Water and Acid-Base Equilibria PDF

For Your Eyes Only WATER and ACID-BASE EQUILIBRIA WATER A. Introduction - in living cells, most chemical reactions occur in an aqueous environment (60% of the mass of a living cell is water) 1. Water is Essential to Biochemistry...

For Your Eyes Only WATER and ACID-BASE EQUILIBRIA WATER A. Introduction - in living cells, most chemical reactions occur in an aqueous environment (60% of the mass of a living cell is water) 1. Water is Essential to Biochemistry a. Biological macromolecules assume specific shapes in response to the chemical and physical properties of water b. Biological molecules undergo chemical reactions in an aqueous environment c. Water is a key reactant in many reactions, usually in the form of H + and OH– ions d. Water is oxidized in photosynthesis to produce molecular oxygen, O2, as part of the process that converts the sun’s energy into usable chemical form - expenditure of that energy under aerobic conditions leads to the reduction of O 2 back to water e. Bathes the cells f. Dissolves and transports compounds in the blood g. Provides a medium for movement of molecules into and throughout cellular compartments h. Separates charged molecules i. Dissipates heat B. Dipole Moment 1. Water Molecule - uncharged - forms hydrogen bonds with other polar molecules a. Oxygen Atom - has two unshared electrons that form an electron dense cloud around it - the cloud lies above and below the plane formed by the water molecule b. Covalent Bond Formed Between the Hydrogen and Oxygen Atoms - the shared electrons are attracted toward the oxygen atom (uneven distribution of electrons), thus giving the oxygen atom a partial negative charge and the hydrogen atom a partial positive charge  oxygen side of the molecule is much more electronegative than the hydrogen side - near the oxygen atom is slightly negative - near the hydrogen atom is slightly positive  the molecule is dipolar and is said to have dipole moment 2. Function - measures the degree of polarity - increases as the difference in electronegativity increases C. Hydrogen Bonds - electrostatic interactions between hydrogen atom of one molecule and the more electronegative atom of an acceptor molecule (ex: oxygen, nitrogen) 1. Water is Strongly Hydrogen Bonded - each water molecule participating in four hydrogen bonds with its neighbors - two in which it donates - two in which it accepts  For Your Eyes Only 2. Strongly Hydrogen Bonded Character of Water - responsible for many of its characteristic properties a. High Heat of Fusion - allows water to act as a heat sink (greater heat loss is necessary for the freezing of water compared to other substances of similar molecular mass) b. High Heat of Vaporization - relatively more heat must be input to vaporize water compared to other substances of similar molecular mass c. Ability to Dissolve Most Polar Compounds - water is an excellent solvent of polar and ionic substances due to its property of surrounding polar molecules and ions with oriented shells of water, thereby attenuating the electrostatic forces between these molecules and ions i. Hydrophobic Effect - tendency of water to minimize its contact with nonpolar (hydrophobic) molecules - largely driven by the increase in entropy caused by the necessity for water to order itself around nonpolar molecules  nonpolar molecules to aggregate reducing the surface area that water must order itself about d. Open Structure - makes ice less dense than liquid water  make ice float  insulate the water beneath it  inhibit total freezing of large bodies of water D. Dissociation of Water 1. Ionization of Water - water molecules have limited tendency to reversibly dissociate (ionize)  hydrogen ion or proton (H+) and a hydroxyl ion (OH-) - concentration of hydrogen ions determines the acidity of the solution, which is expressed in terms of pH 2. Alternative Description to Above Reaction - the proton is never free and binds to a water molecule to form H 3O+ (hydronium ion) 2H2O  H3O+ + OH-  For Your Eyes Only 3. pH of Water a. Dissociation of Water Molecules - into H+ and OH- - very slight b. Hydrogen Ion Concentration of Pure Water - only 0.0000001 M, or 10-7 mol/L - denoted by the term pH  pH of pure water is 7 4. Tendency of Water to Dissociate K = 1.8 x 10-16 mol/L at 25oC 5. Dissociation Constant for Water (Kd) - expresses the relationship between the hydrogen ion concentration [H +], the hydroxide ion concentration [OH-], and the concentration of water [H2O] at equilibrium - water dissociates to such a small extent  [H2O] is essentially constant at 55.5 M 6. Ion Product of Water (Kw) - multiplication of the Kd for water (approximately 1.8 x 10-16 M) by 55.5 M Kw = [H+] [OH-] = 1 x 10-14 mol2/ L2 - ion product of water, [H+][OH-], is constant for all aqueous solutions, even those that contain dissolved acids (proton donors) or dissolved bases (proton acceptor) - if H+ (proton) is added to pure water  [OH-] decreases in order that the product [H+][OH-] will remain constant - if OH is added  [H+] will have to decrease - E. Water as a Solvent 1. Water as Biologic Medium - intracellularly and extracellularly - attributed to - ability to form hydrogen bonds - dipole moment - acts as a solvent for substances the body need - K+ - glucose - adenosine triphosphate (ATP) - proteins 2. Any molecule capable of forming hydrogen bonds can do so with water  dissolve readily in water a. Polar Organic Molecules and Inorganic Salts - can readily dissolve in water because water also forms hydrogen bonds and electrostatic interactions with these molecules b. Organic Molecules containing a high proportion of electronegative atoms (generally oxygen or nitrogen) - soluble in water because these atoms participate in hydrogen bonding with water molecules  For Your Eyes Only c. Chloride, Bicarbonate, and Other Anions - are surrounded by a hydration shell of water molecules arranged with their hydrogen atoms closest to the anion d. Inorganic Cations (Na+, K+) - oxygen atom of water molecules interacts with inorganic cations to surround them with a hydration shell 3. Dipoles - interact with positively and negatively charged ions F. Fluid Compartments in the Body a. Total Body Water - is roughly - 50 to 60% of body weight in adults - 75% of body weight in children b. Body Water Differences - obese people tend to have a lower percentage of body water than thin people - women tend to have a lower percentage than men - older people have a lower percentage than younger people 1. Intracellular Compartment - 60% of the total body water  For Your Eyes Only 2. Extracellular Compartments - 40% of the total body water - includes: i. Interstitial Fluids - fluid in the tissue spaces, lying between cells ii. Blood Plasma iii. Lymph iv. Transcellular water - small, specialized portion of extracellular water that includes - gastrointestinal secretions - urine - sweat - fluid that has leaked through capillary walls because of such processes as increased hydrostatic pressure or inflammation G. Colligative Properties of Solutions - effects of solutes on the solvent - decrease of freezing point - increase of boiling point - decrease of vapour pressure - endow properties of osmotic pressure H. Water and Thermal Regulation 1. Structure of Water - allows it to resist temperature change 2. Heat of Fusion - high - large drop in temperature is needed to convert liquid water to the solid state of ice 3. Thermal Conductivity of Water - high - facilitates heat dissipation from high energy-using areas such as the brain into the blood and the total body water pool 4. Heat Capacity and Heat of Vaporization - remarkably high - liquid water is converted to a gas and evaporates from the skin  cooling effect 5. Responses a. To the Input of Heat - decreasing the extent of hydrogen bonding between water molecules b. To Cooling - increasing the bonding between water molecules  For Your Eyes Only I. Electrolytes J. Osmolality and Water Movement 1. Water Distributes Between the Different Fluid Compartments - according to the concentration of solutes, or osmolality, of each compartment 2. Fluid Osmolality - proportionate to the total concentration of all dissolved molecules - ions - organic metabolites - proteins - expressed as milliosmoles (mOsm)/kg water 3. Semipermeable Cellular Membrane - separates the extracellular and intracellular compartments - contains a number of ion channels through which water can freely move, but other molecules cannot - water can freely move through the capillaries separating the interstitial fluid and the plasma  water will move from a compartment with a low concentration of solutes (lower osmolality) to one with a higher concentration