BIO 201: Biochemistry
BIO 201: Biochemistry BIO 201
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This 18 page Class Notes was uploaded by ASUNursing19 on Thursday February 18, 2016. The Class Notes belongs to BIO 201 at Arizona State University taught by Dr. Penkrot in Winter 2016. Since its upload, it has received 32 views. For similar materials see Human Anatomy/Physiology I in Biology at Arizona State University.
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Date Created: 02/18/16
Chemistry and Physiological Reactions Body is made up of many chemicals Chemistry underlies all physiological reactions: o Movement, digestion, pumping of heart, nervous system Chemistry can be broken down into: o Basic chemistry o Biochemistry Part 1 Basic Chemistry Matter MATTER is anything that has mass and occupies space o Matter can be measured, seen, smelled, and/or felt o Weight is mass plus the effects of gravity States of matter o Matter can exist in three possible states: Solid: definite shape and volume Liquid : changeable shape; definite volume Gas : changeable shape and volume Liquids and gases are both often referred to as "fluids" Energy ENERGY is the capacity to do work or put matter into motion Energy does not have mass, nor does it take up space The greater the work done, the more energy it uses up o Body uses energy that is stored in chemical bonds Kinetic versus potential energy o Energy exists in two possible forms: KINETIC ENERGY : energy in action POTENTIAL ENERGY: stored (inactive) energy o Energy can be transformed from potential to kinetic energy Stored energy can be released, resulting in action Forms of energy o Chemical energy: Stored in bonds of chemical substances o Electrical energy: Results from movement of charged particles o Mechanical energy: Directly involved in moving matter o Radiant or electromagnetic energy : Travels in waves (example: heat, visible light, ultraviolet light, and X rays) Energy form conversions o Energy may be converted from one form to another Example: turning on a lamp converts electrical energy to light energy o Energy conversion is inefficient Some energy is "lost" as heat, which can be partly unusable energy o Usually, converting from type of energy to another energy can result in loss; it is insufficient 2.2 Atoms and Elements All matter is composed of lements o Elements are substances that cannot be broken down into simpler substances by ordinary chemical methods Four elements make up 96% of body: o Carbon, oxygen*(2/3 of body weight is oxygen), hydrogen, and nitrogen o 9 elements make up 3.9% of the body o 11 elements may up <0.01% Periodic table lists all known elements All elements are made up of atoms , which are: o Unique building blocks for each element o Smallest particles of an element with properties of that element o What give each element its particular physical & chemical properties Atomic symbol o One or twoletter chemical shorthand for each element Example: "O" for oxygen, "C" for carbon Some symbols come from Latin names: "Na" (natrium) is sodium; "K" (kalium) is potassium Metals Little Tip If something serves no biological purpose in the body, it is probably not a good idea to consume large amounts of it! o If your body doesn't use it already, there's probably a reason Heavy intake of nonnutrients can be harmful o Example: Argyria Nothing is completely harmless if consumed in large amounts Argyria Silver is not a necessary micronutrient Chromium Chromium picolinate o Cr deficiency is extremely rare o Used for weight loss, to treat depression o More likely to cause DNA damage and mutations than to improve health Cancer Protons, Neutrons, and Electrons Protons determine atomic number Neutrons determine nuclear mass and stability Electrons can be added or lost easily o Basis of chemical properties o Ability to gain, lose, or share electrons determines an element's reactivity Structure of Atoms Atoms are composed of three subatomic particles: o Protons Carry a positive charge (+) Weigh an arbitrary 1 atomic mass unit (1 amu) o Neutrons Have no electrical charge (0) Also weigh 1 amu o Electrons Carry a negative charge () Are so tiny they have virtually no weight (0 amu) Number of positive protons is balanced by number of negative electrons, so atoms are electrically neutral Protons and neutrons are found in a centrally located ucleus ; electrons orbit around the nucleus Chemists devise models of how subatomic particles