remember metabolism is just those chemical reactions in our body…so we have to understand basic...
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Human Anatomy and Physiology
Basic Chemistry Unit 2Remember Metabolism is just those chemical
reactions in our body…so we have to understand Basic Chemistry
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Matter vs. Energy
Matter can store energy through atomic bonds
We understand that matter can release energy…and that moving matter really fast (energy) can create new matter (particle accelerator)
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Matter…a practical working definition
Anything that has mass and occupies space
States of matter:1. Solid—definite shape and volume2. Liquid—definite volume, changeable shape3. Gas—changeable shape and volume
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Energy can be converted from one form to another, but always loses heat between conversions
Potential
Kinetic
Chemical
Electrical
Mechanical
Radiant or Electromagnetic
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Composition of Matter Elements
◦ Cannot be broken down by ordinary chemical means
◦ Each has unique properties: Physical properties
Are detectable with our senses, or are measurable
Chemical properties How atoms interact (bond) with each other
Atoms◦ Unique building blocks for each element◦ Atomic symbol: one- or two-letter chemical shorthand
for each element Au Fe Mg K Cl F Sn Cu
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Elements of the Human Body
Oxygen (O) Carbon (C) Hydrogen (H) Nitrogen (N)
Minor elements: About 3.9% of body mass◦ Calcium (Ca), phosphorus (P), potassium (K), sulfur
(S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe)
Trace elements: < 0.01% of body mass◦ Part of enzymes, e.g., chromium (Cr), manganese
(Mn), and zinc (Zn)
About 96% of body mass
These are major elements
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Your guide to atoms. Make sure you can read it!
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So the basic unit of matter is…the ATOM
Be able to describe and draw an atom…including:
Protons Neutrons Electrons Ions Isotopes
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Atomic Structure Neutrons
No charge Mass = 1 atomic mass unit (amu)
Protons Positive charge Mass = 1 amu
Electrons◦ Orbit nucleus◦ Equal in number to protons in atom◦ Negative charge ◦ 1/2000 the mass of a proton (0 amu)
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Figure 2.1
(a) Planetary model (b) Orbital model
Helium atom
2 protons (p+)2 neutrons (n0)2 electrons (e–)
Helium atom
2 protons (p+)2 neutrons (n0)2 electrons (e–)
Nucleus Nucleus
Proton Neutron Electroncloud
Electron
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Figure 2.2
Proton
Neutron
Electron
Helium (He)(2p+; 2n0; 2e–)
Lithium (Li)(3p+; 4n0; 3e–)
Hydrogen (H)(1p+; 0n0; 1e–)
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Identifying Elements Atomic number = number of protons in
nucleus Mass number = mass of the protons and
neutrons◦ Mass numbers of atoms of an element are not all
identical◦ Isotopes are structural variations of elements that
differ in the number of neutrons they contain Atomic weight = average of mass numbers
of all isotopes
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Figure 2.3
Proton
Neutron
Electron
Deuterium (2H)(1p+; 1n0; 1e–)
Tritium (3H)(1p+; 2n0; 1e–)
Hydrogen (1H)(1p+; 0n0; 1e–)
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Radioisotopes Spontaneous decay (radioactivity) Similar chemistry to stable isotopes Can be detected with scanners Valuable tools for biological research and
medicine Cause damage to living tissue:
◦ Useful against localized cancers◦ Radon from uranium decay causes lung cancer
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Ionic Bonds Ions are formed by transfer of valence shell
electrons between atoms◦ Anions (– charge) have gained one or more
electrons◦ Cations (+ charge) have lost one or more
electrons Attraction of opposite charges results in an
ionic bond
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Molecules and Compounds Most atoms combine chemically with other
atoms to form molecules and compounds◦ Molecule—two or more atoms bonded together
(e.g., H2 or C6H12O6) Compound—two or more different kinds of atoms
bonded together (e.g., C6H12O6) Elements- one or more of the same atom bonded
together
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Chemical Bonds Electrons occupy up to seven electron shells
(energy levels) around nucleus Octet rule: Except for the first shell which is
full with two electrons, atoms interact in a manner to have eight electrons in their outermost energy level (valence shell)
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Figure 2.5a
Helium (He)(2p+; 2n0; 2e–)
Neon (Ne)(10p+; 10n0; 10e–)
2e 2e8e
(a) Chemically inert elements
Outermost energy level (valence shell) complete
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Figure 2.5b
2e4e
2e8e
1e
(b) Chemically reactive elementsOutermost energy level (valence shell) incomplete
Hydrogen (H)(1p+; 0n0; 1e–)
Carbon (C)(6p+; 6n0; 6e–)
1e
Oxygen (O)(8p+; 8n0; 8e–) Sodium (Na)
(11p+; 12n0; 11e–)
2e6e
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Be able to draw and predict bonding for the following atoms:◦ Carbon◦ Hydrogen◦ Calcium◦ Oxygen◦ Fluorine◦ Lithium◦ Sodium◦ Chlorine◦ Helium
Valence Electrons and Bonding
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Ionic
Covalent
Hydrogen
Three important Bonds
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Figure 2.6a-b
Sodium atom (Na)(11p+; 12n0; 11e–)
Chlorine atom (Cl)(17p+; 18n0; 17e–)
Sodium ion (Na+) Chloride ion (Cl–)
Sodium chloride (NaCl)
+ –
(a) Sodium gains stability by losing one electron, and chlorine becomes stable by gaining one electron.
