Transcript Document

PowerPoint® Lecture Slides
prepared by
Barbara Heard,
Atlantic Cape Community
College
CHAPTER
2
Chemistry Comes
Alive: Modified by Dr. Par
Mohammadian
© Annie Leibovitz/Contact Press Images
© 2013 Pearson Education, Inc.
Overview
• Part 1: Basic Chemistry
– Matter
– Energy
– Atoms & Elements
– Chemical Bonds
– Functional Groups
• Part 2: Biochemistry
– Bio-macromolecules:
• Carbohydrates
• Lipids
• Proteins
– Enzymes
• Nucleic acids
Body Chemistry
A general understanding of chemistry is
necessary for understanding human
physiology.
Why? Name specific examples!
– Physiological processes are based on
chemical interactions.
Matter
• Matter—anything that has mass & occupies
space
• 3 states of matter
– Solid—definite shape and volume
– Liquid—changeable shape; definite volume
– Gas—changeable shape and volume
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Energy
• Definition:
– Capacity to do work or put matter into motion
• Types of energy
– Kinetic—energy in action
Name an example!
– Potential—stored (inactive) energy
• Energy can be transferred from potential
to kinetic energy
Forms of Energy
Name an example for each!
• Chemical energy
– Stored in bonds of chemical substances
• Electrical energy
– Results from movement of charged particles
• Mechanical energy
– Directly involved in moving matter
• Radiant or electromagnetic energy
– Travels in waves (e.g., visible light, UV light,
and x-rays)
Composition of Matter: Elements
• Elements
Examples?
– Matter is composed of elements
– Elements cannot be broken into simpler
substances by ordinary chemical methods
• Atoms
– Building blocks for each element
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Major Elements of the Human Body
• Four elements make up 96.1% of body mass
Element
Carbon
Hydrogen
Oxygen
Nitrogen
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Atomic symbol
C
H
O
N
Lesser Elements of the Human Body
9 elements make up 3.9% of body mass
Element
Calcium
Phosphorus
Potassium
Sulfur
Sodium
Chlorine
Magnesium
Iodine
Iron
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Atomic symbol
Ca
P
K
S
Na
Cl
Mg
I
Fe
Atoms – Subatomic particles
• An atom is the smallest unit
of an element. It has:
– A nucleus with positively charged
protons and uncharged neutrons
– Orbiting electrons with negative
charges
– An atomic mass equal to the
number of protons plus the number
of neutrons
– An atomic number equal to the
number of protons
Electron Orbitals
• Orbitals (or shells) are energy levels that surround
the nucleus of an atom.
• Electrons fill the shells, starting with the one
closest to the nucleus.
– The first shell holds 2 electrons.
– Each shell thereafter holds 8 electrons. (Non-biological
elements fill distant shells that hold more than 8.)
– Atoms are most stable when the outer shell is filled.
Electrons in unfilled outer shells participate in bonding;
they are called valence electrons.
Figure 2.1 Two models of the structure of an atom.
Nucleus
Nucleus
Helium atom
Helium atom
2 protons (p+)
2 neutrons (n0)
2 electrons (e–)
2 protons (p+)
2 neutrons (n0)
2 electrons (e–)
Planetary model
Proton
Neutron
Electron
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Electron
cloud
Orbital model
Identifying Elements
• Different elements contain different
numbers of subatomic particles
– Hydrogen has 1 proton, 0 neutrons, and 1
electron
– Lithium has 3 protons, 4 neutrons, and 3
electrons
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Identifying Elements:
• Atomic number = # of protons in nucleus
– Written as subscript to left of atomic symbol
• Ex. 3Li
• Mass number
– Total # of protons and neutrons in nucleus
• Total mass of atom
– Written as superscript to left of atomic symbol
• Ex. 7Li
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Ions and Isotopes
• An atom that loses or
gains electrons
becomes an ion:
– Cation; charge: +
– Anion; charge: • An atom that loses or
gains protons
becomes a different
element.
