AP Biology Unit 1 Study Guide PDF

Chapter 1.3 Notes Science = way of knowing - An approach to understanding the natural world Inquiry = the search for information and explanations of natural phenomena a. Making observations i. Help reveal valuable information about the natural world ii. Tools can be used...

Chapter 1.3 Notes Science = way of knowing - An approach to understanding the natural world Inquiry = the search for information and explanations of natural phenomena a. Making observations i. Help reveal valuable information about the natural world ii. Tools can be used: microscopes, precision thermometers, or high-speed cameras b. Formulating logical explanations (hypotheses) c. Testing the hypotheses i. May require a revision of original hypothesis or the formation of a completely new hypothesis → further testing Hypothesis = an explanation based on observations and assumptions that leads to a testable prediction 1. You can never test all possible explanations 2. You can never prove that a hypothesis is true a. You only prove that it is not false b. Ex. replacing the bulb fixed the lamp, supports the burnt-out bulb hypothesis but could also be because old bulb not screwed properly 3. You can only test natural hypotheses a. Supernatural and religious matters are outside bounds of science Experiment = a scientific test, often done under controlled conditions Data = recorded observations - The information that scientific inquiry is based on 1) Qualitative: descriptions of what is observed 2) Quantitative: expressed as numerical measurements - Can be organized into tables of graphs - Use statistics to analyze the data → significant or random fluctuations Inductive reasoning = a type of logic in which generalizations are derived from a large number of specific observations - Ex. “All organisms are made of cells” → from 2 centuries of microscope observation - Deductive reasoning = a type of logic in which specific results are predicted from a general premise - “if…then” format - Ex. “If the burnt-out bulb hypothesis is correct, then the lamp should work when you replace the bulb with a new one” Realistic Scientific Process Not linear Involves backtracking, repetition, and feedback 4 components: - exploration/discovery - gathering and interpreting data - societal benefits/outcomes - community analysis and feedback Controlled experiment = designed to compare an: a. Experimental group: a set of subjects that has/receives the specific factor being tested b. Control group: a set of subjects that lack the specific factor i. The groups should be identical except for the specific factor Variables = a factor that varies in the experiment a. Independent variable: factor manipulated by experimenter i. Graphed on x-axis b. Dependent variable: factor being measured that is affected by the independent variable i. Graphed on y-axis Theory = an explanation that is broader in scope than a hypothesis, generates new hypotheses, and is supported by a large body of evidence - Ex. “Evolutionary adaptations arise by natural selection” - Theory in everyday language = untested speculation Chapter 2.1 Notes Matter = anything takes up space and has mass, made up of elements Element = a substance that cannot be broken down to other substances by chemical reactions a. 90+ elements occurring in nature i. 20-25% are essential elements: needed for organism to live healthily and reproduce 1. Humans: 25 elements 2. Plants: 17 elements b. 4 elements: oxygen (O), carbon (C), hydrogen (H), and nitrogen (N) i. Makes up 96% of living matter ii. 7 secondary elements: Calcium (Ca), Phosphorus (P), Potassium (K), Sulfur (S), Sodium (Na), Chlorine (Cl), Magnesium (Mg) → remaining 4% c. Trace elements = an element indispensable for life but required in extremely minute amounts Compound = substance consisting of two or more different elements combined in a fixed ratio - Emergent properties: a compound has chemical and physical characteristics different from those of its constituent elements - Ex. sodium chloride (NaCl), sodium (pure sodium = metal) and chlorine (pure chlorine = poisonous gas) in a 1:1 ratio - Ex. water (H20), hydrogen and oxygen in a 2:1 ratio Chapter 2.2 Notes Atom = smallest unit of matter that still retains the properties of an element - Composed of subatomic particles, 3 main kinds: 1. Neutron: no electrical charge (neutral), found in nucleus of an atom 2. Proton: single positive electrical charge, found in the nucleus