SEM 1 Biochemistry Lectures 1-15 PDF

BIOCHEMISTRY WATER 2023 - Semester 1. Lectures 1 -13 combined Lecture contents 1. 2. 3. 4. 5. 6. Why water important to biochemistry Uses of water Physics & chemistry of water Unique physical properties of water Molecular structure of water Noncovalent Bonding in water 1. 2. 3. 7. 8. Ionic int...

BIOCHEMISTRY WATER 2023 - Semester 1. Lectures 1 -13 combined Lecture contents 1. 2. 3. 4. 5. 6. Why water important to biochemistry Uses of water Physics & chemistry of water Unique physical properties of water Molecular structure of water Noncovalent Bonding in water 1. 2. 3. 7. 8. Ionic interactions Hydrogen Bonds van der Waals Forces Thermal Properties of Water Solvent Properties of Water 1. 2. Hydrophilic, hydrophobic, and amphipathic molecules Osmotic pressure 9. Ionization of Water 1. Acids, bases, and pH 2. Buffers 3. titration Why water is important to biochemistry n n n n n n More than 70% earth’s surface covered with water The substance that make possible life on earth Solvent & substrate for many cellular reaction Transports chemicals from place to place Helps to maintain constant body temperature Cell components and molecules (protein, polysaccharides, nucleic acid, membranes) assume their shape in response to water USES OF WATER? Introduction Physic and chemistry of water n n Water is the chemical substance with chemical formula H2O: one molecule of water has two hydrogen atoms covalently bonded to a single oxygen atom. Water is a tasteless, odorless liquid at ambient temperature and pressure, and appears colorless in small quantities, although it has its own intrinsic very light blue hue. n n Oxygen attracts electrons much more strongly than hydrogen, resulting in a net positive charge on the hydrogen atoms, and a net negative charge on the oxygen atom. The presence of a charge on each of these atoms gives each water molecule a net dipole moment. Introduction Unique physical properties of water n n n n n n Exist in all three physical states of matter: solid, liquid, and gas. Has high specific heat Water conducts more easily than any liquid except mercury Water has a high surface tension Water is a universal solvent Water in a pure state has a neutral pH Molecular Structure of Water n n Tetrahedral geometry The oxygen in water is sp3 hybridized. Hydrogens are bonded to two of the orbitals. Consequently the water molecule is bent. The H-O-H angle is 104.5o. n n The bent structure indicate water is polar bcoz linear structure is nonpolar. Phenomenon where charge is separated to partial –ve charge and partial +ve charge is called dipoles. n Water is a polar molecule. • A polar molecule is one in which one end is partially positive and the other partially negative. • Oxygen is more electronegative than hydrogen, so oxygen atom bears a partial –ve charge, hydrogen atoms are partial +ve charge n n Molecules eg water, in which charge is separated are called dipoles. Molecular dipoles will orient themselves in the direction opposite to that of the field when subjected to an electric field. Noncovalent Bonding n n n n Usually electrostatic They occur between the positive nucleus of one atom and the negative electron clouds of another nearby atom Relatively weak, easily disrupted Large no. of noncovalent interactions stabilize macromolecules n Types of noncovalent bonding : 1)Ionic interactions 2)Hydrogen bonding 3)Van der Waals forces -Dipole-dipole -Dipole-induced dipole -Induced dipole-induced dipole Typical “Bond” Strengths Type kJ/mol Covalent >210 Noncovalent Ionic interactions strongest non-covalent Hydrogen bonds van der Waals 4-80 12-30 weakest Hydrophobic interactions 0.3-9 3-12 covalent > ionic > H-bond >hydrophobic > VDW 1) Ionic Interactions n n n Interaction occur between charged atoms or group. Oppositely charged ions are attracted to each other. (eg. NaCl) ions with similar charges eg K+ and Na+ will repel each other n n n n In proteins, certain amino acid side chains contain ionizable groups. Glutamic acid ionized as –CH2CH2COOLysine ionized as -CH2CH2CH2CH2NH3+ Attraction between +ve and –ve charged amino acid side chains forms a salt bridge (-COO-+H3N-) CH2CH2COO - + Salt bridge H3N CH2CH2 2) Hydrogen bonding d+ H d- d+ H Hydrogen bonding is a weak attraction between an electronegative atom (O,N,F) in one molecule and a hydrogen atom in another molecule. H *Has both electrostatic (ionic) and covalent character. O H d+ d- O H d+ H d+ O d- d+ n n Water molecule form hydrogen bond with one another Four hydrogen bonding attraction are possible for each molecule: H O H H O H H H O H O H H O H *2 through the hydrogen *2 through the nonbonding electron pairs n n n The resulting intermolecular hydrogen bond acts as bridge between water molecules. Large no. of intermolecular bond (in liquid/solid states of water),the molecules become large, dynamic. This explain why water have high boiling & melting point. 3)Van Der Waals Forces n n n Force between molecules Occur between permanent and/or induced dipoles 3 types of van der waals forces : - Dipole-dipole interactions - Dipole-induced dipole interactions - Induced dipole-induced dipole interactions a) Dipole-dipole interaction n Occur between molecules containing electronegative atoms, cause positive end of one molecule is directed toward negative end of another n eg. Hydrogen bonds are strong type of dipole-dipole interaction d+ C d- O d+ C d- O b)Dipole-induced dipole interaction n A permanent dipole induces a transient dipole in a nearby molecule by distorting its electron distribution n eg. Carbonyl-containing molecule is weakly attracted to hydrocarbon n Weaker than dipole-dipole interaction d+ C d- H d+ O H H H d- c)Induced dipole-induced dipole interactions n Forces between nonpolar molecules n Because of the constant motion of electron, an atom/molecule can develop a temporary dipole (induced dipole) when the electron are distributed unevenly around nucleus n Neighboring atom can be distorted by the appearance of the temporary dipole which lead to an electrostatic interaction between them n Also known as London dispersion forces n eg. Stacking of base ring in DNA molecule Thermal Properties of Water n n n Hydrogen bonding keeps water in the liquid phase between 0oC and 100oC. Liquid water has a high: Heat of vaporization - energy to vaporize one mole of liquid at 1 atm Heat capacity - energy to change the temperature by 1oC Water plays an important role in thermal regulation in living organisms. Relationship between temperature and hydrogen bond n n n n n Max number of hydrogen bonds form when water has frozen into ice. Hydrogen bonds is approximately 15% break when ice is warmed. Liquid water consists of continuously breaking and forming hydrogen bonds. As the tempt rise, the broken of hydrogen bonds are accelerating. When boiling point is reached, the water molecules break free from one another and vaporize. Solvent properties of water n n n n Water is an ideal biological solvent Water easily dissolves a wide variety of the constituents of living organisms. Water also unable to dissolve some substances This behavior is called hydrophilic