• An atom that gains or
loses neutrons
becomes an isotope
of the same element.
Radioisotopes
• Valuable tools for biological research and
medicine
– Share same chemistry as their stable isotopes
– Most used for diagnosis
• All damage living tissue
– Some used to destroy localized cancers
– Radon from uranium decay causes lung
cancer
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Free Radicals
•Free radicals: Unstable; have at least one
unpaired electron; caused in our body by natural
(sun) and/or manufactured
•Sources (microwave, TV); e.g peroxide ion .O2•Can contribute to diseases (cancer) and aging!
•Antioxidants (e.g. Vit C and E) prevent damage
by releasing electrons without becoming free
radicals!
Combining Matter:
Molecules and Compounds
• Most atoms chemically combined with
other atoms to form molecules and
compounds
– Molecule
• Two or more atoms bonded together (e.g., H2 or
C6H12O6)
• Smallest particle of a compound with specific
characteristics of the compound
– Compound
• Two or more different kinds of atoms bonded
together (e.g., C6H12O6 , but not H2)
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Mixtures
• Most matter exists as mixtures
– Two or more components physically
intermixed
• Three types of mixtures
– Solutions
– Colloids
– Suspensions
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Types of Mixtures: Solutions
• Homogeneous mixtures
• Most are true solutions in body
– Gases, liquids, or solids dissolved in water
– Usually transparent, e.g., atmospheric air or saline
solution
• Solvent
– Substance present in greatest amount
– Usually a liquid; usually water
• Solute(s)
– Present in smaller amounts
• Ex. If glucose is dissolved in blood, glucose is
solute; blood is solvent
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Concentration of True Solutions
• Can be expressed as
– Percent (%) of solute in total solution
(assumed to be water) Prepare 0.9% NaCl solution!
• Parts solute per 100 parts solvent
– Milligrams per deciliter (mg/dl)
– Molarity, or moles per liter (M)
• 1 mole of an element or compound = Its atomic or
molecular weight (sum of atomic weights) in grams
• 1 mole of any substance contains 6.02  1023
molecules of that substance (Avogadro’s number)
Use the example in the book (pg 30) to explain!
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Colloids and Suspensions
• Colloids (AKA emulsions)
– Heterogeneous mixtures, e.g., cytosol
– Large solute particles do not settle out
• Suspensions
– Heterogeneous mixtures, e.g., blood
– Large, visible solutes settle out
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Figure 2.4 The three basic types of mixtures.
Solution
Colloid
Suspension
Solute particles are very tiny,
do not settle out or scatter light.
Solute particles are larger than
in a solution and scatter light;
do not settle out.
Solute particles are very large,
settle out, and may scatter light.
Solute
particles
Solute
particles
Solute
particles
Example
Example
Example
Mineral water
Jello
Blood
Plasma
Settled red
blood cells
Unsettled Settled
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Mixtures versus Compounds
• Mixtures
– No chemical bonding between components
– Can be separated by physical means, such as
straining or filtering
– Heterogeneous or homogeneous
• Compounds
– Chemical bonding between components
– Can be separated only by breaking bonds
– All are homogeneous
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Chemical Bonds
• A molecule forms when electrons of several
atoms interact to form chemical bonds.
– The number of bonds an atom can form is
determined by the number of valence electrons.
• Hydrogen has one electron; it needs one more to fill
the inner shell so that it can form one bond.
• Carbon has 6 electrons; 2 fill the inner shell and 4 are
in the next shell. It needs 4 more electrons so that it
can form 4 bonds.
Figure 2.5a Chemically inert and reactive elements.
Chemically inert elements
Outermost energy level (valence shell) complete
2e
Helium (He)
(2p+; 2n0; 2e–)
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8e
2e
Neon (Ne)
(10p+; 10n0; 10e–)
Figure 2.5b Chemically inert and reactive elements.
