Transcript Chapter 2 - Dr. Eric Schwartz

Chapter 02
Lecture Outline*
Chemical Composition
of the Body
Eric P. Widmaier
Boston University
Hershel Raff
Medical College of Wisconsin
Kevin T. Strang
University of Wisconsin - Madison
*See PowerPoint Image Slides for all
figures and tables pre-inserted into
PowerPoint without notes.
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Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Atoms: the Subunits of Elements
• Atoms are made of protons, neutrons, and
electrons.
• Each element has an atomic number.
– Equal to the number of protons contained
in the atom
• Each element has an atomic weight.
• Hydrogen, oxygen, carbon, nitrogen account
for >99% of the atoms in the human body.
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3
Components of Atoms
• The chemical properties of atoms can be
described in terms of three subatomic
particles—protons, neutrons, and electrons.
• The protons and neutrons are confined to a
very small volume at the center of an atom
called the atomic nucleus.
• The electrons revolve in orbitals at various
distances from the nucleus.
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Figure 2-1
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Atomic Number
• Each chemical element contains a specific
number of protons, and it is this number that is
known as the atomic number.
• Example: hydrogen has an atomic number of 1,
so it has a single proton.
• Because an atom is electrically neutral, the
atomic number is also equal to the number of
electrons in the atom.
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Atomic Weight
• Usually, the number of neutrons in the nucleus of an atom is
equal to the number of protons.
• However, many chemical elements can exist in multiple forms,
called isotopes, which differ in the number of neutrons they
contain.
• For example, the most abundant form of the carbon atom, 12C,
contains 6 protons and 6 neutrons, and has an atomic number
of 6 and an atomic weight of 12.
• The radioactive carbon isotope 14C contains 6 protons and 8
neutrons, giving it an atomic number of 6 but an atomic weight
of 14.
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Atomic Weight
• One gram atomic mass of a chemical element
is the amount of the element, in grams, equal
to the numerical value of its atomic weight.
Thus, 12 g of carbon is 1 gram atomic mass of
carbon.
• One gram atomic mass of any element
contains the same number of atoms (6 x1023;
Avogadro’s constant).
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Atomic Weight
• The atomic weight scale indicates an atom’s
mass relative to the mass of other atoms.
(comparison done to carbon)
• Because the atomic weight scale is a ratio of
atomic masses, it has no absolute units. The
unit of atomic mass is known as a dalton.
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Ions
• If an atom gains or loses one or more electrons, it acquires a net electric
charge and becomes an ion.
• Hydrogen atoms and most mineral and trace element atoms readily form
ions.
• Ions that have a net positive charge are called cations.
Examples: Ca2+, Na+
• Ions that have a net negative charge are called anions.
Example: Cl• Because of their charge, ions are able to conduct electricity when
dissolved in water; consequently, the ionic forms of mineral elements are
collectively referred to as electrolytes.
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Table 2-2
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Molecules
• Two or more atoms bonded together form
a molecule.
• Molecules can be represented by their
component atoms.
• Example: C6H12O6
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Chemical Bonds
Chemical bonds between atoms in a molecule form
when electrons transfer from the outer energy shell
of one atom to that of another, or when two atoms
with partially unfilled electron orbitals share
electrons.
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Covalent Bonds are the Strongest
Chemical Bonds
Fig. 2-2 14
Polar Covalent Bonds
• Electrons are not always shared equally
between two atoms, but instead reside close to
one atom of the pair.
• This atom thus acquires a slight negative
charge, while the other atom becomes slightly
positive.
• Such bonds are known as polar covalent bonds.
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Polar Covalent Bonds
• For example, the bond between hydrogen and oxygen in a hydroxyl group
(—OH) is a polar covalent bond in which the oxygen is slightly negative and
the hydrogen slightly positive.
• Atoms of oxygen, nitrogen and sulfur, which have a relatively strong
attraction for electrons, form polar bonds with hydrogen atoms (Table 2-3).
• One of the characteristics of polar bonds is that molecules that contain such
bonds tend to be more soluble in water than molecules containing the other
major type of covalent bond.
