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Chapter 2
The Chemical Level of Organization
Lecture Outline
Principles of Human Anatomy and Physiology, 11e
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INTRODUCTION
• Since chemicals compose your body and all body activities
are chemical in nature, it is important to become familiar
with the language and fundamental concepts of chemistry.
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Chapter 2
The Chemical Level of Organization
• Matter
– elements
– atoms and molecules
• Chemical bonds
• Chemical energy
• Chemical reactions
• Inorganic compounds
• Organic compounds
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Basic Principles
• Chemistry is the science of the structure and interactions of
matter.
• Matter is anything that occupies space and has mass.
– Mass is the amount of matter a substance contains
– weight is the force of gravity acting on a mass.
• Describe two ways that you could change your weight.
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HOW MATTER IS ORGANIZED
• Chemical Elements
– All forms of matter are composed of chemical elements
which are substances that cannot be split into simpler
substances by ordinary chemical means.
– Elements are given letter abbreviations called chemical
symbols.
• Oxygen (O), carbon (C), hydrogen (H), and nitrogen
(N) make up 96% of body weight.
• These elements, together with calcium (Ca) and
phosphorus (P) make up 98.5% of total body weight.
– Trace elements are present in tiny amounts
• copper, tin, selenium & zinc
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Review
• Table 2.1 lists the major and trace elements of the human
body.
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Structure of Atoms
• Units of matter of all chemical elements are called atoms.
An element is a quantity of matter composed of atoms of the
same type.
• Atoms consist of a nucleus, which contains positively
charged protons and neutral (uncharged) neutrons, and
negatively charged electrons that move about the nucleus in
energy levels (Figure 2.1).
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Structure of Atoms
• Atoms are the smallest units of matter that
retain the properties of an element
• 3 types of subatomic particles
– protons, neutrons and electrons
• Nucleus: protons (p+) & neutrons (neutral
charge)
• Electrons (e-) surround the nucleus as a
cloud (electron shells are designated
regions of the cloud)
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Electron Shells
• Most likely region of the electron
cloud in which to find electrons
• Each electron shell can hold only
a limited number of electrons
– first shell can hold only 2 electrons
– 2nd shell can hold 8 electrons
– 3rd shell can hold 18 electrons
– higher shells (up to 7) hold many more electrons
• Number of electrons = number of protons
• Each atom is electrically neutral; charge = 0
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Structure of Atoms
• Electrons revolve around the nucleus of an atom tending to
spend most of the time in specific atomic regions, called
shells (Figure 2.1a).
– Each shell can hold a certain maximum number of
electrons.
– The first shell, the one nearest the nucleus, can hold a
maximum of 2 electrons; the second shell, 8; the third
shell;18,the fourth shell, 18; and so on (Figure 2.1b).
– The number of electrons in an atom of a neutral element
always equals the number of protons.
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Atomic Number and Mass Number
• Atomic Number
– The number of protons in the nucleus of an atom
– The number of protons in the nucleus makes the atoms
of one element different from those of another as
illustrated in Figure 2.2.
– Since all atoms are electrically neutral, the atomic
number also equals the number of electrons in each
atom.
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Atomic Number & Mass Number
• Atomic number is number of protons in the nucleus. .
• Mass number is the sum of its protons and neutrons.
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Atomic Number and Mass Number
• The mass number of an atom is the sum of the numbers of
protons and neutrons.
– Different atoms of an element that have the same
number of protons but different numbers of neutrons are
called isotopes.
• Isotopes
– Stable isotopes do not change their nuclear structure
over time.
– Certain isotopes called radioactive isotopes are unstable
because their nuclei decay to form a simpler and thus
more stable configuration.
– Radioactive isotopes can be used to study both the
structure and function of particular tissues.
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Atomic Mass
• Mass is measured as dalton (atomic mass unit)
– neutron has mass of 1.008 daltons
– proton has mass of 1.007 daltons
– electron has mass of 0.0005 dalton
• Atomic mass (atomic weight) is close to the mass number
of its most abundant isotope.
