2.
Objectives <ul><li>Structure and polarity of the water molecule </li></ul><ul><li>Solubility of compounds in water </li></ul><ul><li>Properties of acids, bases and buffers </li></ul><ul><li>Describe how buffers maintain pH stability </li></ul>
3.
Water: Structure and Properties <ul><li>Polarity of water: Important for its shape and binding to other water molecules – An asymmetrical distribution of charge. </li></ul><ul><li>Cohesiveness of water: liquid vs solid </li></ul><ul><li>Melting point, boiling point and critical points </li></ul><ul><li>Density of the water molecule </li></ul><ul><li>Solubility of the water molecule </li></ul>
4.
Water is a dipolar Molecule The approximate shape and charge distribution of water. Martin Chaplin, http://www.lsbu.ac.uk/water/molecule.html The two hydrogen atoms have partial positive ( δ + ) charges and the oxygen molecule has a partial negative (2 δ - ) charge. The result of unequal sharing is two electrical dipoles, one along each H-O bond.
5.
H-Bonding between water molecules <ul><li>In its liquid form, each water molecule is hydrogen-bonded to about 3.4 other neighboring water molecules, while in its solid form, it is bound to 4. </li></ul><ul><li>Liquid water is partially structured, while solid water has a regular crystalline structure. </li></ul>http://courses.cm.utexas.edu/jrobertus/ch339k/overheads-1.htm <ul><li>The presence of partial positive and negative charges on the H and O atoms, respectively, leads to an electrostatic attraction between the O atom of one H 2 O molecule and the H of another: This forms a hydrogen bond between the two. </li></ul>
6.
Water is found in 3 forms on planet Earth <ul><li>Water is found in gas liquid and solid on earth. Ie. Earth’s temperature is at the “triple point for water. </li></ul><ul><li>Most molecules become more dense when in solid form. This is not true for water. This property is VERY important for the development of life on earth. </li></ul>Gas Liquid Solid
8.
Water: Cohesion and Adhesion <ul><li>Water is VERY Cohesive (attracted to other water molecules) and also adhesive (attracted to other molecules (polar molecules)). </li></ul><ul><li>Water has the highest Surface Tension of all molecules, except for mercury. </li></ul>http://www.martin-waugh.com/wp-content/uploads/2007/08/water-drop-1b.jpg
9.
Solubility of Molecules in Water <ul><li>Water’s dipolar properties make it very likely that other dipoles (other molecules that are polar or charged) will dissolve in it. </li></ul><ul><li>It is very unlikely that nonpolar molecules (strings of carbons etc.) will dissolve in it. </li></ul><ul><li>These properties are also VERY important for life on earth: The shape of proteins, DNA, RNA, etc. </li></ul>Hemoglobin
12.
Amphipathic Compounds <ul><li>Compounds that contain both polar (or charged) and nonpolar regions. </li></ul><ul><li>Polar or charged region dissolves in water, but the nonpolar regions avoid contact with water molecules. </li></ul>
13.
<ul><li>The nonpolar regions cluster to allow the smallest nonpolar area possible exposed to the aqueous solvent. </li></ul><ul><li>This forms a stable structure called a micelle. </li></ul><ul><li>These structures form on their own and have the highest entropy possible. </li></ul>
14.
Weak Interactions among Biomolecules in Aqueous Solvent <ul><li>4 types of interactions: Hydrogen bonds, Ionic interactions, Hydrophobic interactions, Van der Waals interactions. </li></ul><ul><li>Hydrogen bonds occur between neutral groups and peptide bonds. </li></ul><ul><li>Ionic interactions can be either attractive (negative to positive) or repulsive(neg to neg, or pos to pos). </li></ul><ul><li>Hydrophobic interactions are responsible for holding the nonpolar regions of molecules together; these interactions result from the system achieving the greatest thermodynamic stability. </li></ul><ul><li>Van der Waals interactions occur whtn two uncharged atoms are brought into very close proximity. The two electron clouds influence each other resulting in a transient attractive force. </li></ul>
15.
