Jaymee B. Quindara RN

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Water

Source: Lehninger+ Harper

Overview

Most abundant subs 70% of the wt of most orgs Ionization products H+ & OH- greatly influence the structure, self assembly & properties of all cellular components & macromolecules in the cell. Noncovalent interactions responsible for the strength and specificity of recognition among biomolecules are decisively influenced by the solvent properties of H2O (hydrogen bonding)

Unique set of properties- H2Os dipolar structure & exceptional capacity for forming H-bonds H2Os intrxn w/ solvated biomolecule affects the latters structure Excellent nucleophile- provides pair of e- to form covalent bonds (Lewis base) Reactant or product in many metab rxn Buffer system (e.g. HCO3) keeps ECF btwn 7.35-7.45 Arterial blood- pH Venous blood- CO2 Hypothalamus- reg of water balance (ADH)

Weak Interactions in aq System

Hydrogen bonds - unique properties of H2O. Dissolved polar biomolecules (water-solute) more favourable intrxn than water-water intrxn Nonpolar biomolecules interfere water-water interactions but unable to form water-solute interactions. Weak intrxns have significant influence on the 3D structures of macromolecules.

Hydrogen Bonding Gives Water its Unusual Properties

Properties: High melting point High boiling point High heat of vaporization Due to attraction btwn H2O molecules Each H atom shares an e- pair w/ O at the center. Geometry dictated by the shape of the outer e- orbitals of the O atom w/c is similar to the sp3 bonding orbitals of C. Water geometry- roughly tetrahedron / slightly skewed tetrahedron w/ angle 104.5oC bc of crowding by the nonbonding e- H2O dipole- charge distributed asymmetrically about its structure O more electronegative attracting e- from H resulting to a dipole along each side of the H-O bonds H partial + (+) O partial (2-) O attracted to H atom of another molecule Hydrogen bond Hydrogen bond- relatively weak w/ BDE of 23kJ/mol // Covalent O-H bond 470 k J/mol. // Covalent C-C bond 348 k J/mol Bond dissociation energy (BDE)- energy req to break a bond H bond- 10% covalent & 90% electrostatic At room temp, the thermal energy of an aq sol same magnitude req to break H bond. H bond lifetime 1-20 picoseconds H-O breaking then reformation w/ in 0.1 ps. Flickering clusters

Harper: H-bond half life of 1 s 4.5 kcal/mol req to rupture H-bond in liquid H2O 5% less to break OH bond

H bond form bridges btwn solutes that allow larger molecules to interact w/ each other over several nanometers w/o physical contact. H2O nearly tetrahedral arrangement allows each H2O molecule to form H bond w/ 4 neighbouring H2O molecule. Liquid H2O- 3.5 bonds Ice- 4 bonds ice crystal lattice structure make H2O less dense Breaking of H bond of the crystal lattice of ice requires much energy high melting point of H2O Solid- Liquid H= +5.9 k J/mol Liquid-Gas H= + 44.O k J/mol During melting/ evaporation the entropy of the aq system . At room temp, melting and evaporation occur spontaneously since the tendency of H2O molecules to associate through H bonds is outweighed by the energetic push towards entropy G must be negative for a rxn to be spontaneous Since H (enthalpy) is positive for melting and evaporation, it is the in entropy (S) that makes G negative and drives these changes.

Water forms H Bonds w/ Polar Solutes

Readily form btwn electronegative atom (the H acceptor, usually O or N) and an H atom covalently bonded to another electronegative atom (the H donor) in the same or another molecule. H atoms covalently bonded to C atoms do not participate in H bonding bc H-C bonding is weakly polar. Uncharged but polar biomolecules such as sugar dissolve in H2O bc of the stabilizing effect of H bond btwn OH- grp or COH grp of the sugar & the Polar H2O molecules. R-OH, R-COH/ RCOR, & compounds w/ N-H bond all forms H bond and tend to be soluble in H2O H bond are the strongest when oriented to maximized electrostatic interaction H & 2 atoms that share it are in straight line OR acceptor is in line w/ the covalent bond btwn the donor atom & H (+) charge of H directly btwn the two (-) H bonds are highly directional & capable of holding 2 H bonded molecules or grps in a specific geometric arrangement