Chemically reactive elements
• Valence shell
not full
• Tend to gain,
lose, or share
electrons (form
bonds) with
other atoms to
achieve stability
Outermost energy level (valence shell) incomplete
1e
Hydrogen (H)
(1p+; 0n0; 1e–)
6e
2e
Oxygen (O)
(8p+; 8n0; 8e–)
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4e
2e
Carbon (C)
(6p+; 6n0; 6e–)
1e
8e
2e
Sodium (Na)
(11p+; 12n0; 11e–)
Types of Chemical Bonds
• Three major types
– Ionic bonds
– Covalent bonds
– Hydrogen bonds
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Ionic Bonds
• Ions
– Atom gains or loses electrons and becomes charged
• # Protons ≠ # Electrons
• Transfer of valence shell electrons from one atom
to another forms ions
– One becomes an anion (negative charge)
• Atom that gained one or more electrons
– One becomes a cation (positive charge)
• Atom that lost one or more electrons
• Attraction of opposite charges results in an ionic
bond
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Figure 2.6a–b Formation of an ionic bond.
+
Sodium atom (Na)
(11p+; 12n0; 11e–)
Chlorine atom (Cl)
(17p+; 18n0; 17e–)
Sodium gains stability by losing
one electron, and chlorine becomes
stable by gaining one electron.
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Sodium ion (Na+)
—
Chloride ion (Cl–)
Sodium chloride (NaCl)
After electron transfer,
the oppositely charged ions
formed attract each other.
Figure 2.6c Formation of an ionic bond.
•
Most ionic compounds
are salts
– When dry salts
form crystals
instead of
individual
molecules
– Example is NaCl
(sodium chloride)
Cl–
Na+
Large numbers of Na+ and Cl– ions
associate to form salt (NaCl) crystals.
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Covalent Bonds
• Formed by sharing of two or more valence
shell electrons
• Allows each atom to fill its valence shell at
least part of the time
Reacting atoms
Resulting molecules
or
+
Nitrogen atom
Nitrogen atom
Formation of a triple covalent bond: Two
nitrogen atoms share three electron pairs.
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Structural formula
shows triple bond.
Molecule of nitrogen gas (N2)
Figure 2.7a Formation of covalent bonds.
Reacting atoms
Resulting molecules
+
or
Structural formula
shows single bonds.
Carbon atom
Hydrogen atoms
Formation of four single covalent bonds:
Carbon shares four electron pairs with
four hydrogen atoms.
Reacting atoms
Molecule of methane gas (CH4)
Resulting molecules
+
Oxygen atom
Oxygen atom
Formation of a double covalent bond: Two
oxygen atoms share two electron pairs.
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or
Structural formula
shows double bond.
Molecule of oxygen gas (O2)
Nonpolar Covalent Bonds
• Electrons shared equally
• Produces electrically balanced, nonpolar
molecules such as CO2
Carbon dioxide (CO2) molecules are linear and symmetrical. They are nonpolar.
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Polar Covalent Bonds
• Unequal sharing of electrons produces polar (AKA dipole) molecules such
as H2O
• Small atoms with six or seven valence shell electrons are electronegative,
e.g., oxygen
• Most atoms with one or two valence shell electrons are electropositive,
e.g., Na
–
+
+
V-shaped water (H2O) molecules have two poles of charge—a slightly
more negative oxygen end (–) and a slightly more positive
hydrogen end (+).
Hydrogen Bonds
•
+
–
– Not true bond
Hydrogen bond
– Common between dipoles
(indicated by
such as water
dotted line)
+
–
–
+
–
+
+
Attractive force between
electropositive hydrogen of one
molecule and an electronegative
atom of another molecule
– Also act as intramolecular
bonds, holding a large
molecule in a threedimensional shape
+
–
The slightly positive ends (+) of the water molecules become aligned with
the slightly negative ends (–) of other water molecules.
Functional Groups
Classes of molecules are named after their functional group.
