Chapter 1: An Introduction to Physiology

1. Chapter 1: An Introduction to Physiology -physiology: the study of the functions of living things 1. skeletal 2. articular 3. muscular 4. digestive 5. respiratory 6. urinary 7. reproductive 8. circulatory 9. nervous 10. integumentary 11. endocrine -the human body is comprised of 11 major organ systems: • - in studying these systems we can use two approaches: • teleological • -emphasizes the purpose of a body process • -explains a function in terms of meeting a bodily need • -emphasizes the WHY • e.g. “why do I shiver?” – to warm up because • shivering generates heat • mechanistic • -emphasizes the underlying mechanism by which • this process occurs • -view the body as a machine whose actions can be explained • in terms of cause and effect • -emphasize the HOW • e.g. “why do I shiver?” – detection of body temperature by sensory • receptors leads to activation of the somatic division of the nervous • system and trigger the involuntary contraction of skeletal muscles

2. -physiology is closely related to anatomy – because structure and function are closely related -physiological mechanisms are made possible because of the structure and design of a body part -some relationships are obvious – e.g. structure of the elbow as a hinge joint -others are more subtle – e.g. interface between the air and the blood in the lungs -air sac structure + capillary bed - 300 million air sacs + associated capillaries provides a total surface area for gas exchange = tennis court

3. Life: Levels of Organization • Atoms • Molecules • Macromolecules • Organelles • Cells • Tissues • Organs • Organ systems • Organism

4. Organizational Levels 1. chemical or molecular: 4 major elements within the body -99% of the total number of atoms within the body -C, N, O and H -molecular composition - 67% of our bodies is water

5. Atom = smallest unit of an element that still retains the chemical & physical properties of that element i.e. really, really, really tiny thing! -composed of: protons = one positive charge, 1 atomic mass unit (1.673x10-24g) electrons = one negative charge, no mass (9.109x10-28g) neutrons = no charge, 1 atomic mass unit -elements are grouped on a Periodic Table of Elements -the elements are grouped according to physical and chemical characteristics -on the chart each element is associated with a letter, an atomic number & an atomic mass -each atom is comprised of a nucleus of protons and neutrons + orbiting electrons

6. Periodic Table of Elements IA IIA IIIA IVA VA VIA VIIA VIII http://www.chemicalelements.com/ http://periodic.lanl.gov/default.htm

7. Li 7 3 atomic symbol atomic mass (weight) e.g. # protons (e-) = 6 # pr+6 + #No 6 = 12 C 12 6 atomic number e.g. # protons (e-) = 3 # pr+3 + #No 4 = 7 Atomic mass = number of protons + neutrons Atomic number = number of protons when the element is electrically neutral ** when neutral, the number of protons and electrons are equal

9. Isotope: • same atomic # (same pr+, same e-) • differs only in # of neutrons 12C13C14C pr+: 6 6 6 e-: 6 6 6 No: 6 7 8** radioactive Radioactive isotope uses: -radioactive isotopes have a high neutron to proton ratio -the nuclear “glue” within the nucleus is not strong enough to hold the nucleus together = radioactivity 1. carbon dating - 14C 2. radioactive imaging - e.g. PET scanning -use of FDG – radioactive glucose tracer -18F radioactive isotope (2-fluoro-deoxy-glucose) 3. cancer treatment - 60Co

10. Chemical bonds -forces holding atoms together = chemical bonds -different kinds of chemical bonds – but all involve the electrons of atoms Two types of bonds: 1. Ionic 2. Covalent

11. Electron Configurations • Bed check for electrons • description on how are electrons organized around the nucleus of protons and neutrons? • Bohr model: Nils Bohr proposed electrons orbit around the atom’s nucleus in specific energy levels or orbits (shells) • these shells have a specific energy level – closer the electron is to the nucleus the less energy it needs to “orbit” • this model only works for smaller atoms • larger atoms are described by quantum mechanics – orbitals have energy, momentum/shape, spin and magnetic characteristics • comprised of subshells • 1st shell – closest to the nucleus only holds 2 electrons (s subshell only) • 2nd shell can hold 8 (s and p subshells – 2 + 6 electrons) • 3rd holds 18 (s, p and d subshells – 2 + 6 + 10 electrons) • 4th holds 18

