1Water him no get enemy. Fela kuti

Water, acids, bases and buffers

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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 intracellular and 60% extracellular

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• 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 average 70kg man

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• 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 (δ+), while its two unshared electron pairs constitute a region of local negative charge (δ-)

• The hydrogen nuclei on one molecule of water interacts with the lone pair on an oxygen atom on another water molecule

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The tetrahedral structure of the water molecule

Hydrogen bonding between water molecules

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• 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 O—H bond

• However, The cumulative effect of many hydrogen bonds is equivalent to the stabilizing effect of covalent bonds

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• 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 in hydrogen bonding

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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

• Water’s 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+ 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 is largely responsible for the ease of dissolving salts such as NaCl in water

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• In thermodynamic terms, formation of the solution occurs with a favorable free-energy change: ΔG =Δ H - T Δ S, where Δ H has a small positive value and T Δ S a large positive value; thus ΔG is negative

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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 hemoglobin and myoglobin ; CO2 is either carried as it is by hemoglobin or is changed to the soluble form –bicarbonate, HCO3

-

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Non-polar solutes in water • Non-polar compounds such as benzene and hexane

are hydrophobic—they 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 (ΔH) for dissolving these solutes is generally small

• Hydrophobic solutes, however, offer no such compensation, and their addition to water may therefore result in a small gain of enthalpy; the breaking of hydrogen bonds between water molecules takes up energy from the system

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• 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: ΔG=ΔH - TΔS, where ΔH has a positive value, TΔS has a negative value, and ΔG is positive

• This unfavorable state is relieved when the non-polar solutes coalesce to form droplets

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• 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 of the solvation spheres is the strong tendency of water molecules to form hydrogen bonds among themselves; non-polar molecules are excluded because they cannot form hydrogen bonds

15Formation of an oil-droplet in an aqueous solution

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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

• These stable structures of amphipathic compounds in water, called micelles, may contain hundreds or thousands of molecules

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• 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 while their hydrophobic fatty acyl side chains cluster together, excluding water

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Liposomes are formed through the sonication of a solution of amphipathic molecules.

They have a potential for

drug delivery

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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 and the phosphorus atoms of phosphoesters

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• 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 earth’s water would be steam, making life unlikely

• However, water actually melts at 0 0C and boils at +100 0C; consequently, it is a liquid over most of the wide range of temperatures found on the earth’s surface

• Hydrogen bonding is responsible for this behavior of water

• Energy is required to break hydrogen bonds

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• 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 water’s temperature is substantially higher than expected

• One consequence of water’s high heat of vaporization (the energy required to vaporize 1 mole of a substance at 1 atm) and high heat capacity (the energy that must be added or removed to change the temperature by one degree Celsius) is that water acts as an effective modulator of climatic (and body) temp.

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• Water can absorb solar heat and release it slowly

• Water’s high heat capacity, coupled with the high water content found in most organisms helps maintain an organism’s 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 by metabolic processes

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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, for example, as does an equal number of moles of a non-dissociating solute such as glucose

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• 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 + and Cl-) is dissolved in 1 kg of water, the solution is 2-osmolal

• Measurement of colligative properties is useful in estimating solute concentrations in biological fluids. For example, in blood plasma, the normal total concentration of solutes is remarkably constant (275-295 milliosmolal).

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• 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

• Cells also actively pump out ions such as Na+ into the interstitial fluid to stay in osmotic balance with their surroundings

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• 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+, K+, Cl- and HCO3

- • In addition, the intracellular fluid and the plasma

contain non-diffusible proteins

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• 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 lower in the protein-containing compartment

• Gibbs-Donnan equilibrium is established when the altered distribution of cations and anions results in electrochemical equilibrium

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Semi-permeable membrane

Distribution of inorganic ions in the absence of non-diffusible ions

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More Cl- leaves I to balance charges

Distribution of inorganic ions in the presence of non-diffusible ions

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• 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+ thereby regulating cell volume

• About 1/3 of ATP in an animal cell is used to power Na+-K+ pumps; in nerve cells, which use Na+ and K+ gradients to propagate electrical signals, up to 2/3 of the ATP is used to power these pumps

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Dissociation of Water and the pH Scale• Acids are compounds that donate a hydrogen ion

(H+) 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+ (hydronium), but nevertheless are usually represented simply as H+.Water itself is neutral, neither acidic nor basic

• For the dissociation of water: where the brackets represent

molar concentrations and K is the dissociation constant

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• 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

• KW is known as the ion product of water• Since the concentrations of [H+] [OH_] in pure

water are equal:[H+]= [OH_] = 10-7 M

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• 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

• Strong acids/bases completely dissociate in water; weak acid/bases dissociate only partially

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• 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.

