1. Chapter 2 The Chemical Level of Organization

2. HOW MATTER IS ORGANIZED • Chemical Elements • substances that cannot be split into simpler substances • 112 elements • O, C, H, N, Ca, and P make up 98.5% of total body weight • Trace elements are present in tiny amounts • copper, tin, selenium & zinc

3. Structure of Atoms • All elements = atoms of same type • Subatomic particles • Nucleus • protons (p+) • neutrons (n0) • Electrons (e-) move about nucleus in energy levels • In neutral atom, # e- = # p+ • Atomic Number (Z) • # protons in nucleus • Identifies atom

4. ISOTOPES • Atoms of an element w/ same # of protons but different # of neutrons • Isotopes • Stable isotopes do not change nuclear structure over time • Radioactive isotopes • Unstable  nuclei decay to form simpler & more stable configuration • Assessment of internal abnormalities

5. Ions & Molecules • Ions form when an atom gives up or gains electrons • (+) or (-) charge due to unequal # of p+ and e- • Goal: atomic stability • Molecule results from two or more atoms sharing electrons • Ex: H2, N2, O2, CO2, H2O

6. Free Radicals • Electrically charged atom/molecule w/ unpaired electron • Unstable & highly reactive  chain reactions • Can become stable • giving up an electron • taking an electron from another molecule • Antioxidants inactivate oxygen-derived free radicals • Ex: superoxide radical = oxygen w/ extra electron • can induce tissue damage if chain rxn allowed to propagate

7. Free Radicals & Your Health • Possible sources: absorption of UV energy in sunlight, x-rays, breakdown of harmful substances, & normal metabolic reactions • Cancer, diabetes, Alzheimer, atherosclerosis and arthritis • Dietary antioxidants: vitamins C and E, selenium & beta-carotene (precursor to vitamin A)

8. CHEMICALBONDS • Forces of attraction holding atoms of compound together • Valence e- determine: • type of bonding • chemical stability • 8 e- in outer shell = stable • <8 e- in outer shell  gain/lose/share e- • octet rule

9. Ionic Bonds • Loss or gain of valence electron results in ion formation • Oppositely charged ions attracted to one another • Cations • Anions • Electrolytes • Ex: NaCl in Fig 2.4

10. Covalent Bonds • Formed from sharing one, two, or three pairs of valence e- • Strongest chemical bonds in the body • Single, double, or triple covalent bonds • Bond polarity • Nonpolar covalent bond • Equal sharing of electrons • Polar covalent bond • Unequal sharing of electrons • Electronegativity difference • N—H & C—O

11. Polar Covalent Bonds • Unequal sharing of electrons between atoms • Different centers of positive & negative charge • In a water molecule, O attracts H electrons more strongly • Oxygen has greater EN (indicated by negative delta sign) • Overall polarity of molecule is in direction of oxygen • See Fig 2.6

12. HydrogenBonds • Special polar covalent bonds btwn H atom & electronegative atom • N…H or O…H • Very weakintermolecular bonds • Cohesive properties of water • Occur between δ+ H of one H2O & δ- O of another H2O (See Fig 2.7)

13. Chemical Reactions • New bonds form and/or old bonds are broken • Metabolism = sum of all chemical reactions in the body • Law of conservation of mass • Total mass of reactants equals total mass of products

14. Forms of Energy • Energy = capacity to do work • Kinetic energy = energy of motion • Temperature • Potential energy = energy stored by matter due to its position • Chemical energy • Energy: • Conserved in rxn • May be converted

15. Energy Transfer • Exergonic reaction • bond broken has more energy than one formed • extra energy is released • usually as heat • catabolism of food molecules • Endergonic reaction • requires energy be added to form a bond • usually from a molecule of ATP • EX: building proteins from amino acids

16. Energy Transfer in Chemical Reactions • In living systems, ender- & exergonic reactions occur together • Coupled reactions essential to metabolism • energy released from one reaction drives another • Ex: glucose breakdown releases energy, which is used to build ATP molecules • Ex: ATP fuels transport across membranes, muscle contraction & nerve impulses

