2. WATER : THE SOLVENT FOR BIOCHEMICAL REACTIONS

Transcription

1 2. WATER : THE SOLVENT FOR BIOCHEMICAL REACTIONS

2 2.1 Water and Polarity Both geometry and properties of molecule determine polarity Electronegativity - The tendency of an atom to attract electrons to itself in a chemical bond.

3 2.1 Water and Polarity Both geometry and properties of molecule determine polarity The difference in electronegativity partial positive and negative charge, pictured as δ + & δ - Polar Nonpolar - O-H bonds in water - CH 4, CO 2 δ- 2δ+ δ- O C O

4 2.1 Water and Polarity Solvent Properties of Water The polar nature of water determines its solvent properties. Hydrophilic ( water-loving, from the Greek) Hydrophobic ( water-hating, from the Greek) Amphipathic

5 2.1 Water and Polarity Solvent Properties of Water Hydrophilic *Hydration shells surrounding ions in solution *ion-dipole and dipole-dipole interaction

6 2.1 Water and Polarity Solvent Properties of Water Hydrophilic - Electrostatic attraction between unlike charges - Lowering energy, ion-dipole and dipole-dipole interaction - Ionic compounds with full charges, polar compounds with partial charges

7 2.1 Water and Polarity Solvent Properties of Water Hydrophobic - Hydrocarbons are nonpolar. - Tend not to dissolve in water - The permanent dipole of the water molecule can induce a temporary dipole in the nonpolar molecule. A dipole-induced dipole interaction - hydrophobic interactions: Interactions between nonpolar molecules.

8 2.1 Water and Polarity Solvent Properties of Water Amphipathic - A single molecule may have both polar (hydrophilic) and nonpolar (hydrophobic) portions - e.g., palmitic acid

9 2.1 Water and Polarity Solvent Properties of Water Amphipathic - This compound in the presence of water tends to form structures called micelles. - Van der Waals interaction: Very weak interaction between short-lived temporary dipoles and the dipoles they induce

10 2.1 Water and Polarity

11 2.2 Hydrogen Bonds Noncovalent interaction Bonding of electrostatic origin A special case of dipole-dipole interaction When hydrogen is covalently bonded to an electronegative atom such as oxygen or nitrogen, it has a partial positive charge due to the polar bond. This partial positive charge on hydrogen can interact with an unshared pair of electrons on another electronegative atom. Hydrogen-bond donor and hydrogen-bond acceptor

12 2.2 Hydrogen Bonds Hydrogen bonding sites in HF, H 2 O, and NH 3

13 2.2 Hydrogen Bonds Liquid water and ice crystals - Hydrogen bonding between water molecules in the regular lattice structure of the ice crystal.

14 2.2 Hydrogen Bonds Liquid water and ice crystals Liquid water - Hydrogen bonds are constantly breaking and new ones are constantly forming. - A cluster can break up and re-form in to seconds in water at 25. Ice crystal - A more-or-less-stable arrangement - Lower density, because the fully hydrogen-bonded array is less densely packed ice cubes and icebergs in water

15 2.2 Hydrogen Bonds Hydrogen bonds are much weaker than normal covalent bonds.

16 2.2 HYDROGEN BONDS Both the melting point and the boiling point of water are significantly higher than would be predicted for a molecule of this size.

17 2.2 Hydrogen Bonds Hydrogen bonding also plays a role in the behavior of water as a solvent. Alcohol Ketone Amine

18 2.2 What Is a Hydrogen Bond? Biologically Important Hydrogen Bonds Other Than to Water Molecules The hydrogen bonds have a vital involvement in stabilizing the three-dimensional structures of biologically important molecules, including DNA, RNA, and proteins.

19 2.2 HYDROGEN BONDS How basic chemistry affects life: The importance of the hydrogen bond

20 2.3 Acids, Bases and ph Acids and bases - A acid: a molecule that acts as a proton (hydrogen ion) donor - A base: a molecule that acts as a proton acceptor - Acid strength: amount of hydrogen ion released when a given amount of acid is dissolved in water - Acid dissociation constant, K a HA H + + A - Acid Conjugative base K a = [H+ ][A - ] [HA]

21 2.3 Acids, Bases and ph The acid-base reaction is a proton-transfer reaction. HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) Acid Base Conjugate Conjugate acid to H 2 O base to HA

22 2.3 Acids, Bases and ph Important part of water in biological process. - Central role of water as a solvent - Self-dissociation of water to hydrogen ion and hydroxide ion is small. The water molecule is itself part of a cluster of such molecules.