Chemical Reactions
(reactants) A+B -> AB (product)
Synthesis reactions
Smaller particles are bonded
together to form larger,
more complex molecules.
Example
Amino acids are joined together to
form a protein molecule.
Decomposition reactions
Bonds are broken in larger
molecules, resulting in smaller,
less complex molecules.
Exchange reactions
Bonds are both made and broken
(also called displacement reactions).
Example
Glycogen is broken down to release
glucose units.
Example
ATP transfers its terminal phosphate
group to glucose to form glucosephosphate.
+
Amino acid
molecules
Glycogen
Adenosine triphosphate (ATP)Glucose
Protein
molecule
+
Glucose
molecules
Adenosine diphosphate
(ADP)
Glucosephosphate
Energy Flow in Chemical Reactions
• All chemical reactions are either exergonic
or endergonic
– Exergonic reactions—net release of energy
• Products have less potential energy than reactants
– Endergonic reactions—net absorption of
energy
• Products have more potential energy than
reactants
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Reversibility of Chemical Reactions
• All chemical reactions are theoretically
reversible
– A + B  AB
– AB  A + B
• Chemical equilibrium occurs if neither a
forward nor reverse reaction is dominant
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Rate of Chemical Reactions
• Affected by
–  Temperature   Rate
–  Concentration of reactant   Rate
–  Particle size   Rate
– Catalysts:  Rate without being chemically
changed or part of product
• Enzymes are biological catalysts
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Part 2: Biochemistry
Classes of Compounds
• Inorganic compounds
• *Water, **salts, and many acids and bases
• Do not contain carbon
• Organic compounds
• Carbohydrates, fats, proteins, and nucleic
acids
• Contain carbon, usually large, and are
covalently bonded
• Both equally essential for life
*part of hydrolysis and dehydration synthesis reactions
**Common salts in body: NaCl, CaCO3, KCl, calcium phosphates; Ions (electrolytes) conduct electrical currents in solution
Dehydration synthesis
Monomers are joined by removal of OH from one monomer
and removal of H from the other at the site of bond formation.
+
Monomer 2
Monomer 1
Monomers linked by covalent bond
Hydrolysis
Monomers are released by the addition of a water molecule,
adding OH to one monomer and H to the other.
Monomer 1
+
Monomers linked by covalent bond
Example reactions
Dehydration synthesis of sucrose and its breakdown by hydrolysis
Water is
released
+
Water is
consumed
Glucose
Fructose
Sucrose
Monomer 2
Acids, Bases, and pH
• Sometimes a solution has more H+ ions than OH- ions
ready to release: acid, pH < 7.
• Acids are proton donors
– Release H+ (a bare proton) in solution
– HCl H+ + Cl–
• Sometimes a solution has more OH- ions than
H+ ions ready to release: base (alkaline), pH > above
7. Often called a proton acceptor (Take up H+ from solution)
• NaOH Na+ + OH–
OH– accepts an available proton (H+)
OH– + H+ H2O
Some Important Acids and Bases in Body
• Important acids
– HCl, HC2H3O2 (HAc), and H2CO3
• Important bases
– Bicarbonate ion (HCO3–) and ammonia
(NH3)
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Concentration
(moles/liter)
[OH−]
[H+] pH
pH = - log
[H+]
• Pure water has a
concentration of 10-7, so
the pH is 7. A pH 6
solution actually has 10
times the number of H+
ions.
H+
100
10−14
14
1M Sodium
hydroxide (pH=14)
10−1
10−13
13
Oven cleaner, lye
(pH=13.5)
10−2
10−12
12
10−3
10−11
11
10−4
10−10
10
10−5
10−9
9
10−6
10−8
8
10−7
10−7
7 Neutral
10−8
10−6
6
10−9
10−5
5
10−10
10−4
4
10−11
10−3
3
10−12
10−2
2
10−13
10−1
1
10−14
100
0
Increasingly basic
From 0 to 14:
• 0: the strongest acid
• 14: strongest base.
