds
1.
Two or more atoms may combine to form a molecule.
2.
A molecular formula represents the numbers and kinds of atoms in a
molecule.
3.
If atoms of the same element combine, they produce molecules of that
atom
a.
oxygen – O2
b.
nitrogen – N2
4.
If atoms of different elements combine, they form molecules of substances
called compounds.
a.
water – H2O
b.
glucose – C6H12O6
E.
Bonding of Atoms
1.
2.
The electrons of an atom are arranged in orbits, shells, or energy levels
around the central nucleus; See Fig 2.3, page 55.
A characteristic number of electrons fill each shell:
a.
2 electrons fill the first shell (closest to nucleus);
b.
8 electrons fill the second shell;
c.
8 electrons fill the third shell.
Example 1:
Sodium (Na):
Atomic Number = 11
# protons = 11
# electrons = 11
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2.2
STRUCTURE OF MATTER
E.
Bonding of Atoms
Example 2:
3.
Atomic Number= 17
# protons= 17
# electrons= 17
The way in which atoms react with one another (i.e. their chemical
properties) is based on the electrons in their outermost shell = VALENCE
ELECTRONS
a.
b.
c.
4.
Chlorine(Cl):
The outermost shell of an atom is called its valence shell.
Na has ________ valence electrons;
S has _________ valence electrons.
Summary/Overview:
Example 1:
Fluorine has an Atomic Number of 9. Draw an atom of
fluorine. How and why will fluorine react?
Example 2:
Argon has an Atomic Number of 18. Draw an atom of
argon. How and why will argon react?
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2.2
STRUCTURE OF MATTER
E.
Bonding of Atoms
5.
Atoms form bonds with other atoms to fill their outermost or valence
electron shell (energy level).
a.
6.
"Octet Rule" = except for the first energy level (which contain 2
electrons), atoms react with other atoms so they will have 8
electrons in their valence shell.
Ionic Bonds: See Fig 2.4, page 57.
a.
Ions = atoms that have lost or gained electrons to fill their valence
shell.
b.
cation = a positively charged ion (Na+)
c.
anion = a negatively charged ion (Cl-)
d.
An attraction exists between oppositely charged ions and an ionic
bond results.
Na+ (which is now the cation) has donated its outer electron to Cl- (which
now becomes the anion). Salts, such as table salt or sodium chloride, are
held together by ionic bonds.
7.
Covalent Bonds:
See Fig 2.5, page 58.
a.
A covalent bond is formed by the equal sharing of electrons
between atoms.
b.
a very strong bond
c.
Examples: See Fig 2.6, page 58.
o
H2 (molecular hydrogen)
o
O2 (molecular oxygen)
o H2O (a molecule of water; See Fig 2.7, page 58),
however see below.
d.
Polar Bonds: See Fig 2.8a, page 42.
○
A covalent bond is formed by the unequal sharing of
electrons between atoms.
○
strong bond
○
results in molecules that are polar
o
one end of the molecule is slightly positive, one end
of the molecule is slightly negative
○
within H20 (water) molecules
8.
Hydrogen Bonds: See Fig 2.8b, page 58.
a.
A hydrogen bond is a weak bond formed between hydrogen atoms
(that are covalently bonded to another atom) and another atom.
b.
Examples include interaction between water molecules and DNA
chains.
c.
These bonds are easily broken and put back together.
.
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2.2
STRUCTURE OF MATTER
CHEMICAL BOND SUMMARY TABLE (Keyed at the end of the outline)
TYPE OF BOND
DEFINITION
DESCRIPTION
EXAMPLE
1.
2.
2a.
3.
F.
Chemical Reactions
1.
Definition:
A chemical reaction occurs whenever chemical bonds are formed,
rearranged or broken.
2.
Three Types:
a.
Synthesis = the building of a large molecule (polymer) from
smaller building blocks (monomers);
o
constructive, anabolic reactions;
o
Bonds are formed which now hold chemical energy
o
Water is usually removed from building blocks to form
bond (DEHYDRATION);
Energy

A + B → A—B

H20
o
o
b.
Example = the building of a large protein (polymer) from
many smaller amino acids (monomers).
Decomposition = breaking a large molecule (polymer) down into
its building blocks (monomers);
o
destructive, catabolic, "digestive" reactions;
o
Bonds are broken
(EXERGONIC);
o
Water is used to break bonds (HYDROLYSIS);
o
H20