are put together o Planetary model o Orbital model Planetary model : simplified and outdated because it incorrectly depicts electrons in orbits fixed circular paths o Still useful for illustrations Orbital model : current model used that depicts rbitals, probable regions where an electron is most likely to be located (rather than fixed orbits) o Shading in regions of greatest electrons density results in an electron cloud around nucleus o Useful for predicting chemical behavior of atoms Identifying Elements Atomic number o Equals number of protons in nucleus o Written as subscript to left of atomic symbol Example: 3i Mass number o Equals total number of protons and neutrons in nucleus Total mass of atom o Written as superscript to left of atomic symbol 7 Example Li Isotopes o Structural variations of same element o Atoms contain same number of protons but differ in the number of neutrons they contain Atomic numbers are same, but mass numbers different Atomic weight o Average of mass numbers of all isotope forms of an atom Radioisotopes Radioisotopes are isotopes that decompose to more stable forms o Atom loses various subatomic particles Sometimes loss results in an isotope becoming a different element o As isotope decays, subatomic particles that are being given off release some energy This energy is referred to as radioactivity Can be detected and measured with scanners Radioisotopes are a valuable tool for biological research and medicine o Share same chemistry as their stable isotopes so will be taken up by body Can then be used for diagnosis of disease All radioactivity can damage living tissue o Some types can be used to destroy localized cancer o Some types cause cancer Radon from uranium decay causes lung cancer Ionizing Radiation Radiation transfers energy Ionizing radiation alters the electron cloud of atoms Classification of Ionizing Radiation Three major types: Alpha: α o Helium nucleus Beta: β o Electron Gamma: γ o Photon Other types exist The smaller the particle, the higher the energy… Medical Imaging Ionizing radiation o XRay o CT Scan o PET Scan Magnetic radiation o MRI 2.3 Combining Matter Molecules vs. Compounds Most atoms chemically combine with other atoms to form molecules and compounds o Molecule: general term for 2 or more atoms bonded together o Compound : specific molecule that has 2 or more different kinds of atoms bonded together Example: C 6 O12 6 Molecules with only one type of atom (H or O ) are just callmolecules 2 2 Mixtures Both molecules and compounds are different from mixtures Most matter exists as ixtures : two or more components that are physically intermixed Three basic types of mixtures: o Solutions o Colloids o Suspensions Solutions o Are homogeneous mixtures, meaning particles are evenly distributed throughout Solvent: substance present in greatest amount (does the dissolving) Usually a liquid, such as water Solute(s): substance dissolved in solvent Present in smaller amounts Example: blood sugar glucose is solute, and blood (plasma) is solvent True solutions are usually transparent Example: air (gas solution), salt solution, sugar solution Most solutions in body are true solutions of gases, liquids, or solids dissolved in water Concentration of true solutions o Three common ways to express concentrations: 1. Percent of solute in total solution How many parts of solute are in 100 total parts of solution Solvent is usually water Example: 10 parts salt to 90 parts water is a 10% salt solution 2. Milligrams per deciliter (mg/dl) Deciliter equals 1/100th of a liter Example: normal fasting blood glucose levels are around 80 mg/dl 3. Molarity (M) is number of moles of solute per liter of solvent (water) 1 mole of a compound is equal to its molecular weight (sum of atomic weights) in grams Example: glucose (C H6O 12h6s a molecular weight of 180.12 amu, so 180.12 grams of glucose added to enough H O to2make 1 liter is a 1 M solution of glucose 1 mole of any substance always contains 6.02 x 1023 molecules of that substance (Avogadro's number ) Molarities in the body are so small (can be 0.0001 M), they are expressed in millimoles (mM) so 1000 mM = 1M Colloids o Also known as emulsions ; areheterogeneous mixtures, meaning that particles are not evenly distributed throughout mixture Can see large solute particles in solution, but these do not settle out, but may separate Gives solution a cloudy or milky look o Some undergo solgel (solution to gel) transformations