17. Activation Energy • Energy needed to break bonds & begin reaction (Fig 2.9) • Increasing probability of collision increases chance for reaction • Increasing concentration & temperature are ways of overcoming Ea, thus ↑ chances for collision • more particles are in a given space • particles move more rapidly

18. Factors Influencing Chemical Rxns • Concentration • Temperature • Catalysts • speed up chemical reactions by lowering amount of energy needed to get reaction started (activation energy, Ea) • do not alter difference in potential energy between the reactants & products • orient colliding particles • unchanged at end of reaction  often re-used many times • relevance??? Biological enzymes are catalysts!

19. Effectiveness of Catalysts Difference in PE

20. Types of Chemical Reactions • Synthesis • > two atoms/ions/molecules combine to form new & larger molecules • anabolic reactions (bonds are formed) A + B  AB • generally endergonic • Decomposition • a molecule is broken down into smaller parts • catabolic reactions (bonds are broken) AB  A + B • usually exergonic

21. Reversible Reactions • Chemical reactions can be reversible • Indicated by the 2 arrows pointing in opposite directions between the reactants and the products AB A + B

22. Water • Most important & abundant inorganic compound in all living systems • Polarity makes it a good solvent  almost “universal” solvent • Hydrophiliccompounds • Usually are polar • Dissolve in water • Hydrophobic compounds • Usually nonpolar • Do not dissolve in water • Excellent medium for metabolic reactions of the body

23. Water as a Solvent • Polar covalent bonds (hydrophilic vs. hydrophobic) • Dissolves or suspends many substances • Each water molecule interact w/ 4 ormore neighboring ions/molecules • Hydration spheres • Fig. 2.11 shows how water’s shape makes it such an effective solvent.

24. Water in Chemical Reactions • Hydrolysis: add’n of water breaks molecules apart • Dehydration synthesis • two simple molecules join together • eliminate a molecule of water in process • High heat capacity • Resists changes in temperature  maintain body temp • Due to hydrogen bonding • High heat of vaporization • amount of heat needed to change from liquid to gas • evaporation of water from skin removes lots of heat  why sweat cools you

25. Water as a Lubricant • Major component of mucus & other lubricating fluids • mucus in respiratory and digestive systems • synovial fluid in joints • serous fluids in chest and abdominal cavities • organs slide past one another • Found wherever friction needs to be reduced or eliminated

26. Inorganic Acids, Bases & Salts • Dissociate into ions in water • Acids: H+ + A- HCl  H+ + Cl- • Bases: OH- + cation NaOH  Na+ + OH- • Acid + base  salt & H20 • HCl + NaOH  NaCl + H2O • Salts dissociate into cations & anions in water • metal and nonmetal ions: NaCl + H2O  Na+ + Cl- • not H+ or OH- !! • Electrolytes • important salts in body (Na, Cl, K) • carry electric current (in nerve or muscle)

27. Acid-Base Balance & pH • pH: measure of [H+] in moles/liter (M) • pH scale: 0-14 • pH = 7  neutral  [H+] = [OH-] • pH < 7  acidic  [H+] > [OH-] • pH > 7  alkaline  [H+] < [OH-] • A solution’s acidity or alkalinity is based on the pH scale • Biochemical reactions are very sensitive to even small changes in pH • pH of blood is 7.35 to 7.45

28. The Concept of pH • pH is a logarithmic scale—it is NOT linear! • Therefore—each unit in scale means 10-fold Δ in [H+] • Ex: a change of two pH units represents 100-fold diff in [H+]  • pH 1 contains 10-1 M H+ & pH 3 contains 10-3 M H+ • the diff in H+ ion concentration is 100—not 2! • Ex: pH 8 vs. pH 11 • pH 8 = 10-8 M H+ & pH 11 = 10 -11 M H+ • pH 8 is 1000x more acidic than pH 11 (even tho both are basic!)

