Chapter 2: Water: The Medium of Life PDF

Chapter 2 Water: The Medium of Life  Outline 2.1 What are the properties of water? 2.2 What is pH? 2.3 What are buffers, and what do they do? 2.4 What properties of water give it a unique role in the environment?  Water: The Medium of Life ▪ Life originated, evolved, and thrives in t...

Chapter 2 Water: The Medium of Life  Outline 2.1 What are the properties of water? 2.2 What is pH? 2.3 What are buffers, and what do they do? 2.4 What properties of water give it a unique role in the environment?  Water: The Medium of Life ▪ Life originated, evolved, and thrives in the seas.. Image source: https://kidspressmagazine.com/science-for-kids/misc/misc/evolution-land-animals.html  Water: The Medium of Life ▪ Water and its ionization products, hydrogen ions and hydroxide ions, are critical determinants of the structure and function of many biomolecules, including amino acids and proteins, nucleotides and nucleic acids, and even phospholipids and membranes. Water: The Medium of Life  Water: The Medium of Life ▪ A difference in the concentration of hydrogen ions on opposite sides of a membrane represents an energized condition essential to biological mechanisms of energy transformation. Source: https://www.thoughtco.com/electron-transport-chain-and-energy-production-4136143  2.1 What Are the Properties of Water? ▪ Water has a substantially higher boiling point, melting point, heat of vaporization, and surface tension (anomalously high for a substance of this molecular weight that is neither metallic nor ionic); intermolecular forces of attraction between H2O molecules are high.  2.1 What Are the Properties of Water? ▪ its maximum density is found in the liquid (not the solid) state, and it has a negative volume of melting (that is, the solid form, ice, occupies more space than does the liquid form, water) 2.1 What Are the Properties of Water? Permanent dipoles - occur when two atoms in a molecule have substantially different electronegativity: One atom attracts electrons more than another, becoming more negative, while the other atom becomes more positive. 2.1 What Are the Properties of Water? https://d2gne97vdumgn3.cloudfront.net/api/file/nq42iyAQduHiTj2FmCty http://cdni.wired.co.uk/620x413/g_j/htwoo.jpg ▪ Hydrogen bonding in water is key to its properties. ▪ The solvent properties of water derive from its polar nature 2.1 What Are the Properties of Water? http://www.scottsmithonline.com/interests/medicalschool/biology/110a/Midterm1Materials/Notes/graphics/figure%2002-15.jpg H-bonding is cooperative: H2O molecule serving as an H-bond donor becomes a better H-bond acceptor 2.1 What Are the Properties of Water? http://www.lancoude.com/upload/3/11/31165a8c0882b284f9f836d9545a083d_thumb.gif Hydrogen bonding in ice.  2.1 What Are the Properties of Water? ▪ THE SOLVENT PROPERTIES OF WATER DERIVE FROM ITS POLAR NATURE Because of its highly polar nature, water is an excellent solvent for various compounds http://www.mun.ca/biology/desmid/brian/BIOL2060/BIOL2060-02/02_10.jpg “hydration shells” surrounding ions  2.1 What Are the Properties of Water? (1) Water has a high dielectric constant. Water’s ability to surround ions in dipole interactions and diminish their attraction for each other is a measure of its dielectric constant, D. The attractions between the water molecules interacting with, or hydrating, ions are much greater than the tendency of oppositely charged ions to attract one another. 2.1 What Are the Properties of Water? (2) Water Forms H Bonds with Polar Solutes excellent solvent properties of water stem from its ability to readily form hydrogen bonds with the polar functional groups on these compounds, such as hydroxyls (-OH), amines (-NH2) , and carbonyls (-C=O) 2.1 What Are the Properties of Water? (3) Hydrophobic Interactions - apparent affinity of nonpolar structures for one another Because nonpolar solutes must occupy space, the random H-bonded network of water must reorganize to accommodate them. The water molecules participate in as many H-bonded interactions with one another as the temperature permits. Consequently, the