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Atomic Number and Mass Number
• Atomic Mass
– The atomic mass, also called the atomic weight, of an
element is the average mass of all its naturally occurring
isotopes and reflects the relative abundance of isotopes
with different mass numbers.
– The mass of a single atom is slightly less than the sum of
the masses of its neutrons, protons, and electrons
because some mass (less than1%) was lost when the
atom’s components came together to form an atom.
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Ions, Molecules, & Compounds
• Ions are formed by ionization
– an atom that gave up or gained an electron
– written with its chemical symbol and (+) or (-)
• Molecule
– atoms share electrons
– written as molecular formula showing the number
of atoms of each element (H2O)
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Ions, Molecules, Free Radicals, and Compounds
• If an atom either gives up or gains electrons, it becomes
an ion - an atom that has a positive or negative charge due
to having unequal numbers of protons and electrons.
• When two or more atoms share electrons, the resulting
combination is called a molecule (Figure 2.3a).
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Free Radicals
• A free radical is an electrically charged atom or group of atoms with an
unpaired electron in its outermost shell (Figure 2.3b).
• Unstable and highly reactive; can become stable
– by giving up an electron
– taking an electron from another molecule (example: breaking apart
important body molecules)
• Antioxidants are substances that inactivate oxygen-derived free radicals
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Free Radicals & Your Health
• Produced in your body by absorption of energy in
ultraviolet light in sunlight, x-rays, by breakdown of
harmful substances, & during normal metabolic
reactions
• Linked to many diseases -- cancer, diabetes,
Alzheimer, atherosclerosis and arthritis
• Damage may be slowed with antioxidants such as
vitamins C and E, selenium & beta-carotene
(precursor to vitamin A)
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CHEMICAL BONDS
• The atoms of a molecule are held together by forces of
attraction called chemical bonds.
• The likelihood that an atom will form a chemical bond with
another atom depends on the number of electrons in its
outermost shell, also called the valence shell.
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CHEMICAL BONDS
• An atom with a valence shell holding eight electrons (2
electrons for hydrogen and neon) is chemically stable, which
means it is unlikely to form chemical bonds with other
atoms.
• To achieve stability, atoms that do not have eight electrons
in their valence shell (or 2 in the case of H and He) tend to
empty their valence shell or fill it to the maximum extent.
• octet rule.
– Atoms with incompletely filled outer shells tend to
combine with each other in chemical reactions to
produce a chemically stable arrangement of eight
valence electrons for each atom.
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Ionic Bonds
• When an atom loses or gains a valence electron, ions are
formed (Figure 2.4a).
– Positively and negatively charged ions are attracted to
one another.
– Cations are positively charged ions that have given up
one or more electrons (they are electron donors).
– Anions are negatively charged ions that have picked up
one or more electrons that another atom has lost (they
are electron acceptors).
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Ionic Bonds
• When this force of attraction holds ions having opposite
charges together, an ionic bond results.
– Sodium chloride is formed by ionic bonds (Figure 2.4)
• In general, ionic compounds exist as solids but some may
dissociate into positive and negative ions in solution. Such a
compound is called an electrolyte.
• Table 2.2 lists the names and symbols of the most common
ions and ionic compounds in the body.
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The Ionic Bond in Sodium Chloride
• Sodium loses an electron to become
Na+ (cation)
• Chlorine gains an electron to become
Cl- (anion)
• Na+ and Cl- are attracted to each other
to form the compound sodium chloride
(NaCl) -- table salt
• Ionic compounds generally exist as
solids
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Covalent Bonds
• Covalent bonds are formed by the atoms of molecules
sharing one, two, or three pairs of their valence electrons.
– Covalent bonds are common and are the strongest
chemical bonds in the body.
– Single, double, or triple covalent bonds are formed by
sharing one,two, or three pairs of electrons, respectively
(Figures 2.5a – d).
• Covalent bonds may be nonpolar or polar.
– In a nonpolar covalent bond, atoms share the electrons
equally; one atom does not attract the shared electrons
more strongly than the other atom (Figures 2.5a –.d).