29. Maintaining pH: Buffer Systems • pH in body maintained fairly constant by buffer systems • Buffers resist Δ in pH even when acid/base added • consist of a weak acid & a weak base • convert strong acids/bases into weak acids/bases • Ex: carbonic acid-bicarbonate buffer system in blood • HCO3- acts as weak base • H2CO3 acts as weak acid • H2CO3 ↔ H+ + HCO3- • H2CO3  H+ + HCO3- (in presence of XS base) • H2CO3 H+ + HCO3- (in presence of XS acid)

30. ORGANIC COMPOUNDS:Carbon and Its Functional Groups • Carbon forms bonds w/ itself • Large complex molecules of varying shapes • Most compounds do not dissolve easily in water • useful for building body structures • C compounds held together by covalent bonds • 4 valence e-  forms 4 bonds • Decompose easily • good source of energy • Functional groups have distinct chemical properties when attached to organic molecule

31. Functional Groups • Many different functional groups can attach to carbon skeleton • Very large molecules = macromolecules • Isomers have the same molecular formulas but different structures (glucose & fructose are both C6H12O6) • STRUCTURAL FORMULA OFGLUCOSE (Fig 2.14) C6H12O6 ISOMERS

32. Carbohydrates (CHO) • Primary energy source in humans • Include sugars, starches, glycogen, and cellulose • Used to generate ATP • Structural building blocks (DNA) • Structurally, one H2O molecule/C atom • Function as food reserves • glycogen stored in liver & muscle • Divided into three major groups based on size: • Mono-/di-/polysaccharides

33. SUGARS: Monosaccharides • Names of sugars generally end in “-ose” • Monosaccharides • 3-7 carbon atoms • Monomers for building large CHO molecules in body • Ex: glucose (a hexose) is main energy-supplying compound in body • Humans absorb only 3 simple sugars without further digestion in small intestine • glucose found in syrup or honey • fructose found in fruit • galactose found in dairy products

34. SUGARS: Disaccharides • Formed from two monosacch. by dehydration synthesis • glucose + fructose  sucrose (table sugar) • glucose + glucose  maltose • glucose + galactose  lactose (milk sugar) • Can be split back into simple sugars by hydrolysis • Figure 2.15

35. Polysaccharides • Polymers of up to hundreds of monosaccharides • Primary polysaccharide in humans = glycogen • Stored in liver or skeletal muscles • Hydrolyzed in response to ↓ blood sugar  glucose released into blood (from liver only) • Cellulose • Plant polysaccharide • Not digestible by humans  “fiber”

36. Lipids • Contain carbon, hydrogen & oxygen • Fewer oxygens than CHO (not 2:1 H:O ratio) • Nonpolar covalent bonds • Hydrophobic • Insoluble in polar solvents such as water (plasma) • Only very short-chain fatty acids dissolve in plasma • Increase solubility by forming lipoproteins  “cholesterol”

37. LIPIDS: Triacylglycerols (TAG) • TAG (also called triglycerides) are what we call “fat” • Most plentiful lipids in the body  provide protection, insulation, and energy • Found in fats and oils • Fats = solid @ room temperature • Oils = liquid @ room temperature • Most concentrated form of energy • 9 Calories/gram • Proteins & carbs have only 4 Cal/gram! • Unlimited storage capacity in body  adipose tissue • ANY excess food energy is stored as fat • All TAG contain glycerol backbone & three fatty acids

38. Saturation of Fatty Acids • Determined by number of single or double covalent bonds • Saturated FA contain single covalent bonds & maximum possible # of H atoms • Saturated fats = TAG w/ only saturated fatty acids • Ex: lard, tallow • Unsaturated FA lack some H atoms due to presence of > 1 double bond • Monounsaturated fatty acids have one double bond • olive oil, canola oil, & avocados (yum!!) • Polyunsaturated fatty acids contain > 2 double bonds • corn, safflower, soybean oils • Double bonds form kink in structure of fatty acid • fluid rather than solid

