Chemistry of Life PDF

Chemistry of Life © 2019 McGraw-Hill Education. All rights reserved. Authorized only for instructor use in the classroom. No reproduction or further distribution permitted without the prior written consent of McGraw-Hill Education. Biological organization is the hierarchy of complex biological st...

Chemistry of Life © 2019 McGraw-Hill Education. All rights reserved. Authorized only for instructor use in the classroom. No reproduction or further distribution permitted without the prior written consent of McGraw-Hill Education. Biological organization is the hierarchy of complex biological structures and systems that define life. Each level in the hierarchy represents an increase in organizational complexity, with each "object" being primarily composed of the previous level's basic unit. ©McGraw-Hill Education Life can be organized in a hierarchy of levels. Organization: the levels of biological organization extends within (micro) and beyond (macro) the individual as follows: atoms and molecules→ cells → tissues → organs → organ systems → organisms → populations →communities → ecosystems → biosphere ©McGraw-Hill Education 4 Organization of Life • Organism • Organ systems Organism • Organs • Tissues • Cells Organ system Tissue Organ • Organelles Cell Macromolecule Organelle • Molecules • Atoms © 2019 McGraw-Hill Education Atom Molecule Building blocks of Life Matter has mass and occupies space • 3 forms of matter: • Solid (e.g., bone) • Liquid (e.g., blood) • Gas (e.g., oxygen) • Atom is the smallest particle exhibiting chemical properties of an element • 92 naturally occurring elements make up matter • Organized in periodic table of elements © 2019 McGraw-Hill Education 5 Periodic Table of Elements © 2019 McGraw-Hill Education 6 Most Common Elements of the Human Body © 2019 McGraw-Hill Education 7 Building blocks of Life Components of an atom Atoms composed of three subatomic particles • Neutrons • Mass of one atomic mass unit (amu) • No charge • Protons • Mass of one amu • Positive charge of one (+1) • Electrons • 1/1800th of the mass of a proton or neutron • Negative charge of one (−1) • Located at varying distance from the nucleus in regions called orbitals © 2019 McGraw-Hill Education 8 9 Atomic structure © 2019 McGraw-Hill Education Periodic Table Periodic table • Chemical symbol • Unique to each element • Usually identified by first letter, or first letter plus an additional letter, e.g., C is carbon • Atomic number • Number of protons in an atom of the element • Located above symbol name • Elements arranged by anatomic number within rows • Average atomic mass • Mass of both protons and neutrons • Shown below the element’s symbol on the table © 2019 McGraw-Hill Education 10 11 Building blocks of Life Determining the number of subatomic particles • Proton number = atomic number • Neutron number = atomic mass − atomic number • Neutron number = (p + n) − p • Neutron number of Na = 23 − 11 = 12 • Electron number = proton number © 2019 McGraw-Hill Education Building blocks of Life 12 Diagramming atomic structures An atom has “shells” of electrons surrounding the nucleus • Each shell has given energy level • Each shell holds a limited number of electrons • Innermost shell: two electrons, second shell up to eight • Shells close to the nucleus must be filled first © 2019 McGraw-Hill Education Figure 2.2b 13 Isotopes Isotopes are different atoms of the same element • Same number of protons and electrons; different number of neutrons • Identical chemical characteristics; different atomic masses E.g., carbon exists in three isotopes • Carbon-12, with 6 neutrons • Most prevalent type • Carbon-13, with 7 neutrons • Carbon-14, with 8 neutrons Figure 2.3 © 2019 McGraw-Hill Education Isotopes Weighted average of atomic mass for all isotopes is the average atomic mass Radioisotopes • Contain excess neutrons, so unstable • Lose nuclear components in the form of high energy radiation • Alpha particles • Beta particles • Gamma rays © 2019 McGraw-Hill Education 14 Isotopes Physical half-life • The time for 50% of radioisotope to become stable • Can vary from a few hours to thousands of years Biological half-life • The time required for half of the radioactive material from a test to be eliminated from the body © 2019 McGraw-Hill Education 15 Clinical