Review on Biology Camble Biology Chapters 2-5
Review on Biology Camble Biology Chapters 2-5 1441
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Bio1441 Chapters 25 Review Chapter 2: The Chemical Context of Life I. Matter consists of chemical elements in pure form and in combinations called compounds a. Elements and compounds i. Element: substance that cannot be broken down into other substances by chemical rxns (92 that occur in nature) ii. Compound: substance made of 2 or more different elements combined in a fixed ratio 1. Ex: NaCl—1:1, pure Na is metal (corrosive) and Cl is a gas (toxic)—combined they make table salt (emergent properties) b. The elements of life 1. Of the 92 in nature, approx. 25% are essential for life (humans 25, plants 17) 2. Humans: C, H, O, and N make up 96.3% of living matter; Ca, P, K, S, Cl, Mg = 3.7% ii. Trace elements: needed in small quantities 1. Ex: Fe needed in all lifeforms 2. Ex: vertebrates need iodine b/c it is essential to hormones produced by the thyroid gland (goiter—enlarged gland due to iodine deficiency) c. Evolution of tolerance to toxic elements i. Ex: serpentine plants—can survive in enviro with heavy cobalt, chromium, and nickel deposits II. An element’s properties depend on the structure of its atoms i. Atom: smallest unit of matter that still retains the properties of an element 1. Charge is neutral (protons = electrons) unless indicated otherwise b. Subatomic particles i. Neutrons: no charge, found in nucleus, approx. 1 Dalton ii. Protons: positive, found in nucleus, approx. 1 Dalton iii. Electrons: negative, surround nucleus in a cloud and travel at speed of light (centrifugal th force), approx. 1/2000 Dalton c. Atomic number and atomic mass i. Atomic number: number of protons (and electrons if neutral), written as subscript to the leftdown of an element (Ex: H2) ii. Mass number: sum of protons and neutrons, written as subscript to the leftup of the element ( 23 Na—mass number is 23 [protons + neutrons] and atomic number [protons] is 11 11 so 2311= number of neutrons = 12 iii. Atomic mass: total mass (only slightly more than mass number because electron mass is negligible) d. Isotopes: different number of neutrons of same element i. Ex: carbon12 ( C) which is 99% of carbon in nature and has 6 neutrons; C (6 12 6 neutrons) and C (7 neutrons) are stable isotopes (their nuclei don’t tend to lose subatomic particles—a process called decay); C (8 neutrons) is radioactive (unstable) 1 Bio1441 Chapters 25 Review ii. Radioactive isotope: one in which the nucleus decays spontaneously, giving off particles and energy; when radioactive decay leads to a change in number of protons, it transforms the atom to an atom of another element (neutron splits into 1 electron and 1 proton) 1. Ex: C (6 protons) decays and becomes N (7 protons) e. Radioactive tracers: radioactive isotopes usually used as diagnostic tools in medicine i. Cells use radioactive atoms same as nonradioactive isotopes of same element; radioactive isotopes are incorporated into biologically active molecules which are then used as tracers to track atoms during metabolism 1. Ex: certain kidney disorders diagnosed by injecting small amount of radioactively labeled substances into blood and then analyzing tracer molecules excreted in urine ii. Radiometric dating: measure radioactive decay in fossils to date life 1. Parent isotope decays into daughter isotope as a fixed rate—halflife of isotope (time it takes for 50% of parent isotope to decay); halflife is not affected by temperature, pressure, or any environmental variable f. Energy levels of electrons i. Atoms are mostly empty space; electrons travel at speed of light; when 2 atoms approach during reaction, nuclei don’t come close enough to interact ii. Energy: capacity to cause change 1. Potential energy: energy that matter possesses because of its location or structure; matter tends to move toward the lowest possible PE a. Ex: H O2in a reservoir on a hill has PE because of its altitude; when H 2 is released, energy can be used to move blades and generate electricity; bottom = lease PE b. Electrons have PE due to distance from nucleus; takes work to move a given electron farther from the nucleus so the farther from the nucleus it is, the more PE it has i. Changes in PE in electron occur in fixed amounts (ball bouncing downstairs)—electron can only exist at certain energy level, not in between them iii. Electron shells: average distance from nucleus and energy level 1. Electron can move from shell to shell BUT only by absorbing/losing amount of energy equal to the difference in PE between its position in