B.Sc. Biochem II BPI Unit 1 Water, pH and Buffer

Course: B.Sc. BiochemistrySem IISub: Biophysics and InstrumentationUnit 1.1

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WATER

• Water is the most abundant substance inliving systems, making up 70% or more of theweight of most organisms.

• The first living organisms doubtless arose in anaqueous environment, and served as mediumfor evolution.

• All aspects of cell structure and function areadapted to it.

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• The attractive forces between water molecules andthe slight tendency of water to ionize are of crucialimportance to the structure and function ofbiomolecules.

• The water molecule and its ionization products, H+and OH-, and noncovalent interactions, hydrogenbonds, responsible for the strength and specificity ofbiomolecules

• The solvent properties of water, includes its ability toform hydrogen bonds with itself and with solutes.

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Weak Interactions

• Hydrogen bonds between water moleculesprovide the cohesive forces that make water aliquid at room temperature and that favor theextreme ordering of molecules that is typicalof crystalline water (ice).

• Polar biomolecules dissolve readily in waterbecause they can replace water-waterinteractions with more energetically favorablewater-solute interactions.

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• In contrast, nonpolar biomolecules interfere with water-waterinteractions but are unable to form water-soluteinteractions— consequently, nonpolar molecules are poorlysoluble in water. In aqueous solutions, nonpolar moleculestend to cluster together.

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O

H H

1 molecule of water is

made up of 2 hydrogen atoms

bonded with 1 oxygen atom. The

molecular formula is H2O

STRUCTURE OF WATER

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O

STRUCTURE OF WATER

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The bond that forms water

is a covalent bond

Water is a Polar Molecule-has oppositely charged ends

• Water consists of an oxygen atom bound to two hydrogen atoms by two single covalent bonds.– Oxygen has unpaired &

paired electrons which gives it a slightly negative charge while Hydrogen has no unpaired electrons and shares all others with Oxygen

– Leaves molecule with positively and negative charged ends

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• An electrostatic attraction between theoxygen atom of one water molecule and thehydrogen of another, called a hydrogen bond.

• Hydrogen bonds are relatively weak. Those inliquid water have a bond dissociation energy(the energy required to break a bond) ofabout 23 kJ/mol, compared with 470 kJ/molfor the covalent O-H bond.

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• The nearly tetrahedral arrangement of theorbitals about the oxygen atom allows eachwater molecule to form hydrogen bonds withas many as four neighboring water molecules.

• In liquid water at room temperature andatmospheric pressure, however, watermolecules are disorganized and in continuousmotion, so that each molecule formshydrogen bonds with an average of only 3.4other molecules.

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•In ice, each water

molecule is fixed in space

and forms hydrogen bonds

with a full complement of

four other water molecules

to yield a regular lattice

structure.

•This crystal lattice of ice

makes it less dense than

liquid water, and thus ice

floats on liquid water.

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Water Forms Hydrogen Bonds with Polar Solutes

• Hydrogen bonds are not unique to water.

• They readily form between an electronegativeatom and a hydrogen atom covalently bondedto another electronegative atom.

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• Hydrogen atoms covalently bonded to carbonatoms do not participate in hydrogen bonding,because carbon is only slightly moreelectronegative than hydrogen and thus the C-H bond is only very weakly polar.

• The distinction explains why butanol(CH3(CH2)2CH2OH) has a relatively highboiling point of 117 °C, whereas butane(CH3(CH2)2CH3) has a boiling point of only -0.5 °C.

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• Uncharged but polar biomolecules such as sugars dissolve readily in water because of the stabilizing effect of hydrogen bonds between the hydroxyl groups or carbonyl oxygen of the sugar and the polar water molecules.

• Alcohols, aldehydes, ketones, and compounds containing N-H bonds all form hydrogen bonds with water molecules and tend to be soluble in water.

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Water Interacts Electrostaticallywith Charged Solutes

• Water is a polar solvent.

• It readily dissolves most biomolecules, whichare generally charged or polar compounds

• Compounds that dissolve easily in water arehydrophilic.

• In contrast, nonpolar solvents such aschloroform and benzene are poor solvents forpolar biomolecules but easily dissolve thosethat are hydrophobic—nonpolar moleculessuch as lipids and waxes. 18

• Water dissolves salts such as NaCl byhydrating and stabilizing the Na+ and Cl- ions,weakening the electrostatic interactionsbetween them and thus counteracting theirtendency to associate in a crystalline lattice.

• The same factors apply to chargedbiomolecules, compounds with functionalgroups such as ionized carboxylic acids (-COO-), protonated amines (-NH3+)

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• Water readily dissolves such compounds byreplacing solute-solute hydrogen bonds withsolute-water hydrogen bonds, thus screeningthe electrostatic interactions between solutemolecules.

• The molecules of the biologically importantgases CO2, O2, and N2 are nonpolar makethem very poorly soluble in water.

