Human Anatomy and Physiology Basic Chemistry Unit 2

1. Human Anatomy and PhysiologyBasic Chemistry Unit 2 Remember Metabolism is just those chemical reactions in our body…so we have to understand Basic Chemistry

2. Matter vs. Energy Matter can store energy through atomic bonds We understand that matter can release energy…and that moving matter really fast (energy) can create new matter (particle accelerator)

3. Matter…a practical working definition • Anything that has mass and occupies space • States of matter: • Solid—definite shape and volume • Liquid—definite volume, changeable shape • Gas—changeable shape and volume

4. Energy can be converted from one form to another, but always loses heat between conversions • Potential • Kinetic • Chemical • Electrical • Mechanical • Radiant or Electromagnetic

5. Composition of Matter • Elements • Cannot be broken down by ordinary chemical means • Each has unique properties: • Physical properties • Are detectable with our senses, or are measurable • Chemical properties • How atoms interact (bond) with each other • Atoms • Unique building blocks for each element • Atomic symbol: one- or two-letter chemical shorthand for each element • Au Fe Mg K Cl F Sn Cu

6. Elements of the Human Body • Oxygen (O) • Carbon (C) • Hydrogen (H) • Nitrogen (N) • Minor elements: 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) About 96% of body mass These are major elements

7. Your guide to atoms. Make sure you can read it!

8. So the basic unit of matter is…the ATOM Be able to describe and draw an atom… including: • Protons • Neutrons • Electrons • Ions • Isotopes

9. Atomic Structure • Neutrons • No charge • Mass = 1 atomic mass unit (amu) • Protons • Positive charge • Mass = 1 amu • Electrons • Orbit nucleus • Equal in number to protons in atom • Negative charge • 1/2000 the mass of a proton (0 amu)

10. Nucleus Nucleus Helium atom Helium atom 2 protons (p+) 2 neutrons (n0) 2 electrons (e–) 2 protons (p+) 2 neutrons (n0) 2 electrons (e–) (a) Planetary model (b) Orbital model Proton Neutron Electron Electron cloud Figure 2.1

11. Proton Neutron Electron Hydrogen (H) (1p+; 0n0; 1e–) Helium (He) (2p+; 2n0; 2e–) Lithium (Li) (3p+; 4n0; 3e–) Figure 2.2

12. Identifying Elements • Atomic number = number of protons in nucleus • Mass number = mass of the protons and neutrons • Mass numbers of atoms of an element are not all identical • Isotopes are structural variations of elements that differ in the number of neutrons they contain • Atomic weight = average of mass numbers of all isotopes

13. Proton Neutron Electron Hydrogen (1H) (1p+; 0n0; 1e–) Deuterium (2H) (1p+; 1n0; 1e–) Tritium (3H) (1p+; 2n0; 1e–) Figure 2.3

14. Radioisotopes • Spontaneous decay (radioactivity) • Similar chemistry to stable isotopes • Can be detected with scanners • Valuable tools for biological research and medicine • Cause damage to living tissue: • Useful against localized cancers • Radon from uranium decay causes lung cancer

15. Ionic Bonds • 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

16. Molecules and Compounds • Most atoms combine chemically with other atoms to form molecules and compounds • Molecule—two or more atoms bonded together (e.g., H2 or C6H12O6) • Compound—two or more different kinds of atoms bonded together (e.g., C6H12O6) • Elements- one or more of the same atom bonded together

17. Chemical Bonds • 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)

18. (a) Chemically inert elements Outermost energy level (valence shell) complete 8e 2e 2e Helium (He) (2p+; 2n0; 2e–) Neon (Ne) (10p+; 10n0; 10e–) Figure 2.5a

19. (b) Chemically reactive elements Outermost energy level (valence shell) incomplete 4e 2e 1e Hydrogen (H) (1p+; 0n0; 1e–) Carbon (C) (6p+; 6n0; 6e–) 1e 6e 8e 2e 2e Oxygen (O) (8p+; 8n0; 8e–) Sodium (Na) (11p+; 12n0; 11e–) Figure 2.5b

20. Valence Electrons and Bonding • Be able to draw and predict bonding for the following atoms: • Carbon • Hydrogen • Calcium • Oxygen • Fluorine • Lithium • Sodium • Chlorine • Helium

21. Three important Bonds • Ionic • Covalent • Hydrogen

22. + – 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. (b) After electron transfer, the oppositely charged ions formed attract each other. Figure 2.6a-b

23. Formation of an Ionic Bond • Ionic compounds form crystals instead of individual molecules • NaCl (sodium chloride)

24. CI– Na+ (c) Large numbers of Na+ and Cl– ions associate to form salt (NaCl) crystals. Figure 2.6c

25. Covalent Bonds • Formed by sharing of two or more valence shell electrons • Allows each atom to fill its valence shell at least part of the time

26. Reacting atoms Resulting molecules + or Structural formula shows single bonds. Molecule of methane gas (CH4) Hydrogen atoms Carbon atom (a) Formation of four single covalent bonds: carbon shares four electron pairs with four hydrogen atoms. Figure 2.7a

