Water, acids, bases and buffers Water him no

Water, acids, bases and buffers Water him no

Water, acids, bases and buffers Water him no get enemy. Fela kuti 1

THE BIOLOGICAL IMPORTANCE OF WATER Water is an ideal biological solvent: it dissolves and transports a wide variety of organic and inorganic molecules Water influences the conformations of many biomolecules Water is a reactant or a product in many reactions Water removes excess heat from the body Total body water is roughly 50 to 60% of body weight in adults and 75% of body weight in children Because fat has relatively little water associated with it, obese people tend to have a lower percentage of body water than thin people, women tend to have a lower percentage than men, and older people have a lower percentage than younger people

Approximately 40% of the total body water is 2 intracellular and 60% extracellular The extracellular water includes the fluid in plasma (blood after the cells have been removed) and interstitial water (the fluid in the tissue spaces, lying between cells) Transcellular water is a small, specialized portion of extracellular water that includes saliva, gastrointestinal secretions, ,urine, sweat, cerebrospinal fluid,. Fluid compartments in the body based on an 3 The unique properties of water are due to its

structure Hydrogen bonding A water molecule is an irregular, slightly skewed tetrahedron with oxygen at its center The 1050 angle between the hydrogens differs slightly from the ideal tetrahedral angle, 109.50 Water is a dipole, a molecule with electrical charge distributed asymmetrically about its structure The strongly electronegative oxygen atom pulls electrons away from the hydrogen nuclei, leaving them with a partial positive charge (+), while its two ), while its two unshared electron pairs constitute a region of local negative charge (-)) The hydrogen nuclei on one molecule of water interacts with the lone pair on an oxygen atom on 4

The tetrahedral structure of the water molecule drogen bonding between water molecules 5 Hydrogen bonding favors the self-)association of water molecules into ordered arrays Hydrogen bonding profoundly influences the physical properties of water and accounts for its exceptionally high viscosity, surface tension and boiling point On average, each molecule in liquid water associates through hydrogen bonds with 3.4 others; these bonds are both relatively weak and

transient, with a half-)life of pico seconds In ice, each water molecule forms a hydrogen bond with four other water molecules, giving rise to a crystalline tetrahedral arrangement Rupture of a hydrogen bond in liquid water requires only about 4.5 kcal/mol, less than 5% of the energy required to rupture a covalent OH 6 Hydrogen bonding enables water to dissolve many organic biomolecules that contain functional groups which can participate in hydrogen bonding The oxygen atoms of aldehydes, ketones, and amides, for example, provide lone pairs of electrons that can serve as hydrogen acceptors; alcohols and amines can serve both as hydrogen acceptors and as donors of unshielded hydrogen

atoms for formation of hydrogen bonds Polar groups participating i hydrogen bonding 7 The interaction of water with charged solutes Water has a high dielectric constant; it greatly decreases the force of attraction between charged and polar species relative to water-)free environments with lower dielectric constants Waters strong dipole and high dielectric constant enable water to dissolve large quantities of charged compounds such as salts Water dissolves salts such as NaCl by hydrating

and stabilizing the Na+), while its two and Cl-) ions, weakening the electrostatic interactions between them and thus counteracting their tendency to associate in a crystalline lattice As a salt dissolves, the ions leaving the crystal lattice acquire far greater freedom of motion The resulting increase in entropy of the system is8 In thermodynamic terms, formation of the solution occurs with a favorable free-)energy change: G = H -) T S, where H has a small positive value and T S a large positive value; thus G is negative 9 Non-polar gases and water

The molecules of the biologically important gases O2, CO2 and N2 are non-)polar In O2 and N2, electrons are shared equally by both atoms. In CO2, each C=O bond is polar, but the two dipoles are oppositely directed and cancel each other out The movement of molecules from the disordered gas phase into aqueous solution constrains their motion and the motion of water molecules and represents a decrease in entropy The non-)polar nature of these gases and the decrease in entropy when they enter solution combine to make them very poorly soluble in water O2 is carried by the water soluble proteins 10