View: Medical Imaging of the Thyroid Gland Using Iodine Radioisotopes Radioisotopes introduced into the body during medical procedures • Used by cells in a similar manner to nonradioisotopes • Can trace products of metabolic reactions that use these elements • Thyroid gland darker in areas where less radioactive iodine taken up • Can help locate a nodule © 2019 McGraw-Hill Education 16 Ions and Ionic Compounds Chemical compounds • Stable associations between two or more elements combined in a fixed ratio • Classified as ionic or molecular Ionic compounds are structures composed of ions held together in a lattice by ionic bonds © 2019 McGraw-Hill Education 17 Ions Ions • Atoms with a positive or a negative charge • Produced from loss or gain of one or more electrons • Significant physiological functions • E.g., K+ is used to sports drinks to replace the K+ lost in sweat • E.g., K+ in a large dose is used in some states for lethal injection © 2019 McGraw-Hill Education 18 Ions Losing electrons and the formation of cations Sodium can reach stability by donating an electron • Now satisfies the octet rule • Now has 11 protons and 10 electrons • Charge is +1 Cations are ions with a positive charge © 2019 McGraw-Hill Education 19 Ions Gaining electrons and the formation of anions Chlorine reaches stability by gaining an electron • Now satisfies the octet rule • Now has 17 protons and 18 electrons • Charge is −1 Anions are ions with negative charge Polyatomic ions are anions with more than one atom • E.g., bicarbonate ion and phosphate ion © 2019 McGraw-Hill Education 20 Ionic Bonds 21 Ionic bonds • Cations and anions bound by electrostatic forces • Form salts • E.g., table salt (NaCl) • Each sodium atom loses one outer shell electron to a chlorine atom • Sodium and chlorine ions are held together by ionic bonds in a lattice crystal structure (ionic compound) • E.g., magnesium chloride • Each magnesium atom loses one electron to each of the two chlorine atoms © 2019 McGraw-Hill Education Formation of an Ionic Bond Involving Sodium and Chloride © 2019 McGraw-Hill Education 22 Covalent Bonding, Molecules, and Molecular Compounds Covalently bonded molecule • Electrons shared between atoms of two or more different elements • Termed molecular compounds • E.g., carbon dioxide (CO2), but not molecular oxygen (O2) © 2019 McGraw-Hill Education 23 Chemical Formulas: Molecular and Structural Molecular formula • Indicates number and type of atoms • E.g., carbonic acid (H2CO3) Structural formula • Indicates number and type of atoms • Indicates arrangement of atoms within the molecule • E.g., O=C=O (carbon dioxide) • Allows differentiation of isomers • Same number and type of elements, but arranged differently in space © 2019 McGraw-Hill Education 24 Chemical Formulas: Molecular and Structural Glucose vs. galactose vs. fructose • Same molecular formula • 6 carbon, 12 hydrogen, 6 oxygen • Atoms arranged differently Isomers may have different chemical properties Figure 2.6 © 2019 McGraw-Hill Education 25 Covalent Bonds Covalent bond • Atoms share electrons • Occurs when both atoms require electrons • Occurs with atoms with 4 to 7 electrons in outer shell • Formed commonly in human body using • Hydrogen (H) • Oxygen (O) • Nitrogen (N) • Carbon (C) © 2019 McGraw-Hill Education 26 Covalent Bonds Number of covalent bonds an atom can form Simplest occurs between two hydrogen atoms • Each sharing its single electron Oxygen needs two electrons to complete outer shell • Forms two covalent bonds Nitrogen forms three bonds Carbon forms four bonds © 2019 McGraw-Hill Education 27 Covalent Bonds Single, double, and triple covalent bonds Single covalent bond • One pair of electrons shared • E.g., between two hydrogen atoms Double covalent bond • Two pairs of electrons shared • E.g., between two oxygen atoms Triple covalent bond • Three pairs of electrons shared • E.g., between two nitrogen atoms © 2019 McGraw-Hill Education 28 Single, Double, and Triple Covalent Bonds © 2019 McGraw-Hill Education 29 Covalent Bonds Carbon skeleton formation Carbon • Bonds in straight chains, branched chains, or rings • Carbon present where lines meet at an angle • Additional atoms are hydrogen © 2019 McGraw-Hill Education 30 Covalent Bonds Nonpolar and polar covalent bonds Electronegativity—relative attraction