1. WATER

2. COURSE OUTCOME (C0 2)  CO2: Ability to define and describe the structure, classification and function of biomolecules and water  Terms used in Course Outcome and Teaching  Knowledge: Define, introduce, describe, name, relate, explain, identify and Remember concepts and principles.  Repetition: Repeat and discuss concepts and principles.  Application: Apply, demonstrate, interpret and illustrate concepts and principles.

3. LECTURE CONTENTS 1. Why water is important to biochem ? 2. USES OF WATER 3. PHYSICS & CHEMISTRY OF WATER 4. PHYSICAL PROPERTIES OF WATER 5. Molecular Structure of Water 6. Hydrogen Bonding 7. Polarity of water 8. Noncovalent Bonding 9. Weak, van der Waal’s forces 10. Thermal properties of water 11. Osmosis, reverse osmosis & dialysis 12. Water ionization, pH, titration and buffer 13. Summary 14. The end note Sect

4. Sect 1. Why water is important to biochem ?  More than 70% of the earth’s surface is covered with the molecule of water.  Cell components (protein, poly sacharides, nucleic acid, membranes) asume their shape in response to water  Water acts as a solvent & substrate for many cellular reactions  Water is a common chemical substance that is essential for the survival of all known forms of life. ( In typical usage, water refers only to its liquid form or state, but the substance also has a solid state, ice, and a gaseous state, water vapor. )chemical substance lifeliquidstatesolidicegaseouswater vapor

5. Sect 2 : USES OF WATER  AGRICULTURE  FOR DRINKING  AS SOLVENT  HEAT TRANSFER FLUID  FOOD PROCESSING  INDUSTRIAL APPLICATIONS  AS A SCIENTIFIC STANDARD

6. Sect 3 : Physics & Chemistry of water  Water is the chemical substance with chemical formula H2O: one molecule of water has two hydrogen atoms covalently bonded to a single oxygen atom.chemical substancechemical formula HOmoleculehydrogenatoms covalentlybondedoxygen  Water is a tasteless, odorless liquid at ambient temperature and pressure, and appears colorless in small quantities, although it has its own intrinsic very light blue hue.ambient temperature and pressure  Oxygen attracts electrons much more strongly than hydrogen, resulting in a net positive charge on the hydrogen atoms, and a net negative charge on the oxygen atom.  The presence of a charge on each of these atoms gives each water molecule a net dipole moment.dipole moment

7.  Electrical attraction (hydrogen bonding) between water molecules due to this dipole pulls individual molecules closer together, making it more difficult to separate the molecules and therefore raising the boiling point.  Water can be described as a polar liquid that dissociates disproportionately into the hydronium ion (H3O+(aq)) and an associated hydroxide ion (OH−(aq))hydroniumhydroxide  Water is in dynamic equilibrium between the liquid, gas and solid states at standard temperature and pressure (0°C, 100.000 kPa), and is the only pure substance found naturally on Earth.dynamic equilibriumliquidgassolidstatesstandard temperature and pressure

8. Sect 4 : UNIQUE PHYSICAL PROPERTIES OF WATER  Exist in all three physical states of matter: solid, liquid, and gas.  Has high specific heat  Water conducts more easily than any liquid except mercury  Water has a high surface tension  Water is a universal solvent  Water in a pure state has a neutral pH.pH

9. Sect 5 : Molecular Structure of Water  The oxygen in water is hybridized. Hydrogens are bonded to two of the orbitals. Consequently the water molecule is bent. The H-O-H angle is 104.5 o.

10. Sect 6: Hydrogen Bonding  Water molecules hydrogen bond with one another. Four hydrogen bonding attractions are possible per molecule: 2 through the hydrogens and 2 through the nonbonding electron pairs.

