Unit 2 Notes

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Department
Biological Sciences
Course
BIO SCI 93
Professor
Diane O' Dowd
Semester
Fall

Description
Unit 2 Elements and Compounds  Matter- anything that takes up place and has mass  Element- substance that cannot be broken down to other substances by chemical reactions o 92 in nature  Compound- substance consisting of two or more different elements combined in a fixed ratio The Elements of Life  20-25% of 92 natural elements are essential elements an organism needs to be healthy  O,C,H,N = 96% living matter o Other 4% = Ca, P, K, S  Trace elements- required by an organism in minute quantities o Fe An element’s properties depend on the structure of atoms  Atom- smallest unit of matter that still retains the properties of an element Subatomic Particles  Atomic nucleus- center of an atom, packed with protons and neutrons o Electrons form cloud around nucleus, attracted to positive nucleus  Proton and neutron = 1.7E-24 g o 1 dalton = 1 amu = 1.7E-24 g o Mass of electron = 1/2000 mass of a neutron/proton Atomic Number and Atomic Mass  Atomic number- number of protons  Mass number- sum of protons and neutrons in the nucleus of an atom  Atomic mass- mass number (because electron mass can be disregarded) Isotopes  Isotope- atom of a given element with same number of protons but more neutrons than some atoms of the same element and therefore have greater mass  In nature, an element occurs as a mixture of its isotopes  Radioactive isotope- isotope in which the nucleus decays spontaneously, giving off particles and energy o Leads to change of element if decay changes the number of protons o Radioactive tracers used as diagnostic tools in medicine The Energy Levels of Electrons  Only electrons are directly involved in the chemical reactions between atoms; nuclei can’t touch  Energy- capacity to cause change by doing work  Potential energy- the energy that matter possesses because of its location or structure  The more distant an electron is from the nucleus  the greater its potential energy  Electrons are found in different electron shells, each with a characteristic average distance and energy level o Electron absorbs energy  moves to shell farther out from the nucleus o Electron loses energy  “falls back” to a shell closer to the nucleus; lost energy = heat Electron Distribution and Chemical Properties  Chemical behavior of an atom depends on number of electrons in outermost shell  Valence electrons make a valence shell 1 Electron Orbitals  Orbital- the 3-dimensional space where an electron is found 90% of the time o A component of an electron shell o Maximum of 2 electrons can occupy a single orbital Formation and Function of Molecules Depend on Chemical Bonding Between Atoms  Chemical bonds- attractions concerning valence electrons that keep atoms staying close together Covalent Bonds  Covalent bond- sharing of a pair of valence electrons by two atoms  Two or more atoms constitute a molecule  Single bond- pair of shared electrons  Double bond- two shared pairs of valence electrons  Atom’s valence = atom’s bonding capacity o Gets complicated in the third row of the periodic table (ex. P can have v=3 & v=5)  H & 2 = p2re elements, NOT compounds  Electronegativity- attraction of a particular atom for the electrons of a covalent bond o More electronegative = more strongly it pulls shared electrons toward itself o Nonpolar covalent bond = same electronegativity; polar = unequal electronegativity o Ex. H2O = electrons spend more time near O because O is much more electronegative Ionic Bonds  Ion- charged atom/molecule o Positive = cation; negative = anion o Make an ionic bond  make ionic compounds, or salts (crystals)  Temperature and environment affects the strength of ionic bonds o Ex. Salt crystal dry and salt crystal in water Weak Chemical Bonds  Hydrogen bonds o H-bond- noncovalent attraction between a hydrogen and an electronegative atom o F,O,N  Van der Waals Interactions  - electrons may accumulate by chance in one part of the molecule or another, even in a molecule with nonpolar covalent bonds  results in ever-changing regions of positive and negative charge that enable all atoms and molecules to stick to one another  Occur only when atoms/molecules are very close together  Why a lizard can walk straight up a wall Molecular Shape