the old and new shells a. Ex: light energy excites electron to higher energy shell (photosynthesis) and falls back to old energy shell when it loses energy (released as heat) g. Electron distribution and chemical properties: i. Periodic table arranged in periods that tell number of shells (ex: 3 row = 3 shells) ii. First shell holds up to 2 electrons, all other up to 8 iii. Chemical behavior of atom depends on number of electrons in valence shell (outermost) so atoms with same valence have similar behaviors (Fl and Cl 7) 2 Bio1441 Chapters 25 Review iv. Atom with completed valence shell is unreactive (8 in valence shell or 2 if first shell) – noble gasses (He, Ne, Ar…) – inert h. Electron orbitals: 3D space where electron is found 90% of the time; each orbital holds max 2 electrons i. First electron shell has only one spherical orbital (1s) ii. Second shell has 4 orbitals: 1 large spherical (2s) and 3 dumbbellshaped porbitals (2p), electrons in each have same energy but move in different volumes of space iii. Reactivity of atoms arises from presence of unpaired electrons in one or more orbitals of its valence shell III. The formation and function of molecules depend on chemical bonding between atoms a. Covalent bonds: sharing of a pair of valence electrons by 2 atoms i. Ex: hydrogen atoms (each have 1 electron) overlap 1s orbital and share electrons so each have completed valence shell (2 electrons)—covalent bond froms molecule H (molecula2 formula), H:H (Lewis Dot), H—H (structural formula) ii. Ex: molecules H a2d O ar2 pure elements, NOT compounds (must be made of 2 or more different elements—example: H O) 2 iii. Electronegativity: attraction of an atom for the electron of a covalent bond; the more electronegative the atom is, the more strongly it pulls the shared electrons toward itself 1. Covalent bond between atoms of same element share electrons equally (nonpolar) 2. H O polar because oxygen is more electronegative 2 b. Ionic bonds: 2 atoms so different in attraction for valence electrons that the more electronegative atom takes one electron from the other element—now both are charged (ions); + = cation and = a+ion atract each other to form ionic bond i. Ex: NaCl (Na and Cl) form table salt (compounds from ionic bonds called salts) ii. Does NOT consist of molecules, just ration of elements iii. “ion” applies to atoms and molecules that are charged (ex: NH Cl) 4 c. Weak chemical bonds: needed for emergent properties of life like in cell communication or other temporary interactions) i. Hydrogen bonds: when hydrogen atom covalently bonded to an electronegative atom, H has slightly positive charge that allows it to be attracted to a different electronegative atom nearby—forms Hbond ii. Van der Waals interactions: electron not always equally distributed—even in nonpolar covalent compounds so electrons may accumulate in one place and create negative charge; always changing and allow molecules to stick to one another (ex: gecko climbing) iii. *cumulative effect of weak bonds: reinforce 3D shape of molecule d. Molecular shape and function: i. With 2 atoms of same element, it is always linear but usually more complicated shapes (determined by orbitals) 3 Bio1441 Chapters 25 Review ii. Hybridization of orbitals: single s and 3p orbitals of valence shell involved in covalent bonding combine to form 4 tear dropshaped hybrid orbitals that form tetrahedron (pyramid) iii. Molecular shape determines how biological molecules recognize and respond to one another with specificity; usually bond temporarily by forming weak bonds but only if shapes are complementary (ex: opiates have similar shapes to endorphins so they can latch on to receptors and mimic function but with dangerous effects) IV. Chemical reactions make and break chemical bonds a. Chemical reactions: making/breaking chemical bonds leading to changes in composition of matter (ex: 2H 2 2O 22H O ;2breaks covalent bonds between pure elements to form covalent bonds between different elements to make compounds)—matter is conserved, cannot be created or destroyed b. All chemical rxns are reversible c. Rate of rxn affected by concentration of reactants (greater = more frequently they collide and can react and form products)—same for products (as the accumulate, more collisions) i. Eventually forward and reverse rxns occur at the same rate and relative concentration of reactants and products stop changing—equal point: equilibrium (when it decomposes as quickly as it forms) 1. Both affect each other exactly—dynamic because rxns still going on; no net effect on concentration 2. Reactants and products not necessarily equal