• Consequently, these polar molecules readily dissolve in the blood, interstitial
fluid, and intracellular fluid. Indeed, water itself is the classic example of a
polar molecule, with a partially negatively charged oxygen atom and two
partially positively charged hydrogen atoms.
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Nonpolar Covalent Bonds
• In contrast to polar covalent bonds, bonds between atoms with similar
electronegativities are said to be nonpolar covalent bonds. In such bonds,
the electrons are equally or nearly equally shared by the two atoms.
• Bonds between carbon and hydrogen atoms and between two carbon atoms
are electrically neutral, nonpolar covalent bonds (see Table 2–3).
• Molecules that contain high proportions of nonpolar covalent bonds are
called nonpolar molecules; they tend to be less soluble in water than those
with polar covalent bonds.
• Consequently, such molecules are often found in the lipid bilayers of the
membranes of cells and intracellular organelles. When present in body
fluids such as the blood, they may associate with a polar molecule that
serves as a sort of “carrier” to prevent the nonpolar molecule from coming
out of solution.
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Hydrogen Bonds
• When two polar molecules are in close contact, an electrical attraction may
form between them. For example, the hydrogen atom in a polar bond in one
molecule and an oxygen or nitrogen atom in a polar bond of another
molecule attract each other, forming a type of bond called a hydrogen
bond. Such bonds may also form between atoms within the same molecule.
• Hydrogen bonds are very weak, having only about 4 percent of the strength
of the polar bonds between the hydrogen and oxygen atoms in a single
molecule of water. Although hydrogen bonds are weak individually, when
present in large numbers, they play an extremely important role in
molecular interactions and in determining the shape of large molecules.
• Remember that the shape of large molecules often determines their
functions and their ability to interact with other molecules. For example,
some molecules interact with a “lock-and-key” arrangement that can only
occur if both molecules have the correct shape.
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Hydrogen Bonds Link
Adjacent Water Molecules
Fig. 2-3
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Molecular Shape
Figure 2-5
• Rotation around
chemical bonds
allows different
molecular shapes.
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Ionic Molecules
•
The process of ion formation (ionization) occurs in single atoms and atoms that are
covalently linked in molecules. Within molecules, two commonly encountered
groups of atoms that undergo ionization are the carboxyl group (—COOH) and the
amino group (—NH2).
•
The shorthand formula for only a portion of a molecule can be written as R—
COOH or R—NH2. The carboxyl group ionizes when the oxygen linked to the
hydrogen captures the hydrogen’s only electron to form a carboxyl ion (R—COO–),
releasing a hydrogen ion (H+): R—COOH 12 R —COO– + H+.
•
The amino group can bind a hydrogen ion to form an ionized amino group (R—
NH3+): R—NH2 + H+ 12 R—NH3+.
•
The ionization of each of these groups can be reversed.
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Free Radicals
• Free radicals are atoms or molecules with unpaired electrons in an outermost
orbital. Free radicals are unstable and highly reactive.
• Free radicals are formed by the actions of certain enzymes in some cells,
such as types of white blood cells that destroy pathogens.
• Free radicals are produced in the body following exposure to radiation or
toxin ingestion. These free radicals can do considerable harm to the cells of
the body. For example, oxidation due to long-term buildup of free radicals
has been proposed as one cause of several different human diseases, notably
eye, cardiovascular, and neural diseases associated with aging.
• Examples of biologically important free radicals are superoxide anion,
O2 · –; hydroxyl radical, OH · ; and nitric oxide, NO · .
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Solutions
• Substances dissolved in a liquid are known as solutes.
• The liquid in which solutes are dissolved is the
solvent.
• Solutes dissolve in a solvent to form a solution.
• Water is the most abundant solvent in the body,
accounting for approximately 60 percent of total body
weight.
• However, not all molecules dissolve in water.
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NaCl dissolves in water
Fig. 2-6
• In this example, Na+ & Cl- are the solutes and
water is the solvent.
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Solubility in Water
• Molecules with ionic or polar covalent bonds
have an electrical attraction to water
molecules.