of each atom for electrons • Determines how electrons are shared in covalent bonds • Two atoms of same element have equal attraction for electrons • Resulting bond is nonpolar covalent bond • Sharing of electrons unequally = polar covalent bond © 2019 McGraw-Hill Education 31 Covalent Bonds Nonpolar and polar covalent bonds (continued) In periodic table, electronegativity increases • From left to right across row • From bottom to top in column For 4 most common elements composing living organisms • From least to greatest electronegativity • hydrogen < carbon < nitrogen < oxygen © 2019 McGraw-Hill Education 32 Covalent Bonds Electrons have negative charge • More electronegative atom develops a partial negative charge • Less electronegative atom develops a partial positive charge • Written using Greek delta (δ) followed by superscript plus or minus • Exception to rule of polar bond forming between two different atoms • Carbon bonding with hydrogen © 2019 McGraw-Hill Education 33 Nonpolar, Polar, and Amphipathic Molecules Covalent bonds may be polar or nonpolar • Nonpolar molecules contain nonpolar covalent bonds • E.g., O—O and C—H are nonpolar bonds • Polar molecules contain polar covalent bonds • O—H is a polar bond in the polar molecule water (H2O) • Nonpolar molecules may contain polar covalent bonds, if the polar covalent bonds cancel each other • E.g., carbon dioxide © 2019 McGraw-Hill Education 34 Nonpolar, Polar, and Amphipathic Molecules Amphipathic molecules • Large molecules with both polar and nonpolar regions • E.g., phospholipids © 2019 McGraw-Hill Education 35 36 Nonpolar, Polar, and Amphipathic Molecules © 2019 McGraw-Hill Education Intermolecular Attractions Intermolecular attractions • Weak chemical attractions between molecules • Important for shape of complex molecules • E.g., DNA and proteins • Hydrogen bond • Forms between polar molecules • Attraction between partially positive hydrogen atom and a partially negative atom • Individually weak, collectively strong • Influences how water molecules behave © 2019 McGraw-Hill Education 37 Hydrogen Bonding © 2019 McGraw-Hill Education 38 39 Intermolecular Attractions Other intermolecular attractions • Van der Waals forces • Nonpolar molecules • Electrons orbiting nucleus briefly, unevenly distributed • Induce unequal distribution of adjacent atom of another nonpolar molecule • Individually weak • Hydrophobic interactions • Nonpolar molecules placed in a polar substance • If occurring between parts of large molecule, termed intramolecular attractions © 2019 McGraw-Hill Education Molecular Structure of Water Water • Composes two-thirds of the human body by weight • Polar molecule • One oxygen atom bonded to two hydrogen atoms • Oxygen atom has two partial negative charges • Hydrogens have single partial positive charge • Can form four hydrogen bonds with adjacent molecules • Central to water’s properties © 2019 McGraw-Hill Education 40 41 Properties of Water Phases of water 3 phases of water, depending on temperature: • Gas (water vapor) • Substances with low molecular mass • Liquid (water) • Almost all water in the body • Liquid at room temperature due to hydrogen bonding • Solid (ice) © 2019 McGraw-Hill Education 42 Properties of Water Phases of water (continued ) Functions of liquid water: • Transports • Substances dissolved in water move easily throughout body • Lubricates • Decreases friction between body structures • Cushions • Absorbs sudden force of body movements • Excretes wastes • Unwanted substances dissolve in water are easily eliminated © 2019 McGraw-Hill Education 43 Properties of Water Cohesion, surface tension, and adhesion Cohesion • Attraction between water molecules due to hydrogen bonding Surface tension • Inward pulling of cohesive forces at surface of water • Causes moist sacs of air in lungs to collapse • Surfactant, a lipoprotein, prevents collapse Adhesion • Attraction between water molecules and a substance other than water © 2019 McGraw-Hill Education Properties of Water High specific heat and high heat of vaporization Temperature • Measure of kinetic energy of atoms or molecules within a substance Specific heat • Amount of energy required to increase temperature of 1 gram of a substance by 1 degree Celsius • Water’s value extremely high due to energy needed to