a. Neutron and protons have mass of 1.7 x 10-24 gram or ~1 dalton 3. Electron: single negative electrical charge, move around the nucleus a. Electron mass = 1/2000 of neutron/protons i. Ignore when computing total mass of an atom Atomic nucleus = an atom’s dense central core, consisting of proton and neutrons  - Attraction between negative electrons and positive protons keep electron in vicinity of the nucleus Atomic number = number of protons in the nucleus of an atom - Number is unique to each element, designated by a subscript - All atoms of same element have same number of protons - Atom is usually neutral in electrical charge so # of protons = # of electrons Mass number = total number of protons and neutrons in the nucleus of the atom - # of neutrons = mass number - atomic number Atomic mass = total mass of an atom, mostly concentrated in nucleus Isotope = different atomic forms of same element - Same number of proton but different number of neutrons → differ in atomic mass - Ex. 12C, 13C, and 14C → the first two are stable, the third is unstable - Radioactive isotope: unstable; nucleus decays spontaneously, giving off particles and energy During a chemical reaction, nuclei do not come close enough to react, only electrons are directly involved Energy = capacity to cause change/do work Potential energy = energy that matter possesses because of its location or structure - Matter has natural tendency to move toward lowest possible state of potential energy Electrons of an atom have potential energy from their distance from the nucleus a. Electrons (-) are attracted to the nucleus (+) b. Takes work to move an electron farther from the nucleus i. More distance from nucleus = greater potential energy c. Can only exist at certain energy levels, not in between Electron shells = energy level of electrons at a characteristic average distance from nucleus 1. First shell = closest to nucleus, lowest potential energy a. Holds 2 electrons 2. Second shell = middle, have more energy a. Holds 8 electrons 3. Third (valence) shell = farthest away, highest energy level a. Chemical behavior depends on the number of electrons in this shell i. Valence electrons = electrons in this outermost shell b. Atoms with full valence shells = inert/chemically unreactive Electrons can move from one shell to another by absorbing/losing energy equal to difference in potential energy a. Absorbs energy → moves to shell farther out i. Light energy can excite an electron to a higher energy level b. Loses energy → “falls back” to shell closer i. Lost energy usually released to environment as: 1. Visible light 2. UV radiation Chapter 2.3 Notes Chemical bonds = the attraction between two atoms, from the sharing of valence electrons or the presence of opposite charges on the atoms - Bonded atoms gain complete outer electron shells - Strongest: covalent and ionic bonds in dry ionic compounds - Ionic bonds in aqueous solutions are weak Covalent bond = sharing of a pair of valence electrons by two atoms, strongest in bio a. Double covalent bond = sharing of two pairs of valence electrons by two atoms b. Nonpolar covalent bond = electrons are shared equally between two atoms of similar electronegativity c. Polar covalent bond = covalent bond between atoms that differ in electronegativity i. Shared electrons pulled closer to more electronegative atom 1. Makes that atom slightly negative 2. Other atom slightly positive ii. Ex. H2O → oxygen very electronegative Molecule = 2+ atoms held together by covalent bonds Valence = bonding capacity/number of electrons required to complete the atom’s outermost shell Electronegativity = the attraction of a given atom for the electrons of a covalent bond - More electronegative → more strongly it pulls shared electrons toward it Ion = an atom or group of atoms that has gained or lost one or more electrons → acquires a charge a. Cation: positively charged ion i. Cation and anions attract each other due to opposites attract 1. Attraction = ionic bond *aqueous/dissolved in water  b. Anion: negatively charged ion i. Ex. Sodium atom (Na - 11 electrons) + Chlorine atom (Cl - 17 electrons) 1. Lone valence electron of sodium → chlorine atom a. Both valence shells complete 2. Electron transfer changes charges of each atom: a. Sodium: 11 protons, 10 electrons → charge of 1+ i. Cation b. Chlorine: 17 protons, 18 electrons → charge of 1- i. Anion 3. Form an ionic bond a. Ionic compound sodium chloride = table salt Ionic compound (salts) = compound resulting from the formation of an ionic bond Hydrogen bond = weak chemical bond, non covalent attraction between a hydrogen (slightly +) and electronegative atom - In living organisms, electronegative atom = oxygen or nitrogen Van der Waals interactions = weak attractions between molecules or parts of molecules that result from transient local partial charges - Molecules have ever-changing regions of positive and negative charge (electrons not always symmetrically distributed) - Enables all atoms and molecules to stick together Every molecule has a characteristic size and shape → key to its function - 2 atoms (like H2) are always linear Molecular shape is crucial → determines how biological molecules recognize and respond to one another with specificity - Often bind temporarily through weak interactions - Only works if shapes are complementary Chapter 2.4 Notes Chemical reactions = making and breaking of chemical bonds, which leads to changes in the composition of matter → matter is conserved a. Reactants: starting materials in chemical reactions i. Coefficient indicate number of molecules involved b. Product: material resulting from the chemical reaction i. All atoms in reactants are in the product as well  1. Reactions cannot create or destroy atoms, only rearrange/redistribute the electrons Photosynthesis - occurs within cells of green plant tissues a. 6CO2 + 6H2O —sunlight→ C6H12O6 + 6O2 All chemical reactions are theoretically reversible - Reversibility is indicated by ⇌ symbol Chemical equilibrium = (in a chemical reaction) the state in which the rate of forwards reaction = rate of reverse reaction - Reactions continue in both directions but - The relative concentrations of reactants and products do not change over time → no net effect - Does not mean amount of reactants = products - Means concentration have stabilized at a certain ratio Chapter 2.5 Notes Water is required for life ¾ of Earth’s surface = water ○ Only substance on our planet to exist in all 3 states of matter Cells = 70-95% water Water molecules are polar → unequal sharing of electrons ○ Oxygen: 2 partial negative charges ○ Hydrogen: each hydrogen has one partial positive Held together by a polar covalent bond ○ Hydrogen bond between many water molecules 4 emergent properties of water: Cohesive behavior ○ Water molecules held by multiple hydrogen bonds Cohesion due to hydrogen bonding Allows for transport of water and nutrients against gravity in plant Surface tension (from cohesion): how hard it is to break the surface of a liquid ○ Salt will disrupt surface tension Water molecules will surround polar salt molecules instead of cohesion to itself  Adhesion: clings to other substances (due to hydrogen bonding) Ability to moderate temperature ○ Water absorbs heat from warmer air → releases stored heat to cooler air Absorb/release large amount with only slight change in temperature Energy used to disrupt hydrogen bonds first, before warming the water ○ Heat absorbed when hydrogen bonds break ○ Heat is released when hydrogen bonds form ○ High specific heat (resists temperature change) Keeps temperature fluctuations within limits that permit life Helps maintain homeostasis More moderate along coastlines than inland Calorie = amount of heat it takes to raise the temperature of 1 gram of water by 1°C Also amount of heat released when it cools by 1°C ○ Evaporative Cooling Molecules moving fast enough can overcome their attractions Leave the liquid → enter air as a gas ○ Water has high heat of vaporization As liquid evaporates, remaining liquid surface cools ○ Helps with stabilization of bodies or water and human bodies Expansion upon freezing ○ Ice floats → less dense than liquid water Hydrogen bonds in ice are more “ordered” and locked/stable In water, hydrogen bonds constantly forming and breaking ○ In ice, molecules move too slowly to break hydrogen bonds Extra space in between water molecules when solid ○ Less dense/packed in Insulates the water, allowing underwater life to survive Versatility as a solvent ○ Solution = liquid homogeneous mixture ○ Solvent = dissolving agent Aqueous solution = water is the solvent Solute = substance that is dissolved ○ Hydrophilic = affinity for water ○ Hydrophobic = no affinity for water Lipids → mostly C and H, nonpolar ○ Water is polar, surrounds the compound → hydration shell Form ionic or hydrogen bonds with the molecule  Cannot disrupt covalent bonds Allows individual molecules