• These molecules are able to dissolve in water
and are called hydrophilic (water loving).
• Molecules with nonpolar covalent bonds are
not able to dissolve in water and are called
hydrophobic (water fearing).
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Amphipathic Molecules
• Amphipathic molecules are a special class of molecules that have a polar or
ionized region at one end and a nonpolar region at the opposite end.
• When mixed with water, amphipathic molecules form clusters, with their
polar (hydrophilic) regions at the surface of the cluster where they are
attracted to the surrounding water molecules. The nonpolar (hydrophobic)
ends are oriented toward the interior of the cluster.
• This arrangement provides the maximal interaction between water
molecules and the polar ends of the amphipathic molecules. Nonpolar
molecules can dissolve in the central nonpolar regions of these clusters and
thus exist in aqueous solutions in far higher amounts than would otherwise
be possible based on their low solubility in water.
• The orientation of amphipathic molecules plays an important role in plasma
membrane structure.
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Amphipathic Molecules
Fig. 2-7
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Concentration
• Concentration is expressed as the amount
of solute (based on its molecular weight)
dissolved per liter of solution.
– The molecular weight in grams = 1 mole
– 1 M solution = 1 mole/Liter
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Acids and Bases
• Acids increase the concentration of H+ in
a solution.
• Bases decrease the concentration of H+ in
a solution.
• The amount of H+ in a solution is
expressed as pH.
pH = -log[H+]
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Types of Special Water Reactions
• Hydrolysis is the breakdown of a large
molecule into a small molecule by using water.
• Other types of reactions are:
– Dehydration
– Condensation
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Osmosis
• Water moves between fluid compartments by
the process of osmosis (covered more in
Chapter 4).
• In osmosis, water moves from regions of low
solute concentrations to regions of high solute
concentrations, regardless of the specific type
of solute.
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Terminology
• Organic chemistry is the study of carboncontaining molecules.
• Inorganic chemistry is the study of noncarboncontaining molecules.
• Biochemistry is the chemistry of living
organisms. (can be considered part of organic
chemistry)
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Classes of Organic Molecules
• Carbohydrates
– Monosaccharides
– Disaccharide
– Polysaccharides
• Lipids
–
–
–
–
Fatty Acids
Triglycerides
Phospholipids
Steroids
• Proteins
– Amino Acid Subunits
– Polypeptides
• Nucleic Acids
– DNA
– RNA
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Fig. 2-8
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Carbohydrates
• Monosaccharides are the simplest carbohydrates.
Fig. 2-9 37
Disaccharide Formation
Fig. 2-10
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Glycogen: a Polysaccharide
Fig. 2-11
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How the Body Uses Sugars
• Glycogen exists in the body as a reservoir of available energy
that is stored in the chemical bonds within individual glucose
monomers.
• Hydrolysis of glycogen, as occurs during periods of fasting,
leads to release of the glucose monomers into the blood,
thereby preventing blood glucose from decreasing to
dangerously low levels.
• Glucose is often called “blood sugar” because it is the major
monosaccharide found in the blood.
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Lipids
• Lipids are molecules composed predominantly (but not
exclusively) of hydrogen and carbon atoms.
• These atoms are linked by nonpolar covalent bonds. Thus,
lipids are nonpolar and have a very low solubility in water.
•
• Lipids can be divided into four subclasses: fatty acids,
triglycerides, phospholipids, and steroids.
• Lipids are important in physiology partly because some of
them provide a valuable source of energy. Other lipids are a
major component of all cellular membranes, and still others
are important signaling molecules.
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Lipids
Fig. 2-11
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Fatty Acids
• Fatty acids consist of a chain of carbon and hydrogen atoms
with an acidic carboxyl group at one end.
• When all the carbons in a fatty acid are linked by single
covalent bonds, the fatty acid is said to be a saturated fatty
acid.
• Some fatty acids contain one or more double bonds between
carbon atoms, and these are known as unsaturated fatty acids.
• If one double bond is present, the fatty acid is
monounsaturated, and if there is more than one double bond, it
is polyunsaturated.