11 Non-polar solutes in water Non-)polar compounds such as benzene and hexane are hydrophobicthey are unable to undergo energetically favorable interactions with water molecules, and they interfere with the hydrogen bonding between water molecules All molecules or ions in aqueous solution interfere with the hydrogen bonding of some water molecules in their immediate vicinity, but polar or charged solutes (such as NaCl) compensate for lost water-)water hydrogen bonds by forming new solute-)water interactions; the net change in enthalpy (H) for dissolving these solutes is generally small

Hydrophobic solutes, however, offer no such compensation, and their addition to water may 12 Furthermore, dissolving hydrophobic compounds in water produces a measurable decrease in entropy. Water molecules in the immediate vicinity of a non-)polar solute are constrained in their possible orientations as they form a highly ordered cage-)like shell around each solute molecule The ordering of water molecules reduces entropy. The number of ordered water molecules, and therefore the magnitude of the entropy decrease, is proportional to the surface area of the hydrophobic solute enclosed within the cage of water molecules The free energy change for dissolving a non-)polar

solute in water is thus unfavorable: G=H -) TS, where H has a positive value, TS has a 13 negative value, and G is positive This coalition is called the hydrophobic effect /interaction Hydrophobic interaction refers to the tendency of non-)polar compounds to self-)associate in an aqueous environment This self-)association is driven neither by mutual attraction nor by what are sometimes incorrectly referred to as hydrophobic bonds. Self-) association minimizes energetically unfavorable interactions between non-)polar groups and water A solvation sphere of hydrogen-)bonded water

molecules forms around the hydrophobic molecules Although non-)polar molecules, when in close proximity, are attracted to each other by van der Waals forces, the driving force for in the formation 14 of the solvation spheres is the strong tendency of ormation of an oil-droplet in an aqueous solution 15

Amphipathic molecules in water Amphipathic compounds contain regions that are polar (or charged) and regions that are non-) polar When an amphipathic compound is mixed with water, the polar, hydrophilic region interacts favorably with the solvent and tends to dissolve, but the non-)polar, hydrophobic region tends to avoid contact with the water The non-)polar regions of the molecules cluster together to present the smallest hydrophobic area to the aqueous solvent, and the polar regions are arranged to maximize their interaction with the solvent

16 These stable structures of amphipathic Many biomolecules are amphipathic: proteins tend to fold with the R-) groups of amino acids with hydrophobic side chains in the interior; amino acids with charged or polar amino acid side chains generally are present on the surface in contact with water A similar pattern

prevails in a phospholipid bilayer, where the charged head groups contact water 17 Liposomes are formed through the sonication of a solution of amphipathic molecules. They have a potential for drug delivery

18 19 Water as a Participant in Chemical Reactions Metabolic reactions often involve the attack by lone pairs of electrons residing on electron-)rich

molecules termed nucleophiles upon electron-) poor atoms called electrophiles Nucleophiles and electrophiles do not necessarily possess a formal negative or positive charge; water, whose two lone pairs of electrons bear a partial negative charge, is an excellent nucleophile Other nucleophiles of biologic importance include the oxygen atoms of phosphates, alcohols and carboxylic acids; the sulfur of thiols; the nitrogen of amines; and the imidazole ring of histidine Common electrophiles include the carbonyl carbons in amides, esters, aldehydes, and ketones 20 and the phosphorus atoms of phosphoesters Nucleophilic attack by water generally results in

the cleavage of the amide, glycoside, or ester bonds that hold biopolymers together; this process is termed hydrolysis Conversely, when monomer units are joined together to form biopolymers such as proteins or glycogen, water is a product The Thermal Properties of Water If water followed the pattern of compounds such as hydrogen sulfide, it would melt at -)100 0C and boil at -)910C Under these conditions, most of the earths water would be steam, making life unlikely However, water actually melts at 0 0C and boils at +), while its two 100 0C; consequently, it is a liquid over most of 21 the wide range of temperatures found on the