and hydrophobic properties of water. Hydrophilic molecules n n n n Ionic or polar substances that has an affinity for water In Greek= Hydro, “water” philios, “loving” Water dipole structure and its capacity to form hydrogen bond with electronegative atoms enable water to dissolve ionic and polar substance These substances soluble in water due to 3 kinds of noncovalent bonding : a) ion-dipole b) dipole-dipole c) hydrogen bonding n n n Salts (KCl,NaCl) held together by ionic interactions When ionic compound eg. KCl,NaCl dissolved in water, its ions separate because the polar water molecules attract ions more than the ions attract each other. (ion-dipole interaction) Shells of water mol. cluster around the ions = solvation spheres H O O H H H O K+ O H H O H H H H O H H Cl H H O - H O H Dipole-dipole Interactions n n Organic molecules with ionize group The polar water molecule interacts with carboxyl group of aldehyd & ketones (carbohyd) and hydroxyl group of alcohol H O d- H O H H3C H C d+ O O CH3 H H Dipole-dipole interactions Hydrogen Bonding n A hydrogen attached to an O or N becomes very polarized and highly partial plus. This partial positive charge interacts with the nonbonding electrons on another O or N giving rise to the very powerful H hydrogen bond. O H hydrogen bond R1 O H shown in yellow H O H Hydrophobic molecules n n n Non ionic or nonpolar substance These molecules do not form good attractions with the water molecule. They are insoluble and are said to be hydrophobic (water hating). eg. Hydrocarbon : CH3CH2CH2CH2CH2CH3, hexane n Water forms hydrogen-bonded cagelike structures around hydrophobic molecules, forcing them out of solution. (droplet/into a separate layer) Amphipathic Molecules n n n Amphipathic molecules contain both polar and nonpolar groups. Ionized fatty acids are amphipathic. The carboxylate group is water soluble (hydrophilic) and the long carbon chain is not (hydrophobic). Amphipathic molecules tend to form micelles when mixed with water. n n polar head – orient themselves in contact with water molecules Nonpolar tails – aggregate in the center, away from water Osmotic Pressure n n Osmosis is a spontaneous process in which solvent (eg water) molecules pass through a semi permeable membrane from a solution of lower solute concentration (dilute) to a solution of higher solute concentration (concentrated). Osmotic pressure is the pressure required to stop osmosis (22.4 atm for 1M solution) •Over time, water diffuses from side B (more dilute) to side A (concentrated) A B A B Osmotic Pressure n Osmotic pressure (p) is measured using an osmometer. Osmotic Pressure p = iMRT i = van’t Hoff factor (degree of ionization of solute) M = molarity (concentration of solute in mole/L) R = gas constant (0.082 L.atm/K.mole) T = absolute temp (in Kelvin) Osmolarity = iM (osmol/Liter) n i is the van't Hoff coefficient. n For non-electrolytes (non ionizable solute) n For strong electrolytes i= the number of ions that are produced by the dissociation according to the molecular formula i=1 e.g for NaCl you have 2 ions (1 Na+ and 1 Cl-) so i=2 for CaCl2, 3 ions (1 Ca+2 and 2 Cl-) so i=3 n For weak electrolytes i=(1-a)+na n = the number of ions coming from the 100% dissociation according to the molecular formula a = the degree of dissociation e.g the degree of ionization of 1M CH3COOH solution is 80% a=80%/0.8 , n=2 so, i=(1-0.8) + 2(0.8) =1.8 n n n Osmotic pressure is an important factor affecting cells Cells contain high concentration of solutes – small organic mol., ionic salts, macromolecule Cells may gain or lose water depend on concentration of solute in their environment. Definitions of solutions n n n Isotonic – solutions of equal concentration on either side of the membrane Cells placed in isotonic solution no net movement of water across the membrane Volume of cells are unchanged bcoz water entering & leaving the cell at the same rate. n n n n Hypotonic – solution with a lower solute concentration than the solution on the other side of the membrane Cells placed in hypotonic solution water moves into the cells Cause cells rupture eg. Red blood cells swell & rupture when immersed in pure water (hemolysis) n n n n Hypertonic – solution with higher concentration of solutes than the solution on the other side of the membrane Cells placed in hypertonic solution water moves out the cells Cause cells to shrink eg. Red blood cells shrink when immersed in 3% NaCl solution. (crenation) Water ionization, pH, titration, buffer n n The self-ionization of water is the chemical reaction in which two water molecules react to produce a hydronium (H3O+) and a hydroxide (OH−) ion. Water ionization occurs endothermically due to electric field fluctuations between molecules caused by nearby dipole librations resulting from thermal effects, and favorable localized hydrogen bonding. Water dissociates. (selfionizes) n H2O + H2O = H3O+ + OHn dissociation = ionisation n n Ions may separate but normally recombine within a few min. to seconds. Rarely (about once every eleven hours per molecule at the localized hydrogen bonding arrangement breaks before allowing the separated ions to return, and the pair of ions (H+, OH-) hydrate independently and continue their separate existence. 25°C, or less than once a week at 0°C) Ionization of water n n n may be expressed as Keq = [H3O+][OH-] [H2O]2 The conditions for the water dissociation equilibrium must hold under all situations at 25°C. Kw= [H3O+][OH-] = 1 x 10-14M Pure water ionize into equal amount of [H3O+ ] = [OH-] = 1 x 10-7 M Acids, Bases and pH n When external acids or bases are added to water, the ion product ([H3O+ ][OH-] ) must equal. Kw= [H3O+][OH-] = 1 x 10-14 n The effect of added acids or bases is best understood using the BronstedLowry- theory of acids and bases. Bronstead-Lowry theory n n n Bronsted-Lowry theory is an acid-base theory Acid is a substance that can donate proton (ion H+ donor) acid + base = conjugate base + conjugate acid HCl + H2O = H3O+ + ClAcid Base CA CB C: conjugate (product) A/B n base is a substance that can accept proton RNH2 + H2O = OH- + RNH3+ B A CB CA C: conjugate (product) A/B conjugate base can accept a proton (H+) Measuring Acidity n n n n n Added acids, increase concentration of hydronium ion —> H3O+ In acid solutions [H3O+] > 1 x 10-7 M [OH-] < 1 x 10-7 M Added bases, increase concentration of hydroxide ion. In basic solutions [OH-] > 1 x 10-7 M [H3O+] < 1 x 10-7 M pH scale measures acidity without using exponential numbers. pH Scale n Define: pH = - log(10)[H3O+] 0---------------7---------------14 acidic basic [H3O+]=1 x 10-7 M, pH = ? Strength of Acids n n Strength of an acid is measured by the percent which reacts with water to form hydronium ions. Strong acids (and bases) ionize close to 100%. • eg. HCl, HBr, HNO3, H2SO4 • eg. NaOH, KOH, CaOH Strength of Acids n n Weak acids (or bases) ionize typically in the 1-5% range eg. Organic acid (contain carboxyl groups) CH3COCOOH, pyruvic acid CH3CHOHCOOH, lactic acid CH3COOH, acetic acid Strength of Acids Strength of an acid is also measured by its Ka or pKa values Strength of acid measured by n Dissociation