Download Chapter 2 INTRODUCTION Chapter Overview Basic Principles

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.
Chapter 2
The Chemical Level of Organization
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Basic Principles
Chapter Overview
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• Chemistry is the science of the structure and
interactions of matter
• Matter is anything that occupies space and
has mass
Matter
Chemical bonds
Chemical energy
Chemical reactions
Inorganic compounds
Organic compounds
– 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 is Matter Organized
• Matter is made up of atoms
• An element is a quantity of matter composed
of atoms of the same type
• Atoms join together to form chemicals with
different characteristics
• Chemical characteristics determine
physiology at the molecular and cellular level
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Atomic Particles
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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
• Proton:
– positive, 1 mass unit
• Neutron:
– first shell can hold only 2 electrons
– 2nd shell can hold 8 electrons
– neutral, 1 mass unit
• Electron:
• Number of electrons = number of protons
• Each atom is electrically neutral; charge = 0
– negative, low mass
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The Structure of Atoms
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Particles and Mass
Isotopes
• Atomic number:
• 2 or more elements with equal numbers of
protons but different numbers of neutrons
• 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
– number of protons
• Mass number:
– number of protons plus neutrons
• Atomic weight:
– exact mass of all particles (daltons)
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Chemical Properties of Atoms
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Atomic Number & Mass Number
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Ions
• Atomic number is number of protons in the nucleus.
• Mass number is the sum of its protons and neutrons.
• 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.
– written with its chemical symbol and (+) or (-)
(e.g., Na+, Cl-)
– cations are positive
– anions are negative
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Free Radicals
Free Radicals & Your Health
• A free radical is an electrically charged atom or
group of atoms with an unpaired electron in its
outermost shell
• Unstable and highly reactive; can become stable
• 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’s, atherosclerosis and arthritis
• Damage may be slowed with antioxidants
such as vitamins C and E, selenium & betacarotene (precursor to vitamin A)
– by giving up an electron
– taking an electron from another molecule
– can break apart important body molecules in the process
• Antioxidants are substances that inactivate oxygenderived free radicals
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Chemical Bonds
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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
(valence) shell.
• Atoms with incompletely filled outer shells
tend to combine with each other in chemical
reactions to produce a chemically stable
arrangement of filled outer shells for each
atom.
• Ionic bonds:
– attraction between cations and anions
• Covalent bonds:
– strong electron bonds
• Hydrogen bonds:
– weak polar bonds
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The Ionic Bond in
Sodium Chloride
Covalent Bonds
• Atoms share electrons
to form covalent bonds
• Electrons spend most of
the time between the 2
atomic nuclei
• 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
• Some may dissociate into
positive and negative ions in
solution. Such a compound is
called an electrolyte.
single bond = share 1 pair
double bond = share 2 pair
triple bond = share 3 pair
• Covalent bonds are
common and are the
strongest chemical
bonds in the body.
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Covalent Bonds
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Hydrogen Bonds
Polar Covalent Bonds
• Approximately 5% as
strong as covalent bonds
• Useful in establishing links
between molecules or
between distant parts of a
very large molecule
• Help determine 3-D shape
of large molecules (e.g.,
proteins).
• Give water considerable
cohesion which creates a
very high surface tension
• Unequal sharing of electrons between atoms.
• 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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Chemical Reactions
Forms of Energy
• 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.
• 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 Transfer in Chemical
Reactions
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Energy Transfer in Chemical Reactions
• Reactions in living systems usually involve
both kinds of reactions occurring together.
• 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).
– 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.
• 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.
– 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
Catalysts
• Activation energy is
the energy needed
to initiate a reaction.
• Increases in
concentration or
temperature can help
achieve activation
energy.
• Catalysts are chemical compounds that
speed up chemical reactions by lowering the
activation energy needed for a reaction to
occur.
– 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 reactants
to favor formation of products.
– The catalyst itself is unchanged at the end of the
reaction; it is often re-used many times.
