Transcript Chapter 3

Chapter 2
The Chemical Level of Organization
Lecture Outline
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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.
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HOW MATTER IS ORGANIZED
• Chemical Elements
– Elements are given letter abbreviations called chemical
symbols.
• Oxygen (O), carbon (C), hydrogen (H), and nitrogen
(N) make up 96% of body weight.
– Trace elements are present in tiny amounts
• copper, tin, selenium & zinc
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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
• Nucleus: protons (p+) & neutrons (neutral
charge)
• Electrons (e-) surround the nucleus as a
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
• Number of electrons = number of protons
• Each atom is electrically neutral; charge = 0
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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
– Isotopes
• Isotopes
– Stable isotopes do not change their nuclear structure
over time.
– radioactive isotopes
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Ions, Molecules, Free Radicals, and Compounds
• Ions
• Molecules (Fig 2.3a)
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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
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Ionic Bonds
• When an atom loses or gains a valence electron, ions are
formed (Figure 2.4a).
– Opposites attract
– Cations
– Anions
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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
• Atoms share electrons to
form covalent bonds
• Electrons spend most of the
time between the 2 atomic
nuclei
• Polar covalent bonds
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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
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Hydrogen Bonds
• Weak
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CHEMICAL REACTIONS
• A chemical reaction occurs when new bonds are formed or
old bonds break between atoms (Figure 2.8).
• Law of conservation of mass.
• Metabolism
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Forms of Energy and Chemical Reactions
• Energy
• Kinetic energy
– Temperature is an indirect measure of molecular motion.
• Potential energy
– Chemical energy is a form of potential energy stored in
the bonds of compounds or molecules.
• Law of conservation of energy
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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
• 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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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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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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Reversible Reactions
• Chemical reactions can be reversible.
– Reactants can become products or products can revert
to the original reactants
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Oxidation-Reduction Reactions
• Oxidation is the loss of electrons from a molecule
• Reduction is the gain of electrons by a molecule
• In the body, oxidation-reduction reactions are coupled &
occur simultaneously
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INORGANIC COMPOUNDS AND SOLUTIONS
• Inorganic compounds
• Organic compounds
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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
• Dehydration synthesis
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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.
• Heat of vaporization is also high
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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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Inorganic Acids, Bases & Salts. Acids, bases
and salts always dissociate into ions if they are
dissolved in water (Fig 2.12)
– acids dissociate into H+
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
– 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-]).
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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
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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.
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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
• 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)
• 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
– 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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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
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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
– At room temperature, triglycerides may be either solid
(fats) or liquid (oils).
– Excess dietary carbohydrates, proteins, fats, and oils will
be deposited in adipose tissue as triglycerides.
– Neutral fats composed of a single 3-carbon glycerol
molecule and 3 fatty acid molecules (Figure 2.17).
– 9 calories/gram compared to 4 for proteins &
carbohydrates
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Triglycerides
• 3 fatty acids & one glycerol molecule
• Fatty acids attached by dehydration systhesis
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Chemical Nature of Phospholipids
head
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Phospholipids
• Phospholipids are important membrane components.
• They are amphipathic (Figure 2.18).
• a phosphate group (PO4-3) & glycerol molecule
• forms hydrogen bonds with water
• 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
– 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.
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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)
• Importance of shape
• Denaturation
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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
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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 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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Enzyme Functionality
• Highly specific
• Very efficient
• Under nuclear control
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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)
– Ribonucleic acid (RNA)
– Nucleotides (Figures 2.24a, b).
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DNA Structure
• Contains C, H, O, N and
phosphorus.
• 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
• 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
• Synthesis of ATP
– energy from 1 glucose molecule creates 36
molecules of ATP
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