27. Reacting atoms Resulting molecules + or Structural formula shows double bond. Molecule of oxygen gas (O2) Oxygen atom Oxygen atom (b) Formation of a double covalent bond: Two oxygen atoms share two electron pairs. Figure 2.7b

28. Reacting atoms Resulting molecules + or Structural formula shows triple bond. Molecule of nitrogen gas (N2) Nitrogen atom Nitrogen atom (c) Formation of a triple covalent bond: Two nitrogen atoms share three electron pairs. Figure 2.7c

29. Covalent Bonds • Sharing of electrons may be equal or unequal • Equal sharing produces electrically balanced nonpolar molecules • CO2

30. Figure 2.8a

31. Covalent Bonds • 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

32. Figure 2.8b

33. Figure 2.9

34. Hydrogen Bonds • Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule • Common between dipoles such as water • Also act as intramolecular bonds, holding a large molecule in a three-dimensional shape • http://www.youtube.com/watch?v=lkl5cbfqFRM

35. + – Hydrogen bond (indicated by dotted line) + + – – – + + + – (a) The slightly positive ends (+) of the watermolecules become aligned with the slightlynegative ends (–) of other water molecules. Figure 2.10a

36. (b) A water strider can walk on a pond because of the highsurface tension of water, a result of the combinedstrength of its hydrogen bonds. Figure 2.10b

37. Chemical vs Physical reactions • Chemical bonds are formed, rearranged or broken • Synthesis, Decomposition, Exchange, Redox • Chemical Equations • Reactants • Products • Molecular formulas • Balance atoms • No chemical bonds are changed • Most matter exists as MIXTURES • Solution • Aqueous/ homeogeneous • Solute/ solvent • http://www.youtube.com/watch?v=3G472AA3SEs • Colloid • Medium solute/ gel/ foam • heterogeneous • Suspension • Large solute/ heterogeneous Chemical reactions Physical reactions

38. Concentration of Solutions • Expressed as • Percent, or parts per 100 parts • Milligrams per deciliter (mg/dl) • Ppm: parts per million • Molarity, or moles per liter (M) • 1 mole = the atomic weight of an element or molecular weight (sum of atomic weights) of a compound in grams • 1 mole of any substance contains 6.02  1023 molecules (Avogadro’s number) • What is the molecular weight of a gram of glucose (C6H12O6)? • To make a 1M (one molar) solution of glucose, how much glucose would you weigh to mix with water to make one liter?

39. 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 Mineral water Example Gelatin Example Blood Figure 2.4

40. Chemical Reactions can be expressed as Equations H + H  H2 (hydrogen gas) 4H + C  CH4 (methane) Bonds are always broken, made, rearranged Atoms are always accounted for so formulas must be balanced to work • Chemical equilibrium occurs if neither a forward nor reverse reaction is dominant • Many biological reactions are essentially irreversible due to Energy requirements and/or • Removal of products (reactants) (product)

41. Patterns of Chemical Reactions • Synthesis • Anabolic = forms bonds/ Endergonic (absorbs energy) • A + B → AB Found in rapidly growing tissues • Decomposition • Catabolic = breaks bonds/ Exergonic (releases energy) • AB → A + B Found in metabolic reactions/ cell respiration • Exchange or Displacement • Involves synthesis and decomposition/ atoms rearranged • AB + C → AC + B AB + CD → AD + CB • Oxidation-Reduction (Redox) reactions • Both decomposition and exchange reactions • Electrons move between reactants • Reactant that loses electron is an electron donor/ oxidized • Reactant that gains electron is an electron acceptor/ reduced • LEO the lion goes GER

42. Rate of Chemical Reactions • Rate of reaction is influenced by: •  temperature  rate •  particle size  rate •  concentration of reactant  rate • Catalysts:  rate without being chemically changed • Enzymes are biological catalysts

43. (a) Synthesis reactions Smaller particles are bonded together to form larger, more complex molecules. Example Amino acids are joined together to form a protein molecule. Amino acid molecules Protein molecule Figure 2.11a

44. (b) Decomposition reactions Bonds are broken in larger molecules, resulting in smaller, less complex molecules. Example Glycogen is broken down to release glucose units. Glycogen Glucose molecules Figure 2.11b

45. (c) Exchange reactions Bonds are both made and broken (also called displacement reactions). Example ATP transfers its terminal phosphate group to glucose to form glucose-phosphate. + Glucose Adenosine triphosphate (ATP) + Glucose phosphate Adenosine diphosphate (ADP) Figure 2.11c

46. Oxidation Reduction (Redox) LEO the lion goes GER

47. Common Example of Redox Reactions in Body: Aerobic Cell Respiration Glucose is oxidized to carbon dioxide as it loses hydrogen atoms oxygen is reduced to water as it accepts the hydrogen atoms

48. Classes of Compounds • Inorganic compounds • Water, salts, and many acids and bases • Do not have to contain carbon • Organic compounds • Carbohydrates, fats, proteins, and nucleic acids • Must contain carbon, usually large, and are covalently bonded

49. Water • 60%–80% of the volume of living cells • Most important inorganic compound in living organisms because of its properties

50. Properties of Water • 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 • 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