Water him no get enemy . Fela kuti

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

2. 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

3. 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

4. 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

5. The tetrahedral structure of the water molecule Hydrogen bonding between water molecules

6. 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

7. 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

8. 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

9. 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

10. 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-

12. 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

13. 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

14. 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

15. Formation of an oil-droplet in an aqueous solution

16. 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

17. 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

18. Liposomes are formed through the sonication of a solution of amphipathic molecules. They have a potential for drug delivery

20. Water as a Participant in Chemical Reactions • Metabolic reactions often involve the attack by lone pairs of electrons residing on electron-rich molecules termed nucleophilesupon 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

21. 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

22. 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.

23. 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

24. 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

25. 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).

26. The effects of solutes on colligative properties

27. 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

29. 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

30. 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

31. Semi-permeable membrane Distribution of inorganic ions in the absence of non-diffusible ions

32. More Cl- leaves I to balance charges Distribution of inorganic ions in the presence of non-diffusible ions

33. 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

34. 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

35. 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

36. 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

37. 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, pKais used where • pKa = -log Ka • The stronger the acid the lower its pKavalue • 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

38. 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

39. 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

40. 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 pKaof 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

41. The titration of acetic acid

42. 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

43. 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

44. 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, pKaand the concentration of conjugate acid-base pairs; it is derived as follows:

45. 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; …

46. 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)

47. 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

49. 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 +]

50. 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