(b) After electron transfer, the oppositely charged ions formed attract each other.
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Formation of an Ionic Bond Ionic compounds form crystals instead of
individual molecules◦ NaCl (sodium chloride)
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Figure 2.6c
CI–
Na+
(c) 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
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Figure 2.7a
+
Hydrogenatoms
Carbonatom
Molecule ofmethane gas (CH4)
Structuralformulashows singlebonds.
(a) Formation of four single covalent bonds: carbon shares four electron pairs with four hydrogen atoms.
or
Resulting moleculesReacting atoms
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Figure 2.7b
or
Oxygenatom
Oxygenatom
Molecule ofoxygen gas (O2)
Structuralformulashowsdouble bond.(b) Formation of a double covalent bond: Two
oxygen atoms share two electron pairs.
Resulting moleculesReacting atoms
+
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Figure 2.7c
+ or
Nitrogenatom
Nitrogenatom
Molecule ofnitrogen gas (N2)
Structuralformulashowstriple bond.(c) Formation of a triple covalent bond: Two
nitrogen atoms share three electron pairs.
Resulting moleculesReacting atoms
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Covalent Bonds Sharing of electrons may be equal or
unequal◦ Equal sharing produces electrically balanced
nonpolar molecules CO2
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Figure 2.8a
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Covalent Bonds Unequal sharing by atoms with different
electron-attracting abilities produces polar molecules◦ H2O
Atoms with six or seven valence shell electrons are electronegative, e.g., oxygen
Atoms with one or two valence shell electrons are electropositive, e.g., sodium
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Figure 2.8b
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Figure 2.9
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Hydrogen Bonds Attractive force between electropositive
hydrogen of one molecule and an electronegative atom of another molecule◦ Common between dipoles such as water◦ Also act as intramolecular bonds, holding a large
molecule in a three-dimensional shape◦ http://www.youtube.com/watch?v=lkl5cbfqFRM
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(a) The slightly positive ends (+) of the watermolecules become aligned with the slightlynegative ends (–) of other water molecules.
+
–
–
–– –
+
+
+
+
+
Hydrogen bond(indicated bydotted line)
Figure 2.10a
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Figure 2.10b
(b) A water strider can walk on a pond because of the highsurface tension of water, a result of the combinedstrength of its hydrogen bonds.
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Chemical vs Physical reactions
Chemical reactions Physical reactions
Chemical bonds are formed, rearranged or broken
Synthesis, Decomposition, Exchange, Redox
Chemical Equations◦ Reactants◦ Products◦ Molecular formulas ◦ Balance atoms
No chemical bonds are changed
Most matter exists as MIXTURES◦ Solution
Aqueous/ homeogeneous Solute/ solvent http://www.youtube.com/watch
?v=3G472AA3SEs◦ Colloid
Medium solute/ gel/ foam heterogeneous
◦ Suspension Large solute/ heterogeneous
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Concentration of Solutions Expressed as
◦ Percent, or parts per 100 parts◦ Milligrams per deciliter (mg/dl)◦ Ppm: parts per million◦ Molarity, or moles per liter (M)
1 mole = the atomic weight of an element or molecular weight (sum of atomic weights) of a compound in grams
1 mole of any substance contains 6.02 1023 molecules (Avogadro’s number) What is the molecular weight of a gram of glucose
(C6H12O6)? To make a 1M (one molar) solution of glucose, how much
glucose would you weigh to mix with water to make one liter?