When ice is warmed to its melting point, approximately 15% of the hydrogen bonds break Liquid water consists of ice-)like clusters of molecules whose hydrogen bonds are continuously breaking and forming As the temperature rises, the movement and vibrations of the water molecules accelerate and additional hydrogen bonds are broken When the boiling point is reached, the water molecules break free from one another and vaporize The energy required to raise waters temperature is substantially higher than expected One consequence of waters high heat of vaporization (the energy required to vaporize 1

22 mole of a substance at 1 atm) and high heat Water can absorb solar heat and release it slowly Waters high heat capacity, coupled with the high water content found in most organisms helps maintain an organisms internal temperature The evaporation of water is used as a cooling mechanism, because it permits large losses of heat For example, an adult human may eliminate as much as 1200g of water daily in expired air, sweat and urine The associated heat loss may amount to

approximately 20% of the total heat generated 23 by metabolic processes Colligative Properties Solutes of all kinds alter certain physical properties of the solvent, water: its vapor pressure, boiling point, melting point (freezing point), and osmotic

pressure These are called colligative (tied together) properties, because the effect of solutes on all four properties has the same basis: the concentration of water is lower in solutions than in pure water The effect of solute concentration on the colligative properties of water is independent of the chemical properties of the solute; it depends only on the number of solute particles (molecules, ions) in a given amount of water A compound such as NaCl, which dissociates in solution, has twice the effect on osmotic pressure, 24 Water molecules tend to move from a region of higher water concentration to one of lower water

concentration osmosis When two different aqueous solutions are separated by a semipermeable membrane (one that allows the passage of water but not solute molecules), water molecules diffusing from the region of higher water concentration to that of lower water concentration produce osmotic pressure A solution containing 1 mol of solute particles in 1 kg of water is a 1-osmolal solution When 1 mol of a solute (such as NaCl) that dissociates into two ions (Na +), while its two and Cl-)) is dissolved in 1 kg of water, the solution is 2-)osmolal Measurement of colligative properties is useful in25 ffects

26 of solu tes on colliga tive p ro p If a cell is put in In a hypotonic solution, with lower osmolality than the cytosol, the cell swells as water enters In their natural environments, cells generally contain higher concentrations of biomolecules and ions than their surroundings, so osmotic pressure tends to drive water into cells

If not somehow counterbalanced, this inward movement of water would distend the plasma membrane and eventually cause bursting of the cell (osmotic lysis) In multicellular animals, blood plasma and interstitial fluid are maintained at an osmolality close to that of the cytosol; the high concentration of albumin and other proteins in blood plasma contributes to its osmolality 27 28 Because the effect of solutes on osmolality depends on the number of dissolved particles, not their mass, macromolecules (proteins, nucleic

acids, polysaccharides) have far less effect on the osmolality of a solution than would an equal mass of their monomeric components One effect of storing fuel as polysaccharides (starch or glycogen) rather than as glucose or other simple sugars is prevention of an enormous increase in osmotic pressure within the storage cell The Gibbs-Donnan Equilibrium The three fluid compartments, that is, the intracellular fluid, interstitial fluid and blood plasma each contain diffusible ions such as Na+), while its two , K+), while its two , Cl-) and HCO3-) 29

The negatively charged, non-)diffusible proteins present predominantly in the plasma space will attract positively charged ions and repel negatively charged ions Despite the high permeability of small ions across membranes, a similar concentration of ionic species is not seen The passive distribution of cations and anions is altered to preserve electroneutrality in the compartments The normal difference in concentrations of diffusible ions between the plasma and interstitial compartments is due to the presence of non-) diffusible proteins in plasma The diffusible cation concentration is higher in the compartment containing non-)diffusible, anionic

proteins, whereas diffusible anion concentration is 30 Semi-permeable membrane stribution of inorganic ions in the absence of 31 non-diffusible ions More Cl- leaves I to balance charges Distribution of inorganic ions in the presence non-diffusible ions 32 The existence of ionic asymmetry on the surfaces on the surface of cell membrane results in the

establishment of the electrochemical gradient or membrane potential which provides the means for electrical conduction and active and passive transport A related outcome is that water tends to move from the interstitial space to the plasma (maintaining blood volume) and the intercellular space (causing a constant threat of cellular swelling) Cells must, therefore constantly regulate their osmolality; many animal and bacterial cells pump out inorganic ions such as Na+), while its two thereby regulating cell volume About 1/3 of ATP in an animal cell is used to power 33 +), while its two