12. an atom will always try to complete its outermost shell • basis for bonding reactions • the number of electrons the atom gains or loses to complete its outer shell = valence • chemists really only consider the electrons in the s and p orbitals as valence electrons

14. Molecule: • particle formed by the union of more than one atom • e.g. same kind of atom - O2 • e.g. different types of atoms - H20 Molecules are held together by either covalent or ionic bonds -these bonds form through interactions between the valence electrons

15. 1. Ionic bond: • attraction between 2 oppositely charged atoms (ions) • e.g. Na+ and Cl- NaCl • e.g. Ca2+ and Cl- CaCl2 -positively charged ions = cations -negatively charged ions = anions -these form as one atom transfers electrons to another atom -ions may also be composed of more then one atom = polyatomic -these are treated as the same as monoatomic ions e.g. sulfate SO43-, nitrite NO2-, hydroxide OH-

16. Ionic Reaction

17. -most of your group I and group II metals will form ionic bonds with the group VIIA halogens

18. 2. Covalent Bond: • if it isn’t favorable for an atom to gain or lose an electron • it will have to share it with another • covalent bond = bond in which atoms share electrons • e.g O2, , N3 • e.g H20 • -usually forms when one atom has to lose or gain three or more • electrons • e.g. carbon would have to gain 4 valence electrons to complete its outer shell • nitrogen would have to gain 3 valence electrons • -also form between two identical atoms – e.g. nitrogen, oxygen gas

19. Polar and Non-polar bonds • the sharing of electrons does not have to be equal • nonpolar covalent bond = equal sharing of electrons • e.g. oxygen (O2), methane (CH4) • polar covalent bond = uneven sharing of electrons leading to a slight charge • e.g. water = H20

20. d- O H H d+ d+ • Water: • 60-70% of body weight • covalent bond • POLAR molecule (uneven sharing of • electrons) • polar compounds are attracted to other • the bond between one oxygen and • the hydrogens of adjacent water • molecules = Hydrogen Bond • **HB = occurs between a covalently • bonded hydrogen and negatively • charged atom a distance away

21. Chemical reactions: 3 types: 1. Synthesis - A + B AB (Anabolism reactions) 2. Decomposition - AB A + B (Catabolism reactions) 3. Exchange - AB + CD AD + BC -these equations must be balanced -Law of conservation of Mass or “chemical bookeeping” -i.e. the number of atoms of each element is the same before and after a chemical reaction

22. Chemical reactions • are made up of reactants and creates products • these reactions go on constantly within the human body = metabolism • each reaction involves changes in energy • if the reaction requires energy = endothermic (anabolism) • if it liberates energy = exothermic (catabolism) • atoms, molecules and ions are continuously moving and colliding with one another = kinetic energy • this kinetic energy if big enough (i.e. collision is large enough) can break a bond or cause a new one to form • this collision energy = activation energy • critical to the progression of all chemical reactions in our body • the more often a collision occurs the greater chance a bond will form or break

23. Activation Energy & Catalysts • activation energy = initial “energy investment” required to start a reaction • the reactants must absorb enough energy to cause their chemical bonds to become unstable and created new ones • as these bonds form – energy is released into the environment – if more energy is released than absorbed = heat (exothermic reaction) • two influences on AE – temperature and concentration • concentration – increasing this increases the chance of collision between atoms • temperature – heating a reaction increases the kinetic energy of the reactants – collide more often • catalysts = compounds that lower the activation energy of a reaction

24. Molecules of Life: • the chemicals used in metabolic reactions or those • that are produced by them can be classified into • 2 groups: • 1. Inorganic • 2. Organic

25. Inorganic Compounds • water • oxygen,carbon dioxide • inorganic salts

26. d- O H H d+ d+ Water • major component of blood, plasma, CSF etc… • role in: transporting chemicals • transporting waste products • transporting & absorbing heat • polar molecule - asymmetrical distribution of charge • liquid at room temperature – we can drink it • universal solvent for polar compounds – facilitates most chemical • reactions in the body • -water molecules are cohesive – therefore they cling together • (because of hydrogen bonding) • -this allows the even distribution of dissolved substances • throughout our system – so water is an excellent transport • medium • -the temperature of water rises and falls slowly – prevents • sudden and drastic changes of temperature in our bodies • -water requires high heat to evaporate – it cools our bodies • -frozen water is less dense than liquid water – ice floats • -water freezes from the top down – allows aquatic organisms • to survive winters