A−B → A + B

Energy
o
Example = digesting a large protein we eat into its amino
releasing
chemical
energy
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acid building blocks.
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2.2
STRUCTURE OF MATTER
F.
Chemical Reactions
2.
Three Types:
c.
Exchange Reactions involve degradation followed by synthesis.
A−B + C−D  A + B + C + D  A−C + B−D
*
Many chemical reactions are Reversible Reactions = products can
be changed back to reactants.
A + B  A—B
Chemical Reaction Comparison Table (Keyed at the end of the outline)
SYNTHESIS REACTIONS
DEGRADATION
REACTIONS
GENERAL DESCRIPTION
(Sentence)
DESCRIPTIVE TERMS
BOND FORMATION OR
BREAKING?
IS ENERGY REQUIRED
OR RELEASED?
NAME THAT TERM.
HOW IS WATER
INVOLVED?
NAME THAT TERM.
EXAMPLE IN HUMAN
METABOLISM
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2.2
STRUCTURE OF MATTER
G.
Acids, Bases and Salts
1.
These ions are referred to as electrolytes (charged particles). See Table
2.4, page 59.



a.



b.



c.
2.
Acids dissociate (ionize) in water to form:

a.
b.
c.


Electrolytes must be maintained within a very narrow range in our
blood and tissues (i.e. homeostasis);
Also see Appendix.
Needed for muscle contraction, nerve impulses, bone growth, et
cetera;
Examples include Na+, K+, Cl-, Ca+, PO4-; HCO3-, etc.
a hydrogen cation, H+, and
an anion.
Example = HCl (hydrochloric acid).
H2O

HCl → H+ + Cl-
3.
Bases dissociate (ionize) in water to form:
a
b.
c.
a hydroxyl anion, OH-, and
a cation.
Example = NaOH (sodium hydroxide).
H2O

NaOH → Na+ + OH-
4.
Salts: See Fig 2.9, page 59.
a.
Salts dissociate (ionize) into ions when dissolved in water.
o
an anion is formed and
o
a cation is formed.
o
Example =
NaCl in water.
H2O
↓
NaCl → Na+ + Cl-
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2.2
STRUCTURE OF MATTER
H.
Acid and Base Concentration: See Table 2.5, and Fig 2.10, page 60.
1.


2.
3.
4.
The relative concentrations of hydrogen ions and hydroxyl ions determine
the pH in our blood, fluids, and tissues.
pH in body = [H+] + [OH-] .
pH = -log[H+];
pH Scale ranges from 0 to 14.
0 ---------------------------7---------------------------14
acidic
[H ] > [OH-]
neutral
[H+] = [OH-]
+
5.
basic
[H+] < [OH-]
Physiologic pH = 7.4
a.
pH < 7.4 = acidosis; lethal below 7.0;
b.
pH > 7.4 = alkalosis; lethal above 7.8.
c.
Buffering Systems (Discussed in greater detail in Chapter 21)
Definition:
Buffers (are compounds added to solutions that)
prevent abrupt change in pH.
o
usually weak acids;
o
function by donating H+ when needed and by accepting H+
when in excess;
o
very important in biological systems!
o
Example = the carbonic acid (H2CO3) buffering system.
when pH is rising equation goes to the right