break hydrogen bonds • Contributes to keeping body temperature constant © 2019 McGraw-Hill Education 44 Properties of Water High specific heat and high heat of vaporization (continued ) Heat of vaporization • Heat required for release of molecules from a liquid phase into a gaseous phase for 1 gram of a substance • Water’s value very high due to hydrogen bonding • Sweating cools body • Excess heat dissipated as water evaporates © 2019 McGraw-Hill Education 45 Water as the Universal Solvent Water—solvent of the body Solutes are substances that dissolve in water Water is called universal solvent because most substances dissolve in it • Chemical properties of a substance determine whether it will dissolve or not © 2019 McGraw-Hill Education 46 Water as the Universal Solvent Substances that dissolve in water Polar molecules and ions • Hydrophilic means “water-loving” • Water surrounds substances, forms a hydration shell • Some substances dissolve but remain intact • E.g., glucose and alcohol • Nonelectrolytes remain intact but do not conduct current • Substances dissolve and dissociate (separate) – • NaCl dissociates into Na+ and Cl ions • Acids and bases, such as HCl • Electrolytes can conduct current © 2019 McGraw-Hill Education 47 48 Water as the Universal Solvent Substances that do not dissolve in water Nonpolar molecules • Hydrophobic means “water-fearing” • Hydrophobic exclusion—cohesive water molecules “force out” nonpolar molecules • Hydrophobic interaction—“excluded molecules” • Hydrophobic substances require carrier proteins to be transported within the blood • E.g., fats and cholesterol are unable to dissolve within water © 2019 McGraw-Hill Education 49 Water as the Universal Solvent Substances that partially dissolve in water Amphipathic molecules have polar and nonpolar regions • Polar portion of molecule dissolves in water • Nonpolar portion repelled by water Phospholipid molecules are amphipathic • Polar heads have contact with water • Nonpolar tails group together • Results in bilayers of phospholipid molecules • E.g., membranes of a cell Other amphipathic molecules form a micelle © 2019 McGraw-Hill Education Substance Interaction with Water © 2019 McGraw-Hill Education 50 Water: A Neutral Solvent Water spontaneously dissociates to form ions • Bond between oxygen and hydrogen breaks apart spontaneously • 1/10,000,000 ions per liter • OH group hydroxide ion (OH–) • Hydrogen ion transferred to a second water molecule • Hydronium ion (H3O+) • Equal numbers of positive hydrogen ions and negative hydroxyl ions produced • Water remains neutral H2O + H2O → H3O+ + OH− simplified to H2O → H+ + OH– © 2019 McGraw-Hill Education 51 Acids and Bases Acid dissociates in water to produce H+ and an anion • Proton donor • Increases concentration of free H+ • More dissociation of H+ with stronger acids • E.g., HCl in the stomach • Less dissociation of H+ with weaker acids • E.g., carbonic acid in the blood Substance A (an acid in water) → H+ + Anion © 2019 McGraw-Hill Education 52 53 Acids and Bases Base accepts H+ when added to solution • Proton acceptor • Decreases concentration of free H+ • More absorption of H+ with stronger bases • E.g., ammonia and bleach • Less absorption of H+ with weaker bases • E.g., bicarbonate in blood and in secretions released into small intestine Substance B (a base in water) + H+ → B—H © 2019 McGraw-Hill Education pH, Neutralization, and the Action of Buffers pH is a measure of H+ • Relative amount of H+ in a solution • Range between 0 and 14 The pH of plain water is 7 • Water dissociates to produce 1/10,000,000 of H+ and OH– ions per liter • Equal to 1 × 10–7 or to 0.0000001 pH and H+ concentration are inversely related • Inverse of the log for a given H+ concentration • As H+ concentration increases, pH decreases • As H+ concentration decreases, pH increases © 2019 McGraw-Hill Education 54 pH, Neutralization, and the Action of Buffers Interpreting the pH scale Solutions with equal concentrations of H+ and OH– • Are neutral • Have a pH of 7 Solutions with greater H+ than OH– • Are acidic • Have a pH < 7 Solutions with greater OH– than H+ • Are basic (alkaline) • Have a pH > 7 Moving from one increment to next is a 10-fold change • E.g., a pH of 6 has 10 times greater concentration of H+ than pure water © 2019 McGraw-Hill Education 