Chapter 2: Water
Hydrogen bonds are strongest...
when the hydrogen atom and the two atoms that share it are in a straight line
Delta G, Enthalpy and Entropy
ΔG = ΔH − T ΔS, where ΔG represents the driving force, ΔH the enthalpy change from making and breaking bonds, and ΔS the change in randomness. Because ΔH is positive for melting and evaporation, it is clearly the increase in entropy (ΔS) that makes ΔGnegative and drives these changes. (ΔG) must have a negative value for a process to occur spontaneously:
Keq
A constant, characteristic for each chemical reaction, that relates the specific concentrations of all reactants and products at equilibrium at a given temperature and pressure. Keq = [products]/[reactants] molarity is the unit of concentration used in calculating Keq. The equilibrium constant is fixed and characteristic for any given chemical reaction at a specified temperature. It defines the composition of the final equilibrium mixture, regardless of the starting amounts of reactants and products
Acidosis
A metabolic condition in which the capacity of the body to buffer H+ is diminished; usually accompanied by decreased blood pH.
Alkalosis
A metabolic condition in which the capacity of the body to buffer OH− is diminished; usually accompanied by an increase in blood pH.
Titration curve
A plot of pH versus the equivalents of base added during titration of an acid. At the midpoint of the titration, the concentrations of the proton donor and proton acceptor are equal, and the pH is numerically equal to the pKa.
Conjugate acid-base pair
A proton donor and its corresponding deprotonated species; for example, acetic acid (donor) and acetate (acceptor). Acids may be defined as proton donors and bases as proton acceptors
Condensation reaction
A reaction type in which two compounds are joined with the elimination of water.
Buffers Are Mixtures of Weak Acids and Their Conjugate Bases
A system capable of resisting changes in pH, consisting of a conjugate acid-base pair in which the ratio of proton acceptor to proton donor is near unity (The same) As an example, a mixture of equal concentrations of acetic acid and acetate ion, found at the midpoint of the titration curve in Figure 2-17, is a buffer system.
Cumulative effect
Although these four types of interactions are individually weak relative to covalent bonds, the cumulative effect of many such interactions can be very significant
Henderson-Hasselbalch equation
An equation relating pH, pKa, and ratio of the concentrations of proton-acceptor (A−) and proton-donor (HA) species in a solution:pH = pKa+ log[A−][HA] For example, it shows why the pKa of a weak acid is equal to the pH of the solution at the midpoint of its titration. At that point, [HA] = [A−], and pH = pKa + log 1 = pKa + 0 = pKa
Entropy as driving force in melting/evaporation of water at RT
At room temperature, both the melting of ice and the evaporation of water occur spontaneously; the tendency of the water molecules to associate through hydrogen bonds is outweighed by the energetic push toward randomness
Hydrogen Bond
Bond between a partially negative atom (FON) and a partially positive hydrogen atom A weak electrostatic attraction between one electronegative atom (such as oxygen or nitrogen) and a hydrogen atom covalently linked to a second electronegative atom.
Osmosis official
Bulk flow of water through a semipermeable membrane into another aqueous compartment containing solute at a higher concentration. (less water mols.)
Hydrolysis
Cleavage of a bond, such as an anhydride or peptide bond, by addition of the elements of water, yielding two or more products.
Amphipathic
Containing both polar and non polar regions. hen an amphipathic compound is mixed with water, the polar, hydrophilic region interacts favorably with the water and tends to dissolve, but the nonpolar, hydrophobic region tends to avoid contact with the water. The nonpolar 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 , a phenomenon called the hydrophobic effect.
enzyme-substrate complex.
Disruption of ordered water molecules is part of the driving force for binding of a polar substrate (reactant) to the complementary polar surface of an enzyme: entropy increases as the enzyme displaces ordered water from the substrate and as the substrate displaces ordered water from the enzyme surface
Entropy during melting/evaporation
During melting or evaporation, the entropy of the aqueous system increases as the highly ordered arrays of water molecules in ice relax into the less orderly hydrogen-bonded arrays in liquid water or into the wholly disordered gaseous state
Hydrolases
Enzymes (e.g., proteases, lipases, phosphatases, nucleases) that catalyze hydrolysis reactions.
pH optimum
Enzymes typically show maximal catalytic activity at a characteristic pH, called the pH optimum. On either side of this optimum pH, catalytic activity often declines sharply. Thus, a small change in pH can make a large difference in the rate of some crucial enzyme-catalyzed reactions.