Example: JellO goes from liquid to gel Cytosol of cell is also a solgel type solution Suspensions o Heterogeneous mixtures that contain large, visible solutes that do settle out o Example: mixture of water and sand o Blood is considered a suspension because if left in a tube, the blood cells will settle out Difference Between Mixtures and Compounds Three main differences: o Unlike compounds, mixtures do not involved chemical bonding between components o Mixtures can be separated by physical means, such as straining or filtering; compounds can be separated only by breaking their chemical bonds o Mixtures can be heterogeneous or homogeneous; compounds are only homogeneous 2.4 Chemical Bonds Chemical bonds are "energy relationships" between electrons of reacting atoms o Chemical bonds are not actual physical structures Electrons are the subatomic particles that are involved in all chemical reactions o They determine whether a chemical reaction will take place and if so, what type of chemical bond is formed Role of Electrons in Chemical Bonding Electrons can occupy areas around nucleus called electron shells o Each shell contains electrons that have a certain amount of kinetic and potential energy, so shells are also referred to as rgy levels o Depending of its size, an atom can have up to 7 electron shells o Shells can hold only a specific number of electrons; the shell closest to nucleus is filled first Shell 1 can hold only 2 electrons Shell 2 holds a maximum of 8 electrons Shell 3 holds a maximum of 18 electrons Outermost electron shell is called alence shell o Electrons in valence shell have the most potential energy because they are farthest from nucleus o These are electron that are involved in chemical reactions Octet Rule (rule of eights) o Atoms "desire" 8 electrons in their valence shell Exceptions: smaller atoms (examples: H and He) want only 2 electrons in shell 1 o Desire to have 8 electrons is driving force behind chemical reactions Noble gases already have full 8 valence electrons (or 2 for He) so are not chemically reactive o Most atoms do not have full valence shells Atoms will gain, lose, or share electrons (form bonds) with other atoms to achieve stability to 8 electrons in valence shell Types of Chemical Bonds Three major types of chemical bonds o Ionic bonds o Covalent bonds o Hydrogen bonds (not real chemical bond) Ionic bonds o Ions are atoms that have gained or lost electrons and become charged Number of protons does not equal number of electrons o Ionic bonds involved thetransfer of valence shell electrons from one atom to another, resulting in ions One becomes an anion (negative charge) Atom that gained one or more electrons One becomes a cation (positive charge) Atom that lost one or more electrons o Attraction of opposite charges results in an ionic bond o Most ionic compounds are salts When dry, salts form crystals instead of individual molecules Example: NaCl (sodium chloride = table salt) Covalent bonds o Covalent bonds are formed by sharing of two or more valence shell electrons between two atoms Sharing of 2 electrons results in a single bond Sharing of 4 electrons is a double bond Sharing of 6 electrons is a triple bond o Allows each atom to fill its valence shell at least part of the time o Two types of covalent bonds: Polar and nonpolar covalent bonds o Nonpolar covalent bonds Equal sharing of electrons between atoms Results inelectrically balanced nonpolar molecules such as CO 2 o Polar covalent bonds Unequal sharing of electrons between 2 atoms Results inelectrically polar molecules Atoms have different electronattracting abilities, leading to unequal sharing Atoms with greater electronattracting ability are ctronegative, and those with less are ectropositive H2O is a polar molecule Oxygen is more electronegative, so it exerts a greater pull on shared electrons, giving it a partial negative charge and giving H a partial positive charge Having two different charges is referred to as ole Hydrogen bonds o Attractive force between electropositive hydrogen of one molecule hydrogen of one molecule and an electronegative atom of another molecule Not a true chemical bond, more of a weak magnetic attraction between molecules o Common between dipoles such as water What makes water liquid (plus other unusual attributes) o Also act as intramolecular bonds, holding a large molecule in a threedimensional shape 2.5 Chemical Reactions Chemical reactions occur