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Properties of Water Polar Molecule Cohesion And Adhesion High Specific Heat Density – Greatest At 4oc Universal Solvent Of Life Capillary Action Surface Tension Buoyancy

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Ionization of Water

• Although many of the solvent properties ofwater can be explained in terms of theuncharged H2O molecule, the small degree ofionization of water to hydrogen ions (H+) andhydroxide ions (OH-) must also be taken intoaccount.

• Water molecules have a slight tendency toundergo reversible ionization to yield ahydrogen ion (a proton) and a hydroxide ion,giving the equilibrium

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• Although we commonly show the dissociationproduct of water as H+, free protons do notexist in solution; hydrogen ions formed inwater are immediately hydrated to hydroniumions (H3O+).

• Hydrogen bonding between water moleculesmakes the hydration of dissociating protonsvirtually instantaneous:

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Proton hopping

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• The ionization of water can be measured by itselectrical conductivity; pure water carrieselectrical current as H+ migrates toward thecathode and OH- toward the anode.

• The movement of hydronium and hydroxideions in the electric field is anomalously fastcompared with that of other ions such as Na+,K+, and Cl-.

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• No individual proton moves very far throughthe bulk solution, but a series of proton hopsbetween hydrogen-bonded water moleculescauses the net movement of a proton over along distance in a remarkably short time.

• As a result of the high ionic mobility, acid-basereactions in aqueous solutions are generallyexceptionally fast. Reversible ionization iscrucial to the role of water in cellular function

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Ionization of water inquantitative terms

• The position of equilibrium of any chemicalreaction is given by its equilibrium constant,Keq.

• For the generalized reaction,

• an equilibrium constant can be defined interms of the concentrations of reactants (Aand B) and products (C and D) at equilibrium:

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• Molarity is the unit of concentration used incalculating Keq.

• The equilibrium constant is fixed for any givenchemical reaction at a specified temperature.

• Conversely, we can calculate the equilibriumconstant for a given reaction at a giventemperature if the equilibrium concentrations ofall its reactants and products are known. 29

• The equilibrium constant for the reversible ionization of water is,

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• In pure water at 25° C, the concentration of water is 55.5 M

• which, on rearranging, becomes

• where Kw designates the product (55.5 M)(Keq), the ion product of water at 25 °C.

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• The value for Keq, determined by electrical-conductivity measurements of pure water, is1.8 X10^-16 M at 25° C.

• Substituting this value for Keq gives the valueof the ion product of water:

• Thus the product [H+][OH-] in aqueoussolutions at 25° C always equals 1 X10^-14M^2.

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• When there are exactly equal concentrationsof H+ and OH-, as in pure water, the solution issaid to be at neutral pH.

• At this pH, the concentration of H+ and OH-can be calculated from the ion product ofwater as follows:

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Water as a Reactant

• Water is both the solvent in which metabolicreactions occur and a reactant in manybiochemical processes, including hydrolysis,condensation, and oxidation-reductionreactions.

• Water is not just the solvent in which thechemical reactions of living cells occur; it isvery often a direct participant in thosereactions.

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• The formation of ATP from ADP and inorganicphosphate is an example of a condensationreaction in which the elements of water areeliminated

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• The reverse of this reaction— cleavageaccompanied by the addition of the elementsof water—is a hydrolysis reaction.

• The “metabolic water” formed by oxidation offoods and stored fats is actually enough toallow some animals in very dry habitats tosurvive for extended periods without drinkingwater.

• Green plants and algae use the energy ofsunlight to split water in the process ofphotosynthesis: 36

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The pH Scale Designates theH+ and OH- Concentrations

• The ion product of water, Kw, is the basis for the pH scale.

• The term pH is defined by the expression

• The symbol p denotes “negative logarithm of.”

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• For a precisely neutral solution at 25° C, in which the concentration of hydrogen ions is 1.0 X10^-7 M, the pH can be calculated as follows:

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• Solutions having a pH greater than 7 arealkaline or basic; the concentration of OH- isgreater than that of H+.

• Conversely, solutions having a pH less than 7are acidic.

• Two solutions differ in pH by 1 pH unit meansthat one solution has ten times the H+concentration of the other.

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• The pH of an aqueous solution can beapproximately measured using various indicatordyes, including litmus, phenolphthalein, andphenol red, which undergo color changes.

• Accurate determinations of pH in the laboratoryare made with a glass electrode that is selectivelysensitive to H+ concentration but insensitive toNa+, K+, and other cations.

• In a pH meter the signal from such an electrode isamplified and compared with the signalgenerated by a solution of accurately known pH. 43

• Measurement of pH is one of the most important andfrequently used procedures in biochemistry.

• The pH affects the structure and activity of biologicalmacromolecules; for example, the catalytic activity ofenzymes is strongly dependent on pH.

• Measurements of the pH of blood and urine arecommonly used in medical diagnoses.

• The pH of the blood plasma of people with severe,uncontrolled diabetes, for example, is often below thenormal value of 7.4; this condition is called acidosis. Incertain other disease states the pH of the blood ishigher than normal, the condition of alkalosis.