Harper: R-COH & R-C=OR and R-C=ONH2: O participates in H-bonding by accepting e- ROH & RNH3 : unshielded H serve as H acceptors or donors

Water Interacts Electrostatically with Charged Solutes Water is Polar Solvent Readily dissolved charged/ polar biomolecules hydrophilic. Dissolves many crystalline salts by hydrating their component ions, neutralizing ionic charges & weakens electrostatic charges necessary for lattice formation Dissolve by replacing solute-solute bonds to solute-H2O bond among R-COO-, R-NH3+ and PO4 esters or anhydrides H2O effective in screening the electrostatic interactions btwn dissolved ions bc it has a high dielectric constant (78.5) Dielectric constant- physical property that reflects the number of dipoles in solvent.

F- force/ strength of ionic intrxn depends on charge Q, distance r btwn Q, the dielectric constant F inversely proportional to H2O force of attraction btwn charged & polar species

Entropy Increases as Crystalline Substances Dissolve Substances dissolved leaves the crystal lattice acquire greater freedom of motion in Entropy (S) An in S makes dissolving salts easy. Formation of solution occurs with a favourable free energy change where H has a small positive value, T and S has a large positive value G is negative.

Nonpolar Gas are Poorly Soluble in Water Biologically important gasses CO2, O2 & N2 are nonpolar (equal sharing of e-) Movement of molecules of gas to aq sol constrains their motion and the motion of water S. Some orgs have H2O-soluble carrier CHONs (hemoglobin & myoglobin) that facilitates O2 transport. CO2 in aq sol becomes H2CO3 (carbonic acid) and transported as HCO3- (bicarbonate) ion free or bounded to haemoglobin. In comparison NH3, NO, and H2S are biologically impt gasses that are polar.

Nonpolar Compounds Force Energetically Unfavourable Changes in the Structure of Water Nonpolar compounds are unable to undergo energetically favourable interactions w/ H2O molecules & they interfere with H-bonding among H2O. All molecules in a sol interfere w/ H-bonding of some H2O molecules in their stat envt. However, polar/ charged solutes compensate for the lost of H2O-H2O bonding by forming new Solute-H2O bonding. Small gain in H Nonpolar- no compensation but gain in H. Dissolving nonpolar subs in H2O requires energy from surroundings and produces measurable decrease in S. H2O molecules near nonpolar subs form cage-like shell around each solute molecule. Clathrates- crystalline compound of nonpolar solutes and H2O. Ordering of H2O molecules reduces S. The magnitude of the S is proportional to the surface area of the nonpolar solutes enclosed w/ in H2O molecules H has a positive value and S is negative thus making G positive (Endergonic) Amphipathic compounds- contain regions that are polar and nonpolar. Micelles- stable Amphipathic compounds in H2O which hydrophobic grps are sequestered from H2O (hydrophobic interaction); ordered shell of H2O molecule is & S is . Strength of hydrophobic interaction results from the systems achieving greatest thermodynamic stability by minimizing the number of ordered H2O molecules req to surround hydrophobic portion of solutes. H-bonding btwn H2O and polar solutes cause ordering of H2O molecules but energetic effect is less significant. Force that bind polar substrate to complementary polar surface of enzymes is the in S as the enzyme displaces ordered H2O from the substrate

Harper: CHONs- fold hydrophibic side chains in the interior while charged or polar Aa side chains generally are present on the surface Same goes w/ phospholipid bilayer- phosphatidyl serine or phsophatidyl ethanolamine Maximize formation of energetically favorble charg-dipole, dipole-dipole & H-bonding intrxn btwn polar grps of biomolecules & H2O Minimize unfavorabvle contact btwn H2O & hydrophobic grps.

