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Anatomy & Physiology I Chapter 2 Science that deals with composition and properties of matter Necessary to understand normal and abnormal functioning of body Biochemistry – the study of the molecules that compose living organisms carbohydrates, fats, proteins, and nucleic acids Matter - Anything that has mass and occupies space. All matter is composed of elements. Elements Make up ALL matter Cannot be broken down by ordinary chemical means Each has unique physical and chemical properties 92 occur in nature (24 elements have biological role) Identified by names, chemical symbols or number Described and organized in periodic table Oxygen (O) Carbon (C) Hydrogen (H) Nitrogen (N) About 96% of body mass About 3.9% of body mass: Calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe) Trace Elements (< 0.01% of body mass): Part of enzymes, e.g., chromium (Cr), manganese (Mn), and zinc (Zn) Atoms Unique building blocks for each element Smallest complete units of matter Cannot be broken down or changed by ordinary chemical and physical means Determined by numbers of subatomic particles Protons (positive charge [+]); in nucleus Neutrons (no charge); in nucleus Electrons (negative charge [–]); in concentric clouds that surround nucleus determine the chemical properties of an atom the atom is electrically neutral because number of electrons is equal to the number of protons valence electrons in the outermost shell determine chemical bonding properties of an atom Atoms of different elements contain different numbers of subatomic particles Compare hydrogen, helium and lithium (next slide) Proton Neutron Electron Hydrogen (H) (1p+; 0n0; 1e–) Helium (He) (2p+; 2n0; 2e–) Lithium (Li) (3p+; 4n0; 3e–) Isotopes – varieties of an element that differ from one another only in the number of neutrons and therefore in atomic mass extra neutrons increase atomic weight isotopes of an element are chemically similar have same valence electrons Atomic weight of an element accounts for the fact that an element is a mixture of isotopes 2-10 Hydrogen (1H) (1p+, 0n0, 1e–) Deuterium (2H) (1p+, 1n0, 1e–) Key = Proton = Neutron = Electron Tritium (3H) (1p+, 2n0, 1e–) Isotopes same chemical behavior, differ in physical behavior breakdown (decay) to more stable isotope by giving off radiation Radioisotopes unstable isotopes that give off radiation every element has at least one radioisotope Similar chemistry to stable isotopes Can be detected with scanners Radioactivity radioisotopes decay to stable isotopes releasing radiation we are all mildly radioactive • Ions – charged particles with unequal number of protons and electrons Ionization - transfer of electrons from one atom to another ( stability of valence shell) 11 protons 12 neutrons 11 electrons Sodium atom (Na) 17 protons 18 neutrons 17 electrons Chlorine atom (Cl) Transfer of an electron from a sodium atom to a chlorine atom • Anion – atom that gained electrons (net negative charge) • Cation – atom that lost an electron (net positive charge) • Ions with opposite charges are attracted to each other + – 11 protons 12 neutrons 10 electrons Sodium ion (Na+) 17 protons 18 neutrons 18 electrons Chloride ion (Cl–) Sodium chloride 2 The charged sodium ion (Na+) and chloride ion (Cl–) that result Salts that ionize in water and form solutions capable of conducting an electric current. Electrolyte importance chemical reactivity osmotic effects (influence water movement) electrical effects on nerve and muscle tissue Electrolyte balance is one of the most important considerations in patient care. Imbalances have ranging effects from muscle cramps, brittle bones, to coma and cardiac arrest Chemical particles with an odd number of electrons Produced by normal metabolic reactions, radiation, chemicals Causes tissue damage reactions that destroy molecules causes cancer, death of heart tissue and aging Antioxidants neutralize free radicals in body, superoxide dismutase (SOD) in diet (Selenium, vitamin E, vitamin C, carotenoids) Molecules Formed when two or more atoms unite on the basis of their electron structures Can be made of like atoms or atoms of different elements Compounds Composed of two or more elements Smallest subunits of a compound are molecules Most matter exists as mixtures Two or more components physically intermixed Three types of mixtures Solutions Colloids Suspensions Homogeneous mixtures Usually transparent, e.g., atmospheric air or seawater Solvent