Section 1 Introduction to Biochemical Principles

Chapter 3 Water: The Matrix of Life

Section 3.1: Molecular Structure of Water  Water is essential for life  Water’s important properties include:  Chemical stability  Remarkable solvent properties  Role as a biochemical reactant  Hydration of organism

Section 3.1: Molecular Structure of Water  Water has a tetrahedral geometry  Oxygen is more electronegative than hydrogen (3.5 vs. 2.1) Figure 3.2 Tetrahedral Structure of Water

Section 3.1: Molecular Structure of Water Figure 3.4 Water Molecule Figure 3.3 Charges on a Water Molecule  Larger oxygen atom has partial negative charge (  - ) and hydrogen atoms have partial positive charges (  + )

Section 3.1: Molecular Structure of Water  Bond between oxygen and hydrogen is polar  Water is a dipole because the positive and negative charges are separate Figure 3.5 Molecular Dipoles in an Electric Field

Section 3.1: Molecular Structure of Water  An electron-deficient hydrogen of one water is attracted to the unshared electrons of water forming a hydrogen bond  Can occur with oxygen, nitrogen, and sulfur (i.e., when there is significant difference in EN between H and higher EN atom)  Has electrostatic (i.e., opposite charges) and covalent (i.e., electron sharing) characteristics  Large numbers of hydrogen bonds lead to extended network

Section 3.2: Noncovalent Bonding  Noncovalent interactions are electrostatic  Weak individually, but play vital role in biomolecules because of cumulative effects  Individual H-bonds not strong (~20 kJ/mol), but collectively, they aggregate  Why water has high MP, BP

 Three most important noncoavalent bonds:  Ionic interactions  Van der Waals forces  Hydrogen bonds Section 3.2: Noncovalent Bonding

 Ionic Interactions  Oppositely charged ions attract one another  Biochemistry primarily investigates the interaction of charged groups on molecules, which differs from ionic interactions like those of ionic compounds (e.g., NaCl) Section 3.2: Noncovalent Bonding  Ionized amino acid side chains can form salt bridges with one another

Figure 3.8 Dipolar Interactions  Three types of interactions:  Dipole-dipole  Two permanent dipoles align themselves by charge  (e.g., H-bonds)  Dipole-induced dipole  Permanent dipole induces transient dipole in another molecule by distorting its electron distribution  (e.g., carbonyl induces dipole in aromatic ring pi electrons)  Induced dipole-induced dipole  Transient dipole in 1 molecule induces transient dipole in another molecule  (e.g., Stacking of base rings in DNA: individually weak but collectively strong) Section 3.2: Noncovalent Bonding Van der Waals Forces: Occur between molecules with permanent, and/or induced dipoles

 Water’s melting and boiling points are exceptionally high due to hydrogen bonding  Each water molecule can form four hydrogen bonds with other water molecules  Forms extended network of hydrogen bonds  This explains thermal properties of water: It requires a lot of E input into system to melt ice and boil water, in order to overcome E of H-bonds Section 3.3: Thermal Properties of Water

 Water has an exceptionally high heats of fusion, and vaporization.  Takes a lot of energy to convert water from solid to liquid, and liquid to gas.  Related to H-bonding. Section 3.3: Thermal Properties of Water

 Water has a high heat capacity.  Heat capacity is energy needed to change T of 1 g of substance by 1 ° C.  Water can absorb a lot of E before T increases.  Most living organisms are comprised of ~50-95% water.  The high water content and heat capacity of water help maintain internal T of organisms.  Evaporation of water from body is cooling mechanism.  Human adults eliminate ~1200g of water daily in expired air, sweat, and urine. Section 3.3: Thermal Properties of Water q = Energy in Joules m = mass (g) C = heat capacity ΔT = Change in temperature Q: How much E is required to raise T of 1 Kg of water from 65 ° C to 75 ° C? C of water = J/g· ° C

Figure 3.10 Solvation Spheres  Water is an ideal biological solvent  Dissolves ions, sugars, many amino acids (ionic and polar substances)  Water forms solvation spheres around ions, molecules, thus separating them  Does not dissolve lipids and some amino acids  Allows biological processes (e.g., protein folding and membrane formation) Section 3.4: Solvent Properties of Water