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Covalent Bonds
• Atoms share electrons to
form covalent bonds
• Electrons spend most of the
time between the 2 atomic
nuclei
– single bond = share 1 pair
– double bone = share 2 pair
– triple bond = share 3 pair
• Polar covalent bonds share
electrons unequally between
the atoms involved
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Polar Covalent Bonds
• Unequal sharing of electrons between atoms. (Figure 2.6).
• In a water molecule, oxygen attracts the hydrogen electrons
more strongly
– Oxygen has greater electronegativity as indicated by the
negative Greek delta sign.
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Hydrogen Bonds
• Approximately 5% as strong as covalent bonds
• Useful in establishing links between molecules or
between distant parts of a very large molecule
• Large 3-D molecules are
often held together by a
large number of hydrogen
bonds.
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Hydrogen Bonds
Hydrogen bonds are weak intermolecular bonds; they
serve as links between molecules (usually).
• help determine three-dimensional shape (Figure 2.7)
• give water considerable cohesion which creates a
very high surface tension
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Chemical Reactions
• New bonds form and/or old bonds are broken.
• Metabolism is “the sum of all the chemical reactions in the body.”
• Law of conservation of mass
– The total mass of reactants equals the total mass of the
products. (Count the number of atoms of each element below.)
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CHEMICAL REACTIONS
• A chemical reaction occurs when new bonds are formed or
old bonds break between atoms (Figure 2.8).
– The starting substances of a chemical reaction are
known as reactants.
– The ending substances of a chemical reaction are the
products.
• Remember: In a chemical reaction, the total mass of the
reactants equals the total mass of the products (law of
conservation of mass).
• Metabolism refers to all the chemical reactions occurring in
an organism and involves links between chemical reactions
in different parts of the body, or even different parts of a cell.
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Forms of Energy and Chemical Reactions
• Energy is the capacity to do work.
• Kinetic energy is the energy associated with matter in
motion.
– Temperature is an indirect measure of molecular motion.
• Potential energy is energy stored by matter due to its
position.
– Chemical energy is a form of potential energy stored in
the bonds of compounds or molecules.
• The total amount of energy present at the beginning and
end of a chemical reaction is the same; energy can neither
be created nor destroyed although it may be converted
from one form to another (law of conservation of energy).
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Energy and Chemical Reactions
Example: Chemical reactions involve energy changes.
• forms of energy
– eg., bonds within sugar molecules
• potential energy = stored energy
– eg., molecular vibration measured as temperature
• kinetic energy = energy of motion
• If the chemical bonds within sugar are broken, energy of
sugar can be used to heat the body or create movement.
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Energy Transfer in Chemical Reactions
• An exergonic reaction is one in which the bond being broken
has more energy than the one formed so that extra energy
is released, usually as heat (occurs during catabolism of
food molecules).
• An endergonic reaction is just the opposite and thus
requires that energy be added, usually from a molecule
called ATP, to form a bond, as in bonding amino acid
molecules together to form proteins.
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Energy Transfer in Chemical Reactions
• Reactions in living systems usually involve both
kinds of reactions occurring together.
– exergonic reactions release energy
– endergonic reactions absorb energy
• You will learn of many examples in human
metabolism that involve coupled exergonic and
endergonic reactions; the energy released from
one reaction will drive the other.
– Glucose breakdown releases energy, which is
used to build ATP molecules (that store the
energy for later use in other reactions.)
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Activation Energy
• Atoms, ions & molecules
are continuously moving
& colliding.
• Activation energy is the
collision energy needed
to break bonds & begin a
reaction.
• Increases in concentration &
temperature, increase the
probability of collision
– more particles are in a given space when the concentration
is higher
– particles move more rapidly when temperature is raised
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Activation Energy
• Example:
Firewood does not spontaneously combust. Why?
Give an example of an oxidation reaction that has a
relatively low activation energy.
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Factors that influence the chance that a collision will
occur and cause a chemical reaction include:
• Concentration
• Temperature
• Catalysts are chemical compounds that speed up chemical
reactions by lowering the activation energy needed for a
reaction to occur (Figure 2.10).
– A catalyst does not alter the difference in potential energy
between the reactants and products. It only lowers the
amount of energy needed to get the reaction started.