H2CO3
HCO3-
H+
+
when pH is falling equation goes to the left
carbonic acid
(H+ donor)
bicarbonate ion
(H+ acceptor)
hydrogen ion
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2.3
CHEMICAL CONSTITUENTS OF CELLS
A.
Introduction: Living organisms are composed of both inorganic and organic
matter.
1.
Inorganic Substances are small compounds that do not contain the atoms
C and H; Examples include oxygen, carbon dioxide (CO2) water, salts,
acids & bases.
2.
Organic Substances contains the atoms carbon (and hydrogen). They are
composed of small molecules covalently bonded together to form large
polymers or macromolecules. The four major classes found in cells
include carbohydrates, lipids, proteins, and nucleic acids
B.
Inorganic Substances are small compounds that do not contain the atoms C and
H; Examples include oxygen, carbon dioxide (CO2) water, salts, acids &
bases. See Table 2.6, page 62.
1.
Water is a polar molecule that demonstrates hydrogen bonding and
therefore it possesses very unique characteristics. See Fig 2.7 and 2.8,
page 58.
a.
Water is an excellent solvent (universal?)
o
Many solutes are dissolved in our body's water (i.e. polar
substances dissolve in polar water)
o
Many ionic compounds (i.e. NaCl) dissociate or break apart
in water.
b.
Water participates in many chemical reactions (in our cells and
fluids)
o
Dehydration (synthesis) is when water is removed from
adjacent atoms (of molecules) to form a bond between
them.
o
Hydrolysis (degradation) is when water is used to break
bonds between molecules.
c.
Water is an excellent temperature buffer.
o
absorbs and releases heat very slowly
d.
Water provides an excellent cooling mechanism.
o
It requires a lot of heat to change water from a liquid to a
gas (i.e. high heat of vaporization). If water does change
forms and evaporate, it leaves a cool surface behind.
e.
Water serves as a lubricant
o
mucus, internal organs, and joints
f.
Water is the most abundant component in cells (about 70%).
2.
3.
4.
Oxygen (O2)
a.
gas that is transported in the blood
b.
used to release energy from nutrient molecules
Carbon Dioxide (CO2)
a.
by-product of cellular respiration.
b.
See Fig 2.23b.
Inorganic salts. Many uses (See properties above)
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See Table 2.6, page 62 for Summary of the Inorganic Substances Common in Cells.
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2.3
CHEMICAL CONSTITUENTS OF CELLS
C.
Organic Substances:
1.
contains the atoms carbon (and hydrogen);
2.
small molecules (monomers or building blocks) are covalently bonded
together to form large polymers or macromolecules;
3.
Water is usually involved in the formation and breakage of bonds
between monomers;
4.
*
a.
Dehydration Synthesis = removal of water to form a covalent
bond between monomers;
b.
Hydrolysis = using water to break bonds between monomers.
The four major classes found in cells include:
a.
carbohydrates
b.
lipids
c.
proteins
d.
nucleic acids
See box on page 61 re: many meanings of the word “organic”.
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2.3
CHEMICAL CONSTITUENTS OF CELLS
C.
CARBOHYDRATES (sugars)
1.
contains C, H, and O, in a 1:2:1 ratio (usually);
(Ex: glucose = C6H12O6)
2.
Monomers (building blocks) are monosaccharides;
hexoses = simple 6-C sugars; See Fig 2.11 and 2.23a, page 63.
a.
b.
c.
3.
glucose
fructose
galactose
Polymers are formed by dehydration synthesis.
a.
Disaccharides: 2 monosaccharides covalently bonded together;
See Fig 2.12b, page 63.
o
o
o
b.
maltose = glucose + glucose
lactose = glucose + galactose
sucrose = glucose + fructose
Polysaccharides: many glucose molecules covalently bonded
together; See Fig 2.12c, page 63.
o
o
starch = plant storage carbohydrate;
glycogen = animal storage carbohydrate; stored in liver and
skeletal muscle.
*Polymers are broken down by hydrolysis resulting in monosaccharides.
4.
Function = energy source / energy storage!
*
*
How is the energy that is stored in carbohydrates released?
CELLULAR RESPIRATION OVERVIEW:
oxygen