to be separated from others Acids and Bases Hydrogen ion (H+) is transferred from one water molecule to another, leaves electron behind ○ Transferred as a hydrogen ion (one proton with 1+ charge) Water molecule left behind = hydroxide ion (OH-) Proton binds to other water molecule —> hydronium ion (H3O+) ○ H+ does not exist on its own in aqueous solution Always with water molecule in H3O+ form Reversible reaction Acids increase the H+ concentration in water ○ pH < 7 Weak acid= acids that reversibly release and accept back hydrogen ions Bases reduce the concentration of H+ ○ pH > 7 Each step on pH scale is 10x jump Buffers = minimize changes in concentrations of H+ and OH- in a solution ○ Most contain weak acid and its corresponding base, combine reversibly Internal pH of most living cell = ~7 Either accepting H+ ions or donating H+ ions as needed Ocean Acidification = process by which the pH of ocean is lowered when excess CO2 dissolves in seawater and forms carbonic acid (H2CO3) Reduces carbonate ion concentration ○ Carbonate ions required for calcification Production of calcium carbonate CaCO3 For reef-building corals and animals with shells 25% of human-generated CO2 is absorbed by ocean Chapter 3.1 Notes Carbon can form up to 4 bonds - Organic compounds = contain carbon a. 6 electrons: 2 in first shell, 4 in second shell i. Carbon usually forms single or double covalent bonds  1. When bonding, bond angle toward corners of a tetrahedron a. Two double bonds = four single bonds ii. Valence = # of unpaired electrons/# of possible covalent bonds 1. Hydrogen: 1 2. Oxygen: 2 3. Nitrogen: 3 4. Carbon: 4 Carbon dioxide = inorganic - Very simple and lack hydrogen Carbon chains form skeletons of most organic molecules a. Can be straight, branched, or arranged in closed rings i. Hydrocarbons = organic molecules of only carbon and hydrogen 1. Fats have long hydrocarbon tails a. Serve as stored fuel for plant seeds and animals 2. Also major components of petroleum (a fossil fuel) a. Both of these compounds are hydrophobic i. Nonpolar carbon-hydrogen linkages ii. Hydrocarbons can undergo reactions that release a lot of energy Isomers = compounds with same number of atoms of the same elements but different structures and properties 1. Structural isomers a. Differ in covalent arrangement of their atoms i. # of possible isomers increases as carbon skeletons increases in size b. May also differ in location of double bonds 2. Cis-trans isomers a. Carbons have covalent bonds to same atoms i. Atoms differ in spatial arrangements (due to inflexibility of double bonds) 1. Single bonds can rotate freely b. Cis isomer: Xs on same side of double bonds c. Trans isomer: Xs of opposite sides i. Shape → function, can have large effect on biological activities 3. Enantiomers a. Mirror images of each other but differ in shape due to a asymmetric carbon i. Four atom groups can be arranged around the carbon in two ways 1. “Left hand” and “right hand” versions ii. Usually only one isomer is biologically active 1. Only one has the shape to bind to specific molecules  b. Important in pharmaceutical industry i. Two enantiomers may not be equally effective 1. Can also have different effects a. Emergent properties that depend on specific arrangement of atoms Functional groups = specific configuration of atoms usually attached to carbon skeletons of organic molecules and involved in chemical reactions → 7 chemical groups 1. Hydroxyl (—OH) a. Compound name: Alcohol i. Hydrophilic ii. Makes compounds polar iii. Allows for dehydration reactions 2. Carbonyl (>C=O) a. Compound name: Ketone (if in carbon skeleton) or Aldehyde (if at end of carbon skeleton) i. Hydrophilic ii. Polar iii. Volatile, often reactive 3. Carboxyl (—COOH) a. Compound name: Carboxylic acid, or organic acid i. Hydrophilic ii. Acidic, found on ends of amino acids 1. Makes it low pH 4. Amino ( —NH2) a. Compound name: Amine i. Hydrophilic ii. Basic, found on other ends of amino acids 5. Sulfhydryl (—SH) a. Compound name: Thiol i. Hydrophobic ii. Stabilizes tertiary structure of proteins 6. Phosphate (—OPO32-) a. Compound name: Organic phosphate i. Hydrophilic ii. Energetic and polar 1. Found in ATP, nucleic acids, and phospholipids b. Important in energy transfer i. More complicated organic phosphate = ATP 1. Adenosine attached to 3 phosphate groups ii. ATP —reacts with H2O→ ADP + inorganic phosphate + energy  7. Methyl (—CH3) a. Compound name: Methylated compound i. Not