Hydrogen bonding energetic effect
Hydrogen bonding between water and polar solutes also causes an ordering of water molecules, but the energetic effect is less significant than with nonpolar solutes.
More info about hydrolysis reaction:
Hydrolysis reactions are also responsible for the enzymatic depolymerization of proteins, carbohydrates, and nucleic acids. BREAKING Hydrolysis reactions - are almost invariably exergonic; by producing two molecules from one, they lead to an increase in the randomness of the system. As we shall see, cells circumvent this thermodynamic obstacle by coupling endergonic condensation reactions to exergonic processes, such as breakage of the anhydride bond in ATP.
hypertonic
In a hypertonic solution, one with higher osmolarity than that of the cytosol, the cell shrinks as water moves out. SHRINKS (more water mol inside cell)
Hyper/hypotonic
In a hypertonic solution, one with higher osmolarity than that of the cytosol, the cell shrinks as water moves out. In a hypotonic solution, one with a lower osmolarity than the cytosol, the cell swells as water enters.
hypotonic
In a hypotonic solution, one with a lower osmolarity than the cytosol, the cell swells as water enters. (less water mol. inside cell) 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 movementof water would distend the plasma membrane and eventually cause bursting of the cell (osmotic lysis).
Water as a solvent
It readily dissolves most biomolecules, which are generally charged or polar compounds. Compounds that dissolve easily in water are hydrophilic. Water also dissolves salts like NaCl
Weak Acids and Bases Have Characteristic Acid Dissociation Constants
Ka and Kb weak acids and bases are not completely ionized when dissolved in water.
Proton hopping
No individual proton moves very far through the bulk solution, but a series of proton hops between hydrogen-bonded water molecules causes the net movement of a proton over a long distance in a remarkably short time. High ionic mobility results in PH. OH- also moves rapidly by proton hopping, but in the opposite direction Short "hops" of protons between a series of hydrogen-bonded water molecules result in an extremely rapid net movement of a proton over a long distance. As a hydronium ion (upper left) gives up a proton, a water molecule some distance away (bottom) acquires one, becoming a hydronium ion.
Nonpolar Compounds Force Energetically Unfavorable Changes in the Structure of Water
Nonpolar 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 among water molecules. The free-energy change for dissolving a nonpolar solute in water is thus unfavorable: ΔG = ΔH − T ΔS, where ΔH has a positive value, ΔS has a negative value, and ΔG is positive
Hydrophobic
Nonpolar; describes molecules or groups that are insoluble in water.
Buffer region
One unit of pH up and down. At the midpoint of the buffering region, where the concentration of the proton donor (acetic acid) exactly equals that of the proton acceptor (acetate), the buffering power of the system is maximal; that is, its pH changes least on addition of H+ or OH-. The pH at this point in the titration curve of acetic acid is equal to its pKa.
The Fitness of the Aqueous Environment for Living Organisms
Organisms have effectively adapted to their aqueous environment and, in the course of evolution, have developed means of exploiting the unusual properties of water The high specific heat of water (the heat energy required to raise the temperature of 1 g of water by 1 °C) is useful to cells and organisms because it allows water to act as a "heat buffer," keeping the temperature of an organism relatively constant as the temperature of the surroundings fluctuates and as heat is generated as a byproduct of metabolism.
Osmotic Pressure
Osmotic pressure, Π, measured as the force necessary to resist water movement (Fig. 2-12c), is approximated by the van't Hoff equation: Π = icRT
Participation of water in biological reactions
Participation of water in biological reactions. ATP is a phosphoanhydride formed by a condensation reaction (loss of the elements of water) between ADP and phosphate. R represents adenosine monophosphate (AMP). This condensation reaction requires energy. The hydrolysis of (addition of the elements of water to) ATP to form ADP and phosphate releases an equivalent amount of energy. These condensation and hydrolysis reactions of ATP are just one example of the role of water as a reactant in biological processes.
Hydrogen bonds not just in water
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
Hydrophilic
Polar or charged; describes molecules or groups that associate with (dissolve easily in) water.
Solutes Affect the Colligative Properties of Aqueous Solutions
Solutes of all kinds alter certain physical properties of the solvent, water: its vapor pressure, boiling point, melting point (freezing point), and osmotic pressure.
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 or ions) in a given amount of water.