Examples
Household ammonia
(pH=10.5–11.5)
Household bleach
(pH=9.5)
Egg white (pH=8)
Blood (pH=7.4)
Milk (pH=6.3–6.6)
Increasingly acidic
pH Scale
Black coffee (pH=5)
Wine (pH=2.5–3.5)
Lemon juice; gastric
juice (pH=2)
1M Hydrochloric
acid (pH=0)
pH: Acid-base Concentration
– Relative free [H+] of a solution measured
on pH scale
– As free [H+] increases, acidity increases
• [OH–] decreases as [H+] increases
• pH decreases
– As free [H+] decreases alkalinity increases
• [OH–] increases as [H+] decreases
• pH increases
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Neutralization
• Results from mixing acids and bases
– Displacement reactions occur forming
water and a salt
– Neutralization reaction
• Joining of H+ and OH– to form water neutralizes
solution
Example?
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Buffers
• Buffers stabilize H+ concentration in a solution.
– In blood, two molecules stabilize pH: bicarbonate
ion (HCO3-) and carbonic acid (H2CO3).
HCO3- + H+ ↔ H2CO3
• If blood falls below pH 7.35, the condition is called
acidosis.
• If blood rises above pH 7.45, the condition is called
alkalosis.
Buffers
• Buffers resist abrupt and large swings in pH
– Release hydrogen ions if pH rises
– Bind hydrogen ions if pH falls
• Convert strong (completely dissociated) acids or bases into
weak (slightly dissociated) ones
• Carbonic acid-bicarbonate system (important buffer system of
blood):
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Carbohydrates
• Sugars and starches
Organic compounds:
• Polymers
• Contain C, H, and O [(CH20)n]
• Three classes
– Monosaccharides – one sugar
– Disaccharides – two sugars
– Polysaccharides – many sugars
• Functions of carbohydrates
– Major source of cellular fuel (e.g., glucose)
– Structural molecules (e.g., ribose sugar in RNA)
Carbohydrates
• Monosaccharide: simple sugar, one
carbon ring
– Examples: glucose, fructose,
galactose
• Disaccharide: two monosaccharides
joined by a covalent bond
– Examples: sucrose, maltose,
lactose
• Polysaccharide: several
monosaccharides joined together
– Example: starch (composed of
thousands of glucose molecules)
Carbohydrates
• Glycogen: another
polysaccharide formed to
store sugar in a cell
– Glycogen does not pull in
water via osmosis as simple
sugars do.
• Cellulose: a polysaccharide
made by plants
– Cellulose is not digestible by
humans.
Lipids
• Contain C, H, O (less than in carbohydrates),
and sometimes P
• Insoluble in water
• Main types:
– Triglycerides or neutral fats
– Phospholipids
– Steroids
– Eicosanoids
Figure 2.16a Lipids.
Triglyceride formation
Three fatty acid chains are bound to glycerol by dehydration synthesis.
+
Glycerol
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+
3 fatty acid chains
Triglyceride, or neutral fat
3 water
molecules
Saturated and Unsaturated Fats
• If every carbon on the fatty acid chain shares a
single electron, the fatty acid is saturated.
• If there are double bonds between carbons, the
fatty acid is unsaturated.
• Trans fats – modified
oils – unhealthy
• Omega-3 fatty acids –
“heart healthy”
Ketone Bodies
• Hydrolysis of triglycerides forms free fatty acids in
the blood. These can be used for energy or
converted into ketone bodies by the liver.
– Strict low-carbohydrate diets and uncontrolled diabetes
can result in elevated ketone levels, called ketosis.
– Ketone levels lower pH ->can cause ketoacidosis, which
can lead to coma and death.
Phospholipids
• Modified triglycerides:
– Glycerol + two fatty acids and a phosphorus (P) containing group
• “Head” and “tail” regions have different properties
• Important in cell membrane structure
“Typical” structure of a phospholipid molecule
Two fatty acid chains and a phosphorus-containing group are attached to the glycerol backbone.