Figure 3.11 Diagrammatic View of Structured Water  Structured Water  Water is rarely free flowing  Most of the time, water molecules in cell are non- covalently associated with macromolecules and other cellular components (membranes, proteins, etc.)  A single layer of water molecules on the surface of a biomolecule attracts an second layer, and so on.  Forms complex three- dimensional bridges between cellular components Section 3.4: Solvent Properties of Water

 Sol-Gel Transitions  Cytoplasm has properties of a gel (semi-solid colloidal mixture)  Transition from gel to sol (more liquid state) important in cell movement  Amoeboid motion provides an example of regulated, cellular, sol-gel transitions  Caused by reversible polymerization of G-actin to form F- actin (gel becomes more solid, expands, then gel becomes more liquid, creating a contractile force that pulls the gel forward) Section 3.4: Solvent Properties of Water Figure 3.12 Amoeboid Movement

 Hydrophobic Molecules and the Hydrophobic Effect  Small amounts of nonpolar substances are excluded from the solvation network forming droplets  This hydrophobic effect results from the solvent properties of the water and is stabilized by van der Waals interactions  Water molecule form a cage around the hydrophobic droplets, isolating them.  This effect generates stable lipid membranes and contributes to protein folding. Section 3.4: Solvent Properties of Water Figure 3.13 The Hydrophobic Effect

Section 3.4: Solvent Properties of Water Figure 3.14 Formation of Micelles  Amphipathic Molecules  Contain both polar and nonpolar groups  Amphipathic molecules form micelles when mixed with water  Important feature for the formation of cellular compartments, membranes

 Osmotic Pressure  Osmosis is the spontaneous passage of solvent molecules through a semipermeable membrane  Osmotic pressure is the pressure required to stop the net flow of water across the membrane  Over time, water diffuses from side A (more dilute) to side B (more concentrated)  Osmotic pressure depends on solute concentration Section 3.4: Solvent Properties of Water Figure 3.15 Osmotic Pressure

Section 3.4: Solvent Properties of Water  Can be measured with an osmometer or calculated (=iMRT, p. 89)  Cells may gain or lose water because of the environmental solute concentration  Solute concentration differences between the cell and the environment can have important consequences.  (a) When cells are placed in isotonic solution (i.e., concentrations of solute and water are the same on both sides of selectively permeable membrane), no movement of water is observed in either direction.  (b) A cell in a hypotonic solution (i.e., solution has lower solute concentration), water moves into cells. For example, red blood cells in pure water will absorb water, swell and rupture.  (c) In a hypertonic solution (i.e., higher solute concentration), water moves out of cell and cell shrivels. Figure 3.17 Effect of Solute Concentration on Animal Cells

Section 3.4: Solvent Properties of Water  Most macromolecules have little effect on cellular osmolality, but proteins can, because their ionizable amino acid side chains attract ions of opposite charge.  Size of an ions solvation sphere is inversely related to its charge density. For example, Na+ has a smaller radius than K+, but larger hydration sphere due to charge density. It requires less energy to remove solvation sphere from K+, so there is a tendency for more K+ to accumulate in the cell.

 Intracellular proteins tend to produce a significant net negative charge in cell, which is partially offset by K+ ions.  The extracellular environment has a net positive charge due to Na+ ions.  This asymmetric charge distribution results in formation of an electrical gradient (membrane potential) which provides the means for electrical conduction, active transport, and passive transport Membrane Potential

Self-Ionization of Water Pure water contains H 2 O molecules. In addition, small but equal amounts of H 3 O + and OH - ions are also present. The reason for this is that in one liter of pure water 1.0 x moles of water molecules behave as Brønsted acids and donate protons to another 1.0 x moles of water molecules, which act as Brønsted bases. The reaction is: As a result, absolutely pure water contains 1.0 x mol/L of both H 3 O + and OH -. The term neutral is used to describe any water solution in which the concentrations of H 3 O + and OH - are equal. Thus, pure water is neutral because each of the ions is present at a concentration of 1.0 x M. H 2 O (l) + H 2 O (l) ⇆ H 3 O + (aq) + OH − (aq) K

Self-Ionization of Water Equilibrium in reactions K Equilibrium of water K  K w is ion product of water K w is a constant, but temperature dependent. The concentration of either acid or base may change, but K w does not. If [H 3 O + ] > [OH - ], solution acidic, e.g., [H 3 O + ] = 1.0 x 10 -4, so [OH - ] = 1.0 x at 25°C and 1 atm pressure,