Solutes of all kinds alter certain physical properties of the solvent, water: its vapor pressure, boiling point, melting point (freezing point), and osmotic pressure.
isotonic
Solutions of osmolarity equal to that of a cell's cytosol are said to be isotonic relative to that cell. Surrounded by an isotonic solution, a cell neither gains nor loses water
Conjugate acid-base pairs consist of a proton donor and a proton acceptor.
Some compounds, such as acetic acid and ammonium ion, are monoprotic: they can give up only one proton. Others are diprotic (carbonic acid and glycine) or triprotic (phosphoric acid). The dissociation reactions for each pair are shown where they occur along a pH gradient. The equilibrium or dissociation constant (Ka) and its negative logarithm, the pKa, are shown for each reaction
Hydrophobic effect
The aggregation of nonpolar molecules in aqueous solution, excluding water molecules; caused largely by an entropic effect related to the hydrogen-bonding structure of the surrounding water.
Titration Curves Reveal the pKa of Weak Acids
The amounts of acid and base in titrations are often expressed in terms of equivalents, where one equivalent is the amount of a substance that will react with, or supply, one mole of hydrogen ions in an acid-base reaction.
acid dissociation constant (Ka)
The dissociation constant (Ka) of an acid, describing its dissociation into its conjugate base and a proton Stronger acids, such as phosphoric and carbonic acids, have larger ionization constants; weaker acids, such as monohydrogen phosphate (HPO2−4)(HPO42−), have smaller ionization constants.
Example of condensation
The formation of ATP from ADP and inorganic phosphate is an example of a condensation reaction in which the elements of water are eliminated. The formation of cellular polymers from their subunits would be endergonic and therefore does not occur.
Hydronium ions
The hydrated hydrogen ion (H3O+).
Hydrogen acceptor and hydrogen donor
The hydrogen acceptor is usually oxygen or nitrogen; the hydrogen donor is another electronegative atom.
Ionization of water can be measured by its electrical conductivity
The ionization of water can be measured by its electrical conductivity; pure water carries electrical current as H3O+ migrates toward the cathode and OH- toward the anode.
Nonpolar Gases Are Poorly Soluble in Water
The molecules of the biologically important gases CO2, O2, and N2 are nonpolar. movement of molecules from the disordered gas phase into aqueous solution constrains their motion and the motion of water molecules and therefore represents a decrease in entropy. The nonpolar nature of these gases and the decrease in entropy when they enter solution combine to make them very poorly soluble in water
Number of bonds formed
The nearly tetrahedral arrangement of the orbitals about the oxygen atom (Fig. 2-1a) allows each water molecule to form hydrogen bonds with as many as four neighboring water molecules. In liquid water at room temperature and atmospheric pressure, however, water molecules are disorganized and in continuous motion, so that each molecule forms hydrogen bonds with an average of only 3.4 other molecules. In ice, on the other hand, 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
pKa
The negative logarithm of an equilibrium constant. pKa=−logKa The stronger the tendency to dissociate a proton, the stronger is the acid and the lower its pKa. High Ka, low pKa As we shall now see, the pKa of any weak acid can be determined quite easily.
pH scale
The negative logarithm of the hydrogen ion concentration of an aqueous solution. pH = −log [H+]
Non-covalent bonds
The noncovalent interactions we have described—hydrogen bonds and ionic, hydrophobic, and van der Waals interactions —are much weaker than covalent bonds. Consequently, hydrogen bonds and ionic, hydrophobic, and van der Waals interactions are continually forming and breaking.
Kw (ion product constant for water)
The product of the concentrations of H+ and OH− in pure water; Kw = [H+][OH−] = 1 × 10−14 at 25 °C. Kw=[H+] [OH−] When [H+] is very high, as in a solution of hydrochloric acid, [OH-] must be very low. From the ion product of water we can calculate [H+] if we know [OH-], and vice versa.
Colligative effects
The properties of a solution that depend on the number of solute particles per unit volume; for example, freezing-point depression.
Hydrolysis reaction
The reverse of this condensation reaction—cleavage accompanied by the addition of the elements of water—is a hydrolysis reaction.
The Henderson-Hasselbalch Equation Relates pH, pKa, and Buffer Concentration
The shape of the titration curve of any weak acid is described by the Henderson-Hasselbalch equation, which is important for understanding buffer action and acid-base balance in the blood and tissues of vertebrates.