Example
Phosphatidylcholine
Polar “head”
Nonpolar “tail”
(schematic
phospholipid)
Phosphorus-containing
group (polar “head”)
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Glycerol
backbone
2 fatty acid chains
(nonpolar “tail”)
Steroids
• A steroid is structurally very different
from a triglyceride but nonpolar, so
considered a lipid.
– 3 six-carbon rings + 1 five-carbon
ring + functional groups
• Cholesterol is a steroid used (1) as a
precursor to steroid hormones, such
as testosterone, estrogen, and
aldosterone, and (2) to make
molecules such as vitamin D and bile
salts.
Eicosanoids
• Many different ones
• Derived from a fatty acid (arachidonic acid)
in cell membranes
• Most important eicosanoid
– Prostaglandins
• Role in blood clotting, control of blood pressure,
inflammation, and labor contractions
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Proteins
• Contain C, H, O, N, and sometimes S and P
• Proteins are polymers
• Amino acids (20 types) are the monomers in
proteins
– Joined by covalent bonds called peptide bonds
– Contain amine group and acid group
– Can act as either acid or base
– All identical except for “R group” (in green on figure)
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Amino Acids
• An amino acid has an amino group, a carboxyl
group, and a functional group.
– The functional group is what differentiates
the 20 amino acids.
Amine
group
Acid
group
Generalized
structure of all
amino acids.
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Glycine
is the simplest
amino acid.
Aspartic acid
(an acidic amino
acid) has an acid
group (—COOH)
in the R group.
Lysine
(a basic amino
acid) has an amine
group (—NH2) in
the R group.
Cysteine
(a basic amino acid)
has a sulfhydryl (—SH)
group in the R group,
which suggests that
this amino acid is likely
to participate in
intramolecular bonding.
Figure 2.18 Amino acids are linked together by peptide bonds.
Dehydration synthesis:
The acid group of one amino
acid is bonded to the amine
group of the next, with loss
of a water molecule.
Peptide
bond
+
Amino acid
Dipeptide
Amino acid
Hydrolysis: Peptide bonds
linking amino acids together
are broken when water is
added to the bond.
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Protein Structure
• A chain of amino acids is called a polypeptide
chain.
– primary structure: The chain varies in length
from 3 to 4,500 amino acids.
• Weak hydrogen bonds may form between
neighboring amino acids.
– This may form an alpha helix or a beta fold:
secondary structure.
Protein Structure
• Attraction to amino acids further away
produces bends and folds, creating a
specific 3D shape: tertiary structure
– This structure dictates function.
– Since weak bonds hold tertiary structure
together, a protein is easily denatured
(unfolded) by changes in pH or temperature.
Protein Structure
• Some functional proteins
are composed of multiple
polypeptide chains
covalently bonded
together.
– This is called the
quaternary structure of
the protein.
– Examples are the
hemoglobin in blood
and the hormone
insulin.
Protein Functions
• Structural: collagen fibers in connective
tissues; keratin in skin
• Enzymes: assist every chemical process in
the body
• Antibodies: part of the immune system
• Receptors: receive communication from
other cells for regulation of cell activity
• Carriers: across cell membranes or in blood
Protein Denaturation
• Denaturation
– Globular proteins unfold and lose functional,
3-D shape
• Active sites destroyed
– Can be cause by decreased pH or increased
temperature
• Usually reversible if normal conditions
restored
• Irreversible if changes extreme
– e.g., cooking an egg
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Enzymes
– Proteins that act as biological catalysts
• Regulate and increase speed of chemical reactions
– Lower the activation energy, increase the speed of
a reaction
WITHOUT ENZYME
WITH ENZYME
Less activation
energy required
Energy
Energy
Activation
energy
required
Reactants
Reactants
Product
Progress of reaction
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Product
Progress of reaction
Characteristics of Enzymes
•
Enzymes are specific
– Act on specific substrate
•
Usually end in -ase
•
Often named for the reaction they catalyze
– Hydrolases, oxidases
Figure 2.21 Mechanism of enzyme action.