Section 3.5: Ionization of Water  Acids, Bases, and pH (BL definitions)  An acid is a proton donor  A base is a proton acceptor  Most organic molecules that donate or accept protons are weak acids or weak bases  A deprotonated product of a dissociation reaction is a conjugate base

Section 3.5: Ionization of Water  The pH scale can be used to measure hydrogen ion concentration  pH=-log[H + ] Figure 3.18 The pH Scale and the pH Values of Common Fluids

Section 3.5: Ionization of Water  pK a is used to express the strength of a weak acid  Lower pK a equals a stronger acid  pK a =-logK a  K a is the acid dissociation constant Figure 3.18 The pH Scale and the pH Values of Common Fluids

Section 3.5: Ionization of Water

 Buffers  Regulation of pH is universal and essential for all living things  Certain diseases can cause changes in pH that can be disastrous  Acidosis (blood pH falls below 7.35; blood pH below 7.00 leads to death) and Alkalosis (blood pH above 7.45, leads to convulsions and respiratory arrest)  Buffers help maintain a relatively constant hydrogen ion concentration  Commonly composed of a weak acid and its conjugate base

Section 3.5: Ionization of Water  Buffers Continued  Establishes an equilibrium between buffer’s components  Follows Le Chatelier’s principle  Equilibrium shifts in the direction that relieves the stress Figure 3.19 Titration of Acetic Acid with NaOH

Section 3.5: Ionization of Water  Henderson-Hasselbalch Equation  Establishes the relationship between pH and pK a for selecting a buffer  Buffers are most effective when they are composed of equal parts weak acid and conjugate base  Biological buffers usually contain more conj. base, since biological systems generate acid during metabolism  Best buffering occurs 1 pH unit above and below the pK a pH = pK a + log [A - ] [HA] Henderson-Hasselbalch Equation

Section 3.5: Ionization of Water  Buffering Capacity  Capacity to maintain a certain pH depends on:  Molar concentration of acid-conj. Base pair  Ratio of their concentrations  Increasing concentration increases buffering capacity  Buffer concentration is sum of concentration of weak acid and its conjugate base

Section 3.5: Ionization of Water  Worked Problem 3.5 (Page 91)  Calculate the pH of a mixture of 0.25 M acetic acid (CH 3 COOH) and 0.1 M sodium acetate (NaC 2 H 3 O 2 )  The pK a of acetic acid is 4.76 Solution: pH = pK a + log [acetate] [acetic acid] pH = log [0.1] [0.25] = = 4.36 What happens if we add HCl to a concentration of 0.75 M HCl?

Section 3.5: Ionization of Water  In-Class Problem  Calculate the pH of a mixture prepared by mixing 150 mL of 0.10 M HCl with 300 mL of 0.10 M sodium acetate (NaC 2 H 3 O 2 ), and diluting the mixture to 1.0 L. The pK a of acetic acid is (CH 3 COOH) (NaC 2 H 3 O 2 )

Section 3.5: Ionization of Water  In-Class Problem  Calculate the pH of a mixture prepared by mixing 150 mL of 0.10 M HCl with 300 mL of 0.10 M sodium acetate (NaC 2 H 3 O 2 ), and diluting the mixture to 1.0 L. The pK a of acetic acid is (CH 3 COOH) (NaC 2 H 3 O 2 ) Solution: pH = pK a + log [acetate] [acetic acid] pH = log 5 25 = 4.76 – 0.70 = 4.06

Section 3.5: Ionization of Water Figure 3.20 Titration of Phosphoric Acid with NaOH  Weak Acids with Multiple Ionizable Groups  Each ionizable group can have its own pK a  Protons are released in a stepwise fashion

Section 3.5: Ionization of Water  Physiological Buffers  Buffers adapted to solve specific physiological problems within the body  Bicarbonate Buffer  One of the most important buffers in the blood  CO 2 + H 2 O  H + + HCO 3 - (HCO 3 - is bicarbonate): This is a reversible reaction  Carbonic anhydrase is the enzyme responsible

Section 3.5: Ionization of Water Figure 3.21 Titration of H 2 PO 4 - by Strong Base  Phosphate Buffer  Consists of H 2 PO 4 - /HPO 4 2- (weak acid/conjugate base)  H 2 PO 4 -  H + + HPO 4 2-  Important buffer for intracellular fluids  Protein Buffer  Proteins are a significant source of buffering capacity (e.g., hemoglobin)