H-bonded water network rearranges toward formation of a local cagelike (clathrate) structure surrounding each insoluble solute molecule. More H2O molecules are ordered Less H2O molecules are ordered (lower entropy/disorder) (higher entropy/disorder) In actuality, the “attraction” between nonpolar solutes is an entropy-driven process due to a net decrease in order (net increase in disorder) 2.1 What Are the Properties of Water? (4) Interaction with amphiphilic molecules (Compounds containing both strongly polar and strongly nonpolar groups) micelle 2.1 What Are the Properties of Water? Colligative Properties ▪ The presence of dissolved substances disturbs the structure of liquid water, thereby changing its properties. ▪ The net effect is that solutes, regardless of whether they are polar or nonpolar, fix nearby water molecules in a more ordered array, creating local order. ▪ This influence of the solute on water is reflected in a set of characteristic changes in behavior termed colligative properties, or properties related by a common principle.  2.1 What Are the Properties of Water? Colligative Properties ▪ Freezing point depression ▪ boiling point elevation ▪ vapor pressure lowering ▪ Osmotic pressure effects In effect, by imposing local order on the water molecules, solutes make it more difficult for water to assume its crystalline lattice (freeze) or escape into the atmosphere (boil or vaporize) 2.1 What Are the Properties of Water? ▪ Freezing point depression spraying rock salt or a solution of salt water to prevent ice formation.  2.1 What Are the Properties of Water? ▪ Vapor pressure lowering The vapor pressure of water is the pressure at which the water will transition from a liquid to a gas (vapor). Specifically, the vapor pressure is the point at which the water is in a state of equilibrium, with the same number of water molecules transitioning from liquid to gas and from gas to liquid. 2.1 What Are the Properties of Water? ▪ Vapor pressure lowering The pressure necessary to push water back through the membrane at a rate exactly equaled by the water influx is the osmotic pressure of the solution To minimize the osmotic pressure created by the contents of their cytosol, cells tend to store substances such as amino acids and sugars in polymeric form. For example, a molecule of glycogen or starch containing 1000 glucose units exerts only 1/1000 the osmotic pressure that 1000 free glucose molecules would. -- 1st lec Auto-ionization of Water Source: https://commons.wikimedia.org/wiki/File:Autoionizacion-agua.gif Source:https://chem.libretexts.org/Ancillary_Materials/Exempla rs_and_Case_Studies/Exemplars/Foods/Water_Ionization Equilibrium and Le Chatelier’s Principle Equilibrium and Le Châtelier’s Principle Let us look at the reaction of the weak acid HF and H2O as it proceeds to equilibrium. Initially, only the reactants HF and H2O are present. As F- and H3O+ products build up, the rate of the reverse reaction increases, while the rate of the forward reaction decreases. Equilibrium and Le Châtelier’s Principle Equilibrium and Le Châtelier’s Principle When we alter the concentration of a reactant or product of a system at equilibrium, the rates of the forward and reverse reactions will no longer be equal. We say that a stress is placed on the equilibrium. Le Châtelier’s principle states that when equilibrium is disturbed, the rates of the forward and reverse reactions change to relieve that stress and reestablish equilibrium. Equilibrium and Le Châtelier’s Principle R1 + R2 → P1 + P2 Oxygen–Hemoglobin Equilibrium and Hypoxia Oxygen–Hemoglobin Equilibrium and Hypoxia The transport of oxygen involves an equilibrium between hemoglobin (Hb), oxygen, and oxyhemoglobin (HbO2). When the O2 level is high in the alveoli of the lung, the reaction shifts in the direction of the product HbO2. In the tissues where O2 concentration is low, the reverse reaction releases the oxygen from the hemoglobin. At an altitude of 18 000 ft, a person will obtain 29% less oxygen. When oxygen levels are lowered, a person may experience hypoxia, characterized by increased respiratory rate, headache, decreased mental acuteness, fatigue, decreased physical coordination, nausea, vomiting, and cyanosis (skin, lips or nails turn blue due to a lack of oxygen in the blood). Immediate treatment of altitude sickness includes hydration, rest, and, if necessary, descending to a lower altitude. The adaptation to lowered oxygen levels requires about 10 days. During this time, the bone marrow increases red blood cell production, providing more red blood cells and more hemoglobin. A person living at a high altitude can have 50% more red blood cells than someone at sea level.  