of weak acid : 1) amt which rxts with H2O + HA + H2O = H3O + A to form n Weak acid n conjugate base of HA Strength of weak acid may be determined : Ka = [H3O+][A-] [HA] pKa= -log Ka hyfronium ions 2) its Ka or pKa values Strength of Acids Ka pKa CH3COCOOH CH3CHOHCOOH 3.2x10-3 2.5 1.4x10-4 3.9 CH3COOH 1.8x10-5 4.8 most conc. acid Larger Ka and smaller pKa values indicate stronger acids. Monitoring acidity n The Henderson-Hasselbalch (HH) equation is derived from the equilibrium expression for a weak acid. [A ] conjugate base. since this is the species formed after the acid has lost a proton pH = pKa + log [HA] the acid, before dissociatn (ionisation) HH equation n n The HH equation enables us to calculate the pH during a titration and to make predictions regarding buffer solutions. What is a titration? It is a process in which carefully measured volumes of a base are added to a solution of an acid in order to determine the acid concentration. n n When chemically equal (equivalent) amounts of acid and base are present during a titration, the equivalence point is reached. The equivalence point is detected by using an indicator chemical that changes color or by following the pH of the reaction versus added base, ie. a titration curve. Titration Curve (HOAc with NaOH) acetic acid + sodium hydroxide End point pH Equivalence point NaOH (equivalents added) Titration Curve (HOAc with NaOH) n n At the end point, only the salt (NaOAc) is present in solution. At the equivalence point, equal moles of salt and acid are present in solution. [HOAc] = [NaOAc] pH = pKa Buffer solution n n Buffer : a solution that resists change in pH when small amounts of strong acid or base are added. A buffer consists of: • a weak acid and its conjugate base or • a weak base and its conjugate acid How does buffer work? n n Accepting hydrogen ions from the solution when they are in excess Donating hydrogen ions from the solution when they have depleted Buffer Solutions n n Maximum buffer effect occurs at the pKa for an acid. Effective buffer range is at 1 pH unit above and below the pKa value for the acid or base. eg. H2PO4-/HPO42-, pKa=7.20 buffer range 6.20-8.20 pH Buffer Solutions n n High concentrations of acid and conjugate base give a high buffering capacity. Buffer systems are chosen to match the pH of the physiological situation, usually around pH 7. Physiological buffer n n n n 3 most important buffer in body: Within cells the primary buffer is the phosphate buffer: H2PO4-/HPO42The primary blood buffer is the bicarbonate buffer: HCO3-/H2CO3. Proteins also provide buffer capacity. Side chains can accept or donate protons. (eg. Hemoglobin, serum albumins) n n A zwitterion is a compound with both positive and negative charges. Zwitterionic buffers have become common because they are less likely to cause complications with biochemical reactions. n N-tris(hydroxymethyl)methyl-2aminoethane sulfonate (TES) is a zwitterion buffer example. (HOCH2)3CN+H2CH2CH2SO3- Water : The Medium of Life The End WATER, acids, bases, and buffers I. WATER Water is the solvent of life. - it bathes our cells. -dissolves and transports compounds in the blood - provides a medium for movement of molecules into and throughout cellular compartments - separates charged molecules - dissipates heat - participates in chemical reactions. Most compounds in the body, including proteins, must interact with an aqueous medium function. In spite of the variation in the amount of water, we ingest each day and produce from metabolism, our body maintains a nearly constant amount of water that is approximately 60% of our body weight. A. Fluid Compartments in the Body Total body water is roughly 50 to 60% of body weight in adults and 75% of body weight in children. Approximately 40% of the total body water is intracellular and 60% extracellular. The extracellular water includes the fluid in plasma (blood after the cells have been removed) and interstitial water (the fluid in the tissue spaces, lying between cells). Transcellular water is a small, specialized portion of extracellular water that includes gastrointestinal secretions, urine, sweat, and fluid that has leaked through capillary walls because of such processes as increased hydrostatic pressure or inflammation. B. Hydrogen Bonds in Water The dipolar nature of the water (H2O) molecule allows it to form hydrogen bonds, a property that is responsible for the role of water as a solvent. In H2O, the oxygen atom has two unshared electrons that form an electron dense cloud around it. This cloud lies above and below the plane formed by the water molecule. In the covalent bond formed between the hydrogen and oxygen atoms, the shared electrons are attracted toward the oxygen atom, thus giving the oxygen atom a partial negative charge and the hydrogen atom a partial positive charge. As a result, the oxygen side of the molecule is much more electronegative than the hydrogen side, and the molecule is dipolar. Both the hydrogen and oxygen atoms of the water molecule form hydrogen bonds and participate in hydration shells. A hydrogen bond is a weak non-covalent interaction between the hydrogen of one molecule and the more electronegative atom of an acceptor molecule. The oxygen of water can form hydrogen bonds with two other water molecules, so that each water molecule is hydrogen-bonded to approximately four close neighboring water molecules in a fluid three-dimensional lattice. WATER AS A SOLVENT Polar organic molecules and inorganic salts can readily dissolve in water because water also forms hydrogen bonds and electrostatic interactions with these molecules. Organic molecules containing a high proportion of electronegative atoms (generally oxygen or nitrogen) are soluble in water because these atoms participate in hydrogen bonding with water molecules. Chloride (Cl-), bicarbonate (HCO3-), and other anions are surrounded by a hydration shell of water molecules arranged with their hydrogen atoms closest to the anion. In a similar fashion, the oxygen atom of water molecules interacts with inorganic cations such as Na+and K+ to surround them with a hydration shell. Although hydrogen bonds are strong enough to dissolve polar molecules in water and to separate charges, they are weak enough to allow movement of water and solutes. The strength of the hydrogen bond between two water molecules is only approximately 4 kcal, roughly l/20th of the strength of the covalent O-H bond in the water molecule. Thus, the extensive water lattice is dynamic and has many strained bonds that are continuously breaking and reforming. The average hydrogen bond between water molecules lasts only about 10 psec (1 picosecond is 10"12 sec), and each water molecule in the hydration shell of an ion stays only 2.4 nsec (1nanosecond = 10~9 sec). As a result, hydrogen bonds between water molecules and polar solutes continuously dissociate and reform, thereby permitting solutes to move through water and water to pass through channels in cellular membranes. WATER AND THERMAL REGULATION The structure of water also allows it to resist temperature change. Its heat of fusion is high, so a large drop in