Water and Aqueous Solutions PDF

Water and aqueous solutions Marc A. Ilies, Ph. D. Lehninger - Chapter 2 (p47-73) [email protected]; lab 517, office 517A (Tu, Fr 3-5) For questions, comments please use the discussion tool in Canvas or recitation sess...

Water and aqueous solutions Marc A. Ilies, Ph. D. Lehninger - Chapter 2 (p47-73) [email protected]; lab 517, office 517A (Tu, Fr 3-5) For questions, comments please use the discussion tool in Canvas or recitation session ©MAIlies 2024 1  Water – the solvent of life on Earth - water is the most abundant substance in living systems (~ 70% weight in most organisms) - water is the solvent of life and shaped evolution of life on Earth: the first living organisms on Earth doubtless arose in an aqueous environment, and water continues to play a central role on life on Earth - attractive forces between water molecules through hydrogen bonding and dipole-dipole interactions, the ability of water to dissolve and associate ions through ion-dipole interactions, and the slight tendency of water to ionize are of crucial importance to the structure and function of biomacromolecules - water and its ionizable products H+ (H3O+) and OH- ions influence the properties of all cellular components – proteins, sugars, nucleic acids, lipids, including ionization, self-assembling, etc 2  Structure of water Octet rule dictates that there are four electron pairs around an oxygen atom in water These electrons are in four sp3 orbitals Two of these pairs covalently link two hydrogen atoms to a central oxygen atom The two remaining pairs remain nonbonding (lone pairs) Water geometry is a distorted tetrahedron The electronegativity of the oxygen atom induces a net dipole moment Because of the dipole moment, water can serve as both a hydrogen bond donor and acceptor 3  H-bonding in solid and liquid water Water is associated through H bonds in solid (ice) and liquid state Water has many different crystal forms; the hexagonal ice is the most common Hexagonal ice forms a regular lattice, and thus has a low entropy Hexagonal ice contains more hydrogen bonds/water molecule Thus, ice has lower density than liquid water; Ice floats on top of liquid water When ice melts some of the H-bonds are broken; density increases – max water density at 4 oC; a dynamic structure in which H-bonds are made and broken simultaneously “flickering clusters of water” is observed in liquid state (1 water molecule associated with ~ 3.4 other waters at 25 oC): Hydrogen bonding in ice 4  H-bonding in water: consequences Melting Point, Boiling Point, and Heat of Vaporization of Some TABLE 2-1 Common Solvents Breaking of H bonds Melting point (˚C) Melting point (˚C) Heat of vaporization (J/g)a Water 0 100 2,260 Methanol (CH3OH) –98 65 1,100 Ethanol (CH3CH2OH) –117 78 854 Propanol (CH3CH2CH2OH) –127 97 687 Butanol (CH3(CH2)2CH2OH) –90 117 590 Acetone (CH3COCH3) –95 56 523 Hexane (CH3(CH2)4CH3) –98 69 423 Benzene (C6H6) 6 80 394 Butane (CH3(CH2)2CH3) –135 –0.5 381 Chloroform (CHCl3) –63 61 247 a The heat energy required to convert 1.0 g of a liquid at its boiling point and at atmospheric pressure into its gaseous state at the same temperature. It is a direct measure of the energy required to overcome attractive forces between molecules in the liquid phase. 6  Important properties of water Water has polar covalent bonds, highly individual polar bonds High Polarity: Hydrogen bonding, dissolution of ions, polar molecules High Specific Heat Capacity: Thermal insulation High Heat of Vaporization: Sweating cools the body High Dielectric Constant: Electrical insulation (charge separation) Maximum Density at 4o C: Ice floats Hydrophobic Effects: Micelles, bilayers & amphipathic compounds assemblies, protein folding, enzyme substrate interactions Colligative Properties 7  Water forms hydrogen bonds with polar solutes Hydrogen bonds are formed by many compounds containing highly electronegative elements F, O, N Strength: typically 4–6 kJ/mol for bonds with neutral atoms, and 6–10 kJ/mol for bonds with one charged atom Hydrogen bonds are strongest when the bonded molecules are oriented to maximize electrostatic interaction Ideally the three atoms involved are in a line Glucose in water 8  Polarity of biomolecules - water and biomolecules with functional groups containing O, N are polar, can associate through H-bonding, dipole-dipole interactions → polar biomolecules are soluble in water, (water-compatible, water-loving, hydrophilic) - non-polar molecules (e.g. hydrocarbons) are not very soluble in water (water-incompatible, water-hating, hydrophobic) 9  Water dissolves many salts/inorganic compounds High dielectric constant of H2O reduces attraction between oppositely charged ions in salt crystal; almost no attraction at large (> 40 nm) distances Strong electrostatic interactions between the solvated ions and water molecules (ion-dipole interactions) lower the energy of the system Entropy increases as ordered crystal lattice is dissolved 10  Solubility of polar and non-polar gases in water Non polar = charge Polar unequal charge - solubility of non-polar gasses in H2O increases with the decrease of temperature 11 Nonpolar compounds in water: the hydrophobic effect The hydrophobic effect refers to the tendency of non-polar compounds to aggregate in aqueous solution and exclude the water molecules Entropy decrease hydrophobic The system presents the smallest hydrophobic area to water/aqueous solvent Is one of the main factors behind: – protein folding – protein-protein association – formation of surfactant micelles and lipid membranes – binding of steroid hormones to their receptors Does not arise because of some attractive direct force between two nonpolar molecules 12 Low solubility of hydrophobic solutes can be explained by entropy Bulk water has little order, thus high entropy Water near a hydrophobic solute is forced to adopt a highly ordered state – low entropy Low entropy is not compensated by attractive forces between the non-polar molecules and water ones (of enthalpic nature), therefore the whole process is thermodynamically unfavorable (ΔG > 0), thus hydrophobic solutes have low solubility. l Little order High order ] 13  Origin of the Hydrophobic Effect Consider amphipathic fatty acids in water Fatty acid molecules disperse in the solution; nonpolar tail of each lipid molecule is surrounded by ordered water molecules Entropy of the system decreases System is now in an unfavorable state Nonpolar portions of the amphipathic molecule self- aggregate (self-assemble) so that fewer water molecules are ordered The released water molecules will be more random and the entropy increases All nonpolar groups are sequestered from water, and the released water molecules increase the entropy further Only polar “head groups” are exposed in these supra- molecular assemblies (micelle in this case) and make energetically favorable H-bonds with water molecules 14  Hydrophobic effect and the self-assembly of different amphiphiles into micelles, bilayers, liposomes - the molecular shape of amphiphilic molecules influences the self-assemblies formed: Fatty acids: Phospholipids liposomes: “artificial vesicles” (used for drug delivery e.g. liposomal doxorubicin Myocet®) - Note that the cell membrane is formed because of water environment in which phospholipids are placed; cells cannot exist in another solvent! 15  Hydrophobic effect favors amphipathic ligand binding Binding sites in enzymes and receptors are often hydrophobic, containing ordered water molecules of low entropy Such sites can bind hydrophobic substrates and ligands such as steroid hormones Many drugs are designed to take advantage of the hydrophobic effect Entropy more as binding Occurs, more free energy Delta G (-), Favorable Entropy is less Delta G (+) Unfavorable 17  Noncovalent Interactions in biomolecules Noncovalent interactions do not involve sharing a pair of electrons. Based on their physical origin, one can distinguish: H-bonds and Ionic (Coulombic) Interactions – Electrostatic interactions between permanently charged species, or between the ion and a permanent dipole Dipole Interactions – Electrostatic interactions between uncharged, but polar molecules van der Waals Interactions – Weak interactions between all atoms, regardless of polarity – Attractive (dispersion) and repulsive (steric) component Hydrophobic Effect – Complex phenomenon associated with the ordering of water molecules around nonpolar substances and entropy decrease in the system 18 Important in folding of proteins!  Effects of Solutes on Properties of Water Solutes of all kinds alter Colligative Properties of Water (colligative = tied together), such as boiling point, melting point, vapor pressure, and osmotic pressure (osmolarity). The effect of solutes on colligative properties of water does not depend on the (chemical) nature of the solute, just on the concentration (number of molecules of solute added to a given amount of water, because concentration of water is lower in solutions than in pure water e. g. The effect of adding 1 mol of glucose to 5 mols water in terms of CPW is identical to the addition of 1 mol of ethanol to 5 mols water, as we are adding the same amount of foreign molecules and water will have the same molar concentration in the final solution consequence: Cytoplasm of cells is a highly concentrated solution of different biomolecules, ions and has a high osmotic pressure (the more biomolecules we add the higher the osmotic pressure) Non-colligative Properties of Water: viscosity, surface tension, taste, and color, Depend on the chemical nature of the solute 20  Effects of Solutes on Colligative Properties of Water - Boiling point elevation and freezing point depression: 21  Effects of Solutes on Colligative Properties of Water: Osmosis and osmotic pressure - Osmotic pressure is caused by water passing from a region of high water concentration to a region of low water concentration; measured as the force to resist osmosis: (osmosis = water movement through a semipermeable membrane driven by differences in osmotic pressure (conc of water in two solutions separated by the membrane) 22  Effects of Solutes on Colligative Properties of Water: Osmosis and osmotic pressure Differences in concentration of solutes inside/outside of cell triggers water movement through membrane via osmosis Osmotic burst constitutes a big issue: plasma contains proteins, ions to be isotonic with the cells  hydration of a patient is done with isotonic saline solution, not with pure water! Osmotic pressure depends only on the number (concentration) of solutes, not on their MW; one large biopolymer elicits the same osmotic pressure as a small molecule, therefore large biopolymers (high MW proteins, polysaccharides, nucleic acids) can be stored in large amounts in cells with negligible contribution to total osmotic pressure of the cell; fuel is stored as polysaccharides, not as simple sugars! 23  Ionization of Water  H+ + OH- H2 O  O-H bonds are polar and can dissociate heterolytically Products are a proton (H+) and a hydroxide ion (OH–) Dissociation of water is a rapid reversible process Most water molecules remain un-ionized, thus pure water has very low electrical conductivity (resistance: 18 M cm) The equilibrium is strongly shifted to the left Extent of dissociation depends on the temperature 24  Proton Hydration Protons do not exist free in solution. They are immediately hydrated to form hydronium (oxonium) ions. A hydronium ion is a water molecule with a proton associated with one of the non- bonding electron pairs. Hydronium ions are solvated by nearby water molecules. The covalent and hydrogen bonds are interchangeable. This allows for an extremely fast mobility of protons in water via “proton hopping.” Protons move through water faster than other ions. This high ionic mobility helps to explain why acid- base reactions are so fast. 25  Ionization of Water: Quantitative Treatment Concentrations of participating species in an equilibrium process are not independent but are related via the equilibrium constant: [H + ] [OH-] H2O   H+ + OH- Keq = ———— [H2O ] Keq can be determined experimentally, it is 1.8 10–16 M at 25 C. [H 2O] can be determined from water density, it is 55.5 M. Ionic product of water: + - -14 2 K w = K eq ×[ H 2O] = [ H ][ OH ] = 1×10 M In pure water [H+] = [OH–] = 10–7 M 26  What is pH? pH is defined as the negative pH = -log[H+] logarithm of the hydrogen ion concentration Simplifies equations K w = [ H + ][ OH- ] = 1×10-14 M 2 The pH and pOH must always add to 14 - -log[H + ] - - log[OH ] = +14 In neutral solution, [H+] = [OH–] and the pH is 7 pH + pOH = 14 27 Henderson–Hasselbalch Equation: Derivation [ H + ][ A - ] HA   H+ + A- Ka = [ HA] + [ HA] [H ] = K a [A - ] [ HA] - log[H+] = -logK a -log [ A-] - [A ] pH = pK a + log [ HA] 28 pH scale is logarithmic: 1 unit = 10-fold 29  Dissociation of Weak Electrolytes: Principle O Keq O Weak electrolytes dissociate H3 C + H2O H3 C + H3O+ OH O- only partially in water. K a = K eq ×[ H 2O] Extent of dissociation is determined by the acid [ H + ][ CH3COO- ] dissociation constant Ka. Ka = = 1.74×10-5 M [ CH3COOH] We can calculate the pH if the + [ CH3COOH] Ka is known. But some [ H ] = Ka × [ CH3COO-] algebra is needed! 30  Dissociation of Weak Electrolytes: Example What is the final pH of a solution when 0.1 moles of acetic acid is added to water to a final volume of 1L? O O H3 C Ka We assume that the H3 C + H+ OH O- only source of H+ is 0.1 – x x x the weak acid [ x ][ x ] Ka = = 1.74×10-5 M To find the [H+], a [ 0.1- x] quadratic equation 2 -6 -5 x = 1.74×10 -1.74×10 x must be solved x 2 + 1.74×10-5 x -1.74×10-6 = 0 x = 0.001310, pH = 2.883 31 Dissociation of Weak Electrolytes: Simplification What is the final pH of a solution when 0.1 moles of acetic acid is added to water to a final volume of 1L? O O The equation can be Ka H3 C H3 C + H+ simplified if the amount of OH O- dissociated species is much 0.1 – x x x less than the amount of 0.1 x x undissociated acid [ x ][ x ] Approximation works for Ka = = 1.74×10-5 M [ 0.1] sufficiently weak acids and x 2 = 1.74×10-6 bases Check that x < [total acid] x = 0.00132, pH = 2.880 32  pKa measures acidity pK a = –log K a (strong acid  large Ka  small pKa) 33  Titration curves reveal