H2O (l) + H2O (l) H3O+ (aq) + OH- (aq) Although the equilibrium lies far to the left it is very important to take into consideration, especially for living systems. For pure water [OH– ] = [H+] 2.2 What is pH? We define an aqueous solution as being neutral when the [H+] = [OH-] acidic when [H+] > [OH-] basic when [H+] < [OH-] CHECK THIS OUT ! [H+] = 0.0000001 = 10-7 How can this be abbreviated further? By just describing the power called the POWER OF H pH = –log [H+] = -log (10-7) = 7 Kw is called ionization constant of water and is very small. As with all Kw values, it is temperature dependent. Kw = 1.0 x 10-14 @ 25oC Kw = [H+][OH-] Since [OH– ] = [H+] for pure water, then [H+][OH-] or x2 is: This only means that the neutral value for pH is getting lower, it does not mean that the solution is Kw = x = 1 x 10-7 becoming more acidic as the temperature increase. Kw =(1 x 10-7)(1 x10-7)  Kw = [H+][OH-] or 1.0 x 10-14 = (1.0 x 10-7)(1.0 x10-7) Get the –log of both sides: –log Kw = –log [H+] + –log [OH̶ ] 14.00 = 7.00 + 7.00 pKw = pH + pOH Thus: 14 = pH + pOH Acids and Bases In the early days of chemistry chemists were organizing physical and chemical properties of substances. They discovered that many substances could be placed in two different property categories: Substance A Substance B 1. Sour taste 1. Bitter taste 2. Reacts with carbonates to make CO2 2. Reacts with fats to make soaps 3. Reacts with metals to produce H2 3. Do not react with metals 4. Turns blue litmus red 4. Turns red litmus blue 5. Reacts with B substances to make 5. Reacts with A substances make salt and water salt and water  Arrhenius Definition The Swedish chemist Svante Arrhenius proposed the first definition of acids and bases (Substances A and B became known as acids and bases, respectively.) “Acids are substances that dissociate in water to produce H+ ions and bases are substances that dissociate in water to produce OH- ions” NaOH (aq) Na+ (aq) + OH̶ (aq) Base HCl (aq) H+ (aq) + Cl ̶ (aq) Acid  Arrhenius Definition But what if the acid/base is not dissolved in water? The Arrhenius definition for acids and bases only refers to compounds dissolved in water. Does this mean that acids and bases cannot exist out of water? Not quite, that’s where the Brønsted-Lowry definition comes in.  Brønsted-Lowry Definition Johannes Brønsted and Thomas Lowry revised Arrhenius’s acid-base theory to include other solvents besides water. They defined acids and bases as follows: “An acid is a hydrogen containing species that donates a proton. A base is any substance that accepts a proton.” HCl (aq) + H2O (l) Cl ̶ (aq) + H3O+ (aq) In the above example, what is the Brønsted-Lowry acid? what is the Brønsted-Lowry base?  Brønsted-Lowry Definition Johannes Brønsted and Thomas Lowry revised Arrhenius’s acid-base theory to include other solvents besides water. They defined acids and bases as follows: “An acid is a hydrogen containing species that donates a proton. A base is any substance that accepts a proton.” HCl (aq) + H2O (l) Cl ̶ (aq) + H3O+ (aq) Acid Base Acidic Solution (H+ donor) (H+ acceptor) In the above example, what is the Brønsted-Lowry acid? what is the Brønsted-Lowry base?  