Present in greatest amount, usually a liquid Solute(s) Present in smaller amounts Colloids (emulsions) Heterogeneous translucent mixtures, e.g., cytosol Large solute particles that do not settle out Undergo sol-gel transformations Suspensions: Heterogeneous mixtures, e.g., blood Large visible solutes tend to settle out Solution Colloid Suspension Solute particles are very tiny, do not settle out or scatter light. Solute particles are larger than in a solution and scatter light; do not settle out. Solute particles are very large, settle out, and may scatter light. Solute particles Solute particles Solute particles Example Example Example Mineral water Gelatin Blood Mixtures No chemical bonding between components Can be separated physically, such as by straining or filtering Heterogeneous or homogeneous Compounds Can be separated only by breaking bonds All are homogeneous Electrons occupy up to seven electron shells (energy levels) around nucleus Octet rule: Except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their outermost energy level (valence shell) Stable and unreactive Outermost energy level fully occupied or contains eight electrons (a) Chemically inert elements Outermost energy level (valence shell) complete 8e 2e Helium (He) (2p+; 2n0; 2e–) 2e Neon (Ne) (10p+; 10n0; 10e–) Outermost energy level not fully occupied by electrons Tend to gain, lose, or share electrons (form bonds) with other atoms to achieve stability (b) Chemically reactive elements Outermost energy level (valence shell) incomplete 1e Hydrogen (H) (1p+; 0n0; 1e–) 6e 2e Oxygen (O) (8p+; 8n0; 8e–) 4e 2e Carbon (C) (6p+; 6n0; 6e–) 1e 8e 2e Sodium (Na) (11p+; 12n0; 11e–) Ionic Covalent Hydrogen Ions are formed by transfer of valence shell electrons between atoms Anions (– charge) have gained one or more electrons Cations (+ charge) have lost one or more electrons Attraction of opposite charges results in an ionic bond Sodium atom (Na) (11p+; 12n0; 11e–) Chlorine atom (Cl) (17p+; 18n0; 17e–) + – Sodium ion (Na+) Chloride ion (Cl–) Sodium chloride (NaCl) (a) Sodium gains stability by losing one electron, and chlorine becomes stable by gaining one electron. 1. 2. 3. (b) After electron transfer, the oppositely charged ions formed attract each other. An electron is transferred from the Na atom to Cl atom. The sodium (Na) becomes a cation (Na+) and the chlorine (Cl) becomes an anion (Cl-). The opposite charges attract forming an ionic bond. In step 2, why does sodium become positive and chlorine negative? CI– Na+ (c) Large numbers of Na+ and Cl– ions associate to form salt (NaCl) crystals. Formed by two atoms sharing valence electrons Types of covalent bonds single - sharing of single pair electrons double - sharing of 2 pairs of electrons nonpolar covalent bond shared electrons spend approximately equal time around each nucleus strongest of all bonds polar covalent bond if shared electrons spend more time orbiting one nucleus than they do the other, they lend their negative charge to the area they spend most time Reacting atoms Resulting molecules + Hydrogen atoms or Carbon atom Molecule of methane gas (CH4) Structural formula shows single bonds. Formation of four single covalent bonds: Each hydrogen atom shares its electron with carbon while carbon shares one of its valence electrons with each hydrogen atom. Reacting atoms Resulting molecules + Oxygen atom or Oxygen atom Molecule of oxygen gas (O2) Structural formula shows double bond. Formation of a double covalent bond: Each oxygen atom shares two electrons with the other oxygen atom. Sharing of electrons may be equal or unequal Equal sharing produces electrically balanced nonpolar molecules CO2 Unequal sharing by atoms with different electron-attracting abilities produces polar molecules H2O Atoms with six or seven valence shell electrons are electronegative, e.g., oxygen Atoms with one or two valence shell electrons are electropositive, e.g., sodium Attraction between a slightly positive hydrogen atom in one molecule and a slightly negative oxygen or nitrogen atom in another. Water molecules are weakly attracted to each other by hydrogen bonds very important to physiology protein structure DNA structure++ H O H H H O O H H O H Covalent bond H Hydrogen bond O Water molecule H H Water’s polar covalent bonds and its V-shaped molecule gives water a set of properties that account for its ability to support life. 60%–80% of the volume of living cells Most important inorganic compound in living organisms