concentrations, just stabilized ratio Quiz: 1. In the term trace element, the adjective trace means that a. The element is required in very small amounts 2. Compared with P, the radioactive isotope P has 2 a. One more neutron 3. The reactivity of an atom arises from a. The existence of unpaired electrons in the valence shell 4. Which statement is true of all atoms that are anions? a. The atom has more electrons than protons 5. Which of the following statements correctly describes any chemical rxn that has reached equilibrium? a. The rates of the forward and reverse reactions are equal 6. We can represent atoms by listing the number or protons, neutrons, and electrons; for example— + 0 18 2p , 2n , 2e for Helium. Which of the following represents the O isotope of oxygen? a. 8p , 10n , 8e 7. The atomic number of sulfur is 16. Sulfur combines with hydrogen by covalent bonding to form a compound, hydrogen sulfide. Based on the number of valence electrons in a sulfur atom, predict the molecular formula of the compound. a. H S2 4 Bio1441 Chapters 25 Review 8. What coefficients must be placed in the following blanks so that all atoms are accounted for in the products? C 6 O126___ C H O2+ 6__ CO 2 a. 2; 2 Chapter 3: Water and Life I. The molecule that supports all of life a. Life began in water and evolved for 3 billion years; 75% of earth is water; only substance to exist in 3 states of matter naturally (solid is less dense than liquid form—rare) II. Polar covalent bonds in H O molecules result in hydrogen bonding 2 a. Emergent properties traced to structure and interactions of molecules b. Oxygen is more electronegative than hydrogen so electrons gravitate closer to O—polar covalent bonds so water is a polar molecule c. Properties of oxygen from oppositely charged atoms: slightly positive hydrogen attracted to slightly negative oxygen of nearby molecule and form hydrogen bond (1/20 strength of covalent bond) III. Four emergent properties of water contribute to earth’s sustainability for life: cohesion, ability to moderate temperature, expansion when freezing, versatility as a solvent a. Cohesion: water molecules stay close to each other because of hydrogen bonding i. Contributes to transport of water and dissolved nutrients against gravity in plants (water from roots reaches leaves through network of waterconducting cells that make up the xylem); as water evaporates from the leaf, hydrogen bonds cause water molecules leaving veins to tug on molecules farther down (cohesion)—transpiration ii. Adhesion: clinging of one substance to another; water forms hydrogen bonds with cell wall of plants’ cells iii. Surface tension: measure of how difficult it is to stretch/break the surface of a liquid; at interface between water and air, there is an ordered arrangement of water molecules hydrogen bonded to each other and water below, giving water high surface tension—the invisible film b. Moderation of temperature by water: absorbs heat from air that is warmer and releasing stored heat to cooler air—good heat bank because it can take in a lot of heat with only slight change in its own temperature i. Temperature and heat: anything that moves has kinetic energy (KE)—atoms always moving; faster = more KE 1. Thermal energy: KE associated with random movement of atoms/molecules, not same as temperature 5 Bio1441 Chapters 25 Review a. Temperature: measure of average KE of molecules in body of matter, regardless of volume whereas total thermal energy depends in part on volume b. Ex: coffee pot has higher temperature than cold pool abut pool has greater thermal energy because pool has much greater volume 2. When 2 objects brought together, thermal energy passes from warmer to cooler object until both have same temperature (ice cube cools drink by absorbing heat) a. Heat: thermal energy in transfer from one body of matter to another (unit: calorie—amount of heat it takes to raise temperature of 1g of water by 1 C OR amount of heat it releases to decrease temperature by 1 C; also Joules (J) ii. Evaporative cooling: can happen at any temperature but hotter = faster 1. Heat of vaporization: quantity of heat liquid must absorb for 1g of it to be converted from liquid to gas (water’s is HIGH)—to evaporate 1g of water at 25 C, about 580cal of heat needed—double that of alcohol or ammonia a. Another emergent property of hydrogen bonds that must be broken before molecules can exit to gaseous state 2. High amount of energy required to vaporize water—effects: a. Moderates climate (heat absorbed by tropical seas and released as water circulates poleward in form of rain); steam burns because high heat of