Colligative properties of aqueous solutions <ul><li>Solutes affect the colligative properties of aqueous solutions: </li></ul><ul><ul><li>Vapor pressure </li></ul></ul><ul><ul><li>Boiling point </li></ul></ul><ul><ul><li>Melting (freezing) point </li></ul></ul><ul><ul><li>Osmotic pressure. </li></ul></ul><ul><li>Colligative (“tied together”) properties are so called because the effect of solutes on all four properties has the same basis: the concentration of water is lower in solution than in pure water. </li></ul><ul><li>Independent of the physical properties of the solute; dependent only upon the number of solute particles in a given amount of water. Ie. Any solute will change these properties in the same way. </li></ul>
16.
Osmotic Pressure <ul><li>Osmotic pressure : The hydrostatic pressure produced by a solution in a space divided by a semi-permeable membrane due to a differential in the concentrations of the solutes. </li></ul><ul><li>Water molecules tend to move from a region of higher water concentration to one of lower concentration. </li></ul><ul><li>Isotonic solution : solutions of equal osmotic pressures </li></ul><ul><li>Hypotonic/Hypertonic solution : A solution with a comparatively lower/higher osmotic pressure. </li></ul>
17.
Effect of extracellular osmolarity on water movement across a plasma membrane Effect of extracellular osmolarity on water movement across a man-made, selectively permeable membrane
18.
Ionization of water <ul><li>Water molecules have a slight tendency to reversibly ionize to produce hydrogen ions (protons) and hydroxide ions: generating a reversible equilibrium: H 2 O H + +OH - </li></ul><ul><li>The free proton is generated as part of a hydronium ion, and has negligible existence as H+, therefore the equation is better written: 2H 2 O H 3 O+OH - </li></ul><ul><li>Ionization can be measured as electrical conductivity. Pure water can carry electrical current as H+ move towards the cathode and OH- towards the anode. </li></ul><ul><li>Electrophoresis </li></ul>
19.
<ul><li>Ionization of water is expressed by an equilibrium constant Keq . </li></ul><ul><li>At STP only about 1 in every 10 7 molecules of water is ionized at any instant, the Equilibrium constant for the reversible ionization of water is: </li></ul><ul><li>Keq = [H + ][OH - ]/H 2 O = [H + ][OH - ]/55.5M </li></ul><ul><li>(at 25 ° C the concentration of water is 55.5M.) </li></ul><ul><li>(55.5M)(Keq) = [H + ][OH - ] = Kw (Kw designates the ion product of water at 25 ° C.) </li></ul><ul><li>Thus, Kw = [H + ][OH - ] = (55.5M)(1.8x10 -16 M) = </li></ul><ul><li>1.0x10 -14 M 2 </li></ul>Ionization of water
20.
<ul><li>In a neutral solution, there are equal concentrations of both H + and OH - ions. </li></ul><ul><li>Under these conditions Kw = [H + ][OH - ] = [H + ] 2 . </li></ul><ul><li>Thus, [H + ] = sq. root of 1x10 -14 M 2 = 10 -7 M </li></ul>Ionization of water
21.
<ul><li>Kw is the basis for the pH scale. </li></ul><ul><li>pKw = pH + pOH = 14 </li></ul><ul><li>and pH = Log 1/[H + ] = -log[H + ] </li></ul><ul><li>In a precisely neutral solution at 25 C, in which the concentration of H + ions is 1x10 -7 M </li></ul><ul><li>Thus, pH = log 1/1.0 x 10 -7 = log (1.0 x 10 7 ) </li></ul><ul><li>=log 1.0 + log10 7 = 0+7 = 7 </li></ul>
22.
Acids, Bases and Buffers <ul><li>Definitions: </li></ul><ul><li>Acids = proton donors HA H + + A - </li></ul><ul><li>Bases = proton acceptors B - + H + BH </li></ul><ul><li>A proton donor and its corresponding proton acceptor makes up a conjugate acid-base pair. </li></ul><ul><li>Acetic acid (CH 3 COOH), a proton donor and the acetate anion (CH 3 COO - ), the corresponding proton acceptor constitute a conjugate acid-base pair, related by </li></ul><ul><li>CH 3 COOH CH 3 COO - + H + </li></ul>
23.