because of its properties: solvency cohesion adhesion chemical reactivity thermal stability Solvency - ability to dissolve other chemicals water is called the Universal Solvent Hydrophilic – substances that dissolve in water molecules must be polarized or charged Hydrophobic - substances that do not dissolve in water molecules are non-polar or neutral (fat) Virtually all metabolic reactions depend on the solvency of water A water strider can walk on a pond because of the high surface tension of water, a result of the combined strength of its hydrogen bonds. is the ability to participate in chemical reactions water ionizes into H+ and OH- water ionizes other chemicals (acids and salts) water involved in hydrolysis and dehydration synthesis reactions Occur when chemical bonds are formed, rearranged, or broken Represented as chemical equations Chemical equations contain: Molecular formula for each reactant and product Relative amounts of reactants and products, which should balance H + H H2 (hydrogen gas) (reactants) (product) 4H + C CH4 (methane) Synthesis (combination) reactions Decomposition reactions Exchange reactions Synthesis reactions Smaller particles are bonded together to form larger, more complex molecules. • A + B AB • Always involve bond formation • Anabolic Example Amino acids are joined together to form a protein molecule. Amino acid Molecules (monomers) Protein Molecule (polymer) Decomposition reactions • AB A + B • Reverse synthesis reactions • Involve breaking of bonds Bonds are broken in larger molecules, resulting in smaller, less complex molecules. Example Glycogen is broken down to release glucose units. Glycogen (polymer) • Catabolic Glucose Molecules (monomers) AB + C AC + B Also called displacement reactions Bonds are both made and broken An acid is proton donor (releases H+ ions in water) A base is proton acceptor (accepts H+ ions) releases OH- ions in water pH – a measure derived from the molarity of H+ a pH of 7.0 is neutral pH (H+ = OH-) a pH of less than 7 is acidic solution (H+ > OH-) a pH of greater than 7 is basic solution (OH- > H+ ) Acids and bases are electrolytes pH - measurement of molarity of H+ [H+] on a logarithmic scale a change of one number on the pH scale represents a 10 fold change in H+ concentration a solution with pH of 4.0 is 10 times as acidic as one with pH of 5.0 Our body uses buffers to resist changes in pH slight pH disturbances can disrupt physiological functions and alter drug actions pH of blood ranges from 7.35 to 7.45 deviations from this range cause tremors, paralysis or even death Gastric juice (0.9–3.0) Lemon juice 1M (2.3) Hydrochloric Acid (0) Wine, Bananas, vinegar tomatoes (2.4 -–3.5) (4.7) 3 2 1 0 Bread, black coffee (5.0) 4 Pure water Milk, Egg white (7.0) saliva (8.0) (6.3 -–6.6) Household bleach (9.5) Household ammonia (10.5 - 11.0) Oven cleaner, lye (13.4) 1 M sodium hydroxide (14) 5 6 7 Neutral 8 9 10 11 12 13 14 Acid solutions contain [H+] As [H+] increases, acidity increases Alkaline solutions contain bases (e.g., OH–) As [H+] decreases (or as [OH–] increases), alkalinity increases Neutral solutions: Pure water is pH neutral (contains equal numbers of H+ and OH–) All neutral solutions are pH 7 pH change interferes with cell function and may damage living tissue Slight change in pH can be fatal pH is regulated by kidneys, lungs, and buffers Mixture of compounds that resist pH changes Convert strong acids or bases into weak ones Carbonic acid-bicarbonate system Classes of Compounds 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 bonded Must contain at least one carbon and one hydrogen atom. (CO2 and CO are inorganic because they have no hydrogen atom.) Unique to living systems Include carbohydrates, lipids, proteins, and nucleic acids (Molecules of Life) Many are polymers—chains of similar units (monomers or building blocks) Synthesized by dehydration synthesis Broken down by hydrolysis reactions Macromolecules - very large organic molecules proteins, DNA Polymers – molecules made of a repetitive series of identical or similar subunits (monomers) Monomers - an identical or similar subunits joining monomers to form a polymer dehydration synthesis (condensation) is how living cells form polymers a hydroxyl (-OH) group is removed from one monomer, and a hydrogen (H+) from another producing water as a by-product hydrolysis – opposite of dehydration synthesis; breaks down polymers to monomers a water molecule ionizes into –OH and H+ the covalent bond