1. Biology 102 Lecture 3: Atoms, molecules and life (cont.)

2. Lecture outline 1. Brief review of chemical bonds 2. Molecules: a primer 3. Why is water so important to life?  Basic structure of water molecules  Hydrophilic vs. hydrophobic molecules  Cohesion of water molecules  pH of water: acids, bases and buffers  Role of water in temperature regulation  The magic of ice!

3. 1. Brief review of chemical bonds Atoms react with each other in a way that helps them empty or fill their outer shells Atoms react with each other in a way that helps them empty or fill their outer shells Ionic bonds involve complete electron transfer Ionic bonds involve complete electron transfer Covalent bonds involve electron sharing Covalent bonds involve electron sharing

4. You can tell whether a particular atom tends to empty, fill or share its electrons based on its position on the periodic table!You can tell whether a particular atom tends to empty, fill or share its electrons based on its position on the periodic table!

5. 2. Molecules: a primer Molecule: Two or more atoms of the same or different elements Molecule: Two or more atoms of the same or different elements  If formed from >1 element, called a compound Most biological molecules utilize covalent bonding Most biological molecules utilize covalent bonding  Covalent bonds are stronger than ionic bonds  Most common atoms in biological molecules are C, O, N, P and S Non-polar vs. polar covalent bonds Non-polar vs. polar covalent bonds  Non-polar: equal electron sharing  Polar: unequal electron sharing

6. Non-polar bonds: Non-polar bonds:  Oxygen molecules (O 2 )  Share two electron pairs equally O=O NOTE: Number of bonds is the number of electron pairs that are shared! NOTE: Number of bonds is the number of electron pairs that are shared!  Carbon-carbon bonds (within other molecules)  C-C: sharing of one electron pair (note that this is not a complete molecule) Polar bonds Polar bonds  Carbon-oxygen bonds (within other molecules)  C + -O - : Electrons attracted more strongly by oxygen  (Why?)

7. 3. Why is water so important to life?

8. Basic structure of water molecules Note polarity Note polarity  Why polar? Note hydrogen bonds Note hydrogen bonds  Attractive force between negative O and positive H  NOT between to H atoms!!! Audesirk et al., Fig. 2-5

9. Water molecules stick together Cohesion: tendency of water molecules to stick together Cohesion: tendency of water molecules to stick together  Due to hydrogen bonds  Keeps water fluid at wide range of temperatures  Most similarly-sized molecules are gases at room temperature Surface tension: cohesion of water molecules at the surface Surface tension: cohesion of water molecules at the surface Adhesion: tendency of water molecules to stick to other polar substances Adhesion: tendency of water molecules to stick to other polar substances  “capillary action” is a result of adhesion

10. “Universal solvent”: Water dissolves ionic compounds and polar substances Audesirk et al., Fig. 2-6 Salt crystal dissolved in water Audesirk et al., Fig. 2-7 Glucose molecule dissolved in water Note how negative oxygens of water molecule interact with positively charged ions or parts of polar molecules, and positive hydrogens interact with negative ions or parts of molecules. - +

11. Polar molecules are hydrophilic Polar molecules are hydrophilic  “Hydro”=water, “Philic” = loving  Dissolve in water Non-polar molecules are hydrophobic Non-polar molecules are hydrophobic  “Hydro” = water; “Phobic” = fearing  Those molecules are not dissolved in water, but rather tend to clump together in a hydrophobic interaction  Example: Fats and oils are non-polar  This clustering of hydrophobic molecules (or parts of molecules) is important in the formation of cell membranes

12. pH of water: Acids, bases and buffers Water molecules can break apart to form hydrogen ions (H + ) and hydroxyl ions (OH - ) Water molecules can break apart to form hydrogen ions (H + ) and hydroxyl ions (OH - )  The “shared” electron is completely transferred to the oxygen  Pure water has equal numbers of these two ions Audesirk et al., Ch. 2, p. 30

13. Other substances, called acids and bases, can dissociate in similar ways: Other substances, called acids and bases, can dissociate in similar ways: ACIDS: Dissociate to create excess H + ACIDS: Dissociate to create excess H +  Example: Hydrochloric acid (HCl)  HCl  H + + Cl- BASES: Dissociate to create excess OH - BASES: Dissociate to create excess OH -  Example: Sodium hydroxide (NaOH)  NaOH  Na + + OH - Both acids and bases are corrosive Both acids and bases are corrosive  Can be quite damaging to living systems!

14. Degree of acidity is expressed in pH units: Degree of acidity is expressed in pH units:  pH = -log [H + ]  Example: Pure water has [H+] = 10 -7, so pH = 7  A difference of 1 pH unit corresponds to a tenfold change in [H + ]  Works like the Richter scale for earthquakes!

15. pH of some common substances pH of some common substances

16. Buffers help maintain constant pH in living systems Buffers are able to “accept” H + ions when they are in excess (too acidic) Buffers are able to “accept” H + ions when they are in excess (too acidic) HCO 3 - + H + (excess)  H 2 CO 3 HCO 3 - + H + (excess)  H 2 CO 3 bicarbonate carbonic acid bicarbonate carbonic acid Buffers are able to “give up” H + ions when they are at low levels (too basic) Buffers are able to “give up” H + ions when they are at low levels (too basic) H 2 CO 3  HCO 3 - + H + H 2 CO 3  HCO 3 - + H +

17. Role of water in temperature regulation What is temperature? What is temperature?  Speed of molecular motion Why does temperature regulation matter? Why does temperature regulation matter?  High temperature: damage to proteins  Low temperature: ice crystals rupture cells It takes lots of energy to heat water. It takes lots of energy to heat water.  Why? Water remains liquid at a high range of temperatures. Water remains liquid at a high range of temperatures.  Why?

18. The magic of ice What if ice, like most other solids, sank instead of floating? What if ice, like most other solids, sank instead of floating?