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Figure 2.4
Solution
Soluteparticles
Soluteparticles
Soluteparticles
Solute particles are verytiny, do not settle out orscatter light.
ColloidSolute particles are largerthan in a solution and scatterlight; do not settle out.
SuspensionSolute particles are verylarge, settle out, and mayscatter light.
ExampleMineral water
ExampleGelatin
ExampleBlood
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Chemical Reactions can be expressed as EquationsH + H H2 (hydrogen gas)
4H + C CH4 (methane)
Bonds are always broken, made, rearrangedAtoms are always accounted for so formulas must be
balanced to work Chemical equilibrium occurs if neither a forward nor
reverse reaction is dominant Many biological reactions are essentially irreversible
due to Energy requirements and/or◦ Removal of products
(reactants) (product)
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Synthesis◦ Anabolic = forms bonds/ Endergonic (absorbs energy)◦ A + B → AB Found in rapidly growing tissues
Decomposition◦ Catabolic = breaks bonds/ Exergonic (releases energy)◦ AB → A + B Found in metabolic reactions/ cell respiration
Exchange or Displacement◦ Involves synthesis and decomposition/ atoms rearranged◦ AB + C → AC + B AB + CD → AD + CB
Oxidation-Reduction (Redox) reactions◦ Both decomposition and exchange reactions◦ Electrons move between reactants◦ Reactant that loses electron is an electron donor/ oxidized◦ Reactant that gains electron is an electron acceptor/ reduced
LEO the lion goes GER
Patterns of Chemical Reactions
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Rate of Chemical Reactions Rate of reaction is influenced by:
◦ temperature rate◦ particle size rate ◦ concentration of reactant rate
Catalysts: rate without being chemically changed◦ Enzymes are biological catalysts
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Figure 2.11a
ExampleAmino acids are joined together toform a protein molecule.
(a) Synthesis reactions
Smaller particles are bondedtogether to form larger,more complex molecules.
Amino acidmolecules
Proteinmolecule
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Figure 2.11b
ExampleGlycogen is broken down to releaseglucose units.
Bonds are broken in largermolecules, resulting in smaller,less complex molecules.
(b) Decomposition reactions
Glucosemolecules
Glycogen
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Figure 2.11c
ExampleATP transfers its terminal phosphategroup to glucose to form glucose-phosphate.
Bonds are both made and broken(also called displacement reactions).
(c) Exchange reactions
Glucose Adenosine triphosphate (ATP)
Adenosine diphosphate (ADP)Glucosephosphate
+
+
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Oxidation Reduction (Redox)
LEO the lion goes GER
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Common Example of Redox Reactions in Body:
Aerobic Cell Respiration
Glucose is oxidized to carbon dioxide as it loses hydrogen atoms oxygen is reduced to water as it accepts the hydrogen atoms
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Classes of Compounds Inorganic compounds
Water, salts, and many acids and bases Do not have to contain carbon
Organic compounds Carbohydrates, fats, proteins, and nucleic acids Must contain carbon, usually large, and are
covalently bonded
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Water 60%–80% of the volume of living cells Most important inorganic compound in
living organisms because of its properties
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Properties of Water High heat capacity
◦ Absorbs and releases heat with little temperature change
◦ Prevents sudden changes in temperature High heat of vaporization
◦ Evaporation requires large amounts of heat◦ Useful cooling mechanism
Polar solvent properties◦ Dissolves and dissociates ionic substances◦ Forms hydration layers around large charged
molecules, e.g., proteins (colloid formation)◦ Body’s major transport medium
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Figure 2.12
Water molecule
Ions in solutionSalt crystal
–
+
+
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Properties of Water Reactivity
◦ A necessary part of hydrolysis and dehydration synthesis reactions
Cushioning◦ Protects certain organs from physical trauma,
e.g., cerebrospinal fluid
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Figure 2.14
+
Glucose Fructose
Water isreleased
Monomers linked by covalent bond
Monomers linked by covalent bond
Water isconsumed
Sucrose
(a) Dehydration synthesis
Monomers are joined by removal of OH from one monomerand removal of H from the other at the site of bond formation.
+
(b) Hydrolysis
Monomers are released by the addition of a water molecule, adding OH to one monomer and H to the other.