27. Water: • -excellent solvent for dissolution of polar and ionic substances • e.g. H20 + NaCl • in a solution – a solvent dissolves another substance called • a solute • -water is a versatile solvent because of its polar covalent bonds and its bent • shape which allows it to interact with its neighbours very well hydrophilic • ions and molecules that react with water = • ions and molecules that don’t react = hydrophobic -water + salt: the electronegative O of water attracts the +ve sodium in salt, the electropositive H attracts the –ve chlorine -the salt becomes surrounded by water molecules and the crystal lattice of salt is broken up

28. Solutions, Colloids and Suspensions • solution = homogenous (same) mixtures containing a relatively large amount of one compound (solvent) • e.g. sugar + water • the mixture is the same no matter where you sample it • colloid = solution of larger components called dispersed-phase particles • these particles all carry the same charge (repel each other) • their particles are larger than that of solutions • e.g. plasma proteins within the blood • suspension = solution of larger components called dispersed-phase particles • larger particles than that of colloid • if left undisturbed these particles will settle out to form a solid • e.g. red blood cells within blood • mixture = two or more types of elements or molecules physically blended together without the formation of physical bonds between them • these individual compounds can be separated by physical or chemical means • types of mixtures: combination of solutions, suspensions and colloids

29. Electrolytes • substances that release ions when they react with • water • -these ions will conduct electricity = Electrolytes • -created through the decomposition of ionic substances • e.g. salt • -although polar compounds can also liberate electrolytes • abundant in body fluids • source of ions eg. Na+, Ca2+, K+, Mg2+ • ions play a role in: maintaining water concentration • maintaining pH • bone development • muscle function • nerve function • ions must be maintained in a certain concentration • to maintain homeostasis Inorganic salts: Electrolyte Balance

30. Inorganic Acids & Bases "H+ donators" • 3 types: • 1. release H+ Acids • e.g. HCl H+ + Cl- • 2. release ions to combine with H+Bases • e.g. NaOH Na+ + OH- • 3. acids + bases Salts • e.g. HCl + NaCl H20 + NaCl "H+ acceptors"

31. one must consider acids & bases • in light of how they mix with water • e.g. HCl when mixed in water • dissociates into H+ ions and Cl- • ions • if a base such as NaOH is added – it will • dissociate into Na+ ions and OH- ions • - the Na+ ions will combine with the Cl- ions to • form NaCl, the H+ ions will combine with the • OH- ions to reform water.

32. pH Scale • measures concentration of H+ ions in a solution e.g. pH 6 = 1 x 10-6 pH 7 = 1 x 10-7 pH 8 = 1 x 10-8

33. Buffers: • chemical or compound that keeps the pH of a solution within • a normal range • resists pH change by taking up excess H+ or OH- ions e.g. blood = pH 7.4 “Bicarbonate buffering system” - our blood contains a small amount of carbonic acid H20 + H2CO3 H+ HCO3- H2CO3 carbonic acid excess OH- excess H+

34. H H H H H H H C C C C C C H H H H H H H H H H C C H H C C H C H H H H H C C H H H Organic Compounds • always contain carbon, oxygen and hydrogen • carbon can form 4 covalent bonds with other atoms • e.g. methane • carbon can also form covalent bonds with itself • forming long chain hydrocarbons • or a ring structure symmetrical charge hydrophobic (non-polar)

35. Organic compounds • the skeleton of carbon and hydrogen are frequently combined with other atoms and molecules = functional groups

36. hydrocarbon + carboxyl group “hydrophilic” H H H O carboxyl group = polar group H C C C C O- H H H • these groups confer a specific property to the organic compound • e.g. amino acid vs. nucleotide