Brønsted-Lowry Definition NH3 + H2O NH4+ + OH ̶ Base Acid Basic Solution (H+ acceptor) (H+ donor) Conjugate Pairs In reality, the reaction of HCl with H2O is an equilibrium and occurs in both directions, although in this case the equilibrium lies far to the right. HCl (aq) + H2O (l) Cl ̶ (aq) + H3O+ (aq) Acid Base conjugate base conjugate acid For the reverse reaction Cl ̶ behaves as a Brønsted base and H3O+ behaves as a Brønsted acid. The Cl ̶ is called the conjugate base of HCl. Brønsted-Lowry acids and bases always exist as conjugate acid-base pairs. Their formulas differ by only one proton. Give it a Try! Label the acid, base, conjugate acid, and conjugate base in each reaction: HCl + OH ̶ → Cl ̶ + H2O H2O + H2SO4 → HSO4̶ + H3O+ Strong acids ionize 100% and weak ones do not. A single arrow is used to represent the ionization of strong acids. HCl (g) H+ (aq) + Cl - (aq) double arrows are used to represent ionization of weak acids because an equilibrium is created. HF (g) H+ (aq) + F – (aq) Common Strong Acids and Bases  Strong acids completely dissociate into their ions when they are mixed with water. As the strong acids become more concentrated, they may be unable to fully dissociate. The rule of thumb is that a strong acid is 100% dissociated in solutions of 1.0 M or less.  Strong bases are bases which completely dissociate in water into the cation and OH- (hydroxide ion). The hydroxides of the Group I and Group II metals usually are considered to be strong bases. pH Calculations To calculate pH or pOH, recall that: pH = -log [H3O+], or pOH = -log [OH-] pH + pOH = 14 for aqueous solutions Find the pH of these: 1) 0.15 M solution of Hydrochloric acid 2) 3.00 X 10-7 M solution of Nitric acid Sample Problem: Calculating pH from [H3O+] Aspirin, which is acetylsalicylic acid, was the first nonsteroidal anti-inflammatory drug (NSAID) used to alleviate pain and fever. If a solution of aspirin has a [H3O+] = 1.7 x 10-3 M, what is the pH of the solution? = 2.769551079 pH = 2.77 Calculation of molarity using pH or pOH Molarity of H+ or [H+] = 10–pH Molarity of OH- or [OH-] = 10–pOH The number of decimal places in the log answer is equal to the number of significant figures in the molarity unit Find the Molarity: A solution has a pH of 8.5. What is the Molarity of hydrogen ions in the solution? You only need one piece of information to be able to determine all the values associated with acids and base. Sample Problem: What are the [H3O+] and [OH-] of Diet Coke that has a pH of 3.17? Solution: [H3O+] = 10-3.17 = 6.8 x 10-4 M [OH-] = Kw/[H3O+] = 1.0 x 10-14 M2/6.8 x 10-4 M = 1.5 x 10-11 M Or [OH-] = 10-pOH pOH = pKw – pH = 14.00 – 3.17 = 10.83 [OH-] = 10-10.83 = 1.5 x 10-11 M Reactions of Acids and Bases Acids and Metals Acids React with Carbonates and Bicarbonates https://edu.rsc.org/magnificent-molecules/hydrochloric-acid/3010539.article Reactions of Acids and Bases Acids and Hydroxides: Neutralization 2.3 What are buffers and what do they do? The lungs and the kidneys are the primary organs that regulate the pH of body fluids, including blood and urine. Major changes in the pH of the body fluids can severely affect biological activities within the cells. Buffers are present to prevent large fluctuations in pH. Recall: Strong/Weak Acid A strong acid is one where complete dissociation of the compound occurs. Hydrochloric acid and sulfuric acid are strong acids. A weak acid is one where incomplete dissociation of the compound occurs. Carbonic acid and acetic acid are weak acids. Weak acids dissociate poorly in water; they release protons, but only a small fraction of their molecules dissociate (ionize). 2.3 What are buffers and what do they do?  2.3 What are buffers and what do they do? In a buffer, an acid must be present to react with any OH- that is added, and a base must be available to react with any added H3O+. However, that acid and base must not neutralize each other. Therefore, a combination of an acid–base conjugate pair is used in buffers. A small amount of H3O+ added to the buffer reacts with C2H3O2-, whereas a small amount of OH- added to the buffer neutralizes HC2H3O2. The pH of the solution is maintained as long as the added amounts of acid or base are small compared to the concentrations of the buffer components. Buffering of a Strong Acid Buffering of a Weaker Acid Buffering of a Strong Base Buffer Capacity and Buffer Range Buffer capacity is the ability to resist pH change. The more concentrated the components of a buffer, the greater the buffer capacity. The pH of a buffer is distinct from its buffer capacity. A buffer made of equal volumes of 1.0 M CH3COOH and 1.0 M CH3COO- has the same pH (4.74) as a buffer made of equal volumes of 0.1 M CH3COOH and 0.1 M CH3COO-, but the more concentrated buffer has a much larger capacity for resisting a pH change. 