Similarly, the conjugate of a strong base is a weak acid. The relative strengths of bases are expressed in terms of the pKa of their conjugate acids

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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

• Consider the titration of a 0.1 M solution of acetic acid (for simplicity denoted as HAc) with 0.1 M NaOH at 25 0C

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• Two reversible equilibria are involved in the process:

H2O <--->H++ OH-

HAc <---> H++ 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 OH- combines with the free H+ in the solution to form H2O, to an extent that satisfies the equilibrium relationship of water

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• As free H+ 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 increments of NaOH, the remaining non-dissociated acetic acid is gradually converted into acetate. The end point of the titration occurs at about pH 7.0: all the acetic acid has lost its protons to OH-, to form water and acetate

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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+)

• 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+ or OH- added to the system has much less effect on pH than the same amount added outside the buffer range

• This relatively flat zone is the buffering region of the acetic acid–acetate buffer pair

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• 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 the system is maximal

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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:

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• 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= pKa + 2• If the ratio [A-]/[HA] is 1:10, pH= pKa - 1; …

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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+ concentration of ca. 30%

• The pH value of cytoplasm is slightly lower than that of blood, at 7.0–7.3

• In the lumen of the gastrointestinal tract and in the body’s 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 bases, depending on the state of the metabolism, the pH of urine has a particularly wide range of variation (4.8–7.5)

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• If the H + concentration departs significantly from its normal value, the health and survival of the human body are in jeopardy

• H + is the smallest ion, and it combines with many negatively charged and neutral functional groups

• Changes of [H +], 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 eliminated from the body, buffer systems maintain a fairly constant pH in body fluids

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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 +], significant variations of [H +]in a patient may not be fully appreciated

• For example, if the blood pH decreases from 7.4 to 7.1, [H +] is doubled; or if the pH increases from 7.4 to 7.7, [H +] is halved

• Thus, in clinical situations it is preferable to express [H +] 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 +], which is the mean of the normal range ; the normal range is 7.36-7.44 on the pH scale, or 44-36 nM [H +]

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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 (H2SO4) and hydrochloric acid (HCl) are strong acids

• Organic acids containing carboxylic acid groups (e.g., the ketone bodies acetoacetic acid and β-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 unbuffered body fluids, their pH would be less than 1

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• 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 bicarbonate–carbonic 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

• Under normal metabolic conditions, the body generates more than 13 moles of CO2 per day (approximately 0.5-1 kg)

• About 95% of the CO2 entering the blood diffuses into the red blood cells

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• 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+ and bicarbonate anion, HCO3-

• Although H2CO3 is a weak acid, its dissociation is essentially 100% because of the removal of H+ ions by the buffering action of hemoglobin

• The remainder of the H+ is buffered by phosphate and proteins other than hemoglobin

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• 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

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• 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 small fraction of the concentration of HCO3- and the plasma appears to be poorly protected against an influx of OH- . How is this problem solved?

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• 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 H2CO3 availability is [CO2(d)] + [H2CO3], the so-called total carbonic acid pool, consisting primarily of CO2(d)

• The overall equilibrium for the bicarbonate buffer system is:

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• 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 CO2, and from Ka, the acid dissociation constant for H2CO3:

• Putting this value for [H2CO3] into the expression for the dissociation of H2CO3 gives:

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• 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 H2CO3 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

• It appears that the conjugate acid, because of its small concentration would be overwhelmed by small amounts of alkali

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• However, the acid component is the total carbonic acid pool, that is, [CO2(d)] + [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 CO2(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

solubility coefficient, a = 0.03 mmol of dissolved CO2 per liter of plasma per mm Hg of CO2 pressure

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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 + H+

• 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 system’s capability, the concentrations of H2PO4

- and HPO4-2 in blood are too low (4mEq/L)

to have a major effect • Instead, the phosphate system is an important

buffer in intracellular fluids where its concentration is about 75 mEq/L

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• 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 their low concentrations and pka that is much lower than intracellular pH

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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

• However, protein-bound and dipeptide histidine may be the dominant buffering system in some cells

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• In combination with other amino acids, as in proteins or dipeptides, the imidazole pka 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

Histidine

Imidazole

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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 acids, cysteine and methionine

• It is a strong acid that is dissociated into H+ and sulfate anion (SO4

-2) in the blood and urine

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• 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+) ions a reaction that occurs with a pKa of 9.25

• Ammonia is produced from the catabolism of nitrogen containing biomolecules and kept at very low concentrations in the blood because it is toxic to neural tissues

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• Cells in the kidney generate NH4+ and excrete

it into the urine in proportion to the acidity of the blood

• As the renal tubular cells transport H+ 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 lining