linking one monomer to the other is broken the –OH is added to one monomer the H+ is added to the other Monomers covalently bond together to form a polymer with the removal of a water molecule A hydroxyl group is removed from one monomer and a hydrogen from the next Dimer Monomer 1 Monomer 2 OH O HO H+ + OH– (a) Dehydration synthesis H2 O Splitting a polymer (lysis) by the addition of a water molecule (hydro) a covalent bond is broken All digestion reactions consists of hydrolysis reactions Dimer Monomer 1 OH O H2 O (b) Hydrolysis Monomer 2 H+ + OH– HO Sugars and starches Contain C, H, and O [(CH20)n] Three classes Monosaccharides Disaccharides Polysaccharides Building blocks of carbohydrates: monosaccharides Simple sugars containing three to seven C atoms (CH20)n Glucose – blood sugar; major source of cellular fuel Fructose - found in fruits, vegetables and honey Galactose – component of some glycolipids (i.e. the ABO antigens that determine blood type) Ribose in RNA & Deoxyribose in DNA Double sugars (made of 2 monosaccharides) Maltose – malt sugar (used to make beer) made of 2 glucose monomers Sucrose – table sugar (cane sugar) made of glucose and fructose Lactose – milk sugar made of glucose and galactose (b) Disaccharides Consist of two linked monosaccharides Example Sucrose, maltose, and lactose (these disaccharides are isomers) Glucose Fructose Sucrose Glucose Glucose Maltose Galactose Glucose Lactose Polymers of many simple sugars Starch – storage form of excess glucose in plants (made of several hundred to several thousand glucose units) Glycogen – storage form of excess glucose in animals (made of several thousand glucose units) The liver and skeletal muscles are major depots of glycogen. Glycogen is broken back down into glucose when energy is needed (a process called glycogenolysis). Cellulose – cell walls of plants; provides dietary fiber, roughage (made of approx. 10,000 glucose units) (c) Polysaccharides Long branching chains (polymers) of linked monosaccharides Example This polysaccharide is a simplified representation of glycogen, a polysaccharide formed from glucose units. Glycogen 2-70 Contain C, H, O (less than in carbohydrates), and sometimes P Insoluble in water Main types: Triglycerides (neutral fats) Phospholipids Steroids Eicosanoids Building blocks of lipids: glycerol and fatty acids Chain of 4 to 24 carbon atoms carboxyl (acid) group on one end, methyl group on the other and hydrogen bonded along the sides Classified saturated - carbon atoms saturated with hydrogen unsaturated - contains C=C bonds without hydrogen polyunsaturated – contains many C=C bonds essential fatty acids – obtained from diet, body can not synthesize (saturated) 3 fatty acids covalently bonded to three carbon alcohol, glycerol molecule each bond formed by dehydration synthesis broken down by hydrolysis triglycerides at room temperature when liquid called oils often polyunsaturated fats from plants when solid called fat saturated fats from animals Primary Function - energy storage, insulation and shock absorption (adipose tissue) similar to neutral fat except that one fatty acid replaced by a phosphate group CH3 N+ CH3 Nitrogencontaining group (choline) CH2 structural foundation of cell membrane CH2 O –O Amphiphilic fatty acid “tails” are hydrophobic phosphate “head” is hydrophilic CH3 P O Phosphate group Hydrophilic region (head) O Glycerol O CH2 CH O O C C (CH2)5 CH CH (CH2)5 CH3 (CH2)12 CH3 CH2 O Fatty acid tails Hydrophobic region (tails) Steroid – a lipid with 17 of its carbon atoms in four rings Cholesterol - the ‘parent’ steroid from which the other steroids are synthesized cortisol, progesterone, estrogens, testosterone and bile acids Cholesterol synthesized only by animals especially liver cells 15% from diet, 85% internally synthesized important component of cell membranes one kind of cholesterol does far more good than harm ‘good’ and ‘bad’ cholesterol actually refers to droplets of lipoprotein in the blood complexes of cholesterol, fat, phospholipid, and protein HDL – high-density lipoprotein – “good” cholesterol lower ratio of lipid to protein may help to prevent cardiovascular disease LDL – low-density lipoprotein – “bad” cholesterol high ratio of lipid to protein contributes to cardiovascular disease Simplified structure of a steroid Four interlocking hydrocarbon rings form a steroid. Example Cholesterol (cholesterol is the basis for all steroids formed in the body) 20 carbon compounds derived from a fatty acid