39. Clinical Application • Essential fatty acids (EFA’s) are essential to human health and cannot be made by the human body. They must be obtained from foods or supplements. • ω-3 fatty acids  anti-inflammatory • ω-6 fatty acids  pro-inflammatory • Not all inflammation is bad! • Balance is important • Conjugated fatty acids (CFA’s)  some implications for weight loss… • trans-fatty acids ↑ risk factors for CVD

40. Phospholipids • Important membrane components • Amphipathic • polar head • a phosphate group (PO4-3) & glycerol molecule • forms hydrogen bonds with water • 2 nonpolar fatty acid tails • interact only with lipids • hydrophobic

41. Steroids • Four rings of carbon atoms • Include • cholesterol • important component of cell membranes • starting material for synthesizing other steroids • sex hormones • bile salts • vitamin D • cortisol

42. Other Lipids • Eicosanoids include prostaglandins and leukotrienes. • derived from 20-C fatty acids AA (ω-6) or EPA (ω-3) • prostaglandins have wide variety of functions • modify responses to hormones • contribute to inflammatory response • dilate airways • regulate body temperature • influence formation of blood clots • leukotrienes = allergy & inflammatory responses • PG & LT derived from EPA are biologically inactive • Fatty acids; fat-soluble vitamins (D, E, K); and lipoproteins

43. Proteins • Contain C, H, O, N & sometimes S • 12-18% of body weight • Functions: • Give structure to body (primary role) • Regulate processes • Provide protection • Help muscles contract • Transport substances • Enzymes

44. Proteins • Constructed from combinations of 20 amino acids • dipeptide formed from 2 amino acids joined by peptide bond (covalent bond) • polypeptide chains formed from 10 to 2000 amino acids

45. Amino Acid Structure • Central carbon atom • Amino terminus (NH2) • Carboxyl terminus (COOH) • Side chains (R groups) vary between amino acids • Amino acids identified by side chain

46. Levels of Protein Structure • Primary = sequence of amino acids • Secondary = twisting & folding • Alpha helices • Beta pleated sheets • Tertiary = 3-D shape of folded protein • **Determines function** • Disulfide bridges • Hydrophobic domains in core of folded protein • Quaternary = structure resulting from linkage of 2 polypeptides • Shape influences its ability to recognize & bind other molecules • Denaturation causes loss of characteristic shape and function

47. Bonds of Tertiary Structure • Hydrophobic interaxn on inside of folded protein • Disulfide bridges stabilize • covalent bond btwn S—H groups of 2 cysteine a.a. • H-bonds • Loss of 3-D structure (denaturation)  loss of function • Salts • Heat • Acid

48. Enzymes • Biological catalysts • Names generally end in “ase” • Sucrose is digested by enzyme sucrase • Properties: • Highly specific in terms of substrate & reaction • Highly efficient • Highly regulated by variety of cellular controls • Genes • Active & inactive conformations • Speed up chemical reactions by: • Increasing frequency of collisions • Lowering the activation energy • Properly orienting colliding molecules (Figure 2.23)

49. Enzymes as Catalysts Example: • Normal body temperatures & concentrations are low enough that rxns are effectively blocked by Ea barrier • Lactose reacts very slowly w/ water to yield glc & gal • Lactase (enzyme) orients lactose & water properly • Thousands of lactose/water reactions may be catalyzed by one lactase enzyme • Without lactase, lactose remains undigested in intestines • causes diarrhea and cramping  condition known as lactose intolerance (NOT an allergy!!!)

50. Nucleic Acids: DNA and RNA • Huge organic molecules containing C, H, O, N, P • Deoxyribonucleic acid (DNA) • genetic code inside each cell • regulates most cellular activities • Ribonucleic acid (RNA) • relays instructions from genes in cell’s nucleus • guides assembly of proteins by ribosomes • Basic units of nucleic acids are nucleotides • nitrogenous base • pentose sugar • deoxyribose • ribose • phosphate group (Figures 2.24a,b)