– A catalyst helps to properly orient the colliding particles of
matter so that a reaction can occur at a lower collision
speed.
– The catalyst itself is unchanged at the end of the
reaction; it is often re-used many times.
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Effectiveness of Catalysts
• Catalysts speed up chemical reactions by lowering the
activation energy.
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Catalysts or Enzymes
Example:
• Normal body temperatures and concentrations are low
enough that many chemical reactions are effectively
blocked by the activation energy barrier.
– Lactose typically reacts very slowly with water to break
down into two simple sugars called glucose and
galactose.
– Lactase, an enzyme (catalyst) orients the colliding
particles (lactose and water) properly so that they touch
at the spots that make the reaction happen.
– Thousands of lactose/water reactions may be
catalyzed by one lactase enzyme.
– Without lactase, the lactose will remain undigested in
the intestines and often causes diarrhea and cramping.
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Types of Chemical Reactions
• Synthesis reactions occur when two or more atoms, ions, or
molecules combine to form new and larger molecules.
These are anabolic reactions, meaning that bonds are
formed. (Figure 2.8)
• In a decomposition reaction, a molecule is broken down
into smaller parts. These are catabolic reactions, meaning
that chemical bonds are broken in the process.
• Exchange reactions involve the replacement of one atom or
atoms by another atom or atoms.
• In reversible reactions, end products can revert to the
original combining molecules.
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Synthesis Reactions--Anabolism
• Two or more atoms, ions or molecules combine to
form new & larger molecules
• All the synthesis reactions in the body together are
called anabolism
• Usually are endergonic because they absorb more
energy than they release
• Example
– combining amino acids to form a protein molecule
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Decomposition Reactions--Catabolism
• Large molecules are split into smaller atoms, ions or
molecules
• All decomposition reactions occurring together in the body
are known as catabolism
• Usually are exergonic since they release more energy
than they absorb
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Exchange Reactions
• Substances exchange atoms
– consist of both synthesis and decomposition
reactions
• Example
– HCl + NaHCO3 gives rise to H2CO3 + NaCl
– ions have been exchanged between substances
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Reversible Reactions
• Chemical reactions can be reversible.
– Reactants can become products or products can revert
to the original reactants
• Indicated by the 2 arrows pointing in opposite directions
between the reactants and the products
AB
A + B
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Oxidation-Reduction Reactions
• Oxidation is the loss of electrons from a molecule
(decreases its potential energy)
– acceptor of the electron is often oxygen
– commonly oxidation reactions involve removing a
hydrogen ion (H+) and a hydride ion (H-) from a
molecule
– equivalent to removing 2 hydrogen atoms = 2H
• Reduction is the gain of electrons by a molecule
– increases its potential energy
• In the body, oxidation-reduction reactions are coupled &
occur simultaneously
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INORGANIC COMPOUNDS AND SOLUTIONS
• Inorganic compounds usually lack carbon and are simple
molecules; whereas organic compounds always contain
carbon and hydrogen, usually contain oxygen, and always
have covalent bonds.
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Water
– Water is the most important and abundant inorganic
compound in all living systems.
– An important property of water is its polarity, the uneven
sharing of valence electrons that confers a partial
negative charge near the one oxygen atom and two
partial positive charges near the two hydrogen atoms in
the water molecule (Figure 2.6).
– Water enables reactants to collide to form products
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Water in Chemical Reactions
• Water is the ideal medium for most chemical reactions in
the body and participates as a reactant or product in certain
reactions.
• Hydrolysis breaks large molecules down into simpler ones
by adding a molecule of water.
• Dehydration synthesis occurs when two simple molecules
join together, eliminating a molecule of water in the process.
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Water as a solvent
• In a solution the solvent dissolves the solute.
• Substances which contain polar covalent bonds and
dissolve in water are hydrophilic, while substances which
contain non polar covalent bonds are hydrophobic.
• The polarity of water and its bent shape allow it to interact
with several neighboring ions or molecules. (Figure 2.11)
• Water’s role as a solvent makes it essential for health and
survival.