temperature is needed to convert liquid water to the solid state of ice. The thermal conductivity of water is also high, thereby facilitating heat dissipation from high energy-using areas such as the brain into the blood and the total body water pool. Its heat capacity and heat of vaporization are remarkably high; as liquid water is converted to a gas and evaporates from the skin, we feel a cooling effect. Water responds to the input of heat by decreasing the extent of hydrogen bonding and to cooling by increasing the bonding between water molecules. Osmolality and Water Movement Water distributes between the different fluid compartments according to the concentration of solutes, or osmolality, of each compartment. The osmolality of a fluid is proportionate to the total concentration of all dissolved molecules, including ions, organic metabolites, and proteins (usually expressed as milliosmoles (mOsm)/kg water). The semipermeable cellular membrane that separates the extracellular and intracellular compartments contains a number of ion channels through which water can freely move, but other molecules cannot. II. ACIDS AND BASES Acids- compounds that donate a hydrogen ion (H+) to a solution Bases - compounds (such as the OH- ion) that accept hydrogen ions. Water itself dissociates to a slight extent, generating hydrogen ions (H+),which are also called protons, and hydroxide ions (OH-). The hydrogen ions are extensively hydrated in water to form species such as H3O+, but nevertheless are usually represented as simply H+. Water itself is neutral, neither acidic nor basic. The pH of Water The extent of dissociation by water molecules into H+ and OH- is very slight, and the hydrogen ion concentration of pure water is only 0.0000001 M, or 10-7 mol/L. The concentration of hydrogen ions in a solution is usually denoted by the term pH, which is the negative logI0 of the hydrogen ion concentration expressed in mol/L. Therefore, the pH of pure water is 7. The 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. Because water dissociates to such a small extent, [H2O] is essentially constant at 55.5 M. Multiplication of the Kd for water (approximately 1.8 X 10-16M) by 55.5 M gives a value of approximately 10 -l4 (M)2, which is called the ion product of water (Kw). Because Kw, the product of [H+] and [OH-], is always constant, a decrease of [H+] must be accompanied by a proportionate increase of [OH-]. A pH of 7 is termed neutral because [H+] and [OH-] are equal. acidic solutions have a greater hydrogen ion concentration and a lower hydroxide ion concentration than pure water (pH<7.0) basic solutions have a lower hydrogen ion concentration and a greater hydroxide ion concentration (pH>7.0). Strong and Weak Acids During metabolism, the body produces a number of acids that increase the hydrogen ion concentration of the blood or other body fluids and tend to lower the pH. These metabolically important acids can be classified as weak acids or strong acids by their degree of dissociation into a hydrogen ion and a base (the anion component). Inorganic acids such as sulfuric acid (H2SO4) and hydrochloric acid (HCl) are strong acids that dissociate completely in solution. Organic acids containing carboxylic acid groups (e.g., the ketone bodies (acetoacetic acid and β-hydroxybutyric acid) are weak acids that dissociate only to a limited extent in water. In general, a weak acid (HA), called the conjugate acid, dissociates into a hydrogen ion and an anionic component (A-), called the conjugate base. The name of an undissociated acid usually ends in "ic acid" (e.g., acetoacetic acid) and the name of the dissociated anionic component ends in "ate" (e.g., acetoacetate). The tendency of the acid (HA) to dissociate and donate a hydrogen ion to solution is denoted by its Ka, the equilibrium constant for dissociation of a weak acid. The higher the Ka, the greater is the tendency to dissociate a proton. In the Henderson-Hasselbalch equation, the formula for the dissociation constant of a weak acid is converted to a convenient logarithmic equation. The term pKa represents the negative log of Ka. If the pK., for a weak acid is known, this equation can be used to calculate the ratio of the unprotonated to the protonated form at any pH. From this equation -a weak acid is 50% dissociated at a pH equal to its pKa. Most of the metabolic carboxylic acids have pKas between 2 and 5, depending on the other groups on the molecule. The pKa reflects the strength of an acid. Acids with a pKa of 2 are stronger acids than those with a pKa of 5 because, at any pH, a greater proportion is dissociated. III. BUFFERS A buffer solution = aqueous solution consisting of a mixture of a weak acid and its conjugate base (or vice versa); it prevents changes in the pH of a solution when hydrogen ions or hydroxide ions are added. - their resistance to pH change is caused by the presence of an equilibrium between the acid, HA, and its conjugate base, A−. HA ↔ H+ + A− -the equilibrium is shifted according to the Le Chatelier’s principle (“The Equilibrium Law”): If a system at equilibrium is subjected to some changes in concentration, temperature, volume and pressure, then the system readjusts itself to counteract the effect of the applied change, thus establishing a new equilibrium. A buffer of acetic acid (CH3COOH) and acetate (CH3COO-) - the pH of the solution is graphed as a function of the amount of OH- that has been added. - the OH- is expressed as equivalents of total acetic acid present in the dissociated and undissociated forms. At the midpoint of this curve, 0.5 equivalents of OHhave been added, and half of the conjugate acid has dissociated so that [A-] equals [HA]. This midpoint is expressed in the HendersonHasselbalch equation as the pKa, defined as the pH at which 50% dissociation occurs. Ø As you add more OH- ions and move to the right on the curve, more of the conjugate acid molecules (HA) dissociate to generate H+ ions, which combine with the added OH- ions to form water. => only a small increase in pH results. Ø If you add H+ ions to the buffer at its pKa (moving to the left of the midpoint), conjugate base molecules (A-) combine with the added hydrogen ions to form HA => almost no decrease in pH occurs. Ø 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. Likewise, 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. METABOLIC ACIDS AND BUFFERS An average rate of metabolic activity produces roughly 22,000 mEq acid/day. If all of this acid were dissolved at one time in unbuffered body fluids, their pH would be less than 1. The pH of the blood is normally maintained between 7.36 and 7.44, and intracellular pH at approximately 7.1 (between 6.9 and 7.4). The widest range of extracellular pH over which the metabolic functions of the liver, the beating of the heart, and conduction of neural impulses can be maintained is 6.8 to 7.8. Outside the acceptable range of pH, proteins are denatured and enzymes lose their ability to function. Acid-base homeostasis (acid-base balance) is assured by: • • • • a) buffering systems/agents the bicarbonate-carbonic acid buffer system in extracellular