Lecture 3: Chemistry of Life

1. Lecture 3: Chemistry of Life

2. Lecture 3: Chemistry of Life Goals: 1. Sprint through General Chemistry 2. Whisper past Organic Chemistry 3. Approach Biochemistry cautiously 4. Apply chemistry overview and relate biological chemistry to this course and your life in general Key Terms: Charge, proton, neutron, electron, radioisotope, tracer, chemical bonds a)ionic, b)covalent, c)hydrogen, atom, molecule, pH scale, buffer, basic, acidic, hydrophobic, hydrophillic, acidosis, alkalosis, solute, polar, non-polar. http://pearl1.lanl.gov/periodic/default.htm http://www.chemsoc.org/viselements/pages/pertable_fla.htm

3. Elements • Fundamental forms of matter • Can’t be broken apart by normal means • 92 occur naturally on Earth Less than 12 occur on the exam

4. Most Common Elements in Living Organisms C H O N Carbon Hydrogen Oxygen Nitrogen

5. What Are Atoms? • Smallest particles that retain properties of an element • Made up of subatomic particles: • Protons (+) • Electrons (-) • Neutrons (no charge)

6. Hydrogen and Helium Atoms electron proton neutron HYDROGEN HELIUM Fig. 2.3, p. 22

7. Atomic Number • Number of protons • All atoms of the same element have the same atomic number • Atomic number of hydrogen = 1 • Atomic number of carbon = 6

8. Mass Number Number of protons + Number of neutrons Isotopes vary in mass number (not atomic number or they would be something else)

9. Isotopes Atoms of an element with different numbers of neutrons (different mass numbers) Carbon 12 has 6 protons, 6 neutrons Carbon 14 has 6 protons, 8 neutrons Radioisotopes Have an unstable nucleus that emits energy and particles Radioactive decay transforms radioisotope into a different element Decay occurs at a fixed rate Atomic Mass

10. Radioisotopes as Tracers • Example: Tracer Drug Study • How long does a drug stay in the patient? • Determine dose guidelines • Compound synthesized with a radioisotope • Emissions from the tracer can be detected with special devices • Track levels in the blood, urine and feces • Following movement of tracers is useful in many areas of biology

11. High Sensitivity Very Low Dose

12. Other Uses of Radioisotopes • Drive artificial pacemakers • Biomedical Imaging • Thyroid and bone scans • Radiation therapy Emissions from some radioisotopes can destroy cells. Some radioisotopes are used to kill small cancers.

13. What Determines Whether Atoms Will Interact? The most general of General Chemistry

14. Electrons • Carry a negative charge • Repel one another • Are attracted to protons in the nucleus • Move in orbitals - volumes of space that surround the nucleus y Z X When all p orbitals are full

15. Electron Orbitals • Orbitals can hold up to two electrons • Atoms differ in the number of occupied orbitals • Orbitals closest to nucleus are lower energy and are filled first

16. Shell Model • First shell • Lowest energy • Holds 1 orbital with up to 2 electrons • Second shell • 4 orbitals hold up to 8 electrons CALCIUM 20p+ , 20e-

17. Electron Vacancies • Unfilled shells make atoms likely to react • Hydrogen, carbon, oxygen, and nitrogen all have vacancies in their outer shells CARBON 6p+ , 6e- NITROGEN 7p+ , 7e- HYDROGEN 1p+ , 1e-

18. Chemical Bonds, Molecules, & Compounds • Bond is union between electron structures of atoms • Atoms bond to form molecules • Molecules may contain atoms of only one element - O2 • Molecules of compounds contain more than one element - H2O

19. Only a few atoms, even fewerChemical Bonds Ionic bonds Between metallic and non metallic atoms Easily dissolved by water Covalent Share at least one pair of electrons Polar and non-polar bonds Tight (high energy) bond Hydrogen bonds A hydrogen between atoms Not so tight (low energy) bond: 1/10th covalent

20. 1. Ionic Bonding • One atom loses electrons, becomes positively charged ion • Another atom gains these electrons, becomes negatively charged ion • Charge difference attracts the two ions to each other

21. Ion Formation • Atom has equal number of electrons and protons - no net charge • Atom loses electron(s), becomes positively charged ion • Atom gains electron(s), becomes negatively charged ion

22. Formation of NaCl • Sodium atom (Na) • Outer shell has one electron • Chlorine atom (Cl) • Outer shell has seven electrons • Na transfers electron to Cl forming Na+and Cl- • Ions remain together as NaCl

23. Formation of NaCl 7mm electron transfer SODIUM ATOM 11 p+ 11 e- CHLORINE ATOM 17 p+ 17 e- CHLORINE ION 17 p+ 18 e- SODIUM ION 11 p+ 10 e- Fig. 2.10a, p. 26

24. 2. Covalent Bonding Atoms share a pair or pairs of electrons to fill outermost shell • Single covalent bond • H2 Single bond • Double covalent bond • O2 Double bond • Triple covalent bond • N2 Triple bond

