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Matter—anything that has mass and occupies space Weight—pull of gravity on mass 1 3 states of matter Solid—definite shape and volume Liquid—changeable shape; definite volume Gas—changeable shape and volume 2 Capacity to do work or put matter into motion Types of energy Kinetic—energy in action Potential—stored (inactive) energy Energy can be transferred from potential to kinetic energy Capacity to do work or put matter into motion Types of energy Kinetic—energy in action Potential—stored (inactive) energy Energy can be transferred from potential to kinetic energy 4 Chemical energy Stored in bonds of chemical substances Electrical energy Results from movement of charged particles Mechanical energy Directly involved in moving matter Radiant or electromagnetic energy Travels in waves (e.g., visible light, ultraviolet light, and x-rays) Energy may be converted from one form to another Energy conversion is inefficient Some energy is “lost” as heat (partly unusable energy) 5 Elements Matter is composed of elements Elements cannot be broken into simpler substances by ordinary chemical methods Each has unique properties Physical properties Detectable with our senses, or are measurable Chemical properties How atoms interact (bond) with one another 6 Atoms Unique building blocks for each element Give each element its physical & chemical properties Smallest particles of an element with properties of that element Atomic symbol One- or two-letter chemical shorthand for each element 7 Four elements make up 96.1% of body mass Element Carbon Hydrogen Oxygen Nitrogen 8 9 elements make up 3.9% of body mass Element Calcium Phosphorus Potassium Sulfur Sodium Chlorine Magnesium 9 Iodine Iron 9 Very minute amounts 11 elements make up < 0.01% of body mass Many are part of, or activate, enzymes Chromium Copper 10 Fluorine Manganese Silicon Zinc 10 Atoms are composed of subatomic particles Protons, neutrons, electrons Protons and neutrons found in nucleus Electrons orbit nucleus in an electron cloud Nucleus almost entire mass of the atom Neutrons Carry no charge Mass = 1 atomic mass unit (amu) Protons Carry positive charge Mass = 1 amu Electrons in orbitals within electron cloud Carry negative charge 1/2000 the mass of a proton (0 amu) Number of protons and electrons always equal Planetary model—simplified; outdated Incorrectly depicts fixed circular electron paths But useful for illustrations (as in the text) Orbital model—current model used by chemists Probable regions of greatest electron density (an electron cloud) Useful for predicting chemical behavior of atoms Different elements contain different numbers of subatomic particles Hydrogen has 1 proton, 0 neutrons, and 1 electron Lithium has 3 protons, 4 neutrons, and 3 electrons Atomic number = Number of protons in nucleus Written as subscript to left of atomic symbol Ex. 3Li Mass number Total number of protons and neutrons in nucleus Total mass of atom Written as superscript to left of atomic symbol Ex. 7Li 15 Isotopes Structural variations of atoms Differ in the number of neutrons they contain Atomic numbers same; mass numbers different Atomic weight Average of mass numbers (relative weights) of all isotopes of an atom Atomic number, mass number, atomic weight Give “picture” of each element Allow identification 17 Heavy isotopes decompose to more stable forms Spontaneous decay called radioactivity Similar to tiny explosion Can transform to different element Can be detected with scanners 18 Valuable tools for biological research and medicine Share same chemistry as their stable isotopes Most used for diagnosis All damage living tissue Some used to destroy localized cancers Radon from uranium decay causes lung cancer 19 A molecule is formed when two or more atoms join together chemically. A compound is a molecule that contains at least two different elements. All compounds are molecules but not all molecules are compounds. Textbook: Most atoms chemically combined with other atoms to form molecules and compounds Molecule Two or more atoms bonded together (e.g., H2 or C6H12O6) Smallest particle of a compound with specific characteristics of the compound Compound Two or more different kinds of atoms bonded together (e.g., C6H12O6 , but not H2) 20 Most matter exists as mixtures Two or more components physically intermixed 21 Three types of mixtures Solutions Colloids Suspensions 22 Homogeneous mixtures Most are true solutions in body Gases, liquids, or solids dissolved in water Usually transparent, e.g., atmospheric air or saline solution Solvent Substance present in greatest amount 23 Usually a liquid; usually water Solute(s) Present in smaller amounts Ex. If glucose is dissolved in blood, glucose is solute; blood