+), while its two +), while its two +), while its two Dissociation of Water and the pH Scale Acids are compounds that donate a hydrogen ion (H+), while its two ) to a solution, and bases are compounds (such as the OH-) ion) that accept hydrogen ions Water itself dissociates to a slight extent, generating hydrogen ions , which are also called protons, and hydroxide ions H2O <---> H+ + OH

The hydrogen ions are extensively hydrated in water to form species such as H3O+), while its two (hydronium), but nevertheless are usually represented simply as H+), while its two .Water itself is neutral, neither acidic nor basic For the dissociation of water: where the brackets represent molar concentrations and K 34is Since 1 mole (mol) of water weighs 18 g, 1 liter (L) (1000 g) of water contains 1000/18 = 55.56 mol. Pure water thus is 55.56 molar K can be determined by measurement of the electrical conductivity of pure water, which has the value of 1.8 x 10 -)16 M at 25 indicative of a very small ion concentration, where M (molar) is

the unit of moles per liter Therefore, the concentration of undissociated water is essentially unchanged by the dissociation reaction Substituting for the values of K and [H2O]: _ [H ] [OH ] = 1.8 x10 -)16 M x 55.56 M = 1 x 10-)14 M2 =KW +), while its two KW is known as the ion product of water 35 pH is employed to express proton concentrations

in a convenient form; it is the negative log (to the base ten) of the hydrogen ion concentration: pH=-log[H+] For pure water, pH=-)log [10-)7 ]=7; and pOH =-)log [10-)7 ]=7 A pH of 7 is termed neutral because [H+] and [OH-] are equal. Acidic solutions have a greater hydrogen ion concentration and a lower hydroxide ion concentration (pH<7) than pure water and basic solutions have a lower hydrogen ion concentration and a greater hydroxide ion concentration (pH >7) A decrease in one pH unit reflects a 10-)fold increase in H+ concentration 36

Many biochemicals possess functional groups (carboxyl groups, amino groups, phosphate esters, ) that are weak acids or bases The relative strengths of weak acids and bases are expressed in terms of their dissociation constants For the reaction HA<---> A- +H+ Where Ka is the dissociation constant, HA is the conjugate acid and A-) is the conjugate base Since the numeric values of Ka for weak acids are negative exponential numbers, pKa is used where pKa = -log Ka The stronger the acid the lower its pKa value For any weak acid, its conjugate is a strong base.37

Titration curves reveal the pKa Titration is used to determine the amount of an acid in a given solution A measured volume of the acid is titrated with a solution of a strong base, usually NaOH, of known concentration The NaOH is added in small increments until the acid is consumed (neutralized), as determined with an indicator dye or a pH meter The concentration of the acid in the original solution can be calculated from the volume and concentration of NaOH added A plot of pH against the amount of NaOH added (a titration curve) reveals the pKa of the weak acid

38 Two reversible equilibria are involved in the process: H2O <-)-)-)>H+), while its two +), while its two OH-) HAc <-)-)-)> H+), while its two +), while its two Ac-) The equilibria must simultaneously conform to their characteristic equilibrium constants, which are, respectively, At the beginning of the titration, before any NaOH is added, the acetic acid is already slightly ionized, to an extent that can be calculated from its dissociation constant As NaOH is gradually introduced, the added OH39 -)