vaporization 3. As liquid evaporates, surface of liquid that remains behind cools down— evaporative cooling occurs between hottest molecules (greatest KE) are most likely to leave as gas (average speed of remaining molecules declines because fastest ones are gone) a. Ex: evaporation of water form leaves of plant helps keep tissues in leaves form becoming too warm in sunlight iii. Floating of ice on liquid water: less dense as solid than liquid; solid expands because hydrogen bonds form from water moving too slow to break hydrogen bonds 1. O C—water locked in lattice because water is hydrogen bonded to 4 partners at “arm’s length” a. Importance: if ice sank, eventually ponds, lakes, and oceans would freeze solid, making life impossible (ice insulates liquid below and provides solid habitat for some organisms) c. Water: the solvent of life i. Solution: liquid that is a completely homogenous mixture 1. Solvent: dissolving agent of solution; solute: substance that is dissolved; aqueous solution: where water is the solvent—versatile because of hydrogen bonding and polarity 6 Bio1441 Chapters 25 Review + a. Ex: when NaCl placed in water, H attracted to Cl, separating and surrounding them; sphere of water molecules around each dissolved ion—hydration shell 2. Compound doesn’t have to be ionic to dissolve in water (ex: sugar) because it forms hydrogen bonds with them; even proteins if they have polar/ionic parts a. Blood, sap, etc.—organic compounds dissolved in water ii. Hydrophilic and hydrophobic substances: 1. Hydrophilic: any substance with an affinity for water (doesn’t have to dissolve in water; ex: cotton made of cellulose so has positive and negative regions to make hydrogen bonds; so water adheres to cotton but doesn’t dissolve it— present in cell walls of plants) 2. Hydrophobic: nonionic, nonpolar (cannot form hydrogen bonds) substance that repels water (ex: vegetable oil—nonpolar covalent bonds C—H so equally shared electrons) iii. Solute concentration in aqueous solutions: 1. Molecular mass: the sum of the masses of all the atoms in a molecule a. Ex: C H12 22u11ose) i. 1C = 12 Daltons, 1H = 1 Dalton, 1O = 16 Daltons ii. = 12x12 + 1x22 + 16x11 = 342 Daltons 2. 1 mole: represents an exact number of objects 6.02x10 (Avogadro’s #); 23 there are 6.02x10 Daltons in 1 gram SO to obtain 1 mole of sucrose, we weigh 342 g a. Advantage: a mole of one substance has the same number of molecules as a mole of any other substance i. If the molecular mass of A = 342D and the molecular mass of B = 10D, then both have same number of molecules, NOT mass b. How to make 1L solution with one mole of sucrose: add 346g slowly and stir water then add water until it reaches 1L—1M (molar) solution i. Molarity: (M) number of moles of solute per liter of solution iv. Possible evolution of life on other planets: 800+ planets found but water on few of them (Mars) IV. Acidic and basic conditions affect living organisms: a. Occasionally, hydrogen atom in hydrogen bond between 2 water molecules shifts from one + water to the other; H atom leaves its electron behind and H (proton) is transferred i. Water that lost its H is now OH (hydroxide); proton binds to other water making + H 3 (hydronium) ii. Reversible reaction that reaches equilibrium when water molecules dissociate at the same rate at which they reform iii. Ions at equilibrium (both very reactive); concentration of pure water = 10 M 7 b. Acids and bases: when H = OH, the solution is neutral; more OH = basic, more H = acidic + i. Acid: a substance that increases the hydrogen ion concentration of a solution 7 Bio1441 Chapters 25 Review 1. Ex: HCl added to water results in dissociation of HCl into H and Cl + + ii. Base: substance that reduces hydrogen ion concentration (some accept H like ammonia NH ) 3 + 1. NH : 3nshared electron in nitrogen valence shell attracts H ions from solution NH 3 H NH 4+ + 2. Other bases reduce H concentration indirectly by dissociating to form OH a. Ex: NaOH Na + OH (dissociates completely—strong) + b. Weak acids reversibly release/accept H (carbonic acid into bicarbonate and hydrogen) + H 2O 3 HCO +3H c. pH scale: in any aqueous solution at 25 degrees C, the product of hydrogen and hydroxide 14 ion concentrations is constant at 10 i. [H ] [OH] = 10 ; neutral when equal concentrations – each 10 with pH = 7 7 + ii. pH = log[H ]; pH declines as hydrogen ion concentration increases 1. based on 10s so solution of 3 pH is 1000 times more acidic than pH 6 d. Buffers: human blood and water both = 7 pH but adding acid to both results in greater change in water because presence of buffers in blood + i. A buffer is a