11. Hydrogen Bonding  A hydrogen attached to an O or N becomes very polarized and highly partial plus. This partial positive charge interacts with the nonbonding electrons on another O or N giving rise to the very powerful hydrogen bond. hydrogen bond shown in yellow

12.  Water has an abnormally high boiling point due to intermolecular hydrogen bonding. H bonding is a weak attraction between an electronegative atom in one molecule and an H (on an O or N) in another.

13.  Water molecule with bond and net ( ) dipoles. ++ -- ++

14.  Water is a polar molecule. A polar molecule is one in which one end is partially positive and the other partially negative. This polarity results from unequal sharing of electrons in the bonds and the specific geometry of the molecule. Sect 7. Polarity of water

15. Hydrophillic substances dissolve in water  The polar nature of water makes it an excellent solvent for polar and ionic materials that are water loving (hydophillic)  An ion immersed in a polar solvent such as water attracts the the oppositely charged ends of the solvent dipoles and becomes surrounded by one or more concentric shells of orientated solvent molecules, therby becoming solvated or hydrated when water is the solvent  The bond dipoles of uncharged polar molecules make them soluble in aqueous solutions.

16. HYDROPHOBIC INTERACTIONS  Nonpolar molecules tend to coalesce into droplets in water. The repulsions between the water molecules and the nonpolar molecules cause this phenomenon.  The water molecules form a “cage” around the small hydrophobic droplets.

17. WATER-SOLVENT PROPERTIES  Water dissolves chemicals that have an affinity for it, ie. hydrophilic (water loving) materials. many ionic compounds polar organic compounds  These compounds are soluble in water due to three kinds of noncovalent interactions: 1. ion-dipole 2. dipole-dipole 3. hydrogen bonding

18. Dipole-dipole Interactions  The polar water molecule interacts with an O or N or an H on an O or N on an organic molecule. Dipole-dipole interactions

19. Ion-dipole Interactions  Ions are hydrated by water molecules. The water molecules orient so the opposite charge end points to the ion to partially neutralize charge. The shell of water molecules is a solvation sphere.

20. Nonpolar Molecules  Nonpolar molecules have no polar bonds or the bond dipoles cancel due to molecular geometry.  These molecules do not form good attractions with the water molecule. They are insoluble and are said to be hydrophobic (water hating).  eg.: CH 3 CH 2 CH 2 CH 2 CH 2 CH 3, hexane

21. Nonpolar Molecules-2  Water forms hydrogen-bonded cage like structures around hydrophobic molecules, forcing them out of solution.

22. Amphipathic Molecules  Amphipathic molecules contain both polar and nonpolar groups.  Ionized fatty acids are amphipathic. The carboxylate group is water soluble and the long carbon chain is not.  Amphipathic molecules tend to form micelles, colloidal aggregates with the charged “head” facing outward to the water and the nonpolar “tail” part inside.

23. A Micelle 

24. Sect 8 : Noncovalent Bonding  Ionic interactions  Hydrogen bonding  Van der Waals forces Dipole-dipole Dipole-induced dipole Induced dipole-induced dipole

25. Typical “Bond” Strengths TypekJ/mol Covalent>210 Non-covalent Ionic interactions4-80 Hydrogen bonds12-30 van der Waals0.3-9 Hydrophobic interactions3-12

26.  These forces are electrostatic interactions arising from noncovalent associations between neutral molecules.  These interactions occur among permanent or induced dipoles, where the hydogen bonds form a special kind of dipolar interaction  Interactions among permanent dipoles such as carbonyl groups are much weaker than ionic interactions Sect 9: Weak, van der Waal’s forces

27. Ionic Interactions  Ionic interactions occur between charged atoms or groups.  In proteins, side chains sometimes form ionic salt bridges, particularly in the absence of water which normally hydrates ions. Salt bridge

28.  A permanent dipole can also induce a dipole situation in a neighbouring group by electrostatically distortng its electron distribution.  These dipole-induced interactions are generally much weaker than dipole-dipole interactions.  Non-polar molecules, at any instant, can randomly orientate dipole moment, resulting from the rapid fluctuating motion of their electrons, and polarize and attract neighbouring groups.  These so called dispersion forces are extremely weak and are called London dispersion forces.