Van der Waals Interactions are Weak Interatomic Attractions Two uncharged atoms becomes close together their e- cloud influence each other. Random variation in the positions of the e- around one nucleus may create transient electric dipole that induces opposite electric dipole to nearby atom thus attracting each other. Van der Waals contract- point where net attraction is at max.

Weak Interactions are Crucial to Macromolecular Structure and Fxn Noncovalent interactions are much weaker that covalent bonds. Hydrophobic interactions- association of non polar compounds are much weaker than covalent bonds. Ionic and hydrogen bonds variable in strength depending on the polarity of the solvent and alignment of H-bonded atom but are significantly weaker than covalent bond. Cumulative effects of many intrxns of weak bonds can be very significant. Noncovalent binding of enzyme to its substrate. Contributes to net decrease in the free energy of the system Stability- measured by the equilibrium constant. Stability of a noncovalent interaction can be calculated from the binding energy. The dissociation of two biomolecules that are assoc noncovalently thru multiple weak interactions req all the interaction to be disrupted at the same time. For macromolecules the most stable structure is in w/c weak interactions are maximized. Determine 3D folding of polypeptides Multiple weak interaction Binding of antigen to a specific antibody Enzyme-substrate interaction Hormone-neurotransmitter binding

The complementarity between interacting biomolecules reflects the complementarity and weak interactions between polar, charged, and hydrophobic groups on the surfaces of the molecules.

Harper:

Hydrophobic Intrxns Self association of nonpolar compounds in aq sol to minimize energetically unfavourable intrxns btwn nonpolar grps & H2O Nonpolar grps affect structure of H2O that surrounds them by restricting number of orientation of H2O to permit them to participate in energetically favourable H-bonding Max formation of H-bonds maintained only by order of adjacent H2O molecules S Follows 2nd law of thermodynamics Optimal free energy of C-H H2O mixture is a fxn of maximal H and minimum S. Form micelles & the Phospholipid bilayer

Electrostatic Intrxns Opposite charge btwn biomolecules salt bridges Acts on large distances Facilitates binding of charged molecules & ions to CHONs & NA

Van der Waals Forces Attraction btwn transient dipoles generated by rapid movement of e- on all neutral atoms Act very short distance 2-4

Water as a Nucleophile Nucleophile e- rich molecule Electrophile e- poor molecule Metab rxn transfer of e- from nucleophile to electrophile Lone pair of sp3 e- bear - charge of H2O excellent nucleophile Hydrolysis- results cleavage of amide, glycosidic or ester bonds that hold biopolymers together. Condensation- joining of monomers to form biopolymers produces water (product) Hydrolysis is thermodynamically favoured but amide & phosphodiester bond of polypeptides & oligonucleotides are stable in aq sol. Thermodynamics governing eq of a rxn do not determine rate at w/c it will take place. Enzymes- CHON catalysts used to rate of hydrolytic rxn needed Protease- CHON- Aa Nucleases- phosphodiester bonds in DNA & RNAMany Metabolic Rxn Involve Grp Transfer DG + A = AG + D G frm donor D to acceptor A forming AG Hydrolysis & phosphorolysis of glycogen grp transfer rxn glycosyl grps are transferred to H2O or to orthophosphate. Keq for the hydrolysis of covalent bonds strongly favours the formation of split products. Biosynthesis of macromolecules thermodynamically unfavoured synthesis of covalent bonds- coupled to favoured rxns to favour biopolymer synthesis. Help frm enzyme Newly synthesized polymers do not statly hydrolyzed bc active sites of biosynthetic enzymes sequester substrates in an envt from w/c H2O is excluded.