substance that minimizes changes in H and OH concentrations in solution; it accepts H from solution when excess and donates when depleted (carbonic acid in human blood—see 2b) e. Acidification threat to water quality: carbon dioxide and fossil fuels react with water = more acidic; acid rain with pH = less than 5.2 i. CO d2ssolves in seawater, reacts with water to form carbonic acid—ocean acidification; CO l2vels in air bubbles trapped in ice thousands of years ago—pH dropping and decreased carbonate ion concentration means no calcium carbonate for reefs – loss of habitat Quiz: 1. Which of the following is a hydrophobic material? a. Wax 2. We can be sure that a mole of table sugar and a mole of vitamin C are equal in their a. Number of molecules 3. Measurements show that the pH of a lake is 4.0. what is they hydrogen ion concentration of the lake? a. 10 4 4. What is the hydroxide ion concentration of the lake described in question 3? a. 10 10 5. A slice of pizza has 500kcal. If we could burn the pizza and use all the heat to warm a 50L container of cold water, what would be the approximate increase in the temperature of the water. (Note: a liter of cold water weighs about 1kg.) a. 10 degrees C 8 Bio1441 Chapters 25 Review Chapter 4: Carbon and the Molecular Diversity of Life I. Organic chemistry is the study of carbon compounds. a. Wohler (1828) accidentally made organic compound (urea) b. Organic molecules and the origin of life on Earth: i. Miller (1953) helped bring abiotic synthesis of organic compounds into the context of evolution and concluded that complex organic molecules can spontaneously arise under conditions that were the same as early Earth ii. CHONPS – percentages of each relatively the same in all lifeforms II. Carbon atoms can form diverse molecules by bonding to four other atoms. a. The key to atom’s chemical characteristics is its electron configuration (determines kinds and number of bonds the atom can form) b. Formation of bonds with carbon: C has 6 electrons (2 in first shell and 4 in second); C forms double or single covalent bonds—branches (4); CO : 02C=0 c. Molecular diversity arising from variation in carbon skeletons: i. Hydrocarbons: organic molecules consisting of only carbon and hydrogen (covalent bonds) like oil and fats—long HC tails; undergo reactions that release much energy ii. Isomers: compounds that have the same numbers of atoms of the same elements but different structures i.e. different properties/functions 1. Structural isomers: differ in covalent arrangements of their atoms (ex: straight vs branched); number of possible isomers increases as carbon skeleton increases in size 2. Cistrans isomers: (geometric) carbons have covalent bonds to same atoms, but atoms differ in their spatial arrangements due to the inflexibility of their double bonds a. Single bonds allow atoms they join to rotate without changing compound (double bonds don’t); consider a simple molecule with 2 doublebonded carbons, each with H and X attached i. Both X’s on same side of double bond = cis isomer ii. X’s on opposite sides of double bond = trans isomer b. Subtle difference in chemical structure leads to huge difference in function (ex: chemical in retina) 3. Enantiomers: isomers that are mirror images of each other and that differ in shape due to presence of an asymmetrical carbon (one attached to 4 different atoms/molecules) III. A few chemical groups are key to molecular function: a. Estradiol (estrogen) vs testosterone steroids are organic 4 fused rings BUT different chemicals are attached, affecting molecular shape b. Functional group: chemical groups directly involved in chemical reactions (certain shape/charge) 9 Bio1441 Chapters 25 Review i. hydroxyl, carboxyl, amino, sulfhydryl, phosphate, carbonyl, methyl groups ii. first 6 are chemically reactive and all except sulfhydryl are hydrophilic and thus increase the solubility of organic compounds in water; methyl not reactive Hydroxyl (—OH) Polar due to electronegative oxygen; forms ethanol hydrogen bonds with water, help dissolve compounds like sugars; name: alcohol (specific name ends in ol) Carbonyl (>C=O) Sugars with ketone groups called ketoses; with aldehydes = aldoses; name: ketone (group inside HC chain) and aldehyde (group on end of HC acetone(k); propanol(a) chain) Carboxyl (—COOH) Acts as an acid because combined electronegativity of 2 adjacent oxygen atoms increase dissociation of H (increase number)—very polar bond; name: carboxylic acid/organic acid + Amino (—NH ) 2 Acts as base, can pick up H from surrounding group solution (water, in living organisms); name: anime (amino acids can