Strong acids and strong bases <ul><li>Strong acids and bases are completely dissociated in solution. </li></ul><ul><ul><li>Examples (Acids): HCl, H 2 SO 4 </li></ul></ul><ul><ul><li>(Bases): NaOH, KOH </li></ul></ul>
24.
Weak Acids and Bases <ul><li>These compounds are of more importance in biology. Examples: Acetic acid, bicarbonate (HCO 3 - ). </li></ul><ul><li>Dissociation Constant : Each acid has a tendency to lose its proton in aqueous solution. The stronger the acid, the greater the tendency. </li></ul><ul><li>The tendency of any acid (HA) to lose its proton and form its conjugate base (A - ) is defined by the equilibrium constant (Keq) for the reversible reaction </li></ul><ul><li>HA H + + A - </li></ul>
25.
Dissociation Constant <ul><li>Definition: Dissociation constants are equilibrium constants for ionization reactions , often designated Ka. </li></ul><ul><li>Keq = [H + ][A - ]/[HA] = Ka </li></ul><ul><li>pKa is analogous to pH and is defined by the equation: </li></ul><ul><li>pKa = log 1/Ka = -log Ka </li></ul><ul><li>The stronger the tendency to dissociate a proton, the stronger is the acid and the lower its pKa. </li></ul>
26.
pH, Acids and Buffers <ul><li>Each acid has a tendency to lose its proton in aqueous solution. The stronger the acid, the greater the tendency. </li></ul><ul><li>Buffers are mixtures of weak acids and their conjugate bases. </li></ul><ul><li>The titration curve of a weak acid and its conjugate base shows that these compounds can be used to buffer solutions. </li></ul><ul><li>Buffering is the consequence of two reversible reactions taking place simultaneously and reaching their equilibrium points as governed by their equilibrium constants Kw and Ka. </li></ul>
27.
Titration Curve for Acetic Acid and its Conjugate Base NaOH is added incrementally to HAc, and the pH of the mixture is measured. This value is plotted against the amount of NaOH expressed as the fraction of total NaOH required to convert all the HAc to its protonated form. At the midpoint of titration, the pH is equal to the pKa. The shaded zone is the region useful for buffering power, generally between 10% and 90% titration of a weak acid.
28.
<ul><li>Titration of 0.1M HAc with 0.1 M NaOH at 25 C. two reversible equilibriums are involved: </li></ul><ul><ul><ul><ul><li>H2O H + + OH - Equation 1 </li></ul></ul></ul></ul><ul><ul><ul><ul><li>HAc H + + Ac - Equation 2 </li></ul></ul></ul></ul><ul><li>The equilibria must simultaneously conform to their equilibrium constants. </li></ul><ul><ul><ul><ul><li>Kw = [H + ][OH - ] = 1x10 -14 M 2 </li></ul></ul></ul></ul><ul><ul><ul><ul><li>Ka = [H + ][Ac - ]/[HAc] = 1x10 -5 M </li></ul></ul></ul></ul><ul><li>As NaOH is added, OH - combines with free H + to form H 2 O according to equation 1. As H + is removed, HAc ionizes more to satisfy its equilibrium producing Ac - as NaOH is added. As the titration continues, remaining undissociated HAc is gradually converted to Ac - . </li></ul>
29.