is solvent 23 Can be expressed as Percent of solute in total solution (assumed to be water) Parts solute per 100 parts solvent Milligrams per deciliter (mg/dl) Molarity, or moles per liter (M) 1 mole of an element or compound = Its atomic or molecular weight (sum of atomic weights) in grams 1 mole of any substance contains 6.02 1023 molecules of that substance (Avogadro’s number) 24 Colloids (AKA emulsions) Heterogeneous mixtures, e.g., cytosol Large solute particles do not settle out Some undergo sol-gel transformations e.g., cytosol during cell division Suspensions Heterogeneous mixtures, e.g., blood Large, visible solutes settle out Mixtures No chemical bonding between components Can be separated by physical means, such as straining or filtering Heterogeneous or homogeneous Compounds Chemical bonding between components Can be separated only by breaking bonds All are homogeneous 26 Chemical bonds are energy relationships between electrons of reacting atoms Electrons can occupy up to seven electron shells (energy levels) around nucleus Electrons in valence shell (outermost electron shell) 27 Have most potential energy Are chemically reactive electrons Octet rule (rule of eights) Except for the first shell (full with two electrons) atoms interact to have eight electrons in their valence shell 27 Stable and unreactive Valence shell fully occupied or contains eight electrons Noble gases Valence shell not full Tend to gain, lose, or share electrons (form bonds) with other atoms to achieve stability Three major types Ionic bonds Covalent bonds Hydrogen bonds 30 Ions Atom gains or loses electrons and becomes charged # Protons ≠ # Electrons Transfer of valence shell electrons from one atom to another forms 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 Attraction of opposite charges results in an ionic bond 31 Most ionic compounds are salts When dry salts form crystals instead of individual molecules Example is NaCl (sodium chloride) Formed by sharing of two or more valence shell electrons Allows each atom to fill its valence shell at least part of the time Nonpolar Covalent Bonds Electrons shared equally Produces electrically balanced, nonpolar molecules such as CO2 Polar Covalent Bonds Unequal sharing of electrons produces polar (AKA dipole) molecules such as H2O Atoms in bond have different electron-attracting abilities Small atoms with six or seven valence shell electrons are electronegative, e.g., oxygen Strong electron-attracting ability Most atoms with one or two valence shell electrons are electropositive, e.g., sodium Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule Not true bond Common between dipoles such as water Also act as intramolecular bonds, holding a large molecule in a threedimensional shape Chemical Reactions Occur when chemical bonds are formed, rearranged, or broken Represented as chemical equations using molecular formulas Subscript indicates atoms joined by bonds Prefix denotes number of unjoined atoms or molecules Chemical equations contain Reactants Number and kind of reacting substances Chemical composition of the product(s) Relative proportion of each reactant and product in balanced equations 38 Synthesis (combination) reactions Decomposition reactions Exchange reactions 39 A + B AB Atoms or molecules combine to form larger, more complex molecule Always involve bond formation Anabolic AB A + B Molecule is broken down into smaller molecules or its constituent atoms Reverse of synthesis reactions Involve breaking of bonds Catabolic AB + C AC + B Also called displacement reactions Involve both synthesis and decomposition Bonds are both made and broken Are decomposition reactions Reactions in which food fuels are broken down for energy Are also exchange reactions because electrons are exchanged between reactants Electron donors lose electrons and are oxidized Electron acceptors receive 43 electrons and become reduced C6H12O6 + 6O2 6CO2 + 6H2O + ATP Glucose is oxidized; oxygen molecule is reduced 43 All chemical reactions are either exergonic or endergonic Exergonic reactions—net release of energy Products have less potential energy than reactants Catabolic and oxidative reactions Endergonic reactions—net absorption of energy Products have more potential energy than reactants Anabolic reactions 44 All chemical reactions are theoretically reversible A + B AB AB A + B Chemical equilibrium occurs if neither a forward nor reverse reaction is dominant Many biological reactions are essentially irreversible Due to energy requirements Due to removal of products 45 Affected by Temperature Rate Concentration of reactant Rate Particle size Rate Catalysts: Rate without being chemically changed or part of product Enzymes are biological catalysts 12/10/11 • Study