when chemical bonds are formed, rearranged, or broken These reactions can be written in symbolic forms called chemical equations Chemical equations contain: o Reactants: substances entering into reaction together o Product(s) : resulting chemical end products o Amounts of reactants and products are shown in balanced equations Compounds are represented as molecular formulas o Example H O 2r C H 6 12 6 or 2H 4 o In chemical equations, subscripts indicate how many atoms are joined by bonds, whereas prefix means number of unjoined atoms (example: 4H) Types of Chemical Reactions Three main types of chemical reactions: 1. SYNTHESIS (combination) involve atoms or molecules combining to form larger, more complex molecule Used in anabolic (= building) process A + B > AB 2. DECOMPOSITION reactions involve breakdown of a molecule into smaller molecules or its constituent atoms (reverse of synthesis reactions) Involve catabolic ( = bond breaking) reactions AB > A + B 3. EXCHANGE reactions, also calleddisplacement reactions, involve both synthesis and decomposition Bonds are both made and broken AB + C > AC + B AB + CD > AD + CB In living systems, these reactions are also referred to as uctionoxidation or redox reactions o Atoms are reduced when they gain electrons and oxidized when they lose electrons o Example: C H6O 12 6O >26CO + 62 O + A2P In this example, glucose is oxidized, and oxygen molecule is reduced Energy Flow in Chemical Reactions All chemical reactions are either exergonic or endergonic o Exergonic reactions result in a net release of energy (give off energy) Products have less potential energy than reactants Catabolic and oxidative reactions o Endergonic reactions result in a netabsorption of energy (use up energy) Products have more potential energy than reactants Anabolic reactions Reversibility of Chemical Reactions All chemical reactions are theoretically reversible o A + B < > AB Chemical equilibrium occurs if neither a forward nor a reverse reaction is dominant Many biological reactions are not very reversible o Energy requirements to go backward are too high, or products have been removed Rate of Chemical Reactions The speed of chemical reactions can be affected by: o Temperature : increased temperatures usually increase rate of reaction o Concentration of reactants: increased concentrations usually increase rate o Particle size smaller particles usually increase rate Catalysts increase the rate of reaction without being chemically changed or becoming part of the product o Enzymes are biological catalysts Part 2 Biochemistry Biochemistry is the study of chemical composition and reactions of living matter All chemicals either organic or inorganic o Inorganic compounds Water s, ts , and many acids and bases Do not contain carbon o Organic compounds Carbohydrates, fats, proteins, and nucleic acids Contain carbon, are usually large, and are covalently bonded Both equally essential for life 2.6 Inorganic Compounds Water Most abundant inorganic compound o Accounts for 60%80% of the volume of living cells Most important inorganic compound because of its properties o High heat capacity o High heat of vaporization o *Polar* solvent properties o Reactivity o Cushioning High heat capacity o Ability to absorb and release heat with little temperature change o Prevents sudden changes in temperature High heat of vaporization o Evaporation requires large amounts of heat o Useful cooling mechanism Polar Solvent Properties o Dissolves and dissociates ionic substances o Forms hydration (water) layers around large charged molecules Examples: proteins o Body's major transport medium o Hydrogen bonding Reactivity o Necessary part of hydrolysis and dehydration synthesis reactions Cushioning o Protects certain organs form physical trauma Example: cerebrospinal fluid cushions nervous system organs Salts Salts are ionic compounds that dissociate into separate ions in water o Separate into cations (positively charged molecules) and anions (negatively charged molecules) Not including H and OH ions o All ions are called electrolytes because they can conduct electrical currents in solution o Ions play specialized roles in body functions Example: sodium, potassium, calcium, and iron o Ionic balance is vital for homeostasis o Common salts in body NaCl, CaCO , K3l, calcium phosphates (these salts in particular break apart easily) Acids and Bases Acids and bases are both electrolytes o Ionize and dissociate in