(c) Example reactions
Dehydration synthesis of sucrose and its breakdown by hydrolysis
Monomer 1 Monomer 2
Monomer 1 Monomer 2
+
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Salts Ionic compounds that dissociate in water Contain cations other than H+ and anions
other than OH–
Ions (electrolytes) conduct electrical currents in solution
Ions play specialized roles in body functions (e.g., sodium, potassium, calcium, and iron)
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pH…its all about the ions
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Acids and Bases Both are electrolytes Acids are proton (hydrogen ion) donors
(release H+ in solution) HCl H+ + Cl–
Bases are proton acceptors (take up H+ from solution)◦ NaOH Na+ + OH–
OH– accepts an available proton (H+) OH– + H+ H2O
Bicarbonate ion (HCO3–) and ammonia (NH3)
are important bases in the body
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pH: Acid-Base Concentration pH = the negative logarithm of [H+] in
moles per liter Neutral solutions:
◦ Pure water is pH neutral (contains equal numbers of H+ and OH–)
◦ pH of pure water = pH 7: [H+] = 10 –7 M◦ All neutral solutions are pH 7
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pH: Acid-Base Concentration Acidic solutions
◦ [H+], pH ◦ Acidic pH: 0–6.99◦ pH scale is logarithmic: a pH 5 solution has 10
times more H+ than a pH 6 solution Alkaline solutions
◦ [H+], pH◦ Alkaline (basic) pH: 7.01–14
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Common Substances and their pH value
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Figure 2.13
Concentration(moles/liter)
[OH–]
100 10–14
10–1 10–13
10–2 10–12
10–3 10–11
10–4 10–10
10–5 10–9
10–6 10–8
10–7 10–7
10–8 10–6
10–9 10–5
10–10 10–4
10–11 10–3
10–12 10–2
10–13 10–1
[H+] pHExamples
1M Sodiumhydroxide (pH=14)
Oven cleaner, lye(pH=13.5)
Household ammonia(pH=10.5–11.5)
Neutral
Household bleach(pH=9.5)
Egg white (pH=8)
Blood (pH=7.4)
Milk (pH=6.3–6.6)
Black coffee (pH=5)
Wine (pH=2.5–3.5)
Lemon juice; gastricjuice (pH=2)
1M Hydrochloricacid (pH=0)10–14 100
14
13
12
11
10
9
8
7
6
5
4
3
2
1
0
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Acid-Base Homeostasis pH change interferes with cell function and
may damage living tissue Slight change in pH can be fatal pH is regulated by kidneys, lungs, and
buffers
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Buffers Mixture of compounds that resist pH
changes Convert strong (completely dissociated)
acids or bases into weak (slightly dissociated) ones◦ Carbonic acid-bicarbonate system
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Carbon base◦ Electroneutral! …makes
carbon a small sharer
Functional groups
Monomers and polymers
Organic Compounds
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Carbohydrates
Monosaccharides: CH2O
Disaccharides
FunctionIs Fuel, Easy, AvailableFuel for Our cells
Polysaccharides
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Carbohydrates Sugars and starches Contain C, H, and O [(CH20)n] Three classes
◦ Monosaccharides◦ Disaccharides◦ Polysaccharides
Functions◦ Major source of cellular fuel (e.g., glucose)◦ Oxidation reduction reactions◦ Structural molecules (e.g., ribose sugar in RNA or
glycocalyx)
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Figure 2.15a
ExampleHexose sugars (the hexoses shown here are isomers)
ExamplePentose sugars
Glucose Fructose Galactose Deoxyribose Ribose
(a) MonosaccharidesMonomers of carbohydrates
Simple sugars containing three to seven C atoms (CH20)n
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Figure 2.15b
ExampleSucrose, maltose, and lactose(these disaccharides are isomers)
Glucose Fructose Glucose Glucose Glucose
Sucrose Maltose Lactose
Galactose
(b) DisaccharidesConsist of two linked monosaccharides
Double sugarsToo large to pass through cell membranes
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Figure 2.15c
ExampleThis polysaccharide is a simplified representation of glycogen, a polysaccharide formed from glucose units.