glucose → H20 + CO2

energy (ATP)
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2.3
CHEMICAL CONSTITUENTS OF CELLS
D.
LIPIDS
1.
2.
3.
4.
contain C, H, and O, but much less O than in carbohydrates;
types of lipids:
a.
Fats: See Fig 2.13 and Fig 2.14, page 64. and Fig 2.15a, page 65.
o
monomers (building blocks) = triglycerides (glycerol + 3
fatty acids);
o
saturated vs. unsaturated fats: See Fig 2.13, page 64.
1.
saturated fats:
a. have only single bonds between the carbons in
their fatty acid chains
b. are solid at room temperature
c. are animal fats
d. include bacon grease, lard, butter
e. are nutritionally "BAD" fat
2.
unsaturated fats:
a. have one or more double bond between the
carbons in their fatty acid chains
b. are liquid at RT (oils)
c. are plant fats
d. include corn and olive oil
e. are nutritionally "GOOD" fat
3.
“trans” fats
a. Unsaturated fats that have been solidified by
artificial means
b. Not produced nor maintained in the body
c. Include margarine
d. The WORST type of fat
o
Function = energy storage / energy source
*
See box on page 63 re: A diet rich in fat increases
risk of atherosclerosis.
b.
Phospholipids:
See Fig 2.15b&c, page 65.
o
triglyceride with the substitution of a polar phosphate group
(PO4-) for one fatty acid chain;
o
Function = major cell membrane component.
c.
Steroids:
See Fig 2.16, page 65.
o
four interconnected carbon rings;
o
Example is cholesterol;
o
Function = compose cell membranes; chemical messengers
(hormones).
See Table 2.7, page 65, which compares and contrasts triglycerides,
phospholipids, and steroids.
Solubility
a.
Hydrophilic – soluble in water
b.
Lipophilic – soluble in lipid (also called Hydrophobic)
○ Important in membrane transport discussed in Chapter 3.
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2.3
CHEMICAL CONSTITUENTS OF CELLS
E.
PROTEINS
1.
Monomers = amino acids
a.
Amino Acid Structure (See Fig 2.17, page 66). Draw it below
amino group
b.
2.
carboxyl group
Types of amino acids
o
20 different based on R-groups or side-chains
Polymers are formed by dehydration synthesis between the amino group
of one amino acid and the carboxyl group of a 2nd amino acid.
See Fig 2.18, page 66.
a.
b.
3.
H
|
---- C --|
Rgroup
Bond formed = a peptide bond
Length of amino acid chains may vary:
o
peptide = 2-100 aa's;
o
polypeptide = 100-thousands aa's without a function
o
protein = 100-thousands of aa's with a specific function
Functions of Proteins
a.
structure
o
keratin in hair, nails,and skin
b.
transport
o
hemoglobin
c.
chemical messengers
o
hormones
o
neurotransmitters
d.
movement
o
actin and myosin in muscle
c.
defense
o
antibodies
e.
catalysts
o
ENZYMES
*
Biological catalysts, that increase the rate of chemical
reactions without being consumed by the reaction.
Discussed in greater detail in Chapter 4.
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2.3
CHEMICAL CONSTITUENTS OF CELLS
E.
PROTEINS
4.
Protein Structure (See Fig 2.19, page 67).
a.
Primary (1o) = sequence of amino acids;
b.
Secondary (2o) = twisting of amino acid chain; due to hydrogen
bonding;
c.
Tertiary (3o) = folding of the amino acid chain; due to ionic
bonds, disulfide bridges, and hydrophobic interactions;
d.
Quaternary (4o) = interactions between different amino acid
chains. See Fig 14.9, page 531. Hemoglobin is composed of four
amino acid chain condensed into a functional quaternary structure.
*
See box on page 66 re: faulty chloride pump in cell membrane
proteins that occur in Cystic Fibrosis.
5.
Denaturation of Proteins: the loss of 3-dimensional conformation
(shape) of a protein. This results in loss of function.
a.
Reasons for denaturation:
o
o
o
o
F.
extreme pH values;
●
At what pH do our enzymes work best?
extreme temperature values;
●
At what temperature do our enzymes work best?
harsh chemicals (disrupt bonding);
high salt concentrations.
●
At what osmotic pressure do our enzymes work
best?
NUCLEIC ACIDS
1.
Monomers = nucleotides;
a.
Nucleotide structure = 3 parts:
o
o
o
2.
See Fig 2.20, page 68.
pentose sugar (5-C);
nitrogenous base;

purine (double ring) or

pyrimidine (single ring);
phosphate group.
Polymers are formed by bonding between the sugar of one nucleotide and
the phosphate group of a second nucleotide = sugar/phosphate backbone;
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2.3
CHEMICAL CONSTITUENTS OF CELLS
F.
NUCLEIC ACIDS
3.
Types of Nucleic Acids
a.
DEOXYRIBONUCLEIC ACID = DNA

o
Structure:
See Fig 2.21b, page 68.
1.
Sugar = deoxyribose; See Fig 2.22b, page 68.
2.
Bases = adenine (A), thymine (T), cytosine (C),
guanine (G);
3.
double stranded (resembles ladder); strands held
together by H-bonds between bases on opposite
strands:
a.
A complements T (2 hydrogen bonds);
b.
C complements G (3 hydrogen bonds)
*
Also see Appendix.
4.
double helix (ladder is twisted).
o
Function = genetic material (i.e. genes, chromosomes).
●
DNA directs protein synthesis
o
DNA contains all necessary information needed to
sustain and reproduce life!
b.
Ribonucleic Acid = RNA
o