reactive, serves as recognizable tag ii. Hydrophilic iii. Energy storing and nonpolar Chapter 3.2 Notes Macromolecules = large carbohydrates, proteins, and nucleic acids joined by smaller molecules → usually formed by a dehydration reaction - Not lipids, are comparatively small, not built of monomers Polymer = long molecule consisting of many similar/identical building blocks - Smaller building blocks = monomers - Connected by covalent bonds Enzymes = specialized macromolecules that speed up chemical reactions, typically a protein a. Condensation reaction = two molecules are covalently bonded together, with loss of small molecule i. Dehydration reaction = specifically water molecule is lost 1. Ex. carbohydrate and protein polymers a. (—OH) + (—H) → polymer is lengthened b. Hydrolysis = breaks bonds between two molecules by addition of water i. Disassembles polymers to monomers 1. Ex. digestion a. Enzymes attack polymers → speed up hydrolysis Chapter 3.3 Notes Carbohydrates = a sugar (monosaccharide) or one of its dimers (disaccharides) or polymers (polysaccharides) a. Monosaccharides: simple sugars, monomers for more complex carbohydrates i. Generally some multiple of the unit CH2O 1. Carbon skeleton ranges from 3-7 carbons long a. Hexoses: have 6 carbons  i. Ex. glucose, fructose b. Trioses: 3 carbons c. Pentoses: 5 carbon 2. Carbonyl group can be on end or on interior of a linear sugar a. Makes monosaccharides either aldehydes or ketones ii. In aqueous solutions, glucose molecules form rings 1. Rings = most stable form for sugars under physiological conditions b. Disaccharides: double sugars i. Two monosaccharides joined by a covalent (glycosidic linkage) bond 1. Through dehydration reaction → requires energy a. Ex. glucose + fructose = sucrose c. Polysaccharides: polymers composed of many sugar building blocks i. Joined together by dehydration reactions 1. Can be few hundred to few thousand monosaccharides long a. Through glycosidic linkages ii. Can be storage material 1. Starch a. Storage polysaccharide in plants i. Entirely of glucose monomers joined by glycosidic linkages ii. Stored as granules b. Hydrolyzed when needed → sugar monomers for cells 2. Glycogen a. Storage polysaccharide in animals i. Releases glucose ii. In humans, glycogen stores are depleted in ~1 day b. Extensively coils, extremely branched i. More free ends to breakdown ii. More compact → easier to store iii. Can be building material 1. Cellulose a. Component of tough walls in plants i. Most abundant organic compound 1. Polymer of glucose with 1-4 b. Never branched, molecule is straight, flipped monomers i. Glucose monomers flip upside down every molecule 1. Makes it completely straight ii. Shape gives it strength 1. Hydrogen bonds between strands a. Strand grouped into units called microfibrils  c. Almost all animals and humans cannot digest cellulose i. Helps with digestion 1. No nutrition ii. Cow has cellulose-digesting prokaryotes in its gut 1. Hydrolyze cellulose → glucose → other nutrients 2. Chitin a. Carbohydrate used by arthropods to build exoskeletons i. Also found in fungi for cell walls b. Similar to cellulose i. Glucose monomer has a nitrogen-containing attachment iv. Structure and function determined by sugar monomers and position of glycosidic linkages Chapter 3.4 Notes Lipids → scared of water Fat = glycerol molecule + three fatty acids - Fatty acid has long carbon skeleton (~16-18 carbons) 1. Saturated a. Saturated with max number of hydrogens b. No double bond, flexible so can pack tightly i. Solid at room temperature c. Ex. animal fats, butter 2. Unsaturated a. 1+ double bond b. Causes it to kink i. Almost all are cis bonds ii. Liquid at room temperature 1. Molecules are more spread apart, not as compact, can’t pack together tightly c. If trans: liquid → solid i. Trans fats can contribute to coronary heart disease d. Ex. plant and fish fats, oil Covalent bond = ester linkage in lipids 3. Phospholipids  a. Two fatty acid tails (hydrophobic) + polar head (hydrophilic) i. Make up cell membranes 1. Phospholipid bilayer → naturally form this structure 4. Steroids a. Carbon skeleton of four fused rings i. Can be used as cell signaling b. Cholesterol i. Component in animal cell membranes 1. Is essential, but high levels → heart disease ii. Precursor from which other steroids are synthesized 1. Testosterone → muscle-building, deeper voice, adam’s apple a. Both produced in men and women, but in different levels Major