act as acid OR base because carboxyl and amino groups both present) Sulfhydryl (—SH) Important for maintaining structure of proteins Cysteine (disulfide bridge); two –SH groups can react, (amino acid) forming cross link that helps stabilize protein structure; name: thiol Phosphate (—OPO ) 32 Contributes to neg charge (1 when positioned Glycerol phosphate inside chain of phosphates; 2 when at end); when attached, gives molecule ability to react with water, releasing energy; name: organic phosphate Methyl (—CH ) 3 Affects gene expression when on DNA and sex Methyl cytosine hormones; name: methylated compound IV. ATP: an important source of energy for cellular processes a. Adenosine triphosphate—organic molecule adenosine with 3 phosphates attached 10 Bio1441 Chapters 25 Review Quiz: 1. Organic chemistry is currently defined as a. The study of carbon compounds 2. Which chemical group is most likely to be responsible for an organic molecule behaving as a base? a. Amino 3. Which of the following hydrocarbons has a double bond in its carbon skeleton? a. C H (ration 1:2) 2 4 4. Choose the term that correctly describes the relationship between these two sugar molecules: (same numbers of atoms per element but different shape) a. Structural isomers 5. Identify the asymmetric carbon in this molecule a. One that is bonded to 4 different atoms 6. Which action could produce a carbonyl group? a. The replacement of the –OH of a carboxyl group with hydrogen Chapter 5: I. Macromolecules are polymers built from monomers a. Polymer: long molecule consisting of many similar or identical building blocks linked by covalent bonds—blocks = monomers b. The synthesis and breakdown of polymers: 11 Bio1441 Chapters 25 Review i. Process facilitated by enzymes (specialized proteins that speed up chemical reactions); monomers connected by reaction where covalent bond formed and water molecule lost—dehydration synthesis 1. One monomer provides hydroxyl group (—OH) and other (—H); broken down into monomers by hydrolysis (add water break bond) example: digestion c. The diversity of polymers: small differences between close relatives due to DNA and proteins; arrangement key; there are 20 amino acids but lined up in different order and length II. Carbohydrates serve as fuel and building material: CHO a. Sugars: have formulas with basis CH O (2x: glucose C H O —6a12ca6bonyl group CO and multiple hydroxyl groups OH); depending on location of carbonyl, sugar is either aldose or ketose—glucose is an aldose and fructose is ketose (isomer of glucose); 6carbon sugars called hexoses i. Classification by size of carbon skeleton (37 carbons long); spatial arrangement of their parts around asymmetric carbons (attached to 4 different groups/atoms) and in water, they form rings ii. Disaccharide: 2 monomers joined by glycosidic linkage (covalent bond formed by dehydration synthesis)—ex: maltose, sucrose (glucose and fructose) b. Polysaccharide: monosaccharides joined by glycosidic linkages (storage material hydrolyzed as needed to provide sugar for cells) – others serve as building material for protection (ex: cellulose in cell wall) i. Storage polysaccharides: plants starch (polymer of glucose molecules as granules within cell structures known as plastids—includes chloroplasts) has monomers joined by 14 linkages—some more branched than others 1. Animals glycogen stored in liver/muscle cells—depleted within a day ii. Structural polysaccharides: cellulose important to cell walls—most abundant organic compound on Earth; 2 different ring forms of glucose—1.) starch (αglucose) and 2.) cellulose (βglucose) that differ in placement of hydroxyl group (alpha below and beta above) 1. Parallel cellulose grouped in units called microfibrils 2. Enzymes to hydrolyze alpha and beta (not interchangeable)—animals cannot digest cellulose—it is eliminated in feces and stimulates mucus production; some microorganisms can (bacteria in cow gut) iii. Chitin: carb used by anthropods (insects, spiders, crustaceans, etc.) to build exoskeletons 1. Chemically linked proteins in insects or calcium carbonate encrusted proteins in crabs and other crustaceans 2. Used in cell walls of fungi; also has β linkages like cellulose but with nitrogen containing appendage III. Lipids are a diverse group of hydrophobic molecules: not big enough to be macromolecules 12 Bio1441 Chapters 25 Review a. Mix poorly with water between mostly hydrocarbon regions; important—fats, phospholipids, steroids b. Fats: constructed from glycerol (alcohol with 3 carbons that each have a carboxyl group) and fatty acids (long carbon skeleton with carboxyl group and HC