The Chemical Basis of Life

1. The Chemical Basis of Life Chapter 2

2. Overview • Atoms • Combining Matter • Physically • Chemically • Water • Acids, Bases, and pH • Buffers

3. Matter and Energy Matter: • Occupies space • Has mass: liquid, gas, solid Energy: • Capacity to do work • Measured by effect on matter

4. Chemistry Science of the structure of matter Central to all other sciences Chemistry is part of all living & non-living things

5. Life requires ~25 chemical elements Humans & other living organisms differ from non-living things in elemental composition 96% of body weight made up of C, H, O, N Other 4%: Ca, P, K, S, Na, Cl, Mg & trace elements essential for life (e.g. Fe)

6. Elements Basic units of all matter Can’t be broken down to simpler substances using ordinary chemical methods 112 known elements → periodic table

7. Each element is represented by its atomic symbol (1st letter(s) of element’s name) e.g. carbon = C hydrogen = H oxygen = O

8. In nature, few elements exist in pure form (tend to form compounds) Emergent properties: e.g. NaCl Na (metal) + Cl (poisonous gas) = NaCl (table salt)

9. Atoms Building blocks of elements Unique to each element Give it specific physical and chemical properties Physical properties: Colour, texture, boiling point, melting point, etc. Chemical properties: The ways that atoms interact with other atoms

10. Made up of protons (p+), neutrons (no), electrons (e-) e- p+ no

11. Protons (p+) have positive charge Neutrons (no) have no overall charge Both are heavy particles with approximately same mass Electrons (e-) have negative charge Do not contribute to atomic mass (1/2000th mass of proton) In general, # protons = # electrons No net electrical charge

12. Mass # = p+ + no H a b Atomic # = p+ Generalized Atom Periodic table is ordered by atomic number

13. H He Li 1 4 7 p+ 1 2 3 2 3 1 no 0 2 4 e- 1 2 3 The 3 Smallest Atoms

14. Atomic Mass Approximately equal to mass number (# p+ + # no) because e-s weigh so little In general, atomic weight is about equal to mass # of most abundant isotope e.g. atomic mass of H = 1.008 (indicates that 1H is present in much greater amounts than 2H or 3H forms)

15. Isotopes Different versions of same element Occur with most natural elements Differ in # of neutrons (same atomic # but different mass #) If stable, nucleus remains intact If unstable, is radioactive

16. Radioisotopes Nuclei decompose spontaneously into more stable forms e.g. 14C: half-life of 5700 years ½ atoms turn into 13N Used to date rocks and biological remains Releases particles & energy (breaks chemical bonds in living organisms) Damaging to live tissue but used in biological research & medicine

17. Structure of an Atom Nucleus contains protons & neutrons Electrons move around nucleus = electron cloud Atomic orbitals organized into shells e.g. 11Na

18. Higher-energy shells hold more e-s (2n2) & are located further from nucleus Shells fill up in order of increasing energy e-s can be excited up to higher energy level for brief periods Spontaneously return to lower level while emitting the energy gained via excitation

19. e-s in outer (valence) shell dictate chemical behaviour (these ones interact with those from other atoms) Regardless of # of e-s in each shell, # that can participate in bonding is 8 = octet rule

20. The Octet Rule Atoms want to gain, lose, or share e-s so that have 8 electrons in outer shell Exception = H (only has room in 1st energy level for 2 e-s)

21. Use atomic number to calculate how many e-s are available for bonding • 1st energy level = 2 e-s / shell • 2nd and up = 8 e-s / shell e.g. 6C: Has 4 e-s in outer shell; wants to gain 4 e-s to fill shell for a total of 8 e-s

22. 8O: Has 6 e-s in outer shell Needs 2 7N: Has 5 e-s in outer shell Needs 3

23. 11Na: Has 1 e- in outer shell Needs 7 17Cl: Has 7 e-s in outer shell Needs 1

24. Combining Matter Most atoms do not exist in free state Chemically combine with other atoms to form molecules If atoms are the same = molecule of element e.g. O2 If atoms are different = compound e.g. H2O

25. Molecular Formulas A molecule’s chemical composition is written as a formula Symbols for elements Subscripts for number of atoms of each element e.g. H20 = 2 H, 1 O e.g. 5 H20 = 10 H, 5 O

26. Ways to Represent Compounds e.g. methane (CH4) Structural formula Ball-and-stick model Space-filling model

27. Special Structure: Carbon Ring If icon for ring shows no atoms, assume that C occupies each corner Same goes for 5-carbon rings =