23 2.3 Acids, Bases and ph The degree of dissociation of water - [H K a = + ][OH - ] [H 2 O] The molar concentration of pure water, [H 2 O]= 55.5 M Thus, K a = [H + ][OH - ] 55.5 K a ⅹ55.5 = [H + ][OH - ]= K w Continue

24 2.3 Acids, Bases and ph K w, ion product constant for water - The numerical value of K w can be determined experimentally by measuring the hydrogen ion concentration of pure water. - At 25 in pure water, [H + ]= 10-7 M = [OH - ] Thus, at 25, the numerical value of K w is K w =[H + ][OH - ]= (10-7 )(10-7 )= This relationship is valid for any aqueous solution. Continue

25 2.3 Acids, Bases and ph ph - By exponential notation ph = -log 10 [H + ] - A difference of one ph unit implies a tenfold difference in hydrogen ion concentration. - Pure water with a ph of 7 is neutral, acidic solutions have ph values lower than 7, and basic solutions have ph values higher than 7. pk a = -log 10 K a - Another measure of acid strength - The smaller its value, the stronger the acid. Continue

26 2.3 Acids, Bases and ph

27 2.3 Acids, Bases and ph Monitoring Acidity Equation that connects the K a of any weak acid with the ph of a solution. - In biological practice, it is necessary to control ph for optimum reaction conditions. - The activities of three enzymes affected by ph

28 2.3 Acids, Bases and ph Monitoring Acidity Equation that connects the K a of any weak acid with the ph of a solution. - Derive the involved equation... K a = [H + ][A - ] [HA] log K a = log [H + ] + log -log [H + ] = -log K a + log [A - ] [HA] Then, use the definitions of ph and pk a [A ph = pk a + log - ] [HA] [A - ] [HA] Henderson-Hasselbalch equation Continue

29 2.4 Titration Curves A titration - An experiment in which measured amounts of base are added to a measured amount of acid. - Equivalence point: the point in the titration at which the acid is exactly neutralized 29/43

30 2.4 Titration Curves A way to keep track of protonated and deprotonated forms of acids and their conjugate bases ph < pk a H + on, substance protonated (e.g., HA form) ph > pk a H + off, substance deprotonated (e.g., A- form) [A ph - pk a = log - ] [HA]

31 2.5 Buffers Buffer solution - A mixture of a weak acid and its conjugate base - Tend to resist a change in ph on the addition of moderate amounts of strong acid or base.

32 2.5 Buffers Example 1 - A solution that contains the monohydrogen phosphate and dihydrogen phosphate ions, HPO 4 2- and H 2 PO Henderson-Hasselbalch equation can be used to calculate the HPO 4 2- / H 2 PO 4 - ratio. - [HPO 4 2- ]= M and [H 2 PO 4- ]= 0.1 M conjugate base/weak acid ratio = 0.63 ph = 7.0 If 1.0 ml of 0.10 M HCl is added to 99.0 ml of the buffer, [HPO 4 2- ] + H + [H 2 PO 4- ] Continue

33 2.5 Buffers Example 1 The new ph can be calculated. (pk a 7.20) ph = pka + log [HPO 4 2- ] [H 2 PO 4- ] ph = log new ph = 6.99 Much smaller change than in the unbuffered pure water

34 2.5 Buffers Example 2 - If 1.0 ml of 0.1 M NaOH, H 2 PO OH - HPO 4 2- new ph = 7.01 Much smaller change than in the pure water

35 2.5 Buffers How buffers work? - The ph of a sample being titrated changes very little in the vicinity of the infection point. - Second stage of ionization H 2 PO 4 - H + + HPO 4 2-

36 2.5 Buffers Example 3 - If we choose a ph of 8.2 for a buffer composed of H 2 PO 4 - and HPO 4 2-, ph = pk a + log 8.2 = log [HPO 2-4 ] [H 2 PO 4- ] [HPO 2-4 ] [H 2 PO 4- ] 1 = log [HPO 4 2- ] [H 2 PO 4- ] [HPO 4 2- ] [H 2 PO 4- ] = 10

37 2.5 Buffers When the ph is one unit higher than the pk a, the ratio of the conjugate base form to the conjugate acid form is 10.

38 2.5 Buffers A range of about two units in which the buffer is effective - The H 2 PO 4- /HPO 4 2- pair is suitable as a buffer near ph 7.2, and the CH 3 COOH/CH 3 COO - is suitable near ph 4.76.

39 2.5 Buffers Buffering capacity - Low buffering capacity: a buffer solution with low concentrations of both the acid and base forms - High buffering capacity: a buffer that contains greater amounts of both acid and base - Too high buffering capacity results in high ionic strength which may be harmful to biological systems.

40 2.5 Buffers How We Make Buffers Two forms of the buffer present in the solution at reasonable quantities. - HA and A - are interconverted by adding strong acid or strong base.

41 2.5 Buffers Buffer Systems of Physical Importance Phosphate buffer system (H 2 PO 4- /HPO 4 2- ) in living organisms at ph 7 - Common in laboratory - In most intracellular fluids (Phosphate ion concentration is high) Dissociation of carbonic acid (H 2 CO 3 ) in blood H 2 CO 3 H + + HCO Carbon dioxide can dissolve in water CO 2 (g) + H 2 O(l) H + (aq) + HCO 3- (aq) - CO 2 vs. lactic acid in blood near muscle

42 2.5 Buffers Buffer Systems of Physical Importance Other buffers, zwitterions

43 THE END!! Start by doing what s necessary, then what s possible, and suddenly you are doing the impossible. - St. Francis of Assisi