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Water as a Solvent
• Most versatile solvent known
– polar covalent bonds (hydrophilic versus
hydrophobic)
– its shape allows each water
molecule to interact with 4 or
more neighboring ions/molecules
• oxygen attracts sodium
• hydrogen attracts chloride
• sodium & chloride separate as ionic
bonds are broken
• hydration spheres surround each ion and
decrease possibility of bonds being
reformed
• Water dissolves or suspends many substances
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High Heat Capacity of Water
• Water has a high heat capacity.
• It can absorb or release a relatively large amount of heat
with only a modest change in its own temperature.
• This property is due to the large number of hydrogen ions in
water.
• Heat of vaporization is also high
– amount of heat needed to change from liquid to gas
– evaporation of water from the skin removes large amount
of heat
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Cohesion of Water Molecules
• Hydrogen bonds link neighboring water molecules
giving water cohesion
• Creates high surface tension
– difficult to break the surface of liquid if molecules
are more attracted to each other than to
surrounding air molecules
– respiratory problem causes by water’s cohesive
property
• air sacs of lungs are more difficult to inflate
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Water as a Lubricant
• Water is a major part of mucus and other lubricating fluids.
– mucus in respiratory and digestive systems
– synovial fluid in joints
– serous fluids in chest and abdominal cavities
• organs slide past one another
• It is found wherever friction needs to be reduced or
eliminated
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Solutions, Colloids, and Suspensions
• A mixture is a combination of elements or compounds that
are physically blended together but are not bound by
chemical bonds. Three common liquid mixtures are
solutions, colloids, and suspensions.
• Solution: a substance called the solvent dissolves another
substance called the solute. Usually there is more solvent
than solute in a solution.
• A colloid differs from a solution mainly on the basis of the
size of its particles with the particles in the colloid being
large enough to scatter light.
• Suspension: the suspended material may mix with the
liquid or suspending medium for some time, but it will
eventually settle out.
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Concentration
• The concentration of a molecule is a way of stating the
amount of that molecule dissolved in solution (Table 2.3).
• Percent gives the relative mass of a solute found in a given
volume of solution. (example: 100 ml of a 5% glucose
solution contains 5 grams of glucose. A liter of the same
solution would contain 50 grams of glucose.)
• A mole is the name for the number of atoms in an atomic
weight of that element, or the number of molecules in a
molecular weight of that type of molecule, with the molecular
weight being the sum of all the atomic weights of the atoms
that make up the molecule. (example: a mole of glucose
has a mass of ___g, so a liter of 2 molar glucose would
contain ____x2 grams of glucose.)
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Inorganic Acids, Bases & Salts. Acids, bases
and salts always dissociate into ions if they are
dissolved
– acids dissociate
into in
H+water (Fig 2.12)
and one or more anions
– bases dissociate into OHand one or more cations
– salts dissociate into anions
and cations, none of which
are either H+ or OH• Acid & bases react in the body to form salts
• A salt, when dissolved in water, dissociates into cations and
anions, neither of which is H+ or OH- (Figure 2.12c). Many
salts are present in the body and are formed when acids and
bases react with each other.
– Electrolytes are important salts in the body that carry
electric current (in nerve or muscle)
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Concept of
pH
• pH scale runs from 0 to 14 (concentration of H+ in moles/liter)
• pH of 7 is neutral
(distilled water -- concentration of OH- and H+ are equal)
• pH below 7 is acidic ([H+] > [OH-]).
• pH above 7 is alkaline ([H+] < [OH-]).
• pH is a logarithmic scale
Example: a change of two or three pH units
pH of 1 contains 10x10=100 more H+ than pH of 3
pH of 8 contains 10x10x10=1000 more H+ than pH of 11
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Acid-Base Balance: The Concept of pH
• Biochemical reactions are very sensitive to even small
changes in acidity or alkalinity.
– pH of blood is 7.35 to 7.45
• A solution’s acidity or alkalinity is based on the pH scale
pH 0 (=100 = 1.0 moles H+/L)
pH 7.0 = 10-7 = 0.0000001 moles H+/L = neutrality or equal
numbers of [H+] and [OH-].
pH 14 (= 10-14 = 0.00000000000001 moles H+/L)
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Maintaining pH: Buffer Systems
• The pH values of different parts of the body are maintained
fairly constant by buffer systems, which usually consist of a
weak acid and a weak base.