water ACIDS Are proton donors : they release hydrogen ions (H+), bare protons (have no electrons) in solution o Example: HCl > H Cl Important acids o HCl (hydrochloric acid), HC H 2 3(a2etic acid, abbreviated HAc), and H CO 2 3 (carbonic acid) BASES Are proton acceptors : they pick up H ions in solution + o Example: NaOH > Na + OH When a base dissolves in solution, it releases a hydroxyl ion (OH) o Bicarbonate ion (HCO ) 3nd ammonia (NH ) 3 pH: ACIDBASE CONCENTRATION pH scale is measurement of concentration of hydrogen ions [H ] in a solution The more hydrogen ions in a solu+ion, the more acidic that solution is pH is negative logarithm of [H ] in moles per liter than ranges from 014 pH scale is logarithmic, so each pH unit represents a 10fold difference o Example: a pH 5 solution is 10 times more acidic than a pH 6 solution Acidic solutions have high [H ] but low pH o Acidic pH range is 06.99 + Neutral solutions have equal numbers of H and OH ions o All neutral solutions are pH 7 o Pure water is pH neutral + 7 pH of pure water = pH 7: [H ] = 10 m (perfectly neutral) Alkaline (basic) solutions have low [H ] but high pH o Alkaline pH range is 7.0114 NEUTRALIZATION Neutralization reaction : acids and bases are mixed together o Displacement reactions occur, forming water and a salt NaOH + HCl > NaCl + H O 2 BUFFERS Acidity involves only free H+ in solution, not H bound to anions Buffers resist abrupt and large swings in pH o Can release hydrogen ions if pH rises o Can bind hydrogen ions if pH falls Convert strong acids or bases (completely dissociated) into eak ones (slightly dissociated) o Carbonic acidbicarbonate system (important buffer system of blood) 2.7 Organic Compounds: Synthesis and Hydrolysis Organic molecules contain carbon o Exceptions: CO 2and CO, which are inorganic Carbon is electroneutral o Shares electrons; never gains or loses them o Forms four covalent bonds with other elements o Carbon is essential to living systems Major organic compounds: carbohydrates l, ids p, teins , and ucleic acids Many are polymers o Chains of similar units called nomers (building blocks) Synthesized by dehydration synthesis Broken down by hydrolysis reactions 2.8 Carbohydrates Carbohydrates include sugars and starches Contain C, H, and O o Hydrogen and oxygen are 2:1 ratio Three classes: o Monosaccharides: one single sugar Monomers: smallest unit of carbohydrate o Disaccharides: two sugars o Polysaccharides : many sugars Polymers are made up of monomers of monosaccharides Monosaccharides o Simple sugar containing three to seven carbon atoms o (CH 2) n general formula n = number of carbon atoms o Monomers of carbohydrates o Important monosaccharides Pentose sugars Ribose and deoxyribose Hexose sugars Glucose (blood sugars) Disaccharides o Double sugars o Too large to pass through cell membranes o Important disaccharides Sucrose, maltose, lactose o Formed by dehydration synthesis of two monosaccharides Glucose + fructose > sucrose + water Polysaccharides o Polymers of monosaccharides Formed by dehydration synthesis of many monomers to make long chains of simple sugars o Important polysaccharides Starch: carbohydrate storage form used by plants Glycogen: carbohydrate storage form used by animals o Not very soluble 2.9 Lipids Contain C, H, O, but less than in carbohydrates, and sometimes contain P Insoluble in water Main types: o Triglycerides or neutral fats o Phospholipids o Steroids o Eicosanoids Triglycerides or Neutral Fats o Called fats when solid and oils when liquid o Composed of three fatty acids bonded to a glycerol molecule o Main functions Energy storage Insulation Protection o Stored in fat cells (adipocytes) Triglycerides can be constructed of: o Saturated fatty acids All carbon are linked viasingle covalent bonds resulting in a molecule with the maximum number of H atoms (saturated with H) Solid at room temperature (Example: animal fats, butter) o Unsaturated fatty acids One ore more carbons are linked viadouble bonds , resulting in reduced H atoms (unsaturated) Liquid at room temperature (Example: plant oils, such as olive oil) Trans fats: modified oils; unhealthy Omega3 fatty acids: "heart healthy" Phospholipids o Modified triglycerides Glycerol and two fatty acids plus a phosphoruscontaining group o "Head" and "tail" regions have different properties Head is a polar region and is attracted to water (hydrophilic) Tails are nonpolar and are repelled by water (hydrophobic) o Important