Solutes Affect the Colligative Properties of Aq Sol Solutes alter certain physical properties of their solvent Vapour pressure Boiling point Melting point Freezing point Osmotic pressure. Concentration of H2O in sol is in sol than in pure H2O Effect of solute concentration on the Colligative properties of H2O is independent on the chem prop of solutes H2O molecules diffusing from the region of higher H2O concentration to the region of lower water concentration produce osmotic pressure Osmotic pressure- force that resist water movement. vant Hoff equation: = icRT

R- gas constant 8.315 J/mol T absolute temp ic- osmolarity of the sol, vant Hoff factor i, & solutes molar concentration c vant Hoff factor- measure of the extent to w/c solute dissociates into two or more ionic species.

Osmosis- impt factor in the life of most cells. Permeability of plasma membrane due to CHON channels, aquaporins. Isotonic- both side has same osmolarity Hypotonic- has a lower osmolarity Hypertonic- ECF has higher osmolarity than the cytosol Cells generally contain higher concentration of biomolecules and ions than their surroundings. Counterbalance keep cell from swelling & bursting. Blood plasma and interstitial fld- maintained at an osmolarity close to that of the cytosol. High concentration of albumin contributes to plasma osmolarity. Cell also actively pumps Na+ and other ions into the interstitial fld to stay in osmotic balance w/ ECF. Solutes exert an effect on osmolarity depending on the number of dissolved particles not on their mass. Macromolecules have far less effect on the osmolarity of a sol than would an equal mass of their monomeric components Storing fuels as polysaccharides rather than glucose avoids in osmotic pressure in the storage cell.

Ionization of Water: Weak Acids and Weak Bases

Ionization of H2O can be described by equilibrium constant. Weak acid dissolved in water becomes ionized by donating H+ Weak bases consume H+ and becomes protonated Total hydrogen concentration is expressed as the pH of the solution Water can act as both acid & base

Pure water is slightly ionized H2O molecules have slight tendency to undergo reversible ionization to yield H+ and OH-H2O + H2O H3O+ + OH- H+ do not exist in sol but they are hydrated statly to H3O+ (hydronium ion) Transferred p+ is associated w/ cluster of H2O molecules p+ also exist as multimers H5O2+ and H7O3+ H-bonding btwn H2O molecules makes the hydration of dissociating p+ virtually instantaneous 1g of H2O contains 3.46 x 1022 molecules Probability that H exists as an ion (H+) is 0.01 1 in 100 chance it is an ion (H+) 99 in 100 chances is part of a H2O molecule Actual probability of H atom in pure H2O exist as H+ is 1.8 x 10-9 Probability of it being part of a molecule thus is almost unity. Can also be stated: for every H+ and OH- in pure H2O there are 1.8 Billion or 1.8x109 water molecules Pure H2O carries electric current. H3O+ cathode OH- anode High ionic mobility of H3O+ and OH- results from proton hopping faster than diffusion

High ionic mobility of H+ makes acid-base rxns in aq sol very fast. Reversible ionization plays an impt role in cellular fxn Equilibrium constant Keq

A+B C+D

Equilibrium constant may be approximated by measuring concentrations at equilibrium

The ionization of water is expressed by an Equilibrium Constant Degree of ionization of H2O at eq is small in which at 25oC only about 2 of every 109 molecules in pure water are ionized at any instant. The Keq for the reversible ionization of H2O:

Concentration of H2O at 25oC is 55.5 M grams of H2O in 1L, Concentration of H+ and OH- = 1 x 10-7 M Keq of H2O = 1.8 x 10-16 M at 25oC

Keq (55.5 M) = [H+] [OH-] = Kw

Kw = [H+] [OH-]= (55.5 M) (1.8 x 10-16 M)= 1.0 x 10-14 M2 (ion product of H2O)

When there is exactly equal concentration of H+ and OH- the sol is at neutral pH. Ion product of H2O is constant. H+ or OH- will or

Harper: 1 mole of H2O weighs 18g

H: 1g/mole x 2 = 2O: 16 g/mole x1=118g/mole 1 liter (1000g) of H2O contains 55.56 mole