to achieve an equal osmolality on both sides of the membrane 4. Osmotic Pressure - the force it would take to keep the same amount of water on both sides of the membrane - as water is passed from the blood into the urine to balance the excretion of ions, the blood volume is repleted with water from interstitial fluid - when the osmolality of the blood and interstitial fluid is too high, water moves out of the cell - the loss of cellular water also can occur in hyperglycemia, because the high concentration of glucose increases the osmolality of the blood K. Biomedical Importance of Water 1. Functions of Water a. Principal End-Product of Oxidative Metabolism of Food b. Excellent Nucleophile c. Homeostasis - many of the compounds produced in the body and dissolved in water contain chemical groups that act as acids or bases, releasing or accepting hydrogen ions - hydrogen ion content and the amount of body water are controlled to maintain a constant environment d. Transport of Molecules and Heat 2. Hypothalamus - regulation of water balance 3. States of Water Imbalance - accompanied by Na+ imbalance a. Water Depletion - decreased intake - increased loss  For Your Eyes Only b. Water Excess - increased intake - decreased loss 4. ECF Water and Osmotic Homeostasis a. Factors - osmotic mechanism - non-osmotic mechanisms b. Responses - water conservation by ADH - water acquisition by drinking - > 2% increase of ECF osmolality  thirst 5. Nephrogenic Diabetes Insipidus - inability of renal tubular ADH osmoreceptors to respond to ADH - characteristics - extreme thirst - high water intake - inability to concentrate urine - inability to respond to subtle ECF osmolality changes ACID-BASE EQUILIBRIUM A. Dissociation of a Weak Acid 1. Strong and Weak Electrolytes a. Strong Electrolytes - considered to be fully dissociated in aqueous solutions b. Weak Electrolytes - only partially dissociated (usually 7.0 3. Buffering (refer to titration curve) a. Buffer - mixture of an undissociated acid and its conjugate base (equal concentrations of weak acid and its conjugate base) - causes a solution to resist changes in pH when either H + or OH- is added - if the amount of HA and A- are equal, pH = pKa - if acid is added, A- neutralizes it forming HA - if base is added, HA neutralizes it, in the process being converted to A- b. Behaviour of pH Around Midpoint i. near 50% neutralization, pH changes very slowly (solution is exhibiting buffering) ii. buffering effect is maximal at the midpoint of titration (pH is at the pKa of the acid- conjugate base pair) - a buffer can only compensate for an influx or removal of hydrogen ions within approximately 1 pH unit of its pKa - as the pH of a buffered solution changes from the pKa to one pH unit below the pKa ,the ratio of [A-] to [HA] changes from 1:1 to 1:10 - if more hydrogen ions were added, the pH would fall rapidly because relatively little conjugate base remains - at 1 pH unit above the pKa of a buffer, relatively little undissociated acid remains - more concentrated buffers are more effective simply because they contain a greater total number of buffer molecules per unit volume that can dissociate or recombine with hydrogen ions  For Your Eyes Only c. The Extent to which a Solution Containing a Weak Acid and its Salt can Resist pH Changes Depends on i. Buffer Capacity - the closeness of the buffer pH to pKa - a conjugate acid/base pair can serve as an effective buffer when the pH of the solution is within approximately +/- 1 pH unit of the pKa of the HA ii. Concentration of Buffer Solution REGULATION of BLOOD pH A. Metabolic Acids and Bases 1. Normal Metabolism - generates - CO2 - metabolic acids (lactic acid, ketone bodies) - inorganic acids (sulfuric acid)  increase the hydrogen ion concentration of the blood or other body fluids and tend to lower the pH - average rate of metabolic activity produces roughly 22,000 mEq acid per day - if dissolved at one time in unbuffered body fluids  pH 7.4  alkalosis 2. Change in CO2 levels  respiratory acidosis or alkalosis 3. Change in HCO3-  metabolic acidosis or alkalosis  For Your Eyes Only G. Intracellular pH 1. Major Buffers - phosphate anions - proteins - maintaining a constant pH of intracellular fluids 2. Inorganic Phosphate Anion (H2PO4-) - dissociates to  H+ conjugate base, HPO4-2 - pKa of 7.2 3. Organic Phosphate Anions - ex. glucose 6-phosphate ATP - also act as buffers 4. Intracellular Proteins - contain histidine and other amino acids that can accept protons H. Urinary Hydrogen, Ammonium, and Phosphate Ions 1. Nonvolatile Acids - excreted in the urine - most of the nonvolatile acid hydrogen ion is excreted as undissociated acid that generally buffers the urinary pH between 5.5 and 7.0 (pH of 5.0 is the minimum urinary pH) a. Phosphate and Ammonium Ions b. Uric Acid c. Dicarboxylic Acids d. Tricarboxylic acids - ex: citric acid 2. Sulfuric Acid (H2SO4) - one of the major sources of nonvolatile acid in the body - generated from - sulfate-containing compounds ingested in foods - metabolism of the sulfur-containing amino acids - cysteine - methionine - strong acid - dissociated in the blood and urine into - H+ - sulfate anion (SO4-2) in the blood and urine - urinary excretion helps to remove acid 3. Phosphate Ions - H2PO4- - HPO4-2 - to maintain metabolic homeostasis phosphate excretion (in the urine) = phosphate ingestion - phosphate ingested with food - phosphate anions - organic phosphates (phospholipid)  For Your Eyes Only 4. Ammonium Ions - major contributors to buffering urinary pH, but not blood pH a. Ammonia (NH3) - base that combines with protons to produce ammonium (NH4+) ions NH3 + H+  NH4+ - reaction occurs at pKa 9.25 - produced from amino acid catabolism or absorbed through the intestine - kept at very low concentrations in the blood because it is toxic to neural tissues - kidney cells generate NH4+ and excrete it into the urine in proportion to the acidity (proton concentration) of the blood - as the renal tubular cells transport H+ into the urine, they return bicarbonate anions to the blood I. Hydrochloric Acid - gastric acid - secreted by parietal cells of the stomach into the stomach lumen  strong acidity denatures ingested proteins so they can be degraded by digestive enzymes - neutralized by bicarbonate secreted from pancreatic cells and by cells in the intestinal lining CASE (EMPHYSEMA) A. Patient History A patient with a history of chronic lung disease has suffered from emphysema, which has grown progressively worse over a period of years. The patient experiences chronic shortness of breath. Analysis of the patient’s blood reveals the following: pCO2 = 60 mm Hg [HCO3-] = 34 mM pH = 7.38  For Your Eyes Only B. Diagnosis Because of the chronic lung disease  rate of respiration is depressed, CO2 is not being eliminated from the blood by the lungs  increased pCO2  respiratory acidosis. To compensate for the primary disorder, the kidneys  increase H+ secretion, increased HCO3- reabsorption  increased blood HCO3-  compensatory metabolic alkalosis C. Discussion Chronic condition  kidneys had time to compensate. Blood is slightly below normal (body has not fully compensated for the respiratory acidosis) OTHER APPLICATIONS of the HENDERSON-HASSELBALCH EQUATION A. Preparation of Buffer Solutions of Known pH B. Calculate the Degree of Dissociation of Weak Acid-Conjugate Base Pairs C. Calculate How the pH of Physiologic Solutions (Blood or Intracellular Fluid) Responds to Changes in the Concentration of Weak Acid and/or its Corresponding Salt Form 1. Bicarbonate Buffer System - Henderson-Hasselbalch equation predicts how shifts in [HCO3-] and pCO2 influence pH D. Calculate the Abundance of Ionic Forms of Acidic and Basic Drugs - most drugs are either weak acids or weak bases - acidic drugs (HA) release a proton (H+) causing a charged anion (A-) HA  H+ + A- - weak bases (BA+) can also release H+ producing the uncharged base (B) BH+  B + H+ - drug passes through membranes more readily if it is uncharged - for weak acid - uncharged HA permeates through membranes, A- cannot - for weak base - uncharged form (B) permeates the membrane, BH+ cannot - the effective concentration of the permeable form of each drug at its absorption site is determined by relative concentrations of charged and uncharged forms - the ratio of the 2 forms is determined by the - pH at the site of absorption - strength of the weak or base, which is represented by the pKa of the ionisable group - the Henderson-Hasselbalch equation is useful in determining how much drug is found on either side of a membrane that separates 2 compartments that differ in pH (ex: stomach with pH 1.0-1.5 and blood plasma with pH 7.4 For Your Eyes Only