29. Van der Waals Attractions a. Dipole-dipole b. Dipole-induced dipole c. Induced dipole-induced dipole

30. Sect 10 : Thermal properties of water  Hydrogen bonding keeps water in the liquid phase between 0 o C and 100 o C.  Liquid water has a high: Heat of vaporization - energy to vaporize one mole of liquid at 1 atm Heat capacity - energy to change the temperature by 1 o C  Water plays an important role in thermal regulation in living organisms.

31. Sect 11: OSMOSIS, Reverse Osmosis & Dialysis  Osmosis is a spontaneous process in which solvent (e.g. water) molecules pass through a semi permeable membrane from a solution of lower solute (e.g. chemical) concentration to a solution of higher solute concentration.  Osmosis is the movement of solvent from a region of high concentration (here, pure water) to a region of relatively low concentration (water containing dissolved solute).  Water moves by osmosis and solutes by diffusion.

32. OSMOTIC PRESSURE  Osmotic pressure is the pressure required to stop osmosis or the influx of water (22.4 atm for 1M solution).  Because cells have a higher ion concentration than the surrounding fluids, they tend to pick up water through the semi permeable cell membrane.  The cell is said to be hypertonic relative to the surrounding fluid and will burst (hemolyze) if osmotic control is not effected.

33.  Cells placed in a hypotonic solution will lose water and shrink (crenate).  If cells are placed in an isotonic solution (conc. same on both sides of membrane) there is no net passage of water.

34. Definitions of solutions  Hypotonic solution: A solution with a lower salt concentration than in normal cells of the body and the blood. salt  Hypertonic solution: A solution with a higher salt concentration than in normal cells of the body and the blood. salt  Isotonic solution: A solution that has the same salt concentration as the normal cells of the body and the blood. An isotonic beverage may be drunk to replace the fluid and minerals which the body uses during physical activity. salt

35. OSMOMETER  Osmotic pressure (  ) is measured in an osmometer.

36. Osmotic-pressure formula  iMRT i = van’t Hoff factor (% as ions) M = molarity (mol/L for dilute solutions) R = (normal gas constant expressed in liters and atmospheres) 0.082 L atm/ mol K T = Kelvin temperature Or π= i*C*R*T T= absolute temperature (in Kelvin) R= the gas constant in whatever units you need to express osmotic pressure (e.g. if you want π in atm then R=0.082 L*atm/(mole*K)) C = the concentration of your solute in mole/L

37.  i is the van't Hoff coefficient. For non- electrolytes i=1  For strong electrolytes i= the number of ions that are produced by the dissociation according to the molecular formula e.g for NaCl you have 2 ions (1 Na+ and 1 Cl-) so i=2, for CaCl2, 3 ions (1 Ca+2 from 2 Cl-) so i=3.  For weak electrolytes, if n is the number of ions coming from the 100% dissociation according to the molecular formula and a the degree of dissociation then i=(1-a)+na. E.g. if we assume for CH3COOH a=80% i=(1- 0.8)+2*0.8= 0.2+1.6=1.8

38. Liquids move from high osmotic pressure (high conc. solvent and low conc. solute) to low osmotic pressure (high conc. solute and low conc. of solvent)

39. REVERSE OSMOSIS  Reverse osmosis (RO) is a separation process that uses pressure to force a solution through a membrane that retains the solute on one side and allows the pure solvent to pass to the other side.solution membranesolutesolvent  More formally, it is the process of forcing a solvent from a region of high solute concentration through a membrane to a region of low solute concentration by applying a pressure in excess of the osmotic pressure.osmotic pressure  It is used in water purification and desalination.

40. DIALYSIS  A concentrated solution is separated from a large volume of solvent by a dialysis membrane or bag that is permeable to both water and solutes.  Only small molecules can diffuse through the pores of the membrane.  At equilibrium, the concentrations of small molecules are nearly the same on either side of the membrane, whereas the macromolecules, such as proteins or nucleic acids, remain inside the dialysis bag.