(c) PolysaccharidesLong branching chains (polymers) of linked monosaccharides
Glycogen
Polymers of simple sugars, e.g., starch and glycogen/ Not very soluble
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Lipids: energy storage, structure, insulation
Contain C, H, O (less than in carbohydrates), and sometimes P, Insoluble in water
Triglycerides Steroids
Eicosanoids Lipoproteins
◦ HDL and LDL
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Triglycerides Neutral fats—solid fats and liquid oils Composed of three fatty acids bonded to a
glycerol molecule Main functions
◦ Energy storage◦ Insulation◦ Protection
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Figure 2.16a
Glycerol
+
3 fatty acid chains Triglyceride,or neutral fat
3 watermolecules
(a) Triglyceride formation
Three fatty acid chains are bound to glycerol bydehydration synthesis
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Saturation of Fatty Acids Saturated fatty acids
◦Single bonds between C atoms; maximum number of H
◦Solid animal fats, e.g., butter Unsaturated fatty acids
◦One or more double bonds between C atoms
◦Reduced number of H atoms ◦Plant oils, e.g., olive oil Trans Fats Omega 3 Fatty acids
Omega 6
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Figure 2.16b
Phosphorus-containing
group (polar“head”)
ExamplePhosphatidylcholine
Glycerolbackbone
2 fatty acid chains(nonpolar “tail”)
Polar“head”
Nonpolar“tail”
(schematicphospholipid)
(b) “Typical” structure of a phospholipid molecule
Two fatty acid chains and a phosphorus-containing group areattached to the glycerol backbone.
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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
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Steroids Steroids—basically flat, interlocking four-ring
structure Cholesterol:
◦ vitamin D, steroid hormones, and bile salts◦ Steroid hormones include estrogen, progesterone,
testosterone, adrenocortical hormones (cortisol and aldosterone)
◦ Role in membrane: Limits movement (fluidity) of fatty acid chains Stabilizes membrane Lowers freezing point so allows functional membrane in
colder temperatures
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Figure 2.16c
ExampleCholesterol (cholesterol is thebasis for all steroids formed in the body)
(c) Simplified structure of a steroid
Four interlocking hydrocarbon rings form a steroid.
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Other Lipids in the Body Other fat-soluble vitamins
◦ Vitamins A, E, and K Lipoproteins
◦ Transport fats in the blood◦ HDL: carries cholesterol from cells to liver◦ LDL: carries cholesterol to cells
Eicosanoids◦ Many different types such as prostaglandins and
leukotrienes◦ Derived from a fatty acid (arachidonic acid) in cell
membranes◦ Acts as paracrines
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Protein
Proteome makes up 10-30% of cell mass. While proteins make the basic structure of body, they also have many other vital roles of life.
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Figure 2.17
(a) Generalized structure of all amino acids.
(b) Glycine is the simplest amino acid.
(c) Aspartic acid (an acidic amino acid) has an acid group (—COOH) in the R group.
(d) Lysine (a basic amino acid) has an amine group (–NH2) in the R group.
(e) 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.
Aminegroup
Acidgroup
Polymers of amino acids (20 types)Joined by peptide bonds
Contain C, H, O, N, and sometimes S and P
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Figure 2.18
Amino acid Amino acid Dipeptide
Dehydration synthesis:The acid group of one amino acid is bonded to the amine group of the next, with loss of a water molecule.
Hydrolysis: Peptide bonds linking amino acids together are broken when water is added to the bond.
+
Peptidebond
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Figure 2.19a
(a) Primary structure: The sequence of amino acids forms the polypeptide chain.
Amino acid Amino acid Amino acid Amino acid Amino acid
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Figure 2.19b
a-Helix: The primary chain is coiledto form a spiral structure, which isstabilized by hydrogen bonds.
b-Sheet: The primary chain “zig-zags” backand forth forming a “pleated” sheet. Adjacentstrands are held together by hydrogen bonds.
(b) Secondary structure:The primary chain forms spirals (a-helices) and sheets (b-sheets).
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Figure 2.19c
Tertiary structure of prealbumin(transthyretin), a protein thattransports the thyroid hormonethyroxine in serum and cerebro-spinal fluid.
(c) Tertiary structure: Superimposed on secondary structure. a-Helices and/or b-sheets are folded up to form a compact globular molecule held together by intramolecular bonds.
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Figure 2.19d
Quaternary structure ofa functional prealbuminmolecule. Two identicalprealbumin subunitsjoin head to tail to formthe dimer.
(d) Quaternary structure: Two or more polypeptide chains, each with its own tertiary structure, combine to form a functional protein.