of chemical composition and reactions of living matter • All chemicals either organic or inorganic 47 12/10/11 • Inorganic compounds •Water, salts, and many acids and bases •Do not contain carbon • Organic compounds •Carbohydrates, fats, proteins, and nucleic acids •Contain carbon, usually large, and are covalently 48 bonded • Both equally essential for life 48 12/10/11 • Most abundant inorganic compound – • 60%–80% volume of living cells Most important inorganic compound – Due to water’s properties 49 12/10/11 • • High heat capacity – Absorbs and releases heat with little temperature change – Prevents sudden changes in temperature High heat of vaporization – Evaporation requires large amounts of heat – Useful cooling mechanism 50 12/10/11 • Polar solvent properties – Dissolves and dissociates ionic substances – Forms hydration layers around large charged molecules, e.g., proteins (colloid formation) – Body’s major transport medium 51 12/10/11 • Reactivity – • Necessary part of hydrolysis and dehydration synthesis reactions Cushioning – Protects certain organs from physical trauma, e.g., cerebrospinal fluid 52 12/10/11 • Both are electrolytes – Ionize and dissociate in water • Acids are proton donors – Release H+ (a bare proton) in solution – HCl H+ + Cl– • Bases are proton acceptors 53 – Take up H+ from solution •NaOH Na+ + OH– –OH– accepts an available proton (H+) –OH– + H+ H2O 53 12/10/11 • Ionic compounds that dissociate into ions in water – Ions (electrolytes) conduct electrical currents in solution – Ions play specialized roles in body functions (e.g., sodium, potassium, calcium, and iron) – Ionic balance vital for homeostasis • Contain cations other than H+ 54 and anions other than OH– • Common salts in body – NaCl, CaCO3, KCl, calcium phosphates 54 12/10/11 • pH = negative logarithm of [H+] in moles per liter • pH scale ranges from 0–14 • Because pH scale is logarithmic – – A pH 5 solution is 10 times more acidic than a pH 6 solution Relative free [H+] of a solution measured on pH scale – As free [H+] increases, acidity increases •[OH–] decreases as [H+] increases •pH decreases – As free [H+] decreases alkalinity increases •[OH–] increases as [H+] decreases •pH increases 55 12/10/11 • Acidic solutions [H+], pH – Acidic pH: 0–6.99 • Neutral solutions – Equal numbers of H+ and OH– – All neutral solutions are pH 7 – Pure water is pH neutral •pH of pure water = pH 7: [H+] = 10–7 m 56 • Alkaline (basic) solutions [H+], pH – Alkaline pH: 7.01–14 56 12/10/11 • Results from mixing acids and bases – Displacement reactions occur forming water and a salt – Neutralization reaction •Joining of H+ and OH– to form water neutralizes solution • Acid-base Homeostasis • pH change interferes with cell function and may damage living tissue • Even slight change in pH can be fatal • pH is regulated by kidneys, lungs, and chemical buffers 57 12/10/11 • Acidity reflects only free H+ in solution – Not those bound to anions • Buffers resist abrupt and large swings in pH – Release hydrogen ions if pH rises – Bind hydrogen ions if pH falls • Convert strong (completely dissociated) acids • Carbonic acid-bicarbonate system (important 58 12/10/11 • Molecules that contain carbon – Except CO2 and CO, which are considered inorganic – Carbon is electroneutral •Shares electrons; never gains or loses them •Forms four covalent bonds with other elements • Unique to living systems • Carbohydrates, lipids, proteins, and nucleic acids 59 12/10/11 • Many are polymers – Chains of similar units called monomers (building blocks) • Synthesized by dehydration synthesis • Broken down by hydrolysis reactions 60 12/10/11 • Sugars and starches • Polymers • Contain C, H, and O [(CH20)n] • Three classes • – Monosaccharides – one sugar – Disaccharides – two sugars – Polysaccharides – many sugars Functions of carbohydrates – Major source of cellular fuel (e.g., glucose) – Structural molecules (e.g., ribose sugar in RNA) 61 12/10/11 • • • • Simple sugars containing three to seven C atoms (CH20)n – general formula; n = # C atoms Monomers of carbohydrates Important monosaccharides – Pentose sugars •Ribose and deoxyribose – Hexose sugars •Glucose (blood sugar) 62 12/10/11 • Double sugars • Too large to pass through cell membranes • Important disaccharides – Sucrose, maltose, lactose 63 12/10/11 • Polymers of monosaccharides • Important polysaccharides – • Starch and glycogen Not very soluble 64 12/10/11 • Contain C, H, O (less than in carbohydrates), and sometimes P • Insoluble in water • Main types: – Triglycerides or neutral fats – Phospholipids – Steroids – Eicosanoids 65 12/10/11 • Called fats when solid and oils when liquid • Composed of three fatty acids bonded to a glycerol molecule • Main functions – Energy storage – Insulation – Protection 66 12/10/11 • Saturated fatty acids – Single covalent bonds between C atoms •Maximum number of H atoms – Solid animal fats, e.g., butter • Unsaturated fatty acids – One or more double bonds between C atoms •Reduced number of H atoms 67 – Plant oils, e.g., olive oil – “Heart healthy” • Trans fats – modified oils – unhealthy • Omega-3 fatty acids – “heart healthy” 67 12/10/11 • Modified triglycerides: – Glycerol + two fatty acids and a phosphorus (P) - containing group • “Head” and “tail” regions have different properties • Important in cell membrane structure 68 12/10/11 • Steroids—interlocking four-ring structure • Cholesterol, vitamin D, steroid hormones, and bile salts • Most important steroid – Cholesterol •Important in cell membranes, vitamin D synthesis, steroid hormones, and bile s 69 12/10/11 • Many different ones • Derived from a fatty acid (arachidonic acid) in cell membranes • Most important eicosanoid – Prostaglandins •Role in blood clotting, control of blood pressure, inflammation, and labor contractions 70 12/10/11 • Other fat-soluble vitamins – • Vitamins A, D, E, and K Lipoproteins – Transport fats in the blood 71 12/10/11 • Contain C, H, O, N, and sometimes S an • Proteins are polymers • Amino acids (20 types) are the monom – – – – Joined by covalent bonds called peptide b Contain amine group and acid group Can act as either acid or base All identical except for “R group” (in green 72 12/10/11 73 12/10/11 74 12/10/11 • • Fibrous (structural) proteins – Strandlike, water-insoluble, and stable – Most have tertiary or quaternary structure (3-D) – Provide mechanical support and tensile strength – Examples: keratin, elastin, collagen (single most abundant protein in body), and certain contractile fibers Globular (functional) proteins – Compact, spherical, water-soluble and sensitive to environmental changes – Tertiary or quaternary structure (3-D) – Specific functional regions (active sites) – Examples: antibodies, hormones, molecular chaperones, and enzymes 75 12/10/11 Quaternary multiple separate polypeptides 76 12/10/11 • Denaturation – Globular proteins unfold and lose functional, 3-D shape •Active sites destroyed – Can be cause by decreased pH or increased temperature • Usually reversible if normal conditions restored • Irreversible if changes extreme – e.g., cooking an egg 77 12/10/11 • Globular proteins • Ensure quick, accurate folding and association of other proteins • Prevent incorrect folding • Assist translocation of proteins and ions across membranes • Promote breakdown of damaged or denatured proteins • Help trigger the immune response • Stress proteins – Molecular chaperones produced in response to stressful stimuli, e.g., O2 deprivation – Important to cell function during stress – Can delay aging by patching up damaged proteins and refolding them 78 12/10/11 • Enzymes – Globular proteins that act as biological catalysts •Regulate and increase speed of chemical reactions 79 12/10/11 Lower the activation energy, increase the speed of a reaction (millions of reactions per minute! 80 12/10/11 • • • • Some functional enzymes (holoenzymes) consist of two parts – Apoenzyme (protein portion) – Cofactor (metal ion) or coenzyme (organic molecule often a vitamin) Enzymes are specific – Act on specific substrate Usually end in -ase Often named for the reaction they catalyze – Hydrolases, oxidases 81 12/10/11 82 12/10/11 • Deoxyribonucleic acid (DNA) and ribonucleic acid (RNA) – Largest molecules in the body • Contain C, O, H, N, and P • Polymers – Monomer = nucleotide •Composed of nitrogen base, a pentose sugar, and a phosphate group 83 12/10/11 • Utilizes four nitrogen bases: – Purines: Adenine (A), Guanine (G) – Pyrimidines: Cytosine (C), and Thymine (T – Base-pair rule – each base pairs with its co •A always pairs with T; G always pairs with C • • • • Double-stranded helical molecule (dou Pentose sugar is deoxyribose Provides instructions for protein synthe Replicates before cell division ensuring 84 12/10/11 • • • • Four bases: – Adenine (A), Guanine (G), Cytosine (C), and Uracil (U) Pentose sugar is ribose Single-stranded molecule mostly active outside the nucleus Three varieties of RNA carry out the DNA orders for protein synthesis – Messenger RNA (mRNA), transfer RNA (tRNA), and ribosomal RNA (rRNA) 85 12/10/11 • Chemical energy in glucose captured in this important molecule • Directly powers chemical reactions in cells • Energy form immediately useable by all body cells • Structure of ATP – Adenine-containing RNA nucleotide with two additional phosphate groups 86 12/10/11 • Phosphorylation – Terminal phosphates are enzymatically transferred to and energize other molecules – Such “primed” molecules perform cellular work (life processes) using the phosphat 87