Probability H in pure H+ or as OH- is 1.8x10-9 [H+] & [OH+] = 1.8x10-9 x 55.56 mole = 1x10-7

K= 1.8x10-16 mol/L

K- dissociation constant [H2O]= 55.56 mol/L is constant Used to get ion product Kw constantKw=(K)[H2O] = [H+][OH-]= (1.8x10-16 mol/L)[55.56 mol/L]= 1.0 x 10-14 mol2/L2

The pH scale designates the H+ and OH- Concentrations Kw = 1x10-14 (mol/L)2 for all aq sol even sol of acids & bases basis for the pH scale

pH= -log [H+]

H2O pH is 7.0 at 25oC Neutral Low pH = high [H+] High pH = high [OH-]

Kw= [H+][OH-]= 10-14= log[H+]+ log[OH-]= log 10-14= pH+ pOH = 14

Weak Acids and Bases Have Characteristic Acid Dissociation Constants Strong acids and bases completely ionized in dilute aq sol. Weak acid and bases- ubiquitous in biological systems and play impt roles in metab & its regulation

Acids Proton donors Stronger the acid > tendency to lose p+ Protonated specie

Bases Proton acceptors Unprotonated specie

Conjugate acid base pair Proton donor and corresponding proton acceptor E.g.: Acetate CH3COOH (PD) & Acetate anion CH3COO- (PA)

Tendency of any acid (HA) to lose a p+ and form its conjugate base (A-) is defined by the Keq for the reversible reaction

HA H+ + A-

Ionization constants/ Acid dissociation constants (Ka) Express relative strength of weak acids & bases Equilibrium constant for ionization rxn Strong acids have larger Ka the lower the pKa value and weaker acids has smaller Ka the larger the pKa

pKa= -log KapH= pKa

Ka is negative exponential number pKa is used pKa is r/t Ka as pH is to [H+]

The stronger the tendency to dissociate a p+ the stronger the acid the lower is the pKa acid Weak basestrong conjugate acid Weak acid- strong conjugate base

Polyprotic compounds containing more than 1 dissociable p+ a numerical subscript is assigned to each in order to relative acidity.

RNH3+ RNH2

pKa is the pH at w/c [RNH3+] acid = [RNH2]

[RNH3+] = [RNH2]Ka= [H+] When associated/ protonated and dissociated/ conjugate specie are present at equal concentration the prevailing [H+] = to Ka If log of both side are multiplied by -log :Ka= [H+]

-log Ka= -log [H+]

pKa= -log Ka || pH= -log [H+]

pKa=pH

pKa of an acid grp is the pH at w/c the protonated and unprotonated specie are present at equal concentrations pKa of an acid may be determined by adding 0.5 equivalent of alkali per equivalent of acidthe pH will be the pKa of the acid

Titration cures reveal the pKa of Weak Acids Titration determine amt of acid in a sol w/ strong base (usually NaOH) of known concentration. Base- added in small increments until acid is neutralized Indicator dye pH meter [acid] in the original sol can be calculated from the vol & [NaOH] added Titration curve- plot of pH against amt of base added reveals pKa of a weak acid.

Two reversible equilibrium are involved in the processH2O H+ + OH-

HAc H+ + Ac- Equilibria must simultaneously conform to their characteristic equilibrium constants H2O H+ + OH- || Kw= [H+] [OH-] = 1x10 -14 M2

HAc H+ + Ac- ||

At the beginning of the titration before the base (NaOH) is added, the acid (Acetic acid) is already ionized to an extent that can be calculated from its ionization constant. As base (NaOH) is introduced, its OH- combines w/ the free H+ of the sol H2O to satisfy equilibrium. As free H+ is removed HAc dissociates further to satisfy own equilibrium constant. As titration proceeds more HAc ionizes forming Ac- as the bases (NaOH) is added Midpoint of titration [HAc] proton donor = [Ac-] proton acceptor Impt rel: pH of the equimolar sol of the acid (acetic acid) & conjugate base (acetate) is equal to pKa of the acid As the titration continues by adding further increments of the base (NaOH) nondissociated acid (acetic acid) gradually converted to its conjugate base (acetate). End point- acid lost all its H+ to OH- to form H2O + Conjugate base Throughout the titration the two Equilibria coexist always conforming to its equilibrium constants