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Triglycerides
• Triglycerides (also known as triacylglycerols) constitute the
majority of the lipids in the body.
• Triglycerides form when glycerol, a three-carbon alcohol,
bonds to three fatty acids.
• Triglycerides are found in all cells and comprise part of
cellular membranes, including those of intracellular organelles.
• They are also stored in great quantities in adipose tissue, where
they serve to supply energy to the cells of the body,
particularly during times when a person is fasting or requires
additional energy (maintained exercise, for example).
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Fig. 2-12A
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Fig. 2-12B
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Fig. 2-12C
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Cholesterol
Fig. 2-13
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Phospholipids
• Phospholipids are similar in overall structure to triglycerides,
but the third hydroxyl group of glycerol is linked to phosphate.
• In addition, a small polar or ionized nitrogen-containing
molecule is usually attached to this phosphate.
• These groups constitute a polar (hydrophilic) region at one end
of the phospholipid, whereas the fatty acid chains provide a
nonpolar (hydrophobic) region at the opposite end.
• Therefore, phospholipids are amphipathic. It is this property of
phospholipids that permits them to form the lipid bilayers of
cellular membranes.
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Steroids
• Steroids have a distinctly different structure from those of the
other subclasses of lipid molecules.
• Four interconnected rings of carbon atoms form the skeleton
of every steroid.
• Steroids are NOT water-soluble.
• Examples of steroids are cholesterol, cortisol from the adrenal
glands, and female (estrogen) and male (testosterone) sex
hormones secreted by the gonads.
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Proteins
• Proteins account for about 50 percent of the
organic material in the body (17 percent of the
body weight), and they play critical roles in
almost every physiological process.
• Proteins are composed of carbon, hydrogen,
oxygen, nitrogen, and small amounts of other
elements, notably sulfur. They are
macromolecules, often containing thousands of
atoms.
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Amino Acids
• The subunit monomers of proteins are amino acids.
• Every amino acid except proline has an amino (—NH2) and a
carboxyl (—COOH) group bound to the terminal carbon atom in the
molecule.
• The proteins of all living organisms are composed of the same set of
20 different amino acids, corresponding to 20 different side chains.
• The side chains may be nonpolar (8 amino acids), polar (7 amino
acids), or ionized (5 amino acids).
• The human body can synthesize many amino acids, but several must
be obtained in the diet; these are known as essential amino acids.
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Proteins are Made of Amino Acids
Fig. 2-14 53
Peptide Bonds
Fig. 2-15 54
Fig. 2-16
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Protein Structures
•
•
•
•
Primary Protein Structure
Secondary Protein Structure
Tertiary Protein Structure
Quaternary Protein Structure
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Primary Protein Structure
• Two variables determine the primary structure
of a protein:
(1) The number of amino acids in the chain
(2) The specific type of amino acid at each
position along the chain
• A polypeptide in the primary protein structure
is analogous to a linear string of beads, each
bead representing one amino acid.
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Conformation
• Proteins do not appear in nature like a linear
string of beads on a chain.
• Interactions between side groups of each
amino acid lead to bending, twisting, and
folding of the chain into a more compact
structure.
• The final shape of a protein is known as its
conformation.
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Secondary Protein Structure
•
The attractions between various regions along a polypeptide chain creates
secondary structure in a protein.
•
Because peptide bonds occur at regular intervals along a polypeptide chain, the
hydrogen bonds between them tend to force the chain into a coiled conformation
known as an alpha helix.
•
Hydrogen bonds can also form between peptide bonds when extended regions of a
polypeptide chain run approximately parallel to each other, forming a relatively
straight, extended region known as a beta pleated sheet.
•
The sizes of the side chains and the presence of ionic bonds between side chains
with opposite charges can interfere with the repetitive hydrogen bonding required to
produce these alpha and beta shapes and result in irregular regions called random
coil conformations. These occur in regions linking the more regular helical and
beta pleated sheet patterns.
•
Beta pleated sheets and alpha helices tend to impart upon a protein the ability to
anchor itself into a lipid bilayer.