– convert strong acids or bases into weak acids or bases.
– Example: carbonic acid-bicarbonate buffer system.
• Bicarbonate ions (HCO3-) act as weak bases and
carbonic acid (H2CO3) acts as a weak acid.
• CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3• Table 2.4 shows pH values for certain body fluids compared
to common substances.
– gastric juice 1.2 to 3.0; saliva 6.35 to 6.85; bile 7.6 to 8.6
and blood 7.35 to 7.45
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ORGANIC COMPOUNDS
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Carbon and Its Functional Groups
• The carbon that organic compounds always contain has
several properties that make it particularly useful to living
organisms.
• It can react with one to several hundred other carbon atoms
– forms large molecules of many different shapes.
• Many carbon compounds do not dissolve easily in water
– useful materials for building body structures.
• Carbon compounds are mostly or entirely held together by
covalent bonds and tend to decompose easily
– organic compounds are a good source of energy.
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Carbon & Its Functional Groups
• Many functional groups can attach to carbon skeleton
– esters, amino, carboxyl, phosphate groups (Table 2.5)
• Very large molecules are called macromolecules (or “polymers” if all the
monomer subunits are similar)
• Isomers have the same molecular formulas but different structures
(glucose & fructose are both C6H12O6)
• STRUCTURAL
FORMULA OF
GLUCOSE (Fig 2.14)
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Carbohydrates
• Carbohydrates provide most of the energy needed for life
and include sugars, starches, glycogen, and cellulose.
• Some carbohydrates are converted to other substances
which are used to build structures and to generate ATP.
• Other carbohydrates function as food reserves.
– About 2-3% of total human body weight
• Carbohydrates are divided into three major groups based on
their size: monosaccharides, disaccharides, and
polysaccharides (Table 2.6).
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Carbohydrates
• formed from C, H, and O
– ratio of one carbon atom for each water molecule
• glucose is 6 carbon atoms and 6 water molecules (H20)
• source of energy for ATP formation
• 2-3 % of total body weight
– glycogen is stored in liver and muscle tissue
– sugar building blocks of DNA & RNA
(deoxyribose & ribose sugars)
• Only plants produce starches or
cellulose for energy storage
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Monosaccharides and Disaccharides: Sugars
• Monosaccharides contain from three to seven carbon atoms
and include glucose, a hexose that is the main energysupplying compound of the body.
• Humans absorb only 3 simple sugars without further
digestion in our small intestine
– glucose found syrup or honey
– fructose found in fruit
– galactose found in dairy products
• Disaccharides are formed from two monosaccharides by
dehydration synthesis; they can be split back into simple
sugars by hydrolysis (Figure 2.15). Glucose and fructose
combine, for example, to produce sucrose.
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Disaccharides
• Combining 2 monosaccharides by dehydration synthesis releases a
water molecule.
– sucrose = glucose & fructose
– maltose = glucose & glucose
– lactose = glucose & galactose (lactose intolerance)
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Clinical Application:
• Lactose intolerance is a deficiency of the enzyme lactase.
As a result undigested lactose remains in the feces and
bacterial fermentation of lactose produces gas.
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Polysaccharides
• Polysaccharides are the largest carbohydrates and may
contain hundreds of monosaccharides.
• The principal polysaccharide in the human body is glycogen,
which is stored in the liver or skeletal muscles. (Figure 2.16)
– When blood sugar level drops, the liver hydrolyzes
glycogen to yield glucose which is released from the liver
into the blood
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Lipids
• Lipids, like carbohydrates, contain carbon, hydrogen, and
oxygen; but unlike carbohydrates, they do not have a 2:1
ratio of hydrogen to oxygen.
• They have few polar covalent bonds
– hydrophobic
– mostly insoluble in polar solvents such as water
– combines with proteins (lipoproteins) for transport in
blood
• Table 2.7 summarizes the various types of lipids and
highlights their roles in the human body.