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Catalysts or Enzymes
Effectiveness of Catalysts
• Normal body temperatures and
concentrations are low enough that many
chemical reactions are effectively blocked
by the activation energy barrier.
• Most chemical reactions that sustain life
cannot occur unless the right enzymes
are present
• Enzymes are biological catalysts that
carry out specific chemical reactions in
the body.
– e.g., lactase breaks down lactose into glucose
and galactose.
• Catalysts speed up chemical reactions by
lowering the activation energy.
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Types of Chemical Reactions
Organic and Inorganic Molecules
• Decomposition reaction (catabolism):
• Organic:
AB ! A + B
– molecules based on carbon and hydrogen
• Synthesis reaction (anabolism):
• Inorganic:
A + B ! AB
• Exchange reaction:
– molecules not based on carbon and hydrogen
AB + CD ! AC + BD
• Reversible reaction:
AB " A + B
• Hydrolysis:
A—B—C—D + H2O ! A—B—H + HO—C—D
• Dehydration synthesis (condensation):
A—B—H + HO—C—D ! A—B—C—D + H2O
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Essential Molecules
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Water
• Nutrients:
• Water is the most important and abundant
inorganic compound in all living systems.
• Most of our body weight is water.
• 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 partial positive
charges near the two hydrogen atoms in the
water molecule.
– essential molecules obtained from food
• Metabolites:
– molecules made or broken down in the body
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Properties of Water
Aqueous Solutions
•Polar water molecules form hydration spheres
around ions and small polar molecules to keep
them in solution
• Solubility:
– water’s ability to dissolve a solute in a solvent to
make a solution
• Reactivity:
– most body chemistry uses or occurs in water
• High heat capacity:
– water’s ability to absorb and retain heat
• Lubrication:
– to moisten and reduce friction
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Electrolytes
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Molecules and Water
• Inorganic ions which conduct electricity in
solution
• Electrolyte imbalance seriously disturbs vital
body functions
• Hydrophilic:
– hydro = water, philos = loving
– reacts with water
• Hydrophobic:
– phobos = fear
– does not react with water
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pH
pH Scale
• Has an inverse relationship with H+ concentration:
• pH:
– the concentration of hydrogen ions (H+) in a solution
– more H+ ions mean lower pH
• Neutral pH:
• pH is a logarithmic scale
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– a balance of H and OH
– pure water = 7.0
– pH of 1 contains 100 times more H+ than pH of 3
• Acid (acidic): pH lower than 7.0
– high H+ concentration,
low OH— concentration
• Base (basic): pH higher than 7.0
– low H+ concentration,
high OH— concentration
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Control of pH
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Organic Compounds
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• Acidosis: excess H in body fluid (low pH)
• Usually large and contain carbon,
hydrogen, and oxygen atoms
• 4 major classes
- carbohydrates
- lipids
- proteins
- nucleic acids
– damages cells and tissues
– alters proteins
– interferes with normal physiological functions
• Alkalosis: excess OH— in body fluid (high pH)
• also cause problems, but rarely
• Buffers: control pH
– weak acid/salt compounds (e.g., sodium
bicarbonate)
– neutralizes either strong acid or strong base
– biochemical reactions are very sensitive to even
small changes in acidity or alkalinity
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Carbon
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
• Molecular groups which allow molecules to
interact with other molecules
– 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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Carbohydrates
Monosaccharides
• 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.
• Monosaccharides contain from three to
seven carbon atoms and include glucose, a
hexose that is the main energy-supplying
compound of the body.
• Humans absorb only 3 simple sugars without
further digestion in our small intestine
– About 2-3% of total human body weight
– glucose found in syrup or honey
– fructose found in fruit
– galactose found in dairy products
• Carbohydrates are divided into three major groups
based on their size: monosaccharides,
disaccharides, and polysaccharides
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Disaccharides
Clinical Application:
• Disaccharides are formed from two
monosaccharides by dehydration synthesis; they
can be split back into simple sugars by hydrolysis.
• 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.
– sucrose = glucose & fructose
– maltose = glucose & glucose
– lactose = glucose & galactose (lactose intolerance)
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Polysaccharides
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Lipids = fats
• 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.