and Function  Shapes determined by the positions of the atoms’ orbitals  Shape important in how biological molecules recognize and respond to one another o Ex. Morphine and opiates Chemical reactions make and break chemical bonds  Chemical reaction- making and breaking of chemical bonds, leading to changes in the composition of matter (reactants  products)  Concentration of reactants affects the rate of reaction o Greater concentration of reactions  more frequently they collide with one another and have the opportunity to react and form products 2  Chemical equilibrium- the point at which forward and reverse reaction rates are equal o No net effect on the concentrations of reacts and products Polar covalent bonds in water molecules result in hydrogen bonding  H and O held together by hydrogen bond o Hydrogen bonds, very fragile in liquid form Cohesion of Water Molecules  Cohesion- hydrogen bonds of water collectively hold it together o Contributes to transport of water against gravity in plants  Adhesion- clinging of one substance to another o Adhesion of water to cell walls by hydrogen bonds help counter the downward pull of gravity  Cohesion of water = water’s surface tension, how difficult it is to stretch or break the surface of a liquid Moderation of Temperature by Water  Heat and temperature o Kinetic energy- energy of motion o Heat- form of energy  Amount of heat is a measure of the matter’s total kinetic energy due to motion of its molecules; heat depends in part on the matter’s volume o Temperature- measure of heat intensity that represents the average kinetic energy of the molecules, regardless of volume o Calorie (cal)- amount of heat it takes to raise the temperature of 1g of water by 1°C  1000 cal = kilocalorie (kcal)  1 J = 0.239 cal; 1 cal = 4.184 J  Water’s high specific heat o Specific heat- amount of heat that must be absorbed or lost for 1g of that substance to change its temperature by 1°C  A measure of how well a substance resists changing its temperature when it absorbs or releases heat  Specific heat of water = 1 cal/g x °C  Water has unusually high specific heat • Hydrogen bonding: heat must be absorbed to break H-bonds  Evaporative cooling o Speed of molecular movement varies, and temperature is the average KE of molecules o Heat of vaporization- quantity of heat a liquid must absorb for 1 g of it to be converted from the liquid to the gaseous state o Evaporative cooling- “hottest” molecules most likely to leave as gas  Contributes to stability of temperature in lakes and ponds Floating of Ice on Liquid Water  Water expands when it solidifies  floats Water: The Solvent of Life  Solution- homogenous mixture o Solvent- dissolving agent o Solute- substance that is dissolved  Water is a versatile solvent because of its’ polarity  Hydration shell- sphere of water molecules around each dissolved ion 3  Hydrophilic and hydrophobic substances o Hydrophilic- affinity for water o Colloid- stable suspension of fine particles in a liquid o Hydrophobic- repel water  Oils  Solute concentration in aqueous solutions o Molarity- the number of moles of solute per liter of solution Acidic and basic conditions affect living organisms  Hydrogen ion: H+ , hydroxide ion: OH- , hydronium ion: H3O+ Acids and Bases  Acid- substance that increases the hydrogen concentration of a solution  Base- substance that reduces the hydrogen ion concentration of a solution The pH Scale -14  [H+][OH-] = 10  Acids not only add hydrogen ions to a solution, but also remove hydroxide ions because of the tendency of H+ to combine with OH- to form water o Same with bases: increase OH- concentration but also reduce H+ by formation of water  pH of a solution- the negative logarithm of the hydrogen ion concentration o pH = -log[H+] Buffers  pH of human blood = 7.4  buffer- substance that minimizes changes in the concentrations of H+ and OH= in a solution Organic chemistry is the study of carbon compounds  overall percentages of major elements of life, C,H,O,N,S, and P – almost uniform between organisms  early 1800’s = vitalism, belief in a life force outside the jurisdiction of physical/chemical laws o mechanism, view that physical/chemical laws govern al natural phenomena, including the processes of life  Miller concluded that complex organic molecules could arise spontaneously under conditions thought to have existed