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Lipids = fats
• Formed from C, H and O
– fats
– phospholipids
– steroids
– eicosanoids
– lipoproteins
– some vitamins
• 18-25% of body weight
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Triglycerides
• Triglycerides are the most plentiful lipids in the body and
provide protection, insulation, and energy (both immediate
and stored).
– At room temperature, triglycerides may be either solid
(fats) or liquid (oils).
– Triglycerides provide more than twice as much energy
per gram as either carbohydrates or proteins.
– Triglyceride storage is virtually unlimited.
– Excess dietary carbohydrates, proteins, fats, and oils will
be deposited in adipose tissue as triglycerides.
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Triglycerides
• Neutral fats composed of a single 3-carbon glycerol
molecule and 3 fatty acid molecules (Figure 2.17).
• Very concentrated form of energy
– 9 calories/gram compared to 4 for proteins &
carbohydrates
– our bodies store triglycerides in fat cells if we eat extra
food
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Triglycerides
• 3 fatty acids & one glycerol molecule
• Fatty acids attached by dehydration systhesis
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Saturation of Triglycerides
• Determined by the number of single or double covalent
bonds
• Saturated fats contain single covalent bonds and are
covered with hydrogen atoms----lard
• Monounsaturated are not completely covered with
hydrogen----safflower oil, corn oil
• Polyunsaturated fats contain even less hydrogen atoms---olive and peanut oil
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Clinical Application
• Essential fatty acids (EFA’s) are essential to human health
and cannot be made by the human body. They must be
obtained from foods or supplements.
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Chemical Nature of Phospholipids
head
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Phospholipids
• Phospholipids are important membrane components.
• They are amphipathic, with both polar and nonpolar regions
(Figure 2.18).
– a polar head
• a phosphate group (PO4-3) & glycerol molecule
• forms hydrogen bonds with water
– 2 nonpolar fatty acid tails
• interact only with lipids
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Steroids
• Steroids have four rings of carbon atoms (Figure 2.19a).
• Steroids include
– sex hormone
– bile salts
– some vitamins
– cholesterol, with cholesterol serving as an important
component of cell membranes and as starting material
for synthesizing other steroids.
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Four Ring Structure of Steroids
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Other Lipids
• Eicosanoids include prostaglandins and leukotrienes.
– Lipid type derived from a fatty acid called arachidonic acid
– prostaglandins = wide variety of functions
• modify responses to hormones
• contribute to inflammatory response
• prevent stomach ulcers
• dilate airways
• regulate body temperature
• influence formation of blood clots
– leukotrienes = allergy & inflammatory responses
• Body lipids also include fatty acids; fat-soluble vitamins such as betacarotenes, vitamins D, E, and K; and lipoproteins.
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Proteins
• Proteins give structure to the body, regulate processes,
provide protection, help muscles to contract, transport
substances, and serve as enzymes (Table 2.8).
• Contain carbon, hydrogen, oxygen, and nitrogen
• 12-18% of body weight
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Proteins
• Constructed from combinations of 20
amino acids.
– dipeptides formed from 2 amino acids
joined by a covalent bond called a
peptide bond
– polypeptides chains formed from 10 to
2000 amino acids.
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Amino Acid Structure
•
•
•
•
Central carbon atom
Amino group (NH2)
Carboxyl group (COOH)
Side chains (R groups) vary
between amino acids
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Levels of Structural Organization
• Levels of structural organization include
– primary
– secondary
– tertiary
– quaternary (Figure 2.22)
• The resulting shape of the protein greatly influences its
ability to recognize and bind to other molecules.
• Denaturation of a protein by a hostile environment causes
loss of its characteristic shape and function.
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Formation of a Dipeptide Bond
• Dipeptides formed from 2 amino acids joined by a
covalent bond called a peptide bond
– dehydration synthesis
• Polypeptides chains contain 10 to 2000 amino acids.