See Fig 2.21a, page 68.
Structure
1.
Sugar = ribose; See Fig 2.22a, page 68.
2.
Bases = A,G,C, and uracil (replaces thymine)
3.
single stranded.
4.
Function = transport DNA code during protein
synthesis.
*
Also see Appendix.
See Table 2.8, page 69 for a summary of the four organic compounds in cells.
Interesting Topics
A.
B.
Biomarkers Herald Disease – See introduction on page 51.
CT Scanning and PET Imaging – See From Science to Technology 2.3, pages 70-71.
CHAPTER SUMMARY – see pages 72-73.
CHAPTER ASSESSMENTS – see pages 73-74.
INTEGRATIVE ASSESSMENTS/CRITICAL THINKING– see page 74.
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ORGANIC MOLECULE SUMMARY TABLE (keyed at end of outline)
Organic
Molecule
Composed of
what atoms?
Building
Blocks
(monomers)
Specific types
& functions of
monomers
Specific types
and functions
of polymers
Other
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SUBATOMIC PARTICLE SUMMARY TABLE
SUBATOMIC
PARTICLE
CHARGE
LOCATION
MASS
(WEIGHT)
PROTON
POSITIVE
NUCLEUS
1 amu
NEUTRON
ZERO
(NEUTRAL)
NUCLEUS
1 amu
ELECTRON
NEGATIVE
SHELLS OR
ORBITALS
AROUND
NUCLEUS
0
CHEMICAL BOND SUMMARY TABLE
TYPE OF BOND
DEFINITION
DESCRIPTION
EXAMPLE
1. IONIC
when atoms lose or
gain electrons
becoming ions, and
then oppositely
charged ions are
attracted to one
another
bond is broken by
water
salts, NaCl
2. COVALENT
when 1 or more
pair(s) of electrons
is/are shared equally
by atoms
Very strong bond
the bonds holding
together a molecule
of CO2
2a. POLAR
when 1 or more
pair(s) of electrons
is/are shared
unequally by atoms
Strong bond
the bonds holding
together a molecule
of H2O
3. HYDROGEN
when a (slightly
positive) hydrogen
atom that is already
covalently bonded to
a molecule is
attracted to a slightly
negative atom.
Very weak bond; in
molecules whose
purpose is to easily
break and then come
back together
reactions between
water molecules (i.e.
ice to water to gas);
DNA chains
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Chemical Reaction Comparison Table
SYNTHESIS REACTIONS
DEGRADATION RXN'S
GENERAL
DESCRIPTION
Synthesis involves the
building of a large molecule
(polymer) from smaller
building blocks (monomer).
Degradation involves the
breakdown of a polymer into
individual monomers.
DESCRIPTIVE TERMS
building
constructive
anabolic
breakdown
digestive
decomposition
catabolic
BOND FORMATION OR
BREAKING?
Bonds are formed.
Bonds are broken.
IS ENERGY REQUIRED
OR RELEASED?
Energy is required to form
the bond.
Energy is released when the
bond is broken.
HOW IS WATER
INVOLVED?
NAME THAT TERM.
Water is released when the
bond is formed.
Dehydration
Water is required to break
the bond.
Hydrolysis
EXAMPLE
Building a protein from
individual amino acids;
Building a triglyceride from
glycerol and 3 fatty acids, etc
Breaking a protein into
individual amino acids;
Breaking starch down into
monosaccharides, etc.
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Full file at http://testbank360.eu/solution-manual-holes-human-anatomy-physiology-12th-editionshier
ORGANIC MOLECULE SUMMARY TABLE
Organic
Molecule
Carbohydrates
(sugars)
Lipids (Fats)
Proteins
Nucleic
Acids
Composed of
what atoms?
C, H, 0
C, H, O
C, H, O, N, S
C, H, O, N, P
Building
Blocks
(monomers)
Monosaccharide
or hexoses
Triglycerides:
glycerol and 3
fatty acids
amino acids
nucleotides:
pentose sugar,
phosphate,
nitrogen base
Specific types
& functions of
monomers
glucose,
fructose,
galactose.
TG: energy
20 different
amino acids
N/A
proteins (>100
amino acids);
Many functions:
ENZYMES,
antibodies,
structure,
transport,
chemical
messengers,
storage
DNA: deoxyribonucleic
acid; genetic
material; RNA:
ribonucleic
acid; aids DNA
in protein
synthesis.
energy
Specific types
and functions
of polymers
Other
Information
Disaccharides:
sucrose, lactose,
maltose; energy
_____________
Polysaccharides
Starch (plant);
Glycogen
(animal); energy
storage.
Phospholipid:
cell membrane
component
Steroid: cell
membrane
component and
chemical
messenger (i.e.
cholesterol)
N/A
Saturated (only
single bonds
between C’s in
FA chain) vs.
Unsaturated (at
least 1 double
bond in FA
chain)
DNA controls
cellular activity
by instructing
our cells what
proteins to
make (i.e.
Enzymes
through protein
synthesis).
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