25. Non-polar Covalent Atoms share electrons equally Nuclei of atoms have same number of protons Example: Hydrogen gas (H-H) Polar Covalent Number of protons in nuclei of participating atoms is NOT equal Molecule held together by polar covalent bonds has no NET charge Electrons spend more time near nucleus with most protons Example: Water Electrons more attracted to O nucleus than to H nuclei Two Flavors of Covalent Bonds

26. + Polar Covalent Bonds slight negative charge at this end KEEP YOUR EYE ON THE ELECTRONS molecule has no net charge ( + and - balance each other) O H H slight positive charge at this end

27. Hydrogen Bonding A bond by Hydrogen between two atoms • Important for O and N • Lets two electronegative atoms interact • The H gives one a net + and the other one that is still – is attracted to it. • The H proton becomes “naked” because its electron gets pulled away.

28. - - - + - Covalent Bond Hydrogen Bond KEEP YOUR EYE ON THE ELECTRONS Hydrogen bond figure Like Charge Atoms Repel Each Other Opposite Charge Atoms Attract Each Other

29. Examples of Hydrogen Bonds one large molecule another large molecule a large molecule twisted back on itself

30. Properties of Water Polarity Temperature-Stabilizing Cohesive Solvent

31. Water Is a Polar Covalent Molecule • Molecule has no net charge • Oxygen end has a slight negative charge • Hydrogen end has a slight positive charge O H H

32. Liquid Water + H + H + _ O H O + _ + H +

33. Hydrophilic & HydrophobicSubstances • Hydrophilic substances • Polar • Hydrogen bond with water • Example: sugar • Hydrophobic substances • Non-polar • Repelled by water • Example: oil

34. Temperature-Stabilizing Effects • Water absorbs a lot more heat than other liquids, such as oil, before its temperature rises. • Why? • Heat is Vibration! • Molecules with lots of vibrational energy feel hot. • Much of the added energy disrupts hydrogen bonding rather than increasing the movement of molecules

35. Evaporation of Water • Large energy input can cause individual molecules of water to break free into air • As molecules break free, they carry away some energy (lower temperature) • Evaporative water loss is used by mammals to lower body temperature

36. Why Ice Floats • In ice, hydrogen bonds lock molecules in a lattice • Water molecules in lattice are spaced farther apart then those in liquid water • Ice is less dense than water

37. Water Cohesion • Hydrogen bonding holds molecules in liquid water together • Creates surface tension • Allows water to move as continuous column upward through stems of plants

38. Water Is a Good Solvent • Ions and polar molecules dissolve easily in water • When solute dissolves, water molecules cluster around its ions or molecules and keep them separated

39. Water as a solvent:Spheres of Hydration – – + + + + – + – + Na+ – + + + + – + Cl– – – + – + + + – + + – + Fig. 2.16, p. 29

40. Water • Solvent- polar • Keeps ions in solution • Doesn’t dissolve membranes • Heat management • Loosing heat • Holding heat • Density Changes

41. If it wasn’t ugly enough already: Hydrogen Ions: H+ • Unbound protons • Have important biological effects • Form when water ionizes

42. The pH Scale • Measures H+ concentration of fluid • Change of 1 on scale means 10X change in H+ concentration Highest H+ Lowest H+ 0---------------------7-------------------14 Acidic Neutral Basic

43. Examples of pH Pure water is neutral with pH of 7.0 Acidic Basic (Alkaline)

44. Acids Donate H+ when dissolved in water Acidic solutions have pH < 7 Bases Accept H+ when dissolved in water Acidic solutions have pH > 7 Acids & Bases

45. Buffers Carbonic Acid-Bicarbonate Buffer System Minimize shifts in pH • When blood pH rises, carbonic acid dissociates to form bicarbonate and H+ • H2C03 -----> HC03- + H+ • When blood pH drops, bicarbonate binds H+ to form carbonic acid • HC03- + H+ -----> H2C03 • Acidosis- too much CO2 in blood • Alkalosis- blood pH too low

46. Lecture 2: Chemistry of LifePart 2 Feeling a little burnt out?

47. Demonstration of Chemical Bonds Tests: 1. Water as a solvent 2. Bond strength Predictions: Covalent bonds Ionic bonds Hydrogen bonds Hydrophilic interactions Hydrophobic interactions

48. Hydrogen BondsAliphatic Resin, PVA and Elmer Why does glue work? • Mechanical component • Chemical component Process 1. Infiltrate wood fibers 2. Allow tight contact 3. Remove water (solvent) Demonstration of Hydrogen bond strength

49. Hydrogen BondsAliphatic Resin, PVA and Elmer • Bond Strength: • 3,500 pounds per square inch • Hydrogen bonds form between the wood and glue as the water leaves • Conclusion:

50. Organic Compounds • Hydrogen and other elementscovalently bonded to carbon • Major Classes of Biological Molecules • Carbohydrates • Lipids • Proteins • Nucleic Acids