chain) i. Each glycerol joined to 3 fatty acids by ester linkage (bond formed from dehydration reaction between hydroxyl group and carboxyl group = triacylglycerol ii. If NO DOUBLE BONDS between carbon atoms in chain, then as many H atoms as possible are bonded to the carbon skeleton = saturated fatty acid; straight chains iii. Unsaturated fatty acids: has 1 OR MORE DOUBLE BONDS with one fewer hydrogen atom in each doublebonded carbon—all double bonds in the HC chain are cis bonds (on same side)—causes kink in chain 1. Most animal fats are saturated (solid at room temperature, ex: lard) 2. Most plant/fish fats are unsaturated (liquid at room temperature, ex: fish oil) a. Hydrogenated vegetable oils synthetically solidified by adding H atoms 3. Too many saturated fats cause plaque buildup heart disease (atherosclerosis) a. Unsaturated fats with trans fats even worse iv. Major function of fats = energy storage (HC rich in energy) fat > starch (energy) 1. Humans store fat in adipose cells (also cushion vital organs and insulate the body) c. Phospholipids: 2 fatty acids attached to glycerol and third hydroxyl group of glycerol joined to a phosphate group (negative charge)—usually, additional small charged molecule like choline is linked to phosphate group i. Tails are hydrophobic; head is hydrophilic—when in water, immediately forms bilayer d. Steroids: lipids characterized by carbon skeleton consisting of 4 fused rings—distinguished by different chemical groups attached i. Cholesterol: common in animal cells and precursor from which other steroids (like sex hormones) are synthesized; in vertebrates, synthesized in liver and obtained from diet IV. Proteins include a diversity of structures, resulting in a wide range of functions: a. Account for more than 50% of dry mass of most cells; most structurally sophisticated structure known, 3D each unique; enzymes, defense, storage, transport, communication, movement, support, etc.; all made from 20 amino acids i. Most enzymes are proteins; speed up chemical reactions (catalysts) – selective (lock and key) b. Monomers are amino acids and make up all proteins when joined by peptide bonds – polypeptide; proteins made up of 1 or more polypeptides, folded; Amino acid monomers: organic molecule with both an amino group and a carboxyl group i. R group differs with each amino acid, but rest is constant 13 Bio1441 Chapters 25 Review ii. When in cell at proper pH, amino and carboxyl groups will be ionized iii. Physical and chemical properties of side chain determine shape which determine function of protein (ex: acidic amino acids have side chains with ionized carboxyl group; basic amino acids have ionized amino group in side chain and positive charge) 1. because charged, acidic and basic side chains are hydrophilic c. Polypeptides: when 2 amino acids are positioned so the carboxyl group of one is adjacent to the amino group of the other, dehydration synthesis joins them to form peptide bond (covalent) i. many make polypeptide backbone with different R groups attached—each with unique linear sequence ii. one end of the polypeptide chain has free amino group (Nterminus) while opposite end has free carboxyl group (Cterminus); there are more side chains so chemical nature is still determined by them d. Protein structure and function: functional protein isn’t just a polypeptide chain—it’s 1 or more polypeptides precisely twisted, folded, and coiled into unique shape—all determined by amino acid sequence i. Folding driven by formation of various bonds between parts of the chemical (spherical = globular protein) (long fibers = fibrous protein)—function depends on shape and ability to bind to molecules 1. Ex: antibody (protein) and flu virus; endorphins bind to specific receptor proteins e. Four levels of protein structure: i. Primary: linear chain of amino acids ii. Secondary: regions stabilized by hydrogen bonds between atoms of the polypeptide backbone—αhelix (hbonding every 4 amino acid) or βpleated sheet (spider webs); hydrogen bond between oxygen and hydrogen in backbone iii. Tertiary: 3D shape stabilized by interactions between side chains (Rgroups) 1. Hydrophobic interaction: amino acids with nonpolar side chains cluster together at core of protein (water excludes them)—kept together by van der Waals interactions a. Ionic bonds between positive and negative and hydrogen bonds between polar side chains also help structure 2. Disulfide bridge: very strong covalent bond (S—S) formed when 2 cysteine monomers (each with sulfhydryl groups on their side chains) are brought together by protein folding iv. Quaternary: association of 2 or more