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Ampholytes• Ampholytes are molecules containing both acidic and basic groups.• All of the common amino acids found in proteins are ampholytes

because they contain a carboxyl group (-COOH) that acts as an acid andan amino group (-NH2) that acts as a base.

• As free amino acids, each amino acid has at least two pKa values (somehave more because they have additional acidic or basic groups).

• The titration of an ampholyte generates a more complex plot of pHversus moles of acid (or base) added than are obtained for a simplebuffer with only a single ionizing species because the ionization of eachacidic and basic group of the ampholyte is represented by a step in thetitration curve.

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Buffering against pH Changesin Biological Systems

• Almost every biological process is pH dependent;a small change in pH produces a large change inthe rate of the process.

• The enzymes that catalyze cellular reactions, andmany of the molecules on which they act, containionizable groups with characteristic pKa values.

• The protonated amino and carboxyl groups ofamino acids and the phosphate groups ofnucleotides, for example, function as weak acids;their ionic state depends on the pH of thesurrounding medium. 47

• Cells and organisms maintain a specific andconstant cytosolic pH, keeping biomolecules intheir optimal ionic state, usually near pH 7.

• In multicellular organisms, the pH ofextracellular fluids is also tightly regulated.Constancy of pH is achieved primarily bybiological buffers: mixtures of weak acids andtheir conjugate bases.

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Buffers Are Mixtures of Weak Acidsand Their Conjugate Bases

• Buffers are aqueous systems that tend toresist changes in pH when small amounts ofacid (H+) or base (OH-) are added.

• A buffer system consists of a weak acid (theproton donor) and its conjugate base (theproton acceptor).

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A Simple Expression Relates pH, pKa, and Buffer Concentration

• Henderson-Hasselbalch equation, which is important for understanding buffer action and acid-base balance in the blood and tissues of vertebrates.

• This equation is simply a useful way of restating the expression for the dissociation constant of an acid.

• For the dissociation of a weak acid HA into H+

and A-, the Henderson-Hasselbalch equation can be derived as follows: 51

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• Now invert -log [HA]/[A-], which involves changing its sign, to obtain the Henderson-Hasselbalch equation:

• it shows the pKa of a weak acid is equal to the pH of the solution when [HA] equals [A-], and

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• Henderson-Hasselbalch equation allows to:

(1) calculate pKa, given pH and the molar ratio of proton donor and acceptor;

(2) calculate pH, given pKa and the molar ratio of proton donor and acceptor; and

(3) calculate the molar ratio of proton donor and acceptor, given pH and pKa.

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Physiological importance of buffers in body fluids and tissues

• The intracellular and extracellular fluids ofmulticellular organisms have a characteristicand nearly constant pH.

• The cytoplasm of most cells contains highconcentrations of proteins, which containmany amino acids with functional groups thatare weak acids or weak bases.

• Nucleotides such as ATP, as well as many lowmolecular weight metabolites, containionizable groups that can contribute bufferingpower to the cytoplasm. 55

• Some highly specialized organelles andextracellular compartments have highconcentrations of compounds that contributebuffering capacity: organic acids buffer thevacuoles of plant cells; ammonia buffers urine.

• Two especially important biological buffers arethe phosphate and bicarbonate systems.

• The phosphate buffer system, which acts inthe cytoplasm of all cells, consists of H2PO4

- asproton donor and HPO4

2- as proton acceptor: 56

• The phosphate buffer system is maximallyeffective at a pH close to its pKa of 6.86 andthus tends to resist pH changes in the rangebetween about 5.9 and 7.9.

• It is therefore an effective buffer in biologicalfluids; in mammals, for example, extracellularfluids and most cytoplasmic compartmentshave a pH in the range of 6.9 to 7.4.

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• Blood plasma is buffered in part by thebicarbonate system, consisting of carbonicacid as proton donor and bicarbonate asproton acceptor:

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• This buffer system is more complex than other conjugate acid-base pairs because one of its components, carbonic acid (H2CO3), is formed from dissolved (d) carbon dioxide and water, in a reversible reaction:

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• Carbon dioxide is a gas under normal conditions, and the concentration of dissolved CO2 is the result of equilibration with CO2 of the gas (g) phase:

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• Although many aspects of cell structure andfunction are influenced by pH, it is the catalyticactivity of enzymes that is especially sensitive.

• Enzymes typically show maximal catalytic activityat a characteristic pH, called the pH optimum

• On either side of the optimum pH their catalyticactivity often declines sharply. Thus, a smallchange in pH can make a large difference in therate of some crucial enzyme-catalyzed reactions.

• Biological control of the pH of cells and bodyfluids is therefore of central importance in allaspects of metabolism and cellular activities. 61

References

reading

• Lehninger - Principle of Biochemistry (4th Ed.)

Images

• 1-10: Lehninger - Principle of Biochemistry (4th Ed.)

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