37. Organic substances: 1. carbohydrates 2. lipids 3. proteins 4. nucleic acids • 1. Carbohydrates: • provide energy to cells • supply materials to build certain cell structures • stored as reserve energy supply (humans = glycogen) • water soluble • characterized H - C - OH (ratio C:H 1:2) • e.g. glucose C6H12O6 • sucrose C12H22O12 • classified by size: simple - sugars • complex – polysaccharides • -see Table 2-6 (Tortora) monosaccharides disaccharides

38. A. Simple carbohydrates • monosaccharides = single sugar in which the # of carbon • atoms is low - from 3 to 7 • e.g. pentose - 5 carbon sugar • hexose - 6 carbon sugar • hexose sugars: glucose • galactose • fructose • pentose sugars: ribose • deoxyribose -three ways to represent the structure of glucose 1. Molecular form 2. 3. Simplest form

39. A. Simple carbohydrates • disaccharide = two 6-carbon monosaccharides • -form by a dehydration synthesis reaction • -broken up by a hydrolysis reaction • e.g. glucose + glucose = maltose • e.g. glucose + fructose = sucrose • e.g. glucose + galactose = lactose

40. built of simple carbohydrates • e.g. glycogen • starch • cellulose • multiple, repeating monomers or • “building blocks” polymer B. Complex carbohydrates:

41. Starch & Glycogen • starch = storage form of glucose found in plants • -hydrolyzed into glucose • glycogen = storage form of glucose found in animals • -hydrolyzed into glucose (in liver) Cellulose • polysaccharide found in cell walls in plants • linkage between glucose monomers differs from starch • indigestible

42. Lipids • many types • 1. triglycerides = fats and oils • 2. phospholipids • 3. steroids • cholesterol – animal cell membranes, basis for steroid hormones • bile salts - digestion • vitamin D – calcium regulation • Adrenocorticosteroid hormones • Sex hormones • 4. Eicanosoids • prostaglandins • leukotrienes • 5. Others • fatty acids • carotenes – synthesis of vitamin A • vitamin E – wound healing • vitamin K – blood clotting • lipoproteins – HDL and LDL

43. fatty acid fatty acid fatty acid 2. Lipids • A. Fats • energy supply • most plentiful lipids in your body • composed of C, H and O • “building blocks” = 3 fatty acid chains (hydrocarbons) • 1 glycerol molecule fatty acid portion glycerol portion

44. carboxyl gp • fatty acids - carboxyl at end • -differ in chain length with each fat • -differ in carbon bonding -some fatty acids cannot be made by the body and must be taken in through food = essential fatty acids e.g omega-3 fatty acids -polyunsaturated fatty acids -important in regulating cholesterol levels -lower LDL levels in the blood -increase calcium utilization by body – stronger bones & teeth -reduce inflammation – arthritis -promote wound healing 1. single C bonds - saturated 2. double C bonds - unsaturated monounsaturated: 1 double bond polyunsaturated: 2 or more double bonds

45. similar to fat molecules - glycerol + 2 fatty acids • + a phosphate group • phosphate gp hydrophilic “head” • fatty acid gps hydrophobic “tails” B. Phospholipids • form the majority of the cell • membrane = lipid bilayer

46. C. Steroids • backbone = 4 fused carbon rings • diversity through attached functional groups • e.g. cholesterol • testosterone, estrogen • aldosterone

47. H HO C C N H O H H R 3. Proteins • roles: structural • energy source • chemical messengers • combine with carbohydrates = glycoproteins • receptors • antibodies • metabolic role - enzymes • building blocks = amino acids a.a. = amino group at 1 end, carboxyl at the other between is a single C atom bound to: 1. H atom 2. R group carboxyl gp amino gp

48. some amino acids: asparagine alanine arginine aspartic acid cysteine glutamic acid glycine histidine leucine lysine phenylalanine proline serine thymine tyrosine tryptophan valine

49. 2 a.a. dipeptide 3 a.a. tripeptide 4 or more a.a. polypeptide R R • amino acids joined together by a condensation reaction • forming a peptide bond = between the NH2 of 1 a.a. and • the COOH of the next peptide bond = polar H H HO C C N C C N H O H H O H H d+ d-

50. polypeptides have 4 types of structures or conformations • which affect their ultimate function Protein conformation: 1. primary - a.a. sequence of polypeptides