55 pH © 2019 McGraw-Hill Education 56 pH, Neutralization, and the Action of Buffers Neutralization • When an acidic or basic solution is returned to neutral (pH 7) • Acids neutralized by adding base • E.g., medications to neutralize stomach acid must contain a base • Bases neutralized by adding acid Buffers • Help prevent pH changes if excess acid or base is added • Act to accept H+ from excess acid or donate H+ to neutralize base • Carbonic acid (weak acid) and bicarbonate (weak base) buffer blood pH • Both help maintain blood pH in a critical range (7.35 to 7.45) © 2019 McGraw-Hill Education 57 58 Water Mixtures Mixtures are formed from combining two or more substances Two defining features: • Substances mixed are not chemically changed • Substances can be separated by physical means • E.g., evaporation or filtering © 2019 McGraw-Hill Education Categories of Water Mixtures Three categories of water mixtures: 1. Suspension: material larger in size than 1 mm mixed with water • E.g., blood cells within plasma or sand in water • Does not remain mixed unless in motion • Appears cloudy or opaque; scatters light 2. Colloid: smaller particles than a suspension, but larger than those in a solution • E.g., fluid in cell cytosol and fluid in blood plasma • Remains mixed when not in motion • Scatters light © 2019 McGraw-Hill Education 59 60 Categories of Water Mixtures Three categories of water mixtures (continued ) 3. Solution: homogeneous mixture of material smaller than 1 nanometer • Dissolves in water • Does not scatter light; does not settle if solution not in motion • E.g., sugar water, salt water, blood plasma • Special category of suspension—emulsion • Water and a nonpolar liquid substance • E.g., oil and vinegar salad dressing or breast milk • Does not mix unless shaken © 2019 McGraw-Hill Education Mixtures and an Emulsion © 2019 McGraw-Hill Education 61 Expressions of Solution Concentration Concentration is determined by the amount of solute dissolved in a solution Expressions of concentration: • Mass/volume • Mass of solute per volume of solution • E.g., results from a blood test • Mass/volume percent • Grams of solute per 100 mL solution • E.g., IV solutions © 2019 McGraw-Hill Education 62 Expressions of Solution Concentration Expressions of concentration (continued ) • Molarity • Moles solute/L solution • Alters with changes in temperature • More easily measured in the body than molality • Molality • Moles solute/kg solvent • Does not alter with changes in temperature • Slightly more accurate than molarity © 2019 McGraw-Hill Education 63 Expressions of Solution Concentration Osmoles (osm) is the unit of measurement for the number of particles in a solution • Reflects whether a substance dissolves, or dissolves and dissociates • Can be expressed as osmolarity or osmolality • Osmolarity • Number of particles in a 1 liter solution • Easier to measure • Blood serum expresses as milliosmoles • Osmolality • Number of particles in 1 kg of water • More accurate, but difficult to measure © 2019 McGraw-Hill Education 64 65 Expressions of Solution Concentration Mole = 6.022 × 1023 atoms, ions, or molecules • Mass in grams equal to atomic mass of an element or molecular mass of a compound • E.g., one mole carbon = 12.01 grams To find molecular mass, multiply number of units of each element by average atomic mass and add totals • Some variation due to isotopes © 2019 McGraw-Hill Education 66 © 2019 McGraw-Hill Education Starr-Taggart Figure 2-12 p32 67 • Living things require the pH of the body to stay within a narrow range around 7.35 – 7.45 – Slightly alkaline • Buffers help maintain the pH by taking up excess H+ or OHH20 + CO2 → H2CO3 (carbonic acid) Too basic: H2CO3 (carbonic acid) → H+ + HCO3- (bicarbonate) Too acidic: H+ + HCO3- (bicarbonate) → H2CO3 (carbonic acid) © 2019 McGraw-Hill Education ATP Adenosine triphosphate (ATP) • Nucleotide composed of nitrogenous base adenine, a ribose sugar, and three phosphate groups • Central molecule in transfer of chemical energy within cell • Covalent bonds between last two phosphate groups are unique, energy rich • Release energy when broken Important nucleotide-containing molecules • Nicotinamide adenine dinucleotide • Flavin adenine dinucleotide • Both participate in production of ATP © 2019 McGraw-Hill Education 68