41. KIDNEY DIALYSIS  Reverse osmosis is the technique used in dialysis, which is used by people with kidney failure.dialysis kidney  The kidneys filter the blood, removing waste products (e.g. urea) and water, which is then excreted as urine.urine  A dialysis machine mimics the function of the kidneys. The blood passes from the body via a catheter to the dialysis machine, across an osmotic filter.dialysis

42. Sect 12: Water ionization, pH, titration and buffer  The self-ionization of water is the chemical reaction in which two water molecules react to produce a hydronium (H3O + ) and a hydroxide ion (OH − ).hydroniumhydroxide  Water ionization occurs endothermically due to electric field fluctuations between molecules caused by nearby dipole librations resulting from thermal effects, and favorable localized hydrogen bonding.  Ions may separate but normally recombine within a few min. to seconds. Rarely (about once every eleven hours per molecule at 25°C, or less than once a week at 0°C) the localized hydrogen bonding arrangement breaks before allowing the separated ions to return, and the pair of ions (H+, OH-) hydrate independently and continue their separate existence.

43. Ionization of Water  Water dissociates. (self-ionizes)  H 2 O + H 2 O = H 3 O + + OH - K w = K a [H 2 O] 2 = [H 3 O + ][OH - ]

44. Water Ionization-2  The conditions for the water dissociation equilibrium must hold under all situations at 25 o C. K w = [H 3 O + ][OH - ]=1 x 10 -14  In neutral water, [H 3 O + ] = [OH - ] = 1 x 10 -7 M

45. Water ionization - 3  When external acids or bases are added to water, the ion product ([H 3 O + ][OH - ] ) must equal K w.  The effect of added acids or bases is best understood using the Bronsted- Lowry- theory of acids and bases.

46. Bronsted-Lowry definitions An acid is a substance that can donate a proton A base is a substance that can accept a proton H+ ions (called a protons, since a H+ ion has neither electrons nor neutrons). This definition can be represented by the general chemical reaction A  B + H+ which does not attempt to show electrical charge balance. In this equation - · A is the acid., · B is the base and · H+ (a hydrogen atom without an electron) is a proton.

47. Conjugate acid/base  An acid can donate a proton  An acid (HA) reacts with a base (H 2 O) to form the conjugate base of the acid (A-) and the conjugate acid of base (H 3 O + ) HA + H 2 O = H 3 O + + A - A B CA CB C: conjugate (product) A/B

48. Conjugate base/acid base = proton acceptor RNH 2 + H 2 O = OH - + RNH 3 + B A CB CA

49. Measuring Acidity  Added acids increase the concentration of hydronium ion and bases the concentration of hydroxide ion.  In acid solutions [H 3 O + ] > 1 x 10 -7 M [OH - ] < 1 x 10 -7 M  In basic solutions [OH - ] > 1 x 10 -7 M [H 3 O + ] < 1 x 10 -7 M  pH scale measures acidity without using exponential numbers.

50. pH Scale  Define: pH = - log (10) [H 3 O + ] 0---------------7---------------14 acidic basic [H 3 O + ]=1 x 10 -7 M, pH = ? 7.0

52. Strength of Acids  Strength of an acid is measured by the percent which reacts with water to form hydronium ions.  Strong acids (and bases) ionize close to 100%. e.g.. HCl, HBr, HNO 3, H 2 SO 4

53. Strength of Acids-2  Weak acids (or bases) ionize typically in the 1-5% range. e.g.. CH 3 COCOOH, pyruvic acid CH 3 CHOHCOOH, lactic acid CH 3 COOH, acetic acid

54. Strength of Acids-3  Strength of an acid is also measured by its K a or pK a values. HA + H 2 O = H 3 O + + A - Larger K a and smaller pK a values indicate stronger acids.