fluid; the hemoglobin buffer system in red blood cells; the phosphate buffer system in all types of cells; the protein buffer system of cells and plasma b) respiratory system (quickly) c) renal system (slower) An acid-base nomograph of human serum Amino acids in Proteins Proteins have many functions in the body servings as: • transporters of hydrophobic compounds in the • • • • blood, cell adhesion molecules that attach cells to each other and to the extracellular matrix, hormones that carry signals from one group of cells to another, ion channels through lipid membranes, enzymes that increase the rate of biochemical reactions. The unique characteristics of a protein are dictated by its linear sequence of amino acids, termed its primary structure. The primary structure of a protein determines how it can fold and how it interacts with other molecules in the cell to perform its function. The primary structures of all of the diverse human proteins are synthesized from 20 amino acids arranged in a linear sequence determined by the genetic code. I. GENERAL STRUCTURE OF THE AMINO ACIDS 20 different amino acids are commonly found in proteins. They are all α-amino acids, amino acids in which the amino group is attached to the α-carbon (the carbon atom next to the carboxylate group). The α-carbon has two additional substituents: - a hydrogen atom - an additional chemical group called a side chain (-R) different for each amino acid. The chemical properties of the amino acids give each protein its unique characteristics. Proteins are composed of one or more linear polypeptide chains containing hundreds of amino acids. The names of the different amino acids have been given three-letter and one-letter abbreviations. The three-letter abbreviations use the first two letters in the name plus the third letter of the name or the letter of a characteristic sound, such as trp for tryptophan. The one-letter abbreviations use the first letter of the name of the most frequent amino acid in proteins (such as an "A" for alanine). If the first letter has already been assigned, the letter of a characteristic sound is used (such as an "R" for arginine). Single-letter abbreviations are usually used to denote the amino acids in a polypeptide sequence. CLASSIFICATION OF AMINO ACID SIDE CHAINS The 20 amino acids used for protein synthesis are grouped into different classifications according to the polarity and structural features of the side chains. GENERAL PROPERTIES The α-amino groups all have pK values near 9.4 and are therefore almost entirely in the ammonium ion form below pH 8.0. => In the physiological pH range, both the carboxylic acid and the amino groups of α-amino acids are completely ionized. An amino acid can therefore act either as an acid or a base. Substances with this property are said to be amphoteric and are referred to as ampholytes (amphoteric electrolytes). Molecules that bear charged groups of opposite polarity are known as zwitterions or dipolar ions. The zwitterionic character of the α -amino acids has been established by several methods including spectroscopic measurements and X-ray crystal structure determinations. Because amino acids are zwitterions, their physical properties are characteristic of ionic compounds (solids with a high melting point). In neutral solution In alkaline solution In acidic solution In the solid state Peptide Bonds The α-amino acids polymerize, at least conceptually, through the elimination of a water molecule. The resulting CO — NH linkage is known as a peptide bond. Polymers composed of two, three, a few (3-10), and many amino acid residues (alternatively called peptide units) are known, respectively, as dipeptides, tripeptides, oligopeptides, and polypeptides. These substances, however, are often referred to simply as "peptides." Polypeptides are linear polymers; that is, each amino acid residue is linked to its neighbors in a head-to-tail fashion rather than forming branched chains. Proteins are molecules that consist of one or more polypeptide chains. These polypeptides range in length from ~ 40 to over 4000 amino acid residues. Acid-Base Properties Amino acids and proteins have conspicuous acidbase properties. The α-amino acids have two or, for those with ionizable side groups, three acid-base groups. The titration curve of glycine, the simplest amino acid. At low pH values, both acid - base groups of glycine are fully protonated so that it assumes the cationic form +H3NCH2COOH. In the course of the titration with a strong base, such as NaOH, glycine loses two protons in the stepwise fashion characteristic of a polyprotic acid. The pK values of glycine's two ionizable groups are sufficiently different so that the Henderson-Hasselbalch equation: Consequently, the pK for each ionization step is that of the midpoint of its corresponding leg of the titration curve : - at pH 2.35 the concentrations of the cationic form +H NCH COOH, and the zwitterionic form +H NCH COO-, 3 2 3 2 are equal and similarly. - at pH 9.78 the concentrations of this zwitterionic form +H NCH COO-, and the anionic form, H NCH COO-, are 3 2 2 2 equal. Amino acids never assume the neutral form in aqueous solution. The pH at which a molecule carries no net electric charge is known as its isoelectric point, pI. For the α-amino acids, the application of the Henderson Hasselbalch equation indicates that, to a high degree of precision, where Ki, and Kj are the dissociation constants of the two ionizations involving the neutral species. For monoamino, monocarboxylic acids such as glycine, Ki and Kj represent K1 and K2. The pI value affects the solubility of the molecule: - at pH = pI the molecules have minimum solubility in water or salt solutions => precipitate - at pH ≠ pI the molecule is charged => separation of proteins in electrophoresis Proteins Have Complex Titration Curves The titration curves of the α-amino acids with ionizable side chains, such as that of glutamic acid, exhibit the expected three pK values. However, the titration curves of polypeptides and proteins, rarely provide any indication of individual pK values because of the large numbers of ionizable groups they represent (typically 25% of a protein's amino acid side chains are ionizable). Furthermore, the covalent and three-dimensional structure of a protein may cause the pK of each ionizable group to shift by as much as several pH units from its value in the free α-amino acid as a result of the electrostatic influence of nearby charged groups, medium effects arising from the proximity of groups of low dielectric constant, and the effects of hydrogen bonding associations. The titration curve of a protein is also a function of the salt concentration, as is shown in fig. because the salt ions act electrostatically to shield the side chain charges from one another, thereby attenuating these charge - charge interactions. OPTICAL ACTIVITY The amino acids are, with the exception of glycine, all optically active = they rotate the plane of planepolarized light; L-amino acids. Optically active molecules have an asymmetry such that they are not superimposable on their mirror image in the