Substrates (S)
e.g., amino acids
+
Energy is Water is
absorbed; released.
bond is
formed.
Product (P)
e.g., dipeptide
Peptide
bond
Active site
Enzyme (E)
Enzyme-substrate
complex (E-S)
1 Substrates bind at active 2 The E-S complex
site, temporarily forming an undergoes internal
enzyme-substrate complex. rearrangements that
form the product.
Enzyme (E)
3 The enzyme releases
the product of the
reaction.
Nucleic Acids
• Building blocks: Nucleotides
– Composed of a five-carbon sugar, a
phosphate group, and a nitrogenous base
– Nitrogenous bases fall into two categories:
• Pyrimidine: a single carbon ring + nitrogen
• Purine: 2 carbon rings + nitrogen
Deoxyribonucleic Acid (DNA)
• The sugar in this molecule is
called deoxyribose and can
bind to one of four
nitrogenous bases:
–
–
–
–
Guanine
Thymine
Cytosine
Adenine
DNA Structure
• Deoxyribose bonds with a phosphate group to
form a long chain, which serves as the backbone
of the molecule.
• Each nitrogenous base can form a hydrogen
bond with another to result in a double-stranded
molecule.
– Cytosine can only bind with guanine.
– Tyrosine can only bind with adenine.
• The two chains of DNA are twisted,
forming a double helix.
Ribonucleic Acid (RNA)
• Similar to DNA except:
– Has ribose sugar
instead of
deoxyribose
– Is single-stranded
instead of doublestranded
– Has uracil instead of
thymine
Adenosine Triphosphate (ATP)
• Chemical energy in glucose captured in
this important molecule
• Directly powers chemical reactions in cells
• Energy form immediately useable by all
body cells
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Figure 2.23 Structure of ATP (adenosine triphosphate).
High-energy phosphate
bonds can be hydrolyzed
to release energy.
Adenine
Phosphate groups
Ribose
Adenosine
Adenosine monophosphate (AMP)
Adenosine diphosphate (ADP)
Adenosine triphosphate (ATP)
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Figure 2.24 Three examples of cellular work driven by energy from ATP.
Function of ATP
Solute
game
• Phosphorylation
– Terminal phosphates
are enzymatically
transferred to and
energize other
molecules
– Such “primed”
molecules perform
cellular work (life
processes) using the
phosphate bond
energy
+
Membrane
protein
Transport work: ATP phosphorylates transport proteins,
activating them to transport solutes (ions, for example)
across cell membranes.
+
Relaxed smooth
muscle cell
Contracted smooth
muscle cell
Mechanical work: ATP phosphorylates contractile proteins in muscle cells so the cells can shorten.
+
Chemical work: ATP phosphorylates key reactants, providing
energy to drive energy-absorbing chemical reactions.
© 2013 Pearson Education, Inc.
Case Studies to be discussed:
Ch 2: Critical Thinking and Clinical Application Questions
#2: Some antibiotics act by binding to certain essential enzymes in the target
bacteria. A) How might these antibiotics influence the chemical reactions controlled
by the enzymes? B) What is the anticipated effect on the bacteria? On the person
taking the antibiotic prescription?
#3: Mrs. R., in a diabetic coma, has just been admitted to N hospital. Her blood pH
indicates that she is in severe acidosis, and measures are quickly instituted to bring
her pH back to normal. A) Define pH and note the normal pH of blood. B) Why is
severe acidosis a problem?
#4: Jason (12 yo) was awakened by a loud crash. As he sat up, his fright was
revealed by his rapid breathing (hyperventilation). A) What happens during
hyperventilation?, B) At this point, was his blood pH rising or falling?