2.3 What are buffers and what do they do? CO2 is continually produced as an end product of cellular metabolism. Some CO2 is carried to the lungs for elimination, and the rest dissolves in body fluids such as plasma and saliva, forming carbonic acid H2CO3. As a weak acid, carbonic acid dissociates to give bicarbonate, HCO3-, and H3O+. More of the anion HCO3- is supplied by the kidneys to give an important buffer system in the body fluid—the H2CO3/HCO3- buffer. 2.3 What are buffers and what do they do? 2.3 What are buffers and what do they do? If the CO2 level rises, increasing [H2CO3], the equilibrium shifts to produce more H3O+, which lowers the pH. This condition is called acidosis. A lowering of the CO2 level leads to a high blood pH, a condition called alkalosis. Excitement, trauma, or a high temperature may cause a person to hyperventilate, which expels large amounts of CO2. The kidneys also regulate H3O+ and HCO3-, but they do so more slowly than the adjustment made by the lungs through ventilation.  A. CHEMICAL BUFFER SYSTEMS 1. Bicarbonate Buffer Maintain a 20:1 ratio : HCO3- : H2CO3 (bicarbonate to carbonic acid); H2O + CO2 * H CO HCO3- + H+ 2 3 gas aqueous Lungs eliminate carbon dioxide Kidneys can remove excess non- gas acids and bases HCl + NaHCO3 H2CO3 + NaCl NaOH + H2CO3 NaHCO3 + H2O * Catalyzed by the enzyme carbonic anhydrase (enzyme found in red blood cells, gastric mucosa, pancreatic cells, and renal tubules)  A. CHEMICAL BUFFER SYSTEMS 1. Bicarbonate Buffer – CO2 + H2O → H2CO3 → HCO3- + H+ lowers pH by releasing H+ – CO2 + H2O ← H2CO3 ← HCO3- + H+ raises pH by binding H+ Functions with respiratory and urinary systems – to lower pH, kidneys excrete HCO3- – to raise pH, kidneys excrete H+ and lungs excrete CO2  Limitations of the Carbonic Acid Buffer System Functions only when respiratory system and respiratory control centers are working normally Ability to buffer acids is limited by availability of bicarbonate ions. The addition of a given amount of acid or base to the blood tends to change its H+ concentration considerably. The pKa of carbonic acid is 6.35 while the pH of blood is 7.4, resulting to a weak buffering power.  A. CHEMICAL BUFFER SYSTEMS 2. Phosphate Buffer The phosphate buffer is very effective but not found in high concentrations in extracellular fluid. H2PO4- HPO42- + H+ H2PO4- + OH- HPO42- + H2O H2PO4- HPO42- + H+  A. CHEMICAL BUFFER SYSTEMS 2. Phosphate Buffer Important in the intracellular fluid (ICF) and renal tubules where phosphates are more concentrated and function closer to their optimum pH of 6.8 – constant production of metabolic acids creates pH values from 4.5 to 7.4 in the ICF, avg. 7.0  A. CHEMICAL BUFFER SYSTEMS 3. Protein Buffers a. Amino Acids - free and terminal amino acids – Respond to pH changes by accepting or releasing H+  If acid comes into blood, hydronium ions can be neutralized by the –COO- groups –COO- + H3O+ → - COOH + H2O If base is added, it can be neutralized by the –NH3+ groups –NH3+ + OH- → - NH2 + H2O  A. CHEMICAL BUFFER SYSTEMS 3. Protein Buffers b. Hemoglobin Binds CO2 Binds and transports hydrogen and oxygen Maintains blood pH as hemoglobin changes from oxyhemoglobin to deoxyhemoglobin  Physiologic Buffer Systems Lungs Kidneys B. RESPIRATORY REGULATION Exhalation of carbon dioxide Powerful, but only works with volatile acids Doesn’t affect fixed acids like lactic acid CO2 + H20 H2CO3 H+ + HCO3- Body pH can be adjusted by changing rate and depth of breathing Provide O2 to cells and remove CO2 Ventilation Rates & Effect on pH Balance It’s all about CO2 and the bicarbonate buffering system Increased ventilation rate causes – Removal of CO2 and H2O – Drives this reaction to…? H2O + CO2 H2CO3 HCO3- + H+ – hyperventilation drives the reaction to the left, causing removal of H+, pH goes up – Hypoventilation drives the reaction to the right, causing additional H+, pH goes down  C. RENAL REGULATION Can eliminate large amounts of fixed acid Can also excrete base Can conserve and produce bicarbonate ions Most effective regulator of pH If kidneys fail, pH balance fails Only the kidneys can rid the body of acids generated by cellular metabolism (nonvolatile or fixed acids), while also regulating blood levels of alkaline substances and renewing chemical buffer components.  