55. Strength of Acids-4 K a pK a CH 3 COCOOH 3.2x10 -3 2.5 CH 3 CHOHCOOH 1.4x10 -4 3.9 CH 3 COOH 1.8x10 -5 4.8 Larger K a and smaller pK a values indicate stronger acids.

56. Monitoring Acidity  The Henderson-Hasselbalch (HH) equation is derived from the equilibrium expression for a weak acid.

57. Monitoring Acidity-2  The HH equation enables us to calculate the pH during a titration and to make predictions regarding buffer solutions.  What is a titration? It is a process in which carefully measured volumes of a base are added to a solution of an acid in order to determine the acid concentration.

58. Monitoring Acidity-3  When chemically equal (equivalent) amounts of acid and base are present during a titration, the equivalence point is reached.  The equivalence point is detected by using an indicator chemical that changes color or by following the pH of the reaction versus added base, ie. a titration curve.

59. Glutamic acid titration curve

60. Titration Curve – solved example  0.7 equivalents of NaOH neutralizes 0.7 eq of acid producing 0.7 eq of salt and leaving 0.3 eq of unneutralized acid.  pK a of HOAc is 4.76 30% acid and 70% salt. pH=5.13

61. Buffer Solutions  Buffer : a solution that resists change in pH when small amounts of strong acid or base are added.  A buffer consists of: a weak acid and its conjugate base or a weak base and its conjugate acid

62. Buffer Solutions - 1  High concentrations of acid and conjugate base give a high buffering capacity.  Buffer systems are chosen to match the pH of the physiological situation, usually around pH 7.

63. Buffer Solutions- 2  Within cells the primary buffer is the phosphate buffer: H 2 PO 4 - /HPO 4 2-  The primary blood buffer is the bicarbonate system: HCO 3 - /H 2 CO 3.  Proteins also provide buffer capacity. Side chains can accept or donate protons.

64. Buffer Solutions-3  A Zwitter ion is a compound with both positive and negative charges.  Zwitterionic buffers have become common because they are less likely to cause complications with biochemical reactions.

65. Example of Zwitter ion buffer  N-tris(hydroxymethyl)methyl-2- aminoethane sulfonate (TES) is a example of zwitterion buffer. (HOCH 2 ) 3 CN + H 2 CH 2 CH 2 SO 3 -

66. Buffer Solutions-4  Buffers work by chemically tying up acid and base. Eg.:

67. Solved example on Buffer Calculate the ratio of lactic acid to lactate in a buffer at pH 5.00. The pK a for lactic acid is 3.86 = 13.8

68. Sect 13: SUMMARY  Water is essential for all living things.  Water molecules can form hydrogen bonds with other molecules because they have 2 H atoms that can be donated ans 2 unshared electron pairs that can act as acceptors.  Liquid water is an irregular network of water molecules that each form 4 hydrogen bonds with neighbouring water molecules.

69. Summary contd.  Hydophilic substances such as ions and polar molecules dissolve readily in water.  The hydrophobic effect is the tendency of water to minimizeits contact with nonpolar substances.  Water molecules move from regions of high concentration to regions of low concentration by osmosis.  Solutes move from regions of high conc. to regions of low conc. by diffusion.

70. Summary contd.  Water ionizes to H + (which represents the hydronium ion H 3 O) and OH -.  The concentration of H + ions in solutions is expressed as a pH value.  Acids can donate protons and bases accept protons.  The strength of an acis is expressed as its pK.  Henderson-Hasselbalch equation relates the pH of a solution to the pK and concentration of an acid to its conjugate base.  Buffered solutions resist changes in pH within about one pH unit of the pK of the buffering species.

71. THE END Note & Lao Tzu’s quote on water  “Be careful what you water your dreams with. Water them with worry and fear and you will produce weeds that choke the life from your dream. Water them with optimism and solutions and you will cultivate success. Always be on the lookout for ways to turn a problem into an opportunity for success. Always be on the lookout for ways to nurture your dream.”