• Formed from C, H and O
• Hydrophobic: mostly insoluble in polar solvents
such as water
• Types of lipids
– triglycerides
– phospholipids
– steroids
– eicosanoids
– lipoproteins
– some vitamins
– When blood sugar level
drops, the liver hydrolyzes
glycogen to yield glucose
which is released from the
liver into the blood
• 18-25% of body weight
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Triglycerides
Triglycerides
• Triglycerides are the most plentiful lipids in
the body and provide protection, insulation,
and energy
– 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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Saturation of Triglycerides
Clinical Application
• Determined by the number of single or
double covalent bonds
• Saturated fats contain single covalent
bonds and are saturated with hydrogen
atoms: lard
• Monounsaturated are not completely
saturated with hydrogen: olive and peanut
oil
• Polyunsaturated fats contain even less
hydrogen atoms: safflower oil, corn oil
• 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.
– Omega-3 and omega-6 fatty acids
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Phospholipids
Chemical Nature of Phospholipids
head
• Phospholipids are important membrane
components.
• They are amphipathic, with both polar and
nonpolar regions.
tails
– a polar head
• a phosphate group (PO43-) & glycerol molecule
• forms hydrogen bonds with water
– 2 nonpolar fatty acid tails
• interact only with lipids
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Four Ring Structure of Steroids
Steroids
• Steroids have four rings of carbon atoms
• 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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Other Lipids
Proteins
• Eicosanoids include prostaglandins and
leukotrienes.
• Proteins are the most abundant and
important organic molecules
• Basic elements:
– Lipid type derived from a fatty acid called arachidonic
acid
– prostaglandins = wide variety of functions
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– carbon (C), hydrogen (H), oxygen (O), and
nitrogen (N)
modify responses to hormones
contribute to inflammatory response
prevent stomach ulcers
dilate airways
regulate body temperature
influence formation of blood clots
• Basic building blocks:
– 20 amino acids
– leukotrienes = allergy & inflammatory responses
• Body lipids also include fatty acids; fat-soluble
vitamins such as beta-carotenes, vitamins D, E,
and K; and lipoproteins.
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Proteins
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KEY CONCEPT
Protein Functions
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Support: structural proteins
Movement: contractile proteins
Transport: transport proteins
Buffering: regulation of pH
Metabolic regulation: enzymes
Coordination and control: hormones
Defense: antibodies
• Proteins:
– control anatomical structure and physiological
function
– determine cell shape and tissue properties
– perform almost all cell functions
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Amino Acids
Peptides
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Primary Structure
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Secondary Structure
• Polypeptide:
• Hydrogen bonds form spirals or pleats
– a long chain of amino acids
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Tertiary Structure
Quaternary Structure
• Final protein shape:
• Secondary structure folds into a unique
shape
– several tertiary structures together
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Enzymes
How Enzymes Work
• Enzymes are catalysts:
– proteins that lower the activation energy of a
chemical reaction
– are not changed or used up in the reaction
• Highly specific
• Very efficient
– speed up reaction up to 10 billion times faster
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Enzymes
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Nucleic Acids: DNA and RNA
Protein Denaturation
• Nucleic acids are large organic molecules, which
store and process genetic information at the
molecular level
• 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 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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DNA Structure
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RNA Structure
• 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:
• 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
– nitrogenous base (A-G-TC)
– pentose sugar
– phosphate group
– messenger RNA
– ribosomal RNA
– transfer RNA
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Adenosine Triphosphate (ATP)
Formation & Usage of ATP
• Temporary molecular storage of energy for use in
cellular activities
• Synthesis of ATP
– muscle contraction, transport of substances across cell
membranes, movement of structures within cells and
movement of organelles
– 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
• Consists of 3 phosphate
groups attached to
adenine & 5-carbon
sugar (ribose)
• Hydrolysis of ATP (removal of terminal
phosphate group by enzyme -- ATPase)
– releases energy for cellular processes
– leaves ADP (adenosine diphosphate)
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