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Fibrous and Globular Proteins Fibrous (structural) proteins
◦ Strandlike, water insoluble, and stable ◦ Examples: keratin, elastin, collagen, and certain
contractile fibers Globular (functional) proteins
◦ Compact, spherical, water-soluble and sensitive to environmental changes
◦ Specific functional regions (active sites) ◦ Examples: antibodies, hormones, molecular
chaperones, and enzymes
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Protein Denaturation Shape change and disruption of active sites
due to environmental changes (e.g., decreased pH or increased temperature)
Reversible in most cases, if normal conditions are restored
Irreversible if extreme changes damage the structure beyond repair (e.g., cooking an egg)
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Molecular Chaperones (Chaperonins) Ensure quick and
accurate folding and association of proteins
Assist translocation of proteins and ions across membranes
Promote breakdown of damaged or denatured proteins
Help trigger the immune response
Produced in response to stressful stimuli, e.g., O2 deprivation
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Enzymes
Enzymes are biological catalysts that do not initiate a reaction, but only speed up a reaction. Note the change in configuration as enzyme binds to substrate .
Enzymes speed up reactions by reducing required activation energy. How do enzymes do this?
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Figure 2.20
Activationenergy required
Less activationenergy required
WITHOUT ENZYME WITH ENZYME
Reactants
Product Product
Reactants
Biological catalystsLower the activation energy, increase the speed of a reaction (millions of reactions per minute!)
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Characteristics of Enzymes Often named for the reaction they catalyze;
usually end in -ase (e.g., hydrolases, oxidases)
Chain reactions/ inactive forms Either pure protein or functional enzymes
(holoenzymes)◦Apoenzyme (protein portion) ◦Cofactor (metal ion) or coenzyme (a
vitamin) Copper/ iron B Vitamins
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Holoenzyme
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Figure 2.21
Substrates (S)e.g., amino acids
Enzyme (E)
Enzyme-substratecomplex (E-S)
Enzyme (E)
Product (P)e.g., dipeptide
Energy isabsorbed;bond isformed.
Water isreleased.
Peptidebond
Substrates bindat active site.Enzyme changesshape to holdsubstrates inproper position.
Internalrearrangementsleading tocatalysis occur.
Product isreleased. Enzymereturns to originalshape and isavailable to catalyzeanother reaction.
Active site
+ H2O
1 23
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Proteosomes
Such a strange discovery: a protein that acts like an enzyme. Mad Cow disease is the result of a similar type molecule called a prion.
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Nucleic Acids DNA and RNA
◦ Largest molecules in the body Contain C, O, H, N, and P Building block = nucleotide, composed of N-
containing base, a pentose sugar, and a phosphate group
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Nucleic Acids
DNA (Deoxyribonucleic acid)
RNA(Ribonucleic Acid)
Found in nucleus Genome Replicates (interphase)
and transcribes Monomer of deoxyribose
and 4 complementary bases: adenine/ guanine/ cytosine/ thymine
Double helix Watson and Crick
Produced in nucleus but functions in cytoplasm
Transcription/ translation (protein synthesis)
Different types: mRNA, tRNA, rRNA
Single strand/ straight or folded
Ribose sugar and substitute uracil for thymine
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Figure 2.22
Deoxyribosesugar
Phosphate
Sugar-phosphatebackbone
Adenine nucleotideHydrogenbond
Thymine nucleotide
PhosphateSugar:
Deoxyribose PhosphateSugarThymine (T)Base:
Adenine (A)
Adenine (A)
Thymine (T)
Cytosine (C)
Guanine (G)
(b)
(a)
(c) Computer-generated image of a DNA molecule
Nucleotides: Monomers of Nucleic Acids
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Figure 2.23
Adenosine triphosphate (ATP)
Adenosine diphosphate (ADP)
Adenosine monophosphate (AMP)
Adenosine
Adenine
Ribose
Phosphate groups
High-energy phosphatebonds can be hydrolyzedto release energy.
Adenine-containing RNA nucleotide with two additional phosphate groups
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Adenosine Triphosphate
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Function of ATP Phosphorylation:
◦ Terminal phosphates are enzymatically transferred to and energize other molecules
◦ Such “primed” molecules perform cellular work (life processes) using the phosphate bond energy
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Figure 2.24
Solute
Membraneprotein
Relaxed smoothmuscle cell
Contracted smoothmuscle cell
+
+
+
Transport work: ATP phosphorylates transportproteins, activating them to transport solutes(ions, for example) across cell membranes.
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.
(a)
(b)
(c)