As free H+), while its two is removed, HAc dissociates further to satisfy its own equilibrium constant The net result as the titration proceeds is that more and more HAc ionizes, forming Ac-), as the NaOH is added At the midpoint of the titration, at which exactly 0.5 equivalent of NaOH has been added, one-)half of the original acetic acid has undergone dissociation that the concentration of the proton donor, [HAc], now equals that of the proton acceptor, [Ac-)] At this midpoint, a very important relationship holds: the pH of the equimolar solution of acetic acid and acetate is exactly equal to the pKa of acetic acid (4.76) As the titration is continued by adding further 40

increments of NaOH, the remaining non-)dissociated 41 The titration of acetic aci d

What are Buffers? Buffers are solutions that resist change in pH when small amounts of proton (acid) or hydroxide (base) are added They are either a mixture of a weak acid (HA) and its conjugate base (A-)) or a mixture of a weak base (B) and its conjugate acid (HB+), while its two ) The mixture of equal concentrations of acetic acid and acetate ion, found at the midpoint of the titration curve is a buffer system The titration curve of acetic acid has a relatively flat zone extending about 1 pH unit on either side of its midpoint pH 4.76 In this zone, an amount of H+), while its two or OH-) added to the

system has much less effect on pH than the same 42 amount added outside the buffer range At the midpoint of the buffering region, where the conc. of the proton donor exactly equals that of the

proton acceptor, the buffering power of 43 The Henderson-Hasselbalch Equation The shape of the titration curve of weak acids and bases is described by the Henderson-)Hasselbalch equation This equation relates pH, pKa and the concentration of conjugate acid-)base pairs; it is derived as follows:

44 At the midpoint of titration, the concentrations of proton acceptor and donor are equal; log (1)= 0; pH= pKa If the ratio [A-)]/[HA] is 100:1, pH= pK a +), while its two 2 If the ratio [A-)]/[HA] is 1:10, pH= pKa -) 1; 45

Normal pH values in organisms pH values in the cell and in the extracellular fluids are kept constant within narrow limits In the blood, the pH value normally ranges only between 7.35 and 7.45; this corresponds to a maximum change in the H+), while its two concentration of ca. 30% The pH value of cytoplasm is slightly lower than that of blood, at 7.07.3 In the lumen of the gastrointestinal tract and in the bodys excretion products, the pH values are more variable Extreme values are found in the stomach (ca.2) and in the small intestine (> 8)

Since the kidney can excrete either acids or 46 If the H +), while its two concentration departs significantly from its normal value, the health and survival of the human body are in jeopardy H +), while its two is the smallest ion, and it combines with many negatively charged and neutral functional groups Changes of [H +), while its two ], therefore, affect the charged regions of many molecular structures, such as enzymes, cell membranes and nucleic acids, and dramatically alter physiological activity If the plasma pH reaches either 6.8 or 7.8, death may be unavoidable Despite the fact that large amounts of acidic and basic metabolites are produced and

47 eliminated from the body, buffer systems 48 A More Meaningful Way of Stating the Concentration of Hydrogen Ions

In clinical acid-)base problems, the use of the pH scale has some disadvantages Since the pH is the logarithm of the reciprocal of [H +), while its two ], significant variations of [H +), while its two ]in a patient may not be fully appreciated For example, if the blood pH decreases from 7.4 to 7.1, [H +), while its two ] is doubled; or if the pH increases from 7.4 to 7.7, [H +), while its two ] is halved Thus, in clinical situations it is preferable to express [H +), while its two ] directly as nanomoles per liter in order to better evaluate acid-)base changes and interpret laboratory tests A blood pH of 7.40 corresponds to 40 nM [H +), while its two ], 49 which is the mean of the normal range ; the normal