in cell membrane structure Steroids o Consist of four interlocking ring structures o Common steroids: cholesterol, vitamin D, steroid hormones (i.e.: sex hormones), and bile salts o Most important steroid is cholesterol Is building block for vitamin D, steroid synthesis, and bile salt synthesis Important in cell plasma membrane structure Eicosanoids o Many different ones o Derived from a fatty acid (arachidonic acid) found in cell membranes o Most important eicosanoids are prostaglandins Play a role in blood clotting, control of blood pressure, inflammation*, and labor contractions 2.10 Proteins Comprise 2030% of cell mass Have most varied functions of any molecules o Structural, chemical (enzymes), contraction (muscles) Contain C, H, O, N, and sometimes S and P Polymers of amino acid monomers held together by peptide bonds Shape and function due to four tructural levels Amino Acids and Peptide Bonds All proteins are made from 20 types of amino acids o Amino acids are the monomers; proteins or polypeptides are the polymers o Joined by covalent bonds called peptide bonds o Contain both an amine group and acid group o Can act as either acid or base o Differ by which of 20 different "R groups" is present Structural Levels of Proteins Four levels of protein structure determine shape and function 1. Primary: linear sequence of amino acids (order) 2. Secondary: how primary amino acids interact with each other Alpha ( α helix coils resemble a spring Beta (β) pleated sheets resemble accordion ribbons 3. Tertiary: how secondary structures interact 4. Quaternary: how 2 or more different polypeptides interact with each other Protein Denaturation Denaturation : globular proteins unfold and lose their functional 3D shape o Long, fibrous proteins are more stable o Active sitesbecome deactivated Can be caused by decreased pH (increased acidity) or increased temperature Sometimes reversible if normal conditions restored Irreversible if changes are extreme o Example: cannot undo cooking an egg Enzymes and Enzyme Activity Enzymes : globular proteins that act as biological catalysts o Catalysts regulate and increase speed of chemical reactions without getting used up in the process o Lower the energy needed to initiate a chemical reaction Leads to an increase in the speed of a reaction Allows for millions of reactions per minute Characteristics of enzymes o Most functional enzymes, referred to as oloenzymes , consist of two parts Apoenzyme (protein portion) Cofactor (metal ion) or oenzyme (organic molecule, often a vitamin) o *Enzymes are very specific* Act on a very specific ubstrate o Names usually end in ase and are often named for the reaction they catalyze Example: hydrolases, oxidases 2.11 Nucleic Acids Nucleic acids, composed of C, H, O, N, and P, are the largest molecules in the body Nucleic acid polymers are made up of monomers called ucleotides o Nucleotides are composed of nitrogen base, a pentose sugar, and a phosphate group Two major classes: o Deoxyribonucleic acid (DNA) o Ribonucleic acid (RNA) DNA holds the genetic blueprint for the synthesis of all proteins o Doublestranded helical molecule (uble helix) located in cell nucleus o Nucleotides contain a deoxyribose sugar, phosphate group, and one of four nitrogen bases: Purines: adenine (A), guanine (G) Pyrimidines: cytosine (C), thymine (T) o Bonding of nitrogen base from strand to opposite strand is very specific Follows complementary basepairing rules: A always pairs with T G always pairs with C RNA links DNA to protein synthesis and is slightly different from DNA o Singlestranded linear molecule is active mostly outside nucleus o Contains aribose sugar (not deoxyribose) o Thymine is replaced withuracil o Three varieties of RNA carry out the DNA orders for protein synthesis Messenger RNA (mRNA), transfer RNA (tRNA), and ribosomal RNA (rRNA) 2.12 ATP Chemical energy released when glucose is broken down is captured in P (adenosine triphosphate) ATP directly powers chemical reactions in cells o Offers immediate, usable energy needed by body cells o ATP is the body's energy currency Structure of ATP o Adeninecontaining RNA nucleotide with two additional phosphate groups Terminal phosphate group of ATP can be transferred to other compounds that can use energy stored in phosphate bond to do work o Loss of phosphate group converts ATP to ADP o Loss of second phosphate group converts ADP to AMP
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