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1. Figure 2.1 The hierarchy of biological order from atom to organism

2. Organisms are made of MATTER Matter has mass and occupies space ELEMENT: fundamental substance -- can’t break down to other substances by chemical reaction (“rxn”) e.g., oxygen, carbon, nitrogen, gold… each has a symbol in the Periodic Table

3. COMPOUND (molecule): 2+ elements e.g., NaCl (sodium chloride) Life (on Earth) requires about 24 elements Major elements in most organisms by weight (about 96% of total human weight): oxygen (O), hydrogen (H), carbon (C), nitrogen (N)

4. Table 2.1 Naturally Occurring Elements in the Human Body

5. Trace elements: usually less than 1/100 percent of organism: e.g., copper, cobalt, fluorine, selenium, zinc, iodine, etc. Can contribute to structure, allow enzymes to work, etc.

6. Each element corresponds to an ATOM Atoms are made of subatomic particles: protons, neutrons, and electrons (e-) (these are enough for our purposes) Atom has NUCLEUS made of proton(s) and usually neutrons (except hydrogen -- has NO neutrons, just 1 proton + 1 e-) Electrons “orbit” nucleus

7. Protons have + charge Electrons have - charge Neutrons: neutral Electrons are attracted to + charge of nucleus Mass of atom is almost all due to protons and neutrons (e- weigh almost nothing)

8. Elements are defined by number of subatomic particles: a given element has a specific number of protons in nucleus This is the ATOMIC NUMBER: e.g., hydrogen: 1H, helium: 2He... (see Periodic Table) For most atoms in organisms, # of protons = # of neutrons (so overall charge of atom is neutral)

9. ATOMIC MASS: # of protons + neutrons Atomic number is subscript and mass is superscript e.g., 11H (hydrogen: 1 proton, NO neutrons) 42H (helium: 2 protons + 2 neutrons) 2311Na (sodium: 11 protons + 12 neutrons)

10. Figure 2.10 Electron configurations of the first 18 elements

11. Electrons “orbit” nucleus at specific distances to form “shells” 1st shell has maximum 2 e- 2nd, 3rd etc. have max. 8 e- (we’ll come back to this) ISOTOPE: different forms of same element (same # of protons) but different # neutrons e.g., 146C, 136C, 126C (most abundant)

12. ISOTOPE: different forms of same element (same # of protons) but different # neutrons Some isotopes are unstable (radioactive) and can “decay” (give off subatomic particles and sometimes even gain others as a result) -- carbon 14 is an example Sometimes harmful, but also useful in medicine and in dating age of fossils, rocks, etc. if rate of decay known

13. Figure 2.5 Two simplified models of a helium (He) atom

14. Recall that e- “orbit” nucleus (as a “cloud”) at fixed distances These represent ORBITALS (I will call these “SHELLS” for simplicity) Each shell’s electrons have a set amount of stored potential energy (E = abbreviation for energy) -- E is the ability to do work The further out they are the more E

15. Figure 2.9 Energy levels of an atom’s electrons

16. Figure 2.10 Electron configurations of the first 18 elements

17. Hydrogen: 1 e- ONLY (plus 1 proton in nucleus; NO neutrons) So 1 shell Helium: 2 e- (plus 2 protons + 2 neutrons) Still 1 shell Lithium: 2 e- in 1st shell, 1 e- in 2nd shell Neon: 2 e- in 1st shell, 8 in 2nd (8 e- is max. for shells above the 1st)

18. Figure 2.11 Electron orbitals

19. Outermost shell = VALENCE shell If valence shell is full (2 e- for 1st shell, 8 for others), element is unreactive If not, element “wants” to share or take or give e- to gain stability: reacts with other atoms to do this VALENCE = # of e- element “needs”

20. CHEMICAL BONDING: e- sharing or transfer between atoms to achieve stability (Elements with similar valences have similar chemical properties and reactivity) Major types of bonds involving e- sharing or transfer: COVALENT (polar or nonpolar; e- shared) IONIC: e- transferred (spend all their time orbiting atom that “needs” them)

21. Covalent bonds: 1 e- shared: single bond 2 shared: double 3: triple, etc. If sharing is equal (e- spend same amount of time orbiting each atom), bond is NONPOLAR covalent If unequal (e- attracted more to particular atom): POLAR COVALENT (e.g., H2O)

