WATER By KHAIRUL FARIHAN KASIM

Chapter 3: Outline Molecular Nature of Water Noncovalent Bonding Ionic interactions Hydrogen Bonds van der Waals Forces Thermal Properties of Water Solvent Properties of Water Hydrophilic, hydrophobic, and amphipathic molecules Osmotic pressure Ionization of Water Acids, bases, and pH Buffers Physiological buffers

Water Solvent for all chemical reactions. Transports chemicals from place to place. Helps to maintain constant body temperature. Part of digestive fluids. Dissolves excretion products.

3.1 Molecular Structure of Water The oxygen in water is sp3 hybridized. Hydrogens are bonded to two of the orbitals. Consequently the water molecule is bent. The H-O-H angle is 104.5o.

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.

Water molecule with bond ( ) and net ( ) dipoles. d+ d-

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.

3.2 Noncovalent Bonding Ionic interactions Hydrogen bonding Van der Waals forces Dipole-dipole Dipole-induced dipole Induced dipole-induced dipole

Typical “Bond” Strengths Type kJ/mol Covalent >210 Noncovalent Ionic interactions 4-80 Hydrogen bonds 12-30 van der Waals 0.3-9 Hydrophobic interactions 3-12

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

Hydrogen Bonding In water molecules each of molecules can form hydrogen bonds with four other water molecules. two through the hydrogens and two through the nonbonding electron pairs.

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

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.

3.3 Thermal Properties Hydrogen bonding keeps water in the liquid phase between 0oC and 100oC. Liquid water has a high: Heat of vaporization - energy to vaporize one mole of liquid at 1 atm Heat capacity - energy to change the temperataure by 1oC Water plays an important role in thermal regulation in living organisms.

3.4 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: ion-dipole dipole-dipole hydrogen bonding

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.

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

Hydrogen Bonding hydrogen bond shown in yellow 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

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.: CH3CH2CH2CH2CH2CH3, hexane

Nonpolar Molecules-2 Water forms hydrogen-bonded cagelike structures around hydrophobic molecules, forcing them out of solution.

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.

A Micelle

Osmotic Pressure-2 Osmosis is a spontaneous process in which solvent molecules pass through a semipermeable membrane from a solution of lower solute concentration to a solution of higher solute concentration. Osmotic pressure is the pressure required to stop osmosis.

Osmotic Pressure-3 Osmotic pressure (p) is measured in an osmometer.

Osmotic Pressure-4 = iMRT i = van’t Hoff factor (% as ions) M = molarity (mol/L for dilute solns) R = 0.082 L atm/ mol K T = Kelvin temperature The concentration of a solution can be expressed in terms of osmolarity. The unit is osmol/ liter Osmolarity = iM where i (the vant’t hoff factor) represents the degree of ionization of the solute species. The degree of ionization of a 1 M NaCl solution is 90% with 10% of the NACl existing as ion pairs. i = [Na+] + [Cl-] + [NaCl]unionized i = 0.9 + 0.9 + 0.1 = 1.9

Class discussion Problem 1 When 0.1 g of urea (M.W = 60) is diluted to 100 mL, what is the osmotic pressure of the solution? Assume that the temperature is 25°C

Problem 2 Estimate the osmotic pressure of a solution of 0.1 M NaCl at 25°C. Assume 100% ionization of solute?

Osmotic Pressure-5 Because cells have a higher ion concentration than the surrounding fluids, they tend to pick up water through the semipermeable cell membrane. The cell is said to be hypertonic relative to the surrounding fluid and will burst (hemolyze) if osomotic control is not effected.

Osmotic Pressure-6 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.

Kw = Ka [H2O]2 = [H3O+ ][OH-] 3.5 Ionization of Water Water dissociates. (self-ionizes) H2O + H2O = H3O+ + OH- Kw = Ka [H2O]2 = [H3O+ ][OH-]

Kw= [H3O+][OH-]=1 x 10-14 [H3O+ ] = [OH-] = 1 x 10-7 M Water Ionization-2 The conditions for the water dissociation equilibrium must hold under all situations at 25o. Kw= [H3O+][OH-]=1 x 10-14 In neutral water, [H3O+ ] = [OH-] = 1 x 10-7 M

Water: A/B Properties When external acids or bases are added to water, the ion product ([H3O+ ][OH-] ) must equal Kw. The effect of added acids or bases is best understood using the Lowry-Bronsted theory of acids and bases.