same way that a left hand is not superimposable on its mirror image, a right hand. The two molecules depicted are not superimposable since they are mirror images. This situation is characteristic of substances that contain tetrahedral carbon atoms that have four different substituents. The central atoms in such atomic constellations are known as asymmetric centers, or chiral centers, and are said to have the property of chirality (Greek: cheir, hand). The Cα atoms of all the amino acids, with the exception of glycine, are asymmetric centers. Glycine, which has two H atoms substituent to its Cα atom, is superimposable on its mirror image and is therefore not optically active. An asymmetric (chiral) carbon, linking to four different substituents, can have two configurations, producing a pair of stereoisomers called enantiomers. The number of enantiomers = 2n n = number of the asymmetric carbons Molecules that are nonsuperimposable mirror images are known as enantiomers of one another. The mixture of racemic mixture the 2 enantiomers = Enantiomeric molecules: - have identical chemical and physical properties => - are physically and chemically indistinguishable by most techniques. - rotate plane-polarized light (+/−) by equal amounts but in opposite directions. - with other enantiomers they give different chemical reactions: In drugs for example, often only one of a drug's enantiomers is responsible for the therapeutic effects, meanwhile the other enantiomer is less active, inactive, or sometimes even harmful (responsible for adverse effects). Ex: zoplicone, albuterol, omeprazol, citalopram, etc. The two are enantiomers The two are the same Configuration may also result from the presence of asymmetric carbons. Protein Structure Protein Functions • Three examples of protein functions Alcohol dehydrogenase oxidizes alcohols to aldehydes or ketones – Catalysis: Almost all chemical reactions in a living cell are catalyzed by protein enzymes. – Transport: Some proteins transports various substances, such as oxygen, ions, and so on. Haemoglobin carries oxygen – Information transfer: For example, hormones. Insulin controls the amount of sugar in the blood Amino Acids • When scientists first turned their attention to nutrition, early in the nineteenth century, they quickly discovered that natural products containing nitrogen were essential for the survival of animals. In 1839, the Swedish chemist Jacob Berzelius coined the term protein (Greek: proteios, primary) for this class of compounds. The physiological chemists of that time did not realize that proteins were actually composed of smaller units, amino acids, although the first amino acids had been isolated in 1830. In fact, for many years, it was believed that substances from plants—including proteins—were incorporated whole into animal tissues. This misconception was laid to rest when the process of digestion came to light. After it became clear that ingested proteins were broken down to smaller compounds containing amino acids, scientists began to consider the nutritive qualities of those compounds. Modern studies of proteins and amino acids owe a great deal to nineteenthand early twentieth-century experiments. We now understand that nitrogen containing amino acids are essential for life and that they are the building blocks of proteins. The central role of amino acids in biochemistry is perhaps not surprising: Several amino acids are among the organic compounds believed to have appeared early in the earth’s history (Section 1-1A). Amino acids, as ancient and ubiquitous molecules, have been co-opted by evolution for a variety of purposes in living systems. We begin this chapter by discussing the structures and chemical properties of the common amino acids, including their stereochemistry, and end with a brief summary of the structures and functions of some related compounds. Staphylococcus epidermidis, which grows on human skin, links the amino acid glutamate into long chains. These help protect the bacteria from changes in extracellular salt concentration that normally occur on the skin surface. Pathways to Discovery William C. Rose and the Discovery of Threonine William C. Rose (1887–1985) Identifying the aminonacid constituents of proteins was a scientific challenge that grew out of studies of animal nutrition. At the start of the twentieth century, physiological chemists (the term biochemist was not yet used) recognized that not all foods provided adequate nutrition. For example, rats fed the corn protein zein as their only source of nitrogen failed to grow unless the amino acids tryptophan and lysine were added to their diet. Knowledge of metabolism at that time was limited mostly to information gleaned from studies in which intake of particular foods in experimental subjects (including humans) was linked to the urinary excretion of various compounds. Results of such studies were consistent with the idea that compounds could be transformed into other compounds, but clearly, nutrients were not wholly interchangeable. At the University of Illinois, William C. Rose focused his research on nutritional studies to decipher the metabolic relationships of nitrogenous compounds. Among other things, his studies of rat growth and nutrition helped show that purines and pyrimidines were derived from amino acids but that those compounds could not replace dietary amino acids. To examine the nutritional requirements for individual amino acids, Rose hydrolyzed proteins to obtain their component amino acids and then selectively removed certain amino acids. In one of his first experiments, he removed arginine and histidine from a hydrolysate of the milk protein casein. Rats fed on this preparation lost weight unless the amino acid histidine was added back to the food. However, adding back arginine did not compensate for the apparent requirement for histidine. These results prompted Rose to investigate the requirements for all the amino acids. Using similar experimental approaches, Rose demonstrated that cysteine histidine, and tryptophan could not be replaced by other amino acids. From preparations based on hydrolyzed proteins, Rose moved to mixtures of pure amino acids. Thirteen of the 19 known amino acids could be purified, and the other six synthesized. However, rats fed these 19 amino acids as their sole source of dietary nitrogen lost weight. Although one possible explanation was that the proportions of the pure amino acids were not optimal, Rose concluded that there must be an additional essential amino acid, present in naturally occurring proteins and their hydrolysates but not in his amino acid mixtures. After several years of effort, Rose obtained and identified the missing amino acid. In work published in 1935, Rose showed that adding this amino acid to the other 19 could support rat growth. Thus, the twentieth and last amino acid, threonine, was discovered. Experiments extending over the next 20 years revealed that 10 of the 20 amino acids found in proteins are nutritionally essential, so that removal of one of these causes growth failure and eventually death in experimental animals. The other 10 