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Tertiary Protein Structure
• Once secondary structure has been formed,
associations between additional amino acid
side chains become possible.
• These interactions fold the polypeptide into its
final three-dimensional conformation, making
it a functional protein.
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Factors that Determine Tertiary Structure
• Five major factors determine the tertiary structure of a
polypeptide chain once the amino acid sequence (primary
structure) has been formed:
1. Hydrogen bonds between portions of the chain or with
surrounding water molecules
2. Ionic bonds between polar and ionized regions along the
chain
3. Attraction between nonpolar (hydrophobic) regions
4. Covalent disulfide bonds linking the sulfur-containing side
chains of two cysteine amino acids
5. van der Waals forces
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Quaternary Protein Structure
• Some proteins are composed of more than one polypeptide
chain and are said to have quaternary structure.
• They are known as multimeric (“many parts”) proteins.
• The same factors that influence the conformation of a single
polypeptide also determine the interactions between the
polypeptides in a multimeric protein.
• Therefore, the chains can be held together by interactions
between various ionized, polar, and nonpolar side chains, as
well as by disulfide covalent bonds between the chains.
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Polypeptides: Conformations
Fig. 2-17 63
Amino acid interactions
Fig. 2-18
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Multimeric example
• The polypeptide chains in a multimeric protein may be identical or
different.
• For example, hemoglobin, the protein that transports oxygen in the
blood, is a multimeric protein with four polypeptide chains; two of
one kind and two of another.
• Even a single amino acid change resulting from a mutation may
have devastating consequences.
• An example of this is when a molecule of valine replaces a molecule
of glutamic acid in the b chains of hemoglobin. The result of this
change is a serious disease called sickle cell anemia.
• When red blood cells in a person with this disease are exposed to
low oxygen levels, their hemoglobin precipitates. This contorts the
red blood cells into a crescent shape, which makes the cells fragile
and unable to function normally.
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Hemoglobin
Fig. 2-19 66
Nucleic Acids
• Nucleic acids are extremely important because they
are responsible for the storage, expression, and
transmission of genetic information.
• There are two classes of nucleic acids, deoxyribonucleic acid (DNA) and ribonucleic acid (RNA).
• DNA molecules store genetic information coded in
the sequence of their genes, whereas RNA molecules
are involved in decoding this information into
instructions for linking together a specific sequence
of amino acids to form a specific polypeptide chain.
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Nucleotides
Fig. 2-20 68
Bases
Fig. 2-21 69
DNA
• Four different nucleotides are present in DNA,
corresponding to the four different bases that
can be bound to deoxyribose. These bases are
divided into two classes:
1. The purine bases: adenine (A) and guanine
(G), which have double rings of nitrogen
and carbon atoms
2. The pyrimidine bases, cytosine (C) and
thymine (T), which have only a single ring
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DNA
• A DNA molecule consists of two chains of
nucleotides coiled around each other in the form of a
double helix.
• The two chains are held together by hydrogen bonds
between a purine base on one chain and a pyrimidine
base on the opposite chain.
• Specificity is imposed on the base pairings by the
location of the hydrogen-bonding groups in the four
bases. G is always paired with C, and A with T.
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Deoxyribonucleic acid (DNA)
Fig. 2-22
Fig. 2-23
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RNA
• RNA molecules differ in only a few respects from
DNA:
1. RNA consists of a single chain of nucleotides.
2. In RNA, the sugar in each nucleotide is ribose
rather than deoxyribose.
3. The pyrimidine base thymine in DNA is replaced
in RNA by the pyrimidine base uracil (U).
(A–U pairing)
• The other three bases, adenine, guanine, and cytosine,
are the same in both DNA and RNA.
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ATP
• The purine bases are important not only in
DNA and RNA synthesis, but also in a
molecule that serves as the molecular energy
source for all cells.
• In all cells, from bacterial to human, adenosine
triphosphate (ATP) is the primary molecule
that receives the transfer of energy from the
breakdown of fuel molecules—carbohydrates,
fats, and proteins.
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ATP
Fig. 2-24 76