28. Mixtures 2 or more substances No chemical bonding = physical intermixing Living material contains 3 types: Solutions Colloids Suspensions

29. Mixture #1: Solution Homogeneous Transparent Does not settle out Solvent • Present in largest quantity • Usually liquid • Water is body’s principle solvent Solute • Present in smaller quantity

30. Mixture #2: Colloid Heterogeneous Translucent or milky Does not settle out Can undergo sol-gel transformation e.g. cytosol in cells

31. Mixture #3: Suspension Heterogenous Settles out e.g. blood settles out into plasma & cells

32. Inert if outer e- valence shell is filled Do not tend to form bonds e.g. He Reactive if outer shell is not filled React with other atoms to gain / lose / share e-s to fill shells e.g. O Chemical Bonds Attractive forces between atoms

33. Ionic Bonds Transfer of e- s from one atom to another Become ions (charged particles) Gain e- →negative charge = anion Lose e- →positive charge = cation Both become stable & combine to form ionic compound (a.k.a. salt)

34. Salts Release ions other than H+ and OH- Usually form when acids and bases mix Dissociate in water into component ions (electrolytes that can conduct electricity) Important in living organisms: e.g. Na+, K+, Ca2+ used in nerve transmission, muscle contraction e.g. plant cells use salts to take up water from soil

35. Covalent Bonds E- sharing Each atom fills outer shell part of the time Can be single, double, or triple bonds e.g. H2: H-H; O2: O=O; N2: NN Can be polar or non-polar bonds

36. Atoms can be: Electropositive: 1-2 valence shell e-s Tend to lose e-s Electronegative: 6-7 valence shell e-s Tend to attract e-s strongly Electrically balanced

37. O C O Non-polar Covalent Bonds Electrically balanced Equal sharing of e-s

38. Polar Covalent Bonds Unequal e- sharing One element has more protons = stronger pull on e-s = has e-s more of the time = slightly electronegative Results in molecule with + & - charges at either end Often occurs when atoms are of different sizes - O H H + +

39. Hydrogen Bonds Not a true bond = can’t form molecules Attraction between covalently-bound H atom & electronegative atom (can be different molecule or different area of same molecule) e.g. between water molecules, between complementary bases in DNA

40. Mixtures vs. Compounds

41. Water’s Life-Giving Properties The universal solvent Water is important because: • Life originated in it • All known living things depend on water (metabolic processes, respiration, photosynthesis) • Maintains cell structure/shape

42. Characteristics of Water • Polar molecules • Specific heat capacity • Heat of vaporization • Density of water • Cohesion • Adhesion • Surface tension • Good solvent All result from H-bonding

43. - + + Polarity of the Water Molecule One end slightly positive, other slightly negative = no net charge Attracts other water molecules (cohesion) Attracts sugar & other polar (hydrophilic) molecules Repels oil & other non-polar (hydrophobic) molecules

44. Why is Polarity Important? If water were linear (non-polar), not bent (polar): • It would not liquify except at high pressures • It would probably not remain liquid over more than about a 20°C. temperature range • Polarity helps water stay liquid because molecules so strongly attracted to each other • It would dissolve very few other substances • Polarity of water molecules can cause temporary polarity in non-polar molecules; virtually everything will dissolve to a small extent in water In consequence, life could not exist anywhere

45. Water & Heat H-bonds make it difficult to separate molecules H-bonds are constantly forming & breaking When temperature is stable, H bonds form at the same rate that they break

46. Heat of Vaporization When temperature increases: H bonds break & stay broken Individual molecules escape into air = evaporation Heat energy changes liquid H2O into gaseous form High boiling point (100°C)

47. When water cools: H-bonds reform H-bonds release heat energy as temperature drops

48. Specific Heat Capacity = energy required to raise given amount of substance by 1°C Water has high specific heat capacity: At high temperatures, water absorbs heat as H-bonds break (can absorb a lot before temperature measurably rises) As water cools, heat released from formation of H-bonds slows down cooling

49. Water’s high specific heat capacity: • Helps regulate Earth’s climate by buffering large changes in temperature • Helps moderate internal temperature

50. Density of Water Water reaches max. density at 4°C (becomes less dense at lower temps) When temp decreases below 0°C: Molecules don’t move enough to break H-bonds so become locked = ice