on early Earth The Formation of Bonds with Carbon  C has 6 electrons; 2 in first shell and 4 in second  Four single bonds = tetrahedron o Figure 2.17B – 109.5°  Most frequent partners: H, O, N  CO =2inorganic compound because it is very simple and lacks H even though it contains carbon Molecular Diversity Arising from Carbon Skeleton Variation  Carbon chains form the skeletons of organic molecules 4  Hydrocarbons o Hydrocarbons- organic molecules consisting of only carbon and hydrogen o Neither petroleum or fat dissolve in water; both are hydrophobic because the great majority of their bonds are relatively nonpolar carbon-to-hydrogen linkages  Hydrocarbons also can undergo reactions that release a relatively large amount of energy  Isomers o Isomers- compounds with same numbers of atoms of the same elements but different structures and different properties o Structural isomers- differ in the covalent arrangements of their atoms  Figure 4.7A  C5H12  Number of possible isomers increases as carbon skeletons increase in size  Cis-trans isomers (geometric isomers)- carbons have covalent bonds to the same atoms but these atoms differ in their spatial arrangements due to the inflexibility of double bonds o Figure 4.7B: The arrangement with box X’s on the same side of the double bond is called a cis isomer; the arrangement with X’s on opposite sides is called a trans isomer  Enantiomers- isomers that are mirror images of each other and differ in shape due to the presence of an asymmetric carbon, one that is attached to four different atoms or groups of atoms o Usually only one isomer is biologically active because only that form can bind to specific molecules in an organism o Figure 4.7C The Chemical Groups Most Important in the Processes of Life  Functional groups- chemical groups that affect molecular function by being directly involved in chemical reactions  7 chemical groups most important in biological processes: hydroxide, carbonyl, carboxyl, amino, sulfhydryl, phosphate, and methyl groups o All but methyl groups are functional groups; are also hydrophilic and thus increase the solubility of organic compounds in water o Methyl group is not reactive, but serves as a recognizable tag on biological molecules  Chemical groups: o Hydroxyl  Hydrogen atom bonded to oxygen atom, which is bonded to carbon skeleton of the organic molecule (-OH)  Alcohols • Ethanol 5  Polar – electrons spend more time near the electronegative oxygen atom  Can form hydrogen bonds with water molecules to help dissolve organic compounds such as sugars o Carbonyl  Carbon atom joined to an oxygen atom by a double bond (>CO)  Ketone if carbonyl is within carbon skeleton; aldehyde if carbonyl is at end of carbon skeleton  Acetone (ketone)  Propanal (aldehyde)  Ketone and aldehyde may be structural isomers with different properties  Ketone and aldehyde groups are found in sugars, giving rise to two major groups of sugars: ketoses (containing ketone groups) and aldoses (containing aldehyde groups) o Carboxyl  An oxygen atom double-bonded to a carbon atom that is also bounded to an –OH group (-COOH)  Carboxylic acids  Acetic acid  Acts as an acid; can donate an H+ because the covalent bond between oxygen and hydrogen is so polar  Found in cells in the ionized form with a charge of -1 and called a carboxylate ion o Amino  A nitrogen atom bonded to two hydrogen atoms and to the carbon skeleton (-NH2)  Amines  Glycine (both an amine and a carboxylic acid because has both an amino group and a carboxyl group); compounds with both are called amino acids  Acts as a base; can pick up an H+ from the surrounding solution 6  Found in cells in the ionized form with a charge of +1 o Sulfhydryl  Sulfur atom bonded to an atom of hydrogen (-SH)  Thiols  Cysteine  Two sulfhydryl groups can react, forming a covalent bond; “cross-linking”  more stabilized protein structure  Cross-linking of cysteines in hair proteins maintains the curliness or straightness of hair. Straight hair can be “permanently” curled by shaping it around curlers and then breaking and re-forming the cross-linking bonds o Phosphate  Phosphorus atom bonded to four oxygen atoms; one oxygen is bonded to the carbon skeleton and two oxygens carry negative charges (-OPO ) 2- 3  Organic phosphates  Glycerol phosphate  Contributes neg
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