Buffering Against pH Changes in Biological Systems

Almost every biological process is pH dependent Ionic intrxns - stabilize CHON molecule allow an enzyme to recognize & bind substrate Cells orgs maintain specific constant cytosolic pH 7 Keep biomolecules in their optimal ionic state Multicell orgs pH of ECF is tightly regulated

Buffers Are mixture of Weak Acids Their Conjugate Bases Buffers- aq system tend to resist changes in pH when small amt of acid/ base are added Consist of weak acid (p+ donor) its conjugate base (p+ acceptor) Buffering region Flat zone in the titration curve Amt of H+ or OH- added has less effect on pH than the same amt added outside the region Midpoint has maximal buffering power in w/c its pH changes least on addition of H+ or OH- pH at midpoint = pKa Buffering result from 2 reversible Equilibria occurring in a sol of nearly equal concentration of p+ donor & its conjugate p+ acceptor Whenever H+ or OH- is added to a buffer small change in ration of the relative concentration of the weak acid its anion small change in pH Concentration of 1 component balanced by in the other. Sum of the buffer components does not change only the ratio.

Henderson-Hasselbalch Eq Relates pH, pKa and Buffer Concentration Titration of Acetic acid H2PO4- & NH4+ have nearly identical shapes Curve reflect funda law or relationship Shape of titration curve- describe by Henderson-Hasselbalch equation. Impt in understanding buffer action & A-B balance in the blood & tissues Useful way of restating the expression for the ionization constant of an acid

1st Solve for

2nd Take log of both sides

3rd Substitute pH for log and pKa for

4th invert w/c involves changing sign to obtain Henderson-Hasselbalch equation

Shows why pKa of a weak acid is = pH of the sol @ midpoint of titration

[HA] = [A-]

pH= pKa + log 1

pH= pKa + 0

pH= pKa

Allows us to calculate pKa given pH & molar ratio of p+ donor & acceptor. Allows us to calculate pH given pKa & molar ratio of p+ donor & acceptor. Calculate the molar ratio of p+ donor & acceptor given pH & pKa

Harper Has great predictive value in protonic Equilibria When an acid is exactly half neutralized, [A-] = [HA] under these condition = Therefore at half neutralization, pH= pKa When the ratio [A-]/[HA]= 100:1pH= pKa + log + pH= pKa + log 100/1 = pKa + (-1)

Acid strength depends on molecular structure Many acids of biologic interest posses >1 dissociating grp Presence of nearby (-) charge hinders release of p+ from nearby grp raising its pKa Phosphoric & Citric Acid

Effect of adjacent charge w/ distance 2nd pKa for Succinic acid w/c has two Meth grps btwn its RCOOH is 5.6 2nd pKa for Glutaric acid w/c has additional Meth grp is 5.4

pKa Value Depends on Properties of the Medium pKa fxnal grp influenced by surrounding medium Medium may or the pKa depending whether the undissociated acid or its conjugate base is the charged specie

Weak Acids or Bases Buffer Cells & Tissues against pH Changes ICF & ECF have a characteristic & nearly constant pH. 1st line of defence against changes in pH buffer systems Cytoplasm- hi concentration of many Aa w/ fxnal grps that are weak acids or weak bases Histidine pKa = 6.0 side chains exist as protonated or unprotonated near neutral pH Nucleotides such as ATP & low molecular wt metabolites contain ionisable grps that can contribute to buffering power of cytoplasm Highly specialized organelles & extracellular compartments have hi concentration of compounds that contribute buffering capacity. E.g. NH4 buffer urine PO4 buffer systems H2PO4 proton donor H2PO42- proton acceptor H2PO4 H+ + H2PO42- Effective at pH close to pKa of 6.86 resist pH change 5.9 to 7.4 Effective biological buffer ECF & cytoplasmic compartments have pH of 6.9 to 7.4 HCO3- buffer system Buffer of blood plasma H2CO3 proton donor (Carbonic Acid) HCO3- proton acceptor (Bicarbonate)

