1. Chapter 2 Life’s Chemical Basis

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

3. Mass Number Number of protons + Number of neutrons Isotopes vary in mass number

4. 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

5. Radioisotopes • Have an unstable nucleus that emits energy and particles • Radioactive decay transforms radioisotope into a different element • Decay occurs at a fixed rate

6. Radioisotopes as Tracers • Tracer is substance with a radioisotope attached to it • Emissions from the tracer can be detected with special devices • Following movement of tracers is useful in many areas of biology

7. Radioisotopes in Medicine • Positron-Emission Tomography (PET) uses radioisotopes to form images of body tissues • Patient is injected with tracer and put through a PET scanner • Body cells absorb tracer at different rates • Scanner detects radiation caused by energy from decay of the radioisotope, and radiation then forms an image • Image can reveal variations and abnormalities in metabolic activity

8. Other Uses of Radioisotopes • Drive artificial pacemakers • Radiation therapy Emissions from some radioisotopes can destroy cells. Some radioisotopes are used to kill small cancers.

9. ELEMENTS IN LIVING THINGS • 4 major elements in all living things: O, C, H, N (decreasing amounts) • TRACE ELEMENTS—required only in small amounts Ex: iron, iodine, copper

10. What Determines Whether Atoms Will Interact? The number and arrangement of their VALENCE electrons

11. Shell Model • First shell • Lowest energy • Holds 1 orbital with up to 2 electrons • Second shell • 4 orbitals hold up to 8 electrons

12. Electron Vacancies • Unfilled shells make atoms likely to react AND bond

13. Chemical Bonds, Molecules, & Compounds • Bond is union between electron structures of atoms • Atoms bond to form molecules

14. Important Bonds in Biological Molecules • Ionic Bonds • Covalent Bonds • Hydrogen Bonds

15. Ionic Bonding • One atom loses electrons, becomes positively charged ion • Between metals & nonmetals • Another atom gains these electrons, becomes negatively charged ion • This type of bond is LEAST affected by water’s presence

16. Fig. 2-8a(2), p.24

17. Covalent Bonding Atoms SHARE a pair or pairs of electrons to fill outermost shell • Single covalent bond • Double covalent bond • Triple covalent bond

18. Nonpolar Covalent Bonds • Atoms share electrons equally • Example: Hydrogen gas (H-H)

19. Polar Covalent Bonds • Unequal sharing of electrons • Ex: Water - Electrons more attracted to O nucleus than to H nuclei

20. 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 + +

21. Hydrogen Bonding The weakest type of bond (20 x easier to break than a covalent bond); are rapidly broken and made This is what holds water molecules together The more hydrogen bonds in a molecule= the more stable the molecule

22. Hydrophilic & HydrophobicSubstances • Hydrophilic substances • Polar • Likewater (so hydrogen bond to it) • Example: sugar • Hydrophobic substances • Nonpolar • Repelled by water • Example: oil

23. Properties of Water • Polar (dissolves polar solutes)

24. Temperature-Stabilizing Effects • Water can absorb much heat before its temperature rises (has a high specific heat capacity) • How is this important for aquatic organisms?

25. Evaporation of Water • Lots of heat must be added to break H-bonds & vaporize water • As it vaporizes, it carries a lot of heat with it(lower the temperature)—AKA high heat of vaporization • Evaporative water loss is used by mammals to lower body temperature

26. WHY SOLID WATER (ICE )FLOATS IN LIQUID WATER • 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(water contracts when cooled to 4 C, but expands from 4 to O C • This is why lakes freeze from top to bottom (helps insulate aquatic life below)

27. High Cohesion & Adhesion • Involves hydrogen bonds • Cohesion—water molecules stick to each other; Adhesion—water molecules stick to other things • Creates surface tension (Ex: water striders) • Allows water to move as continuous column upward through stems of plants(AKA transpiration)

28. The pH Scale • Measures H+(hydronium ion)concentration • Change of 1 on scale means 10X change in H+concentration Ex: a substance with pH of 2 vs. a pH of 6 is 10,000 more acidic Highest H+ Lowest H+ 0---------------------7-------------------14 Acidic Neutral Basic

29. Acids & Bases • Acids • Donate H+ (hydronium)when dissolved in water • Acidic solutions have pH < 7 • Bases • Accept H+ when dissolved in water; contain hydroxide ion • Acidic solutions have pH > 7

30. More hydronium = lower pH = more acidic

31. Weak and Strong Acids • Weak acids • Reluctant H+ donors • Can also accept H after giving it up • Carbonic acid (H2CO3) is example • Strong acids • Completely give up H+ when dissolved • Hydrochloric acid (HCl) is example

32. ACID RAIN • A coal-burning power plant emits sulfur dioxide, which dissolves in water vapor to form acid rain (pg. 29) Fig. 2-13, p.29

33. Buffer Systems • Minimizes or helps prevent changes in pH • Partnership between weak acid and base it forms when dissolved • Two work as pair to counter shifts in pH

34. Carbonic Acid-Bicarbonate Buffer System—found in our blood & in the ocean • When blood pH rises due to increased carbon dioxide levels, carbonic acid breaks apart to form bicarbonate and H+ H2C03 -----> HC03- + H+ • When blood pH drops, bicarbonate binds H+ to form carbonic acid HC03- + H+ -----> H2C03