Metabolic Acids and Bases During metabolism, the body produces a number of acids that increase the hydrogen ion concentration of the blood or other body fluids and tend to lower the pH These metabolically important acids can be weak

acids or strong acids Inorganic acids such as sulfuric acid (H 2SO4) and hydrochloric acid (HCl) are strong acids Organic acids containing carboxylic acid groups (e.g., the ketone bodies acetoacetic acid and -) hydroxybutyric acid) are weak acids An average rate of metabolic activity produces roughly 22,000 mEq acid per day If all of this acid were dissolved at one time in 50 51 52 Until the acid produced from metabolism can be excreted as CO2 in expired air and as ions (and

unmetabolized organic acids) in the urine, it needs to be buffered in the body fluids The major buffer systems in the body are: the bicarbonatecarbonic acid buffer system, which operates principally in extracellular fluid; the hemoglobin buffer system in red blood cells; the phosphate buffer system in all types of cells; the protein buffer system in cells and plasma and phosphate and ammonia in the urine The Bicarbonate Buffer System The major source of metabolic acid in the body is the gas CO2, produced principally from fuel oxidation in the TCA cycle 53 Under normal metabolic conditions, the body

Within the red blood cells, the enzyme carbonic anhydrase I catalyzes the conversion of most of the CO2 to carbonic acid (H2CO3 ) Carbonic anhydrase II is found in most tissues including the lung, bone and renal tubular cells Carbonic acid is a weak acid that partially dissociates into H+), while its two and bicarbonate anion, HCO3 Although H2CO3 is a weak acid, its dissociation is essentially 100% because of the removal of H+), while its two ions by the buffering action of hemoglobin The remainder of the H+), while its two is 54 hroc

yte mem bran e pr o t ei ns As the concentration of HCO 3-) (i.e., of metabolic CO2) in red blood cells increases, an imbalance occurs between the bicarbonate ion concentrations in the red blood cell and plasma This osmotic imbalance causes a marked efflux of HCO 3-) to plasma and consequent influx of Cl-) from plasma in order to maintain the balance of charges This exchange, known as the chloride shift, takes place through an antiporter known as band

3 protein 55 Once in the extracellular fluid, HCO3-) serves as a major buffer Most of the CO2 produced in the body reaches the lungs carried by the plasma in the form of HCO3-) In the lungs, the events that took place in the erythrocytes are reversed and CO2 is exhaled Buffering capacity is greatest at or near the pKa of the conjugate-)acid base pair The pKa of H2CO3 is 3.8 but it is a good buffer at the blood pH of 7.4; how could this be? The most effective buffers are those that contain

equal concentrations of both components. But at pH 7.4, the concentration of H2CO3 is a very -) 56 Gaseous carbon dioxide from the lungs and tissues is dissolved in the blood plasma, symbolized as CO2(d), and hydrated to form H2CO3: In mammalian body fluids, the equilibrium for the carbonic anhydrase reaction lies far to the left, such that about 500 CO2 molecules are present in solution for every molecule of H2CO3 Because dissolved CO2 and H2CO3 are in equilibrium, the proper expression for H 2CO3

availability is [CO2(d)] +), while its two [H2CO3], the so-)called total carbonic acid pool, consisting primarily of CO 2(d) The overall equilibrium for the bicarbonate buffer 57 system is: An expression for the ionization of H2CO3 under such conditions (that is, in the presence of dissolved CO2) can be obtained from Kh, the equilibrium constant for the hydration of CO 2, and from Ka, the acid dissociation constant for H2CO3: Putting this value for [H2CO3] into the expression for the dissociation of H2CO3 gives:

58 KaKh, the product of two constants, can be defined as a new equilibrium constant, Koverall The value of Kh is 0.003 and Ka is equal to 0.000269. Therefore, Koverall = 8.07 x 10-7 and pKoverall = 6.1 This gives a modified Henderson-)Hasselbalch equation for the bicarbonate buffer system: The concentration gap that existed between H 2CO3 and HCO3-) has been greatly narrowed by usingCO2(d) in the equation But still, 6.1 is more than one unit away from 7.4 and the ratio of conjugate base (bicarbonate) over

conjugate acid (mainly carbondioxide) is 20:1 59 However, the acid component is the total carbonic acid pool, that is, [CO2(d)] +), while its two [H2CO3], which is stabilized by its equilibrium with CO2(g) The gaseous CO2 buffers any losses from the total carbonic acid pool by entering solution as CO 2(d), and blood pH is effectively maintained Thus, the bicarbonate buffer system is an open system In the equilibrium expression for the bicarbonate-) carbonic acid buffer system at pH 7.4, the carbonic acid term can be replaced by a pressure term because the carbonic acid concentration is proportional to the partial pressure of carbondioxide ,PCO2, in the blood