Water: A/B Properties-2 Lowry-Bronsted acid = proton donor HA + H2O = H3O+ + A- A B CA CB C: conjugate (product) A/B

Water: A/B Properties-3 Lowry-Bronsted base = proton acceptor RNH2 + H2O = OH- + RNH3+ B A CB CA

Measuring Acidity Added acids increase the concentration of hydronium ion and bases the concentration of hydroxide ion. In acid solutions [H3O+] > 1 x 10-7 M [OH-] < 1 x 10-7 M In basic solutions [OH-] > 1 x 10-7 M [H3O+] < 1 x 10-7 M pH scale measures acidity without using exponential numbers.

Define: pH = - log(10)[H3O+] 0---------------7---------------14 pH Scale Define: pH = - log(10)[H3O+] 0---------------7---------------14 acidic basic [H3O+]=1 x 10-7 M, pH = ? 7.0

pH Scale - 2 [H3O+]=1 x 10-5 M, pH = ? 5 (acidic) [H3O+]=1 x 10-10 M, pH = ? 10 (basic) What if preexponential number is not 1?

pH Scale - 3 [H3O+]=2.6 x 10-5 M, pH = ? 4.59 (acidic) 8.20 (basic) [H3O+]=7.8 x 10-3 M, pH = ? 2.11 (acidic)

inverse log of negative pH pH Scale-4 pH to [H3O+]? inverse log of negative pH orange juice, pH 3.5. [H3O+]=? [H3O+] = 3.2 x 10-4 M urine, pH 6.2. [H3O+]=? [H3O+] = 6.3 x 10-7 M

Strong acids (and bases) ionize close to 100%. 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%. eg. HCl, HBr, HNO3, H2SO4

Weak acids (or bases) ionize typically in the 1-5% range . Strength of Acids-2 Weak acids (or bases) ionize typically in the 1-5% range . eg. CH3COCOOH, pyruvic acid CH3CHOHCOOH, lactic acid CH3COOH, acetic acid

HA + H2O = H3O+ + A- Strength of Acids-3 Strength of an acid is also measured by its Ka or pKa values. HA + H2O = H3O+ + A- Larger Ka and smaller pKa values indicate stronger acids.

Strength of Acids-4 CH3COCOOH 3.2x10-3 2.5 CH3CHOHCOOH 1.4x10-4 3.9 Ka pKa CH3COCOOH 3.2x10-3 2.5 CH3CHOHCOOH 1.4x10-4 3.9 CH3COOH 1.8x10-5 4.8 Larger Ka and smaller pKa values indicate stronger acids.

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

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.

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.

Titration Curve (HOAc with NaOH) pH moles OH- per mole acid

Titration Curve (HOAc with NaOH) - 2 At the equivalence point, only the salt (NaOAc) is present in solution. At the inflection point, equal moles of salt and acid are present in solution. [HOAc] = [NaOAc] pH = pKa

Class discussion

Titration Curve (HOAc with NaOH) - 3 The pKa for acetic acid is 4.76. 1. Calculate the relative percents of acetic acid and acetate ion when the acid is titrated with 0.7 equivalents of NaOH. 2. Use the Henderson-Hasselbalch equation to calculate the pH at this point.

Titration Curve (HOAc with NaOH) - 4 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. pKa of HOAc is 4.76 30% acid and 70% salt. pH=5.13

Assignment

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

Buffer Solutions-2 Maximum buffer effect occurs at the pKa for an acid. Effective buffer range is +/- 1 of the pK value for the acid or base. eg. H2PO4-/HPO42-, Ka=7.20 buffer range 6.20-8.20 pH

Buffer Solutions-3 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.

Buffer Solutions-4 Within cells the primary buffer is the phosphate buffer: H2PO4-/HPO42- The primary blood buffer is the bicarbonate system: HCO3-/H2CO3. Proteins also provide buffer capacity. Side chains can accept or donate protons.

Buffer Solutions-5 A zwitterion 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.

Buffer Solutions-6 N-tris(hydroxymethyl)methyl-2-aminoethane sulfonate (TES) is a zwitterion buffer example. (HOCH2)3CN+H2CH2CH2SO3-

Buffer Solutions-7 Buffers work by chemically tying up acid and base. Eg.:

Class discussion

Buffer Solutions-8 Calculate the ratio of lactic acid to lactate in a buffer at pH 5.00. The pKa for lactic acid is 3.86 = 13.8

Water : The Medium of Life The End