polypeptides (some proteins only); ex: collagen (40% of proteins in people); hemoglobin (oxygen binding protein of red blood cells) 14 Bio1441 Chapters 25 Review f. Sicklecell disease: a change in primary structure i. One substance in amino acid side chain changes shape ii. Physical and chemical conditions also affect shape (pH, salt concentration, temperature, etc.) 1. Could cause protein to unravel (denaturation) or become inhibited; ex: protein put in nonpolar solution g. Protein folding in the cell: go through many intermediate stages—track them i. Chaperonins: proteins crucial to folding process (assist but don’t determine final structure); keep new polypeptide segregated from disruptive chemical conditions in the cytoplasm when it folds spontaneously ii. Misfolding can cause diseases like cystic fibrosis, Alzheimer’s, Parkinson’s, and mad cow disease iii. A single protein has thousands of atoms; structure can be seen using xray crystallography, nuclear magnetic resonance spectroscopy (NMR), biometrics (to predict shape) V. Nucleic acids store, transmit, and help express hereditary information: a. DNA gives instructions for its own synthesis and direct RNA synthesis—gene expression b. Genetic material inherited from parents; each chromosome has 1 long DNA molecule (with 100s of genes) i. DNA directs synthesis of messenger mRNA which then interacts with proteins to make proteins on ribosomes c. The components of nucleic acids: macromolecules; polynucleotides consist of nucleotides (monomer)—made up of a fivecarbon sugar (pentose), a nitrogenous base, and one or more phosphate group(s) (portion without phosphate called a nucleoside); Nbases tend to take up hydrogen ions from the solution, acting as bases i. Pyrimidines: nitrogen bases with single ring (C, T, U); Purines: nitrogen bases with double ring (A, G) ii. Sugars: deoxyribose lacks oxygen atom on second carbon ring (DNA); (RNA) ribose iii. Phosphate: attached to 5’ end of nucleotide d. Nucleotide polymers: linkage of nucleotides from dehydration synthesis; adjacent nucleotides joined by phosphodiester linkage – phosphate links sugars of 2 nucleotides; results in repeating pattern of sugarphosphate backbone (one end has phosphate attached to 5’carbon end and other end has hydroxyl group on 3’ carbon i. Linear order of basis in a gene specifies the amino acid sequence—primary structure of protein e. The structures of DNA and RNA molecules: i. DNA has 2 polynucleotides (strands) wound to form a double helix and sugar phosphates run opposite 5’3’ and 3’5’ (antiparallel); sugarphosphate on outside and nitrogen base on interior—bases held together by hydrogen bond; AT bond, GC bond—complementary 15 Bio1441 Chapters 25 Review ii. RNA is single strand; basepairing can occur and causes shape iii. tRNA (transfer) brings amino acids to ribosome during polypeptide synthesis; is about 80 nucleotides long and functional shape results from basepairing nucleotides where complementary stretches of molecule can run antiparallel to each other 1. in RNA, thymine not present, A and U bond VI. Genomics and proteomics have transformed biological inquiry and applications: a. From genes to protein shape and function, mapping the biome Quiz: 1. Which of the following categories includes all others in the list? Mono, poly, starch, or carb? a. Carbohydrate 2. The enzyme amylase can break glycosidic linkages between glucose monomers only if the monomers are in the α form. a. Glycogen, starch, and amylopectin (cellulose is beta) 3. Which of the following is true of unsaturated fats? a. They have double bonds in the carbon chains of their fatty acids 4. The structural level of a protein lease affected by a disruption in hydrogen bonding is the a. Primary 5. Enzymes that break down DNA catalyze the hydrolysis of the covalent bonds that join nucleotides together. What would happen to DNA molecules treated with these enzymes? a. The phosphodiester linkages of the polynucleotide backbone would be broken 6. The molecular formula for glucose is C H6O 12W6at would be the molecular formula for a polymer made by linking ten glucose molecules together by dehydration synthesis? a. C H O 60 102 51 7. Which of the following pairs of base sequences could form a short stretch of a normal double helix of DNA? a. 5’ATGC3’ with 5’GCAT3’ Monosaccharides Glycosidic linkage Polysaccharide Carbohydrate Fatty acids Ester linkage Triacylglycerol Lipid Amino acids Peptide bonds Polypeptide Protein Nucleotides Phosphodiester/hydrogen bonds Polynucleotide Nucleic acids 16 Bio1441 Chapters 25 Review Good luck 17
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