the pKas of weak acids - Titration is used to determine the amount of acid present in solution, using a strong base (NaOH) of known concentration; pH is measured in parallel (via pH electrode, etc); pKa is determined [A - ] pH = pK a + log [ HA] Comparison of titration curves for 3 weak acids 34  Buffers are mixtures of weak acids and their anions (conjugate base) Buffers resist change in pH; At pH = pKa, there is a 50:50 mixture of acid and anion forms of the compound Buffering capacity of acid/anion system is the greatest at pH = pKa Buffering capacity is lost when the pH differs from pKa by more than 1 pH unit- Acetic Acid-Acetate as a Buffer System: ► Cells and organisms maintain a specific and constant cytosolic pH, usually near 7; phosphate -based buffers are the main ones used intracellularly, doubled by protein buffers: The particular R groups of important amino acids influence protein activity. They also contribute to the ability of proteins to act as buffers. 35 pH homeostasis in the extracellular space/blood, acidosis/alkalosis Normal blood pH is maintained between 7.35 and 7.45 through the bicarbonate buffer system The lungs regulate the amount of CO2 in the blood the kidneys regulate the bicarbonate. When the pH decreases to below 7.35 an acidotic condition is present. Acidosis means that the hydrogen ions are increased while pH and bicarbonate ions are decreased. A greater number of hydrogen ions are present in the blood than can be absorbed by the buffer systems; Alkalosis results when the pH is above 7.45 (e.g. in hyperventilation due to stress, anxiety). This condition results when the buffer base (bicarbonate ions) is greater than normal and the concentration of hydrogen ions are decreased; 36  Both acidosis and alkalosis can be of two different types: Respiratory: acidosis or alkalosis is caused by various malfunctions of the lungs including hypo or hyperventilation respectively. – Examples causing respiratory acidosis are emphysema, chronic bronchitis, asthma, severe pneumonia, cardiac arrest. – Alveolar hyperventilation resulting in respiratory alkalosis can be caused by anxiety, hysteria and stress, fever and large doses of aspirin, which stimulates the respiratory center in the brainstem, meningitis and pregnancy. Respiratory alkalosis is often produced iatrogenically during mechanical ventilation of patients Metabolic: acidosis or alkalosis is caused by various metabolic disorders which result in an excessive build up or loss of acids or bases. Metabolic acidosis can result in stimulation of chemoreceptors and so increase alveolar ventilation, leading to respiratory compensation, otherwise known as Kussmaul breathing (deep and labored breathing) a specific type of hyperventilation. This is common with diabetic ketoacidosis. Metabolic alkalosis most commonly occurs with profuse vomiting which depletes H+ ions leading to an increase of bicarbonate in the body. Compensatory Mechanisms: - Respiratory acidosis is compensated by the kidney, in effect causing a metabolic alkalosis and also through an increase in breathing frequency, depth - Metabolic alkalosis is compensated by decrease breathing rate, in effect causing respiratory acidosis. Renal compensation cause an increase in bicarbonate excretion. Respiratory alkalosis can be compensated by temporary breathing in a paper bag. 37  pH, drug absorbtion, enzymatic activity - pH inside the GI tract changes from 1.5 (stomach) to 7.6 (small intestine); ionizable drugs will start to be absorbed at different locations in the GI tract: - Different enzymes work at different optimum pH for enzymatic activity; food is processed in different locations of the GI tract: 38  Goals and Objectives Upon completion of this lecture at minimum you should be able to answer the following: ► Structure and of water and its influences on life on Earth, H-bonding in solid and liquid water, consequences; ► Properties of water, H-bond formation with polar solutes, polarity of biomacromolecules ► Dissolution of inorganic (ionic) compounds in water, solubility of gases in water, hydrophobic solutes in water and the hydrophobic effect, origins and consequences of the hydrophobic effect, non-covalent interactions in biomolecules ► Colligative and non-colligative properties of water, osmotic pressure impact on cells, isotonic, hypertonic, hypotonic states ► Ionization of water, proton hydration and transport, ionic product of water, pH, Henderson–Hasselbalch Equation, pH scale, dissociation of weak electrolytes in water with examples, acidity, weak acids, titration curves, buffers and their impact in cells and organisms, acidosis/alkalosis, types, mechanisms of compensation ► Drug ionization and pH dependence of drug absorption, optimum pH for enzyme activity 39