For normal plasma at 370C 60 61 The Phosphate Buffer System Phosphate is an abundant anion in cells, both in inorganic form and as an important functional group on organic molecules that serve as metabolites or macromolecular precursors The inorganic phosphate buffer consists of the weak acid-)conjugate base pair dihydrogen phosphate/hydrogen phosphate H2PO4-) <-)-)-)> HPO4-)2 +), while its two H+), while its two The pka of the system is 7.2 so it would appear

that it is an excellent choice for buffering blood Although the blood pH of 7.4 is well within the buffer systems capability, the concentrations of H2PO4-) and HPO4-)2 in blood are too low (4mEq/L) to have a major effect 62 The titration of p hosphoric Organic phosphate anions, such as glucose 6-) phosphate and ATP, also act as intracellular buffers Although cells contain other weak acids these substances are unimportant as buffers because of 63

their low concentrations and pka that is much lower Protein Buffers (The Histidine System) Histidine is one of the 20 naturally occurring amino acids commonly found in proteins It possesses as part of its structure an imidazole

group, a five-)membered heterocyclic ring possessing two nitrogen atoms. The pKa for dissociation of the imidazole hydrogen of histidine is 6.04 In cells, histidine occurs as the free amino acid, as a constituent of proteins, and as part of dipeptides in combination with other amino acids Because the concentration of free histidine is low and its imidazole pKa is more than 1 pH unit removed from prevailing intracellular pH, its role in intracellular buffering is minor 64 In combination with other amino acids, as in proteins or dipeptides, the imidazole pk a may increase substantially approach the physiological

pH The pka of weak acids can be affected by their environments The main protein in erythrocytes, hemoglobin, uses its histidines to buffer the protons released from carbonic acid and other sources Other cells are endowed with other proteins that assist in intracellular buffering Albumin in the blood also serves as a buffer ol e I z a d

mi Histidine 65 Urinary Buffers The non-)volatile acid that is produced from body

metabolism cannot be excreted as expired CO2 and is excreted in the urine Most of the non-)volatile acid hydrogen ion is excreted as undissociated acid that generally buffers the urinary pH between 5.5 and 7.0; a pH of 5.0 is the minimum urinary pH The acid secretion includes inorganic acids such as phosphate and ammonium ions, as well as uric acid, dicarboxylic acids, and tricarboxylic acids such as citric acid Sulfuric acid is generated from the sulfate-) containing compounds ingested in foods and from metabolism of the sulfur-)containing amino 66 Urinary excretion of phosphate ions helps to

remove acid; to maintain metabolic homeostasis, we must excrete the same amount of phosphate in the urine that we ingest with food as phosphate anions or organic phosphates such as phospholipids Whether the phosphate is present in the urine as H2PO4-) or HPO4 -2 depends on the urinary pH and the pH of blood Ammonium ions are major contributors to buffering urinary pH, but not blood pH Ammonia (NH3) is a base that combines with protons to produce ammonium (NH4+), while its two ) ions a reaction that occurs with a pKa of 9.25 Ammonia is produced from the catabolism of 67

Cells in the kidney generate NH4+), while its two and excrete it into the urine in proportion to the acidity of the blood As the renal tubular cells transport H+), while its two into the _ urine, they return HCO3 anions to the blood Hydrochloric acid (HCl), also called gastric acid, is secreted by parietal cells of the stomach into the stomach lumen, where the strong acidity denatures ingested proteins so they can be degraded by digestive enzymes When the stomach contents are released into the lumen of the small intestine, gastric acid is neutralized by bicarbonate secreted from pancreatic cells and by cells in the intestinal

68 lining

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