amino acids were considered “dispensable” since animals could synthesize adequate amounts of them. Rose’s subsequent work included verifying the amino acid requirements of humans, using graduate students as subjects. Knowing which amino acids were required for normal health— and in what amounts—made it possible to evaluate the potential nutritive value of different types of food proteins. Eventually, these findings helped guide the formulations used for intravenous feeding. McCoy, R.H., Meyer, C.E., and Rose, W.C., Feeding experiments with mixtures of highly purified amino acids. VIII. Isolation and identification of a new essential amino acid, J. Biol. Chem. 112, 283– 302 (1935). [Freely available at http://www.jbc.org.] Amino Acid Structure • KEY CONCEPTS • The 20 standard amino acids share a common structure but differ in their side chains. • Peptide bonds link amino acid residues in a polypeptide. • Some amino acid side chains contain ionizable groups whose pK values may vary. The analyses of a vast number of proteins from almost every conceivable source have shown that all proteins are composed of 20 “standard” amino acids. Not every protein contains all 20 types of amino acids, but most proteins contain most, if not all, of the 20 types. The common amino acids are known as α-amino acids because they have a primary amino group (—NH2) as a substituent of the α carbon atom, the carbon next to the carboxylic acid group (—COOH; Fig. 1). The sole exception is proline, which has a secondary amino group (—NH—), although for uniformity we will refer to proline as an α-imino acid. The 20 standard amino acids differ in the structures of their side chains (R groups). Table 1 displays the names and complete chemical structures of the 20 standard. Amino acid: Basic unit of protein R NH3+ C Amino group H Different side chains, R, determine the COO properties of 20 amino Carboxylic acid group acids. An amino acid 20 Amino acids Glycine (G) Alanine (A) Valine (V) Isoleucine (I) Leucine (L) Proline (P) Methionine (M) Phenylalanine (F) Tryptophan (W) Asparagine (N) Glutamine (Q) Serine (S) Threonine (T) Tyrosine (Y) Cysteine (C) Lysine (K) Arginine (R) Histidine (H) Asparatic acid (D) Glutamic acid (E) White: Hydrophobic, Green: Hydrophilic, Red: Acidic, Blue: Basic Each protein has a unique structure! Amino acid sequence NLKTEWPELVGKSVEEAK KVILQDKPEAQIIVLPVGTI VTMEYRIDRVRLFVDKLD Folding! Amino Acids Are Dipolar Ions The amino and carboxylic acid groups of amino acids ionize readily. The pK values of the carboxylic acid groups (represented by pK1 ) lie in a small range around 2.2, while the pK values of the α-amino groups (pK2) are near 9.4. At physiological pH (∼7.4), the amino groups are protonated and the carboxylic acid groups are in their conjugate base (carboxylate) form (pKR). . An amino acid can therefore act as both an acid and a base. Molecules such as amino acids, which bear charged groups of opposite polarity, are known as dipolar ions or zwitterions. Amino acids, like other ionic compounds, are more soluble in polar solvents than in nonpolar solvents. As we will see, the ionic properties of the side chains influence the physical and chemical properties of free amino acids and amino acids in proteins. Amino Acid Ionization The nonionic and zwitterionic forms of a simple amino acid such as alanine are shown in Fig. The zwitterionic form predominates at neutral pH. The nonionic form does not occur in significant amounts in aqueous solution at any pH. A zwitterion can act as either an acid (proton donor) or a base (proton acceptor). Peptide Bonds Link Amino Acids Amino acids can be polymerized to form chains. This process can be represented as a condensation reaction (bond formation with the elimination of a water molecule), as shown in Fig. The resulting CO—NH linkage, an amide linkage, is known as a peptide bond. Polymers composed of two, three, a few (3–10), and many amino acid units are known, respectively, as dipeptides, tripeptides, oligopeptides, and polypeptides. These substances, however, are often referred to simply as “peptides.” After they are incorporated into a peptide, the individual amino acids (the monomeric units) are referred to as amino acid residues. Polypeptides are linear polymers rather than branched chains; that is, each amino acid residue participates in two peptide bonds and is linked to its neighbors in a head-to-tail fashion. The residues at the two ends of the polypeptide each participate in just one peptide bond. The residue with a free amino group (by convention) is called the amino terminus or Nterminus. The residue with a free carboxylate group (at the right) is called the carboxyl terminus or C-terminus. Proteins are molecules that contain one or more polypeptide chains. Variations in the length and the amino acid sequence of polypeptides are major contributors to the diversity in the shapes and biological functions of proteins. Amino Acid Side Chains Are Nonpolar, Polar, or Charged The most useful way to classify the 20 standard amino acids is by the polarities of their side chains. According to the most common classification scheme, there are three major types of amino acids: (1) those with nonpolar R groups, (2) those with uncharged polar R groups, and (3) those with charged polar R groups. The Nonpolar Amino Acid Side Chains Have a Variety of Shapes and Sizes. Nine amino acids are classified as having nonpolar side chains. The threedimensional shapes of some of these amino acids are shown in Fig. Glycine has the smallest possible side chain, an H atom. Alanine, valine, leucine, and isoleucine have aliphatic hydrocarbon side chains ranging in size from a methyl group for alanine to isomeric butyl groups for leucine and isoleucine. Methionine has a thioether side chain that resembles an n-butyl group in many of its physical properties (C and S have nearly equal electronegativities, and S is about the size of a methylene group). Proline has a cyclic pyrrolidine side group. Phenylalanine (with its phenyl moiety) and tryptophan (with its indole group) contain aromatic side groups, which are characterized by bulk as wel as nonpolarity Isoleucine Alanine Phenylalanine Some amino acids with nonpolar side chains. The amino acids are shown as ball-and-stick models embedded in transparent space-filling models. The atoms are colored according to type, with C green, H white, N blue, and O red. Uncharged Polar Side Chains Have Hydroxyl, Amide, or Thiol Groups. Six amino acids are commonly classified as having uncharged polar side chains . Serine and threonine bear hydroxylic R groups of different sizes. Asparagine and glutamine have amide-bearing side chains of different sizes. Tyrosine has a phenolic group (and, like phenylalanine and tryptophan, is aromatic). Cysteine is unique among the 20 amino acids in that it has a thiol group that can form a disulfide bond with another cysteine through the oxidation of the two thiol groups . Serine Glutamine Some amino acids with uncharged polar side chains. Note the presence of electronegative atoms on the side chains. The thiol groups of two cysteine residues are readily