Entropy Increases as Crystalline Substances Dissolve
This is spontaneous (negative delta G). As a salt such as NaCl dissolves, the Na+ and Cl− ions leaving the crystal lattice acquire far greater freedom of motion. The resulting increase in entropy (randomness) of the system is largely responsible for the ease of dissolving salts such as NaCl in water. In thermodynamic terms, formation of the solution occurs with a favorable free-energy change: ΔG = ΔH − T ΔS, where ΔH has a small positive value and T ΔS a large positive value; thus ΔG is negative.
Water and salts
Water dissolves many crystalline salts by hydrating their component ions. The NaCl crystal lattice is disrupted as water molecules cluster about the Cl−and Na+ ions. The ionic charges are partially neutralized, and the electrostatic attractions necessary for lattice formation are weakened.
Hydrogen Bonding Gives Water Its Unusual Properties
Water has a higher melting point, boiling point, and heat of vaporization than most other common solvents. These unusual properties are a consequence of attractions between adjacent water molecules that give liquid water great internal cohesion.
Water as a reactant
Water is not just the solvent in which the chemical reactions of living cells occur; it is very often a direct participant in those reactions.
Pure water is slightly ionized
Water molecules have a slight tendency to undergo reversible ionization to yield a hydrogen ion (a proton) and a hydroxide ion, giving the equilibrium H2O⇌H++OH−
Osmosis
Water molecules tend to move from a region of higher water concentration to one of lower water concentration, in accordance with the tendency in nature for a system to become disordered. 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 the region of lower water concentration produce osmotic pressure
Glucose oxidation
When glucose is oxidized = form water and CO2. Water is both the solvent in which metabolic reactions occur and a reactant in many biochemical processes, including hydrolysis, condensation, and oxidation-reduction reactions.
Weak acids and bases
When weak acids are dissolved in water, they contribute H+ by ionizing; weak bases consume H+ by becoming protonated.
How a buffer system works?
Whenever H+ or OH− is added to a buffer, the result is a small change in the ratio of the relative concentrations of the weak acid and its anion and thus a small change in pH. The decrease in concentration of one component of the system is balanced exactly by an increase in the other. The sum of the buffer components does not change, only their ratio changes. Each conjugate acid-base pair has a characteristic pH zone in which it is an effective buffer . The H2PO−4/HPO2−4H2PO4−/HPO42− pair has a pKa of 6.86 and thus can serve as an effective buffer system between approximately pH 5.9 and pH 7.9; the NH+4/NH3NH4+/NH3 pair, with a pKa of 9.25, can act as a buffer between approximately pH 8.3 and pH 10.3. + - 1 of pKa is the range of buffer system in pH
Important
endothermic and exothermic CAN be spontaneous ENDERGONIC (gibbs free positive) non spontaneous EXERGONIC (Gibbs free negative) spontaneous To make endergonic reactions occur, couple them!!!!
What bonds are present between water molecules?
hydrogen bonds Each hydrogen atom of a water molecule shares an electron pair with the central oxygen atom. Unequal electron sharing in two electric dipoles in the water molecule. As a result, there is an electrostatic attraction between the oxygen atom of one water molecule and the hydrogen of another called a hydrogen bond. oxygen atom of the upper molecule and a hydrogen atom of the lower one. Hydrogen bonds are longer and weaker than covalent O—H bonds. O-H covalent between water is electrostatic hydrogen bonds
nonpolar solvents
nonpolar solvents such as chloroform and benzene are poor solvents for polar biomolecules but easily dissolve those that are hydrophobic—nonpolar molecules such as lipids and waxes.
pH scale is logarithmic
pH by 1 pH unit means that one solution has ten times the H+concentration of the other
ph and pKa
1.At the midpoint of the titration, at which exactly 0.5 equivalent of NaOH (base) has been added per equivalent of the acid 2.one-half of the original acetic acid has undergone dissociation, 3.so that the concentration of the proton donor, [HAc], now equals that of the proton acceptor, [Ac−]. pH = pKa --------------------------------------------------- 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. pH = pKa at this point 1/2 of original acid has undergone dissociation. Equal concentration of acid and base at this point. Concentration of the proton donor (acetic acid) exactly equals that of the proton acceptor (acetate). pka=pH (equimolar and half dissociation) pH is equal to the pKa when there are equal amounts of protonated and deprotonated forms of the acid The pKa can be determined experimentally; it is the pH at the midpoint of the titration curve for the acid or base