H2CO3 H+ + HCO3-

More complex than other conjugate acid-base pairs bc of H2CO3 formed from dissolved (d) CO2 in H2O in a reversible rxn

CO2 (d) + H2O H2CO3

CO2 is a gas under N conditions & concentration of CO2 (d) is the result of equilibration w/ CO2 of gas (g) phase

CO2 (g) CO2(d)

pH of Bicarbonate system depends on [H2CO3] & [HCO3-] [H2CO3] depends on [CO2](d) [CO2] (d) depends on [CO2] (g) or the partial (p) of CO2 (pCO2) Thus the pH of a bicarbonate buffer exposed to a gas phase is ultimately determined by the [HCO3-] in the (aq) phase and by pCO2 in the (g) phase.

Bicarbonate Buffer System: in depth Effective buffer near pH 7.4 bc H2CO3in the blood is in equilibrium w. A large reserve capacity of CO2 in the airspaces in the lungs Involves 3 reversible equilibria btwn of CO2 in the lungs & H2CO3in the blood

Pathway:Vigorous exercise

Muscles Lactic Acid

H+

Blood

H+ + HCO3-

H2CO3

plasma CO2(d)

pCO2 airspace

CO2 exhaled

Pathway: CHON catabolism

NH3

plasma H+

H2CO3

Dissociation

H+ + HCO3-

CO2 (g) dissolved

CO2 (d)

Rate of respiration can quickly adjust pH to nearly constant Ctrled b Brain stem- E.g. detection of pCO2 or plasma pH deeper more freq breathing pH of plasma 7.4 very little H2CO3 is present in comparison w/ HCO3- Addition of small amt of base (NH3 or OH-) would titrate this H2CO3 exhausting now the buffering capacity. Impt role of Carbonic acid at pKa= 3.57 at 37oC in buffering blood plasma seem inconsistent at pH range 1 unit above and below its pKa Large reservoir of CO2(d) in the blood & its rapid equilibration w/ H2CO3

CO2(d) + H2O H2CO3

Kh as the equilibrium constant for the hydration of CO2

Then to take CO2(d) reservoir into accnt, we express as Kh [CO2(d)] in acid dissociation of

Experimentally derived value of Kh= 3.0x10-3 M Experimentally derived value of Ka= 2.7x10-4 M

Kcombi = (3.0x10-3 M) (2.7x10-4 M) = 8.1x10-7 M2pKcombi = -log KcombipKcombi= 6.1

In medicine, CO2(d) referred as the conjugate acid & use the pKcombi 6.1 to simplify calculate of pH from [CO2(d)]

Henderson Hasselbalch equation

pCO2 is expressed in kilopascals = 4.6 to 6.7 kPa 0.23 is the corresponding solubility for CO2 in H2O 1.2 kPa HCO3- in plasma 24mM

Untreated Diabetes Produce Life-Threatening Acidosis Blood pH 7.35- 7.45 Enzymes that fxn in the blood have evolved to have max activity in the pH range. Catalytic activity of enzymes are sensitive to pH pH optimum- pH in w/c enzymes show max catalytic activity Small change in pH can have a large diff in rate of some enzyme catalyzed rxn Biological ctrl of pH of cells & body flds is impt for metab & cellular activities Changes in pH can have physiological consequences Un-TxDM uses stored fatty acids as primary fuel w/c can lead to accumulation of hi [H2CO3], [-hydroxybutyric acid] & [Acetoacetic acid] 90 mg/mL; Healthy- < 3mg/mL Urinary excretion of 5g/ 24 hr; Healthy