called arachidonic acid hormone-like chemical signals between cells includes prostaglandins – produced in all tissues role in inflammation, blood clotting, hormone action, labor contractions, blood vessel diameter O COOH OH OH 2-79 Polymers of amino acids (20 types) Joined by peptide bonds Contain C, H, O, N, and sometimes S and P Building blocks of proteins: amino acids peptide – any molecule composed of two or more amino acids joined by peptide bonds peptide bond – joins two amino acids formed by dehydration synthesis Peptides named for the number of amino acids dipeptides have 2 tripeptides have 3 oligopeptides have fewer than 10 to 15 polypeptides have more than 15 proteins have more than 50 Amine group Acid group (a) Generalized structure of all amino acids. (b) Glycine is the simplest amino acid. (c) Aspartic acid (d) Lysine (an acidic amino acid) (a basic amino acid) has an acid group has an amine group (—COOH) in the (–NH2) in the R group. R group. (e) Cysteine (a basic amino acid) has a sulfhydryl (–SH) group in the R group, which suggests that this amino acid is likely to participate in intramolecular bonding. • There are 20 naturally occurring amino acids. • All 20 amino acids are identical except for their R group. (notice the structures highlighted green) Dehydration synthesis: The acid group of one amino acid is bonded to the amine group of the next, with loss of a water molecule. Peptide bond + Amino acid Amino acid Hydrolysis: Peptide bonds linking amino acids together are broken when water is added to the bond. Dipeptide Amino acid Amino acid Amino acid Amino acid Amino acid The sequence of amino acids forms the polypeptide chain. Shape change and disruption of active sites due to environmental changes (e.g., decreased pH or increased temperature) Reversible in most cases, if normal conditions are restored Irreversible if extreme changes damage the structure beyond repair (e.g., cooking an egg) Enzymes - proteins that function as biological catalysts permit reactions to occur rapidly at normal body temperature Substrate - substance an enzyme acts upon Naming Convention named for substrate with -ase as the suffix amylase enzyme digests starch (amylose) All enzymes are proteins Substrates (S) e.g., amino acids + Product (P) e.g., dipeptide Energy is absorbed; bond is formed. Water is released. H2O Peptide bond Active site Enzyme (E) Enzyme-substrate complex (E-S) 1 Substrates bind 2 Internal at active site. rearrangements Enzyme changes leading to shape to hold catalysis occur. substrates in proper position. Enzyme (E) 3 Product is released. Enzyme returns to original shape and is available to catalyze another reaction. DNA and RNA Largest molecules in the body Contain C, O, H, N, and P Building block = nucleotide, composed of nitrogenous base, a pentose sugar, and a phosphate group Four bases: adenine (A), guanine (G), cytosine (C), and thymine (T) Double-stranded helical molecule in the cell nucleus Provides instructions for protein synthesis Replicates before cell division, ensuring genetic continuity Phosphate Sugar: Deoxyribose Base: Adenine (A) Thymine (T) Adenine nucleotide Sugar Phosphate Thymine nucleotide Hydrogen bond (a) Sugar-phosphate backbone Deoxyribose sugar Phosphate Adenine (A) Thymine (T) Cytosine (C) Guanine (G) (b) (c) Computer-generated image of a DNA molecule Four bases: adenine (A), guanine (G), cytosine (C), and uracil (U) Single-stranded molecule mostly active outside the nucleus Three varieties of RNA carry out the DNA orders for protein synthesis messenger RNA, transfer RNA, and ribosomal RNA Adenine-containing RNA nucleotide with two additional phosphate groups High-energy phosphate bonds can be hydrolyzed to release energy. Adenine Phosphate groups Ribose Adenosine Adenosine monophosphate (AMP) Adenosine diphosphate (ADP) Adenosine triphosphate (ATP) Phosphorylation: Terminal phosphates are enzymatically transferred to and energize other molecules Such “primed” molecules perform cellular work (life processes) using the phosphate bond energy Solute + Membrane protein (a) Transport work: ATP phosphorylates transport proteins, activating them to transport solutes (ions, for example) across cell membranes. + Relaxed smooth muscle cell Contracted smooth muscle cell (b) Mechanical work: ATP phosphorylates contractile proteins in muscle cells so the cells can shorten. + (c) Chemical work: ATP phosphorylates key reactants, providing energy to drive energy-absorbing chemical reactions.