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Levels of
Structural
Organization
• Primary is unique sequence of amino acids
• Secondary is alpha helix or pleated sheet folding
• Tertiary is 3-dimensional shape of polypeptide chain
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Levels of Structural
Organization
•
•
•
•
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Primary…
Secondary …
Tertiary…
Quaternary is relationship of
multiple polypeptide chains
88
Bonds of Tertiary & Quaternary Structure
• Disulfide bridges stabilize the
tertiary structure of protein
molecules
• Covalent bonds between sulfhydryl
groups of 2 cysteine amino acids
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Protein Denaturation
• The function of a protein depends on its ability to bind to
another molecule
• Hostile environments such as heat, acid or salts will
change a protein’s 3-D shape and destroy its ability to
function
– raw egg white when cooked is vastly different
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Enzymes
• Catalysts in living cells are called enzymes.
• Enzymes are highly specific in terms of the “substrate” with
which they react.
• Enzymes are subject to variety of cellular controls.
• Enzymes speed up chemical reactions by increasing
frequency of collisions, lowering the activation energy and
properly orienting the colliding molecules (Figure 2.23).
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Enzymes
• Enzymes are protein molecules that act as catalysts
• Enzyme = apoenzyme + cofactor
– Apoenzymes are the protein portion
– Cofactors are nonprotein portion
• may be metal ion (iron, zinc, magnesium or
calcium)
• may be organic molecule derived from a vitamin
• Enzymes usually end in suffix -ase and are named for
the types of chemical reactions they catalyze
– oxidase, kinase, and lipase are examples
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Enzyme Functionality
• Highly specific
• Very efficient
– speed up reaction up to 10
billion times faster
• Under nuclear control
– rate of synthesis of enzyme
– inhibitory substances
– inactive forms of enzyme
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Clinical Application
• inherited disorder: galactosemia
– Infant lacks enzyme.
– Galactose accumulates in the blood causing anorexia.
– Treatment is elimination of milk from the diet.
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Nucleic Acids: Deoxyribonucleic Acid (DNA) and
Ribonucleic Acid (RNA)
• Nucleic acids are huge organic molecules that contain
carbon, hydrogen, oxygen, nitrogen, and phosphorus.
• Deoxyribonucleic acid (DNA) forms the genetic code inside
each cell and thereby regulates most of the activities that
take place in our cells throughout a lifetime.
• Ribonucleic acid (RNA) relays instructions from the genes in
the cell’s nucleus to guide each cell’s assembly of amino
acids into proteins by the ribosomes.
• The basic units of nucleic acids are nucleotides, composed
of a nitrogenous base, a pentose, sugar, and a phosphate
group (Figures 2.24a, b).
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DNA Structure
• Contains C, H, O, N and
phosphorus.
• Each gene of our genetic material
is a piece of DNA that controls the
synthesis of a specific protein.
• A molecule of DNA is a chain of
nucleotides.
• A nucleotide includes:
– nitrogenous base (A-G-T-C)
– pentose sugar
– phosphate group
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DNA Fingerprinting
• Used to identify criminal, victim or a child’s parents
– need only strand of hair, drop of semen or spot of blood
• Certain DNA segments are repeated several times
– unique from person to person
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RNA Structure
• Differs from DNA
– single stranded
– ribose sugar not deoxyribose sugar
– uracil nitrogenous base replaces thymine
• Types of RNA within the cell, each with a specific
function
– messenger RNA
– ribosomal RNA
– transfer RNA
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Adenosine Triphosphate (ATP)
• Temporary molecular storage of energy as it is being
transferred from exergonic catabolic reactions to cellular
activities
– muscle contraction, transport of substances across cell
membranes, movement of structures within cells and
movement of organelles
• Consists of 3 phosphate
groups attached to
adenine & 5-carbon
sugar (ribose)
• (Figure 2.25).
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Formation & Usage of ATP
• Hydrolysis of ATP (removal of terminal phosphate
group by enzyme -- ATPase)
– releases energy
– leaves ADP (adenosine diphosphate)
• Synthesis of ATP
– enzyme ATP synthase catalyzes the addition of
the terminal phosphate group to ADP
– energy from 1 glucose molecule is used during
both anaerobic and aerobic respiration to create
36 to 38 molecules of ATP
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