Base Excretion Only regulated by the kidney. Primary base in the body is HCO3-. The kidney can retain or excrete HCO3- as needed. Importance of Renal Regulation ◻ For every hydrogen ion buffered by bicarbonate – a bicarbonate ion is consumed. H2O + CO2 H2CO3 HCO3- + H+ ◻ To maintain the capacity of the buffer system, the bicarbonate must be regenerated ◻ However, when bicarbonate is formed from carbonic acid (CO2 and H2O) equimolar amounts of [H+] are formed Importance of Renal Regulation H2O + CO2 H2CO3 HCO3- + H+ Bicarbonate formation can only continue if these hydrogen ions are removed This process occurs in the cells of the renal tubules where hydrogen ions are secreted into the urine and where bicarbonate is generated and retained in the body Renal Bicarbonate Regeneration Bicarbonate is freely filtered through the glomerulus so plasma and glomerular filtrate have the same bicarbonate concentration approx 4300 mmol of bicarbonate would be filtered in 24 hrs Without re-generation of bicarbonate the buffering capacity of the body would be depleted causing acidotic state In healthy bodies, virtually all the filtered bicarbonate is recovered  Renal Bicarbonate Regeneration Renal Bicarbonate Regeneration involves the enzyme carbonate dehydratase (carbonic anhydrase) Luminal side of the renal tubular cells impermeable to bicarbonate ions Carbonate dehydratase catalyses the formation of CO2 and H2O from carbonic acid (H2CO3) in the renal tubular lumen CO2 diffuses across the luminal membrane into the tubular cells Regulation of Plasma pH - Acidosis Figure 27– 11a  Renal Responses to Acidosis The renal response to acidosis is 3-fold: A. increased reabsorption of the filtered HCO3− B. increased excretion of titratable acids, and C. increased production of ammonia. Although these mechanisms are probably activated immediately, their effects are generally not appreciable for 12–24 hr  Renal Responses to Acidosis A. increased reabsorption of the filtered HCO3− (1) CO2 within renal tubular cells combines 5 3 with water in the presence of carbonic 4 2 anhydrase. (2) The carbonic acid (H2CO3) formed 1 rapidly dissociates into H+ and HCO3−. (3) Bicarbonate ion then enters the bloodstream (4) while the H+ is secreted into the renal tubule, (5) where it reacts with filtered HCO3− to form H2CO3.  Renal Responses to Acidosis B. Increased Excretion of Titratable Acids After all of the HCO3− in tubular fluid is reclaimed, the H+ secreted into the tubular lumen can combine with HPO42− to form H2PO4− (the latter is not readily reabsorbed because of its charge and is eliminated in urine). The net result is that H+ is excreted from the body as H2PO4−, and the HCO3− that is generated in the process can enter the bloodstream.  Renal Responses to Acidosis C. Increased Formation of Ammonia After complete reabsorption of HCO3− and 2 consumption of the phosphate buffer, the NH3/NH4+ pair becomes the most important urinary buffer 3 1 (1) Deamination of glutamine within the mitochondria of proximal tubular cells is the principal source of NH3 production in the kidneys. (2) The ammonia formed is then able to passively cross the cell’s luminal membrane, enter the tubular fluid, and (3) react with H+ to form NH4+. Unlike NH3, NH4+ does not readily penetrate the luminal membrane. Excretion of NH4+ in urine effectively eliminates H+ Regulation of Plasma pH - Alkalosis Figure 27– 11b  Renal Responses to Alkalosis When the body is in alkalosis, tubular cells secrete bicarbonate ions and reclaim hydrogen ions and acidify the blood The mechanism is the opposite of bicarbonate ion reabsorption process