function of fats = energy storage - Compact way for animals to carry their energy stores Cis → same side of double bonds Trans → different side of double bonds Chapter 3.5 Notes Protein = molecule of 1+ polypeptides fold/coiled into a specific 3D structure a. Account for 50+% of dry mass of most cells b. All constructed of same set of 20 amino acids i. Linked in unbranched polymers 1. Bond = peptide bond ii. Polymer: polypeptide Play an important role in many bodily functions: speed up chemical reactions, defense, storage, transport, cellular communication, movement, or structural support. Enzymatic Defensive Storage Transport Acceleration of Protection against Storage of amino Transport of chemical reactions disease acids substances Ex. digestive Ex. antibodies help Ex. casein protein of Ex. hemoglobin enzymes catalyze destroy/inactivate milk stores for baby transports oxygen  hydrolysis in food viruses and bacteria mammals from lungs Hormonal Receptor Contractile/Motor Structural Coordination of Response of cell to Movement of Supports tissues organism’s activities chemical stimuli various structure and other Ex. insulin causes Ex. receptors of Ex. actin and Ex. keratin protein tissues to take up nerve cell detect myosin proteins of hair, horn, glucose signaling molecules contract muscles feather, etc Enzymatic protein regulate metabolism - Act as catalysts: chemical agents that selectively speed up chemical reactions without being consumed/reactant of the reaction Amino acid = organic molecule with amino, carboxyl group, and R group a. Center carbon atom = alpha carbon b. R group (side chain) → determine unique characteristics of the amino acid i. Differs with each amino acid ii. Determines how polypeptide folds, final shape, chemical properties 1. Folding driven by bonds that form between parts of the chain Grouped according to R group 1. Nonpolar side chain → hydrophobic 2. Polar → hydrophilic 3. Acidic amino acids → side chains are negative 4. Basic amino acids → side chains are positive a. Acidic and basic side chains are hydrophilic i. Because they are charged Polypeptides can become linked through dehydration reactions repeated over and over a. Repeating N-C-C pattern → polypeptide backbone i. Range in length from few to 1,000+ amino acids b. One end has free amino group, other end has free carboxyl group 4 levels of protein structure: 1. Primary a. Sequence of amino acids i. Dictated by inherited genetic information, not random 2. Secondary  a. Hydrogen bonds between the atoms of polypeptide backbone i. Between negative oxygens and positive hydrogen + nitrogens ii. Have 2 kinds of secondary structures: 1. Helix = coil held together by hydrogen bonds every 4th amino acid 2. Pleated sheet = two or more segments of polypeptide chain lying side by side, hydrogen bonds between 2 parallel segment a. Shown as folded/flat arrow i. Pointing towards the carboxyl end 3. Tertiary a. Overall shape of polypeptide from R group interactions i. Many types of interactions that contribute to tertiary structure: 1. Hydrophobic interaction = amino acids with hydrophobic (nonpolar) R groups cluster on inside/core of protein a. Out of contact with water i. Held by van der Waals interactions on inside 2. Hydrogen bonds between polar side chains 3. Ionic bonds between positively and negatively charged R groups a. Both 2 and 3 are weak in aqueous environment i. But cumulative effect helps give the shape 4. Covalent bonds = disulfide bridges a. Form between two cysteine monomers with sulfhydryl groups (—SH) on R group i. Sulfur of one bonds to sulfur on second 4. Quaternary a. When a protein consists of 2+ polypeptide chains → one macromolecule i. Aggregation of polypeptide subunits 1. Ex. collagen: 3 identical helical polypeptides a. Accounts for 40% of protein in body 2. Ex. hemoglobin: oxygen-binding protein of red blood cells a. Globular protein, 4 polypeptide subunits i. 2 of helix and 2 of pleated sheet Sickle-cell disease a. Inherited blood disorder → caused by substitution of glutamic acid with valine (6) i. Abnormal red blood cells 1. Aggregate into fibers and deforming into sickle shape a. Angular cells clog tiny blood vessels → block blood flow Denaturation = process in which protein loses its native shape due to disruption of weak chemical bonds and interactions → become inactive a. Can occur if aspects of environment is changed: i. pH, salt concentration, temperature, etc b. Misshapen → biologically inactive