The Hendersson-Hasselbach Equation <ul><li>The Hendersson-Hasselbach equation relates pH, pK and buffer concentration: </li></ul><ul><ul><ul><ul><ul><li>pH = pKa + log [A-]/[HA] </li></ul></ul></ul></ul></ul><ul><li>The derivation: </li></ul><ul><ul><ul><ul><ul><li>Ka = [H+][A-]/[HA]; solving for [H+], [H+] = Ka[HA]/[A-] </li></ul></ul></ul></ul></ul><ul><ul><ul><ul><ul><li>Taking the negative log: -log[H+] = -log Ka – log[HA]/[A-] </li></ul></ul></ul></ul></ul><ul><ul><ul><ul><ul><li>Substituting pH for –log[H+] and pKa for –log[Ka], </li></ul></ul></ul></ul></ul><ul><ul><ul><ul><ul><li>pH = pKa – log[HA]/[A-] </li></ul></ul></ul></ul></ul><ul><li>This converts to the Hendersson-Hasselbach Equation: pH = pKa + log [A - ]/[HA] </li></ul><ul><li>This equation is useful in calculating pKa, pH and the molar ratio of proton donor and acceptor. </li></ul>
30.
Example Problems <ul><li>Calculate the pKa of lactic acid, given that the concentration of lactic acid is 0.010 M and the concentration of lactate is 0.087 M, the pH is 4.80. </li></ul><ul><ul><ul><ul><ul><li>pH = pKa + log [lactate]/[lactic acid] </li></ul></ul></ul></ul></ul><ul><ul><ul><ul><ul><li>pKa = pH – log [lactate]/[lactic acid] = 4.8 – log 0.087/0.010 = 4.80 – log 8.7 = 4.80 – 0.94 = 3.86 </li></ul></ul></ul></ul></ul><ul><li>Calculate thepH of a mixture of 0.1M acetic acid and 0.2 M sodium acetate. The pKa is 4.76. </li></ul><ul><ul><ul><ul><ul><li>pH = pKa + log [acetate]/[acetic acid] = 4.76+log 0.2/0.1 </li></ul></ul></ul></ul></ul><ul><ul><ul><ul><ul><li>pH = 4.76+0.301 = 5.06 </li></ul></ul></ul></ul></ul><ul><li>Caculate the ratio of the concentrations of acetate and acetic acid required in a buffer system of pH 5.30. </li></ul><ul><ul><ul><ul><ul><li>Log[acetate]/[acetic acid] = pH – pKa = 5.30 – 4.76 = 0.54; </li></ul></ul></ul></ul></ul><ul><ul><ul><ul><ul><li>Antilog 0.54 = 3.47 </li></ul></ul></ul></ul></ul>
31.
Weak acids or bases buffer cells and tissues against pH changes <ul><li>The intracellular and extracellular fluids of multicellular organisms have a characteristic and nearly constant pH, maintained by using physiological buffer systems . </li></ul><ul><ul><li>Extracellular: </li></ul></ul><ul><ul><li>Buffering capacity of plasma proteins, especially hemoglobin. Basic and acidic amino acid side chains act as weak acids and bases – a minor, but significant system. </li></ul></ul><ul><ul><li>The Bicarbonate Buffer System: Blood plasma is buffered in part by the bicarbonate system, consisting of carbonic acid (H 2 CO 3 ) as proton donor, and bicarbonate (HCO 3 - ) as proton acceptor…. </li></ul></ul><ul><ul><ul><li>This is the most important extracellular buffering system! </li></ul></ul></ul><ul><ul><ul><li>CO 2 + H 2 O [H 2 CO 3 ] H + + HCO 3 - </li></ul></ul></ul>Titration curve for the carbonic-acid-bicarbonate buffer. Note that the pH of the blood (7.4) lies outside the region of greatest buffering capacity (green).
32.
<ul><ul><li>Intracellular: The cytoplasm of most cells contains a high concentration of proteins, which contain amino acids with functional groups that are weak acids or bases. </li></ul></ul><ul><ul><ul><li>The phosphate buffer system: H 2 PO 4- as proton donor, and HPO 4 2- as proton acceptor. This system uses the second dissociation of phosphoric acid (PK 2 = 6.86). </li></ul></ul></ul><ul><ul><ul><ul><ul><li>H 2 PO 4- H+ + HPO 4 2- </li></ul></ul></ul></ul></ul><ul><ul><ul><li>This buffer system is maximally effective at pH close to its pKa of 6.86, between pH of 5.9 and 7.9 </li></ul></ul></ul>