22. Figure 2.12 Covalent bonding in four molecules

23. Figure 2.12x Methane

24. Figure 2.13 Polar covalent bonds in a water molecule

25. H2O is a POLAR molecule because distribution of e- (and therefore charge) is uneven: e- from each H are strongly attracted to O SO: oxygen part of molecule has negative charge (b/c e- are negative) H parts are positively charged (b/c the e- of each H spends most of its time near O) (will discuss more later)

26. IONIC BONDING: complete e- transfer e.g., NaCl (sodium chloride -- table salt) Valence of Na = 7 1 e- in outermost shell SO to be stable either “wants” 7 more OR can give up the 1 valence e-, lose a shell but now have outermost shell with 8 e- Valence of Cl = 1 -- could be stable if it got just 1 more e-

27. SO: Na can give its 1 valence e- to Cl Now each has 8 e- in outermost (valence) shell They stay together as a salt when dry, but in water (polar molecules), NaCl dissociates into Na+ (cation = positive charge) and Cl- (anion = negative charge) Ion = charged particle; gain or loss of e- determines charge

28. Figure 2.14 Electron transfer and ionic bonding

29. Figure 2.15 A sodium chloride crystal

30. HYDROGEN BONDS: between molecules NOT e- sharing; H with + charge (as in H2O) attracted to - charge of an atom in another molecule e.g., water molecules are attracted to other water molecules; weak bonds Other types of molecule-molecule bonds: VAN DER WAALS FORCES: individual bonds weak, but many together can be strong

31. Figure 2.16 A hydrogen bond

32. Figure 2.17 Molecular shapes due to hybrid orbitals

33. Figure 2.18 Molecular shape and brain chemistry

34. Figure 2.19 A molecular mimic

35. CHEMICAL REACTIONS: (“rxns”): Make/break bonds e.g., 2H2 + O2 2H2O 4 H and 2 O 4 H and 2 O Reactants Products Break nonpolar covalent H-H bond and nonpolar covalent O=O bonds Make new bonds btw H and O

36. Unnumbered Figure (Page 38) Chemical reaction between hydrogen and oxygen

37. Photosynthesis (simplified): (many steps) 6CO2 + 6H2O C6H12O6 + H2O (a sugar) Reactants Products

38. Figure 2.20 Photosynthesis: a solar-powered rearrangement of matter

39. Some rxns go to completion: All reactants are converted to products, no reverse BUT: most rxns are REVERSIBLE: Go back and forth btw reactants and products e.g., 3H2 + N2 2NH3 (ammonia)

40. If concentrations of reactants are high, lots of chances to meet, lots of rxns (Note: abbreviation for concentration is [ ] ) As products accumulate, more chances of going the other way (“backwards”) So for reversible rxns, rate in each direction is concentration-dependent

41. When a balance in forward and reverse rates is achieved, have chemical equilibrium This is a dynamic situation: rxns are still happening but rates in both directions are the same This does NOT mean that concentrations of reactants and products have to be the same; depending on conditions, one or other direction may be favored

42. CHAPTER 3: WATER

43. Figure 3.0 Earth

44. Earth is about 4.6 billion years old Very hot until about 4 billion years ago Cooled, water in atmosphere condensed, precipitated, formed oceans Life arose in water at least 3.6 billion years ago, didn’t move to land until about 1.5 billion years ago

45. Pre-biological molecules formed; triggered by electricity, volcanic activity, UV radiation, etc. (can replicate this in lab)...then Life Early atmosphere had lots of H2O, NH3, CO2, CO, etc. BUT very little O2 O2 accumulated in atmosphere because of photosynthesis (this will become important later)

46. Water is essential to all known life Very unusual molecule, many special properties Recall that H2O is a polar molecule: O end has - charge, H ends have + charge SO: water molecules attract each other, form H bonds; each H2O can form bonds with 4 others

47. Figure 3.1 Hydrogen bonds between water molecules 

48. In liquid, these bonds are weak and transient (temporary): H bonds continuously made and broken (break more easily the hotter it is) Bonding causes cohesion: molecules “stick together” Important in capillary action, e.g., movement of water through plant vessels against force of gravity

49. Figure 3.2 Water transport in plants

50. Causes surface tension: water has a “coating”: explains why you can overfill a glass of water, skip a stone Some organisms can move across water’s surface

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