oxidized to form a covalently linked dimeric amino acid known as cystine. In cystine, the two cysteines are joined by a disulfide bond (Fig.). The disulfide-linked cystine residue is strongly hydrophobic. In proteins, disulfide bonds form covalent links between different parts of a polypeptide chain, or between two different polypeptide chains Charged Polar Side Chains Are Positively or Negatively Charged. Five amino acids have charged side chains . The side chains of the basic amino acids are positively charged at physiological pH values. Lysine has a butylammonium side chain, and arginine bears a guanidino group. Histidine carries an imidazolium moiety. Note that only histidine, with a pKR of 6.04, readily ionizes within the physiological pH range. Consequently, both the neutral and cationic forms occur in proteins. In fact, the protonation–deprotonation of histidine side chains is a feature of numerous enzymatic reaction mechanisms. The side chains of the acidic amino acids, aspartic acid and glutamic acid, are negatively charged above pH 3; in their ionized state, they are often referred to as aspartate and glutamate. Asparagine and glutamine are, respectively, the amides of aspartic acid and glutamic acid. Positively Charged Amino Acids The side-chains of lysine and arginine are fully positively charged at neutral pH. In lysine, a primary amino group is attached to the e carbon of the side-chain. In arginine, the guanidinium group of the side-chain is postively charged. The histidine R group contains an aromatic imidazole group that is partially positively charged at neutral pH . Histidine residues function in many enzymecatalyzed reactions as proton donors and/or acceptors. Negatively Charged Amino Acids The R groups of aspartate and glutamate contain carboxyl groups that are fully negatively charged at neutral pH (pKRs of 3.65 and 4.25). In aspartate, the carboxyl group is attached to the ß carbon of the amino acid backbone. In glutamate, the carboxyl group is attached to the g carbon. Aspartate Lysine Some amino acids with charged polar side chains The foregoing allocation of the 20 amino acids among the three different groups is somewhat arbitrary. For example, glycine and alanine, the smallest of the amino acids, and tryptophan, with its heterocyclic ring, might just as well be classified as uncharged polar amino acids. Similarly, tyrosine and cysteine, with their ionizable side chains, might also be thought of as charged polar amino acids, particularly at higher pH values. In fact, the deprotonated side chain of cysteine (which contains the thiolate anion, S−) occurs in a variety of enzymes, where it actively participates in chemical reactions. Inclusion of a particular amino acid in one group or another reflects not just the properties of the isolated amino acid, but its behavior when it is part of a polypeptide. The structures of most polypeptides depend on a tendency for polar and ionic side chains to be hydrated and for nonpolar side chains to associate with each other rather than with water. This property of polypeptides is the hydrophobic effect in action. As we will see, the chemical and physical properties of amino acid side chains also govern the chemical reactivity of the polypeptide. It is therefore worthwhile studying the structures of the 20 standard amino acids in order to appreciate how they vary in polarity, acidity, aromaticity, bulk, conformational flexibility, ability to cross-link, ability to hydrogen bond, and reactivity toward other groups. The pK Values of Ionizable Groups Depend on Nearby Groups The α-amino acids have two or, for those with ionizable side chains, three acid–base groups. At very low pH values, these groups are fully protonated, and at very high pH values, these groups are unprotonated. At intermediate Ph values, the acidic groups tend to be unprotonated, and the basic groups tend to be protonated. Thus, for the amino acid glycine, below pH 2.35 (the pK value of its carboxylic acid group), the +H3NCH2COOH form predominates. Above pH 2.35, the carboxylic acid is mostly ionized but the amino group is still mostly protonated (+H3NCH2COO−). Above pH 9.78 (the pK value of the amino group), the H2NCH2COO− form predominates. Note that in aqueous solution, the un-ionized form (H2NCH2COOH) is present only in vanishingly small quantities. The pH at which a molecule carries no net electric charge is known as its isoelectric point, pI. For the α-amino acids, pI = 1/2 (pKi + pKj) where Ki and Kj are the dissociation constants of the two ionizations involving the neutral species. For monoamino, monocarboxylic acids such as glycine, Ki and Kj represent K1 and K2. However, for aspartic and glutamic acids, Ki and Kj are K1 and KR, whereas for arginine, histidine, and lysine, these quantities are KR and K2 . Of course, amino acid residues in the interior of a polypeptide chain do not have free α-amino and α-carboxyl groups that can ionize (these groups are joined in peptide bonds. Furthermore, the pK values of all ionizable groups, including the N- and C-termini, usually differ from the pK values listed in Table 3-1 for free amino acids. For example, the pK values of αcarboxyl groups in unfolded proteins range from 3.5 to 4.0 In the free amino acids, the pK values are much lower, because the positively charged ammonium group electrostatically stabilizes the COO– group, in effect making it easier for the carboxylic acid group to ionize. Similarly, the pK values for α-amino groups in proteins range from 7.5 to 8.5. In the free amino acids, the pK values are higher, due to the electron-withdrawing character of the nearby carboxylate group, which makes it more diffi cult for the ammonium group to become deprotonated. In addition, the three-dimensional structure of a folded polypeptide chain may bring polar side chains and the N and C-termini close together. The resulting electrostatic interactions between these groups may shift their pK values up to several pH units from the values for the corresponding free amino acids. For this reason, the pI of a polypeptide, which is a function of the pK values of its many ionizable groups, is not easily predicted and is usually determined experimentally. Titration of Simple Amino Acids The titration curves of simple amino acids such as glycine that have non-dissociable R groups, have two plateaus, which correspond to the dissociation and titration of the a-carboxyl group (pK1, left) and the aamino group (pK2, right). As shown above the curve, the predominant ionic species in solution at low pH is the fully protonated form, +H3N-CH2-COOH (net charge = +1), In between the two plateaus, the zwitterionic form, +H3N-CH2-COO- (net charge = 0) predominates. At the end of the titration, the fully dissociated species H2N-CH2-COO- (net charge -1) predominates. The curve shows that glycine has two regions of buffering power centered ±1 pH unit above pK1 and pK2. The Henderson-Hasselbalch equation can be used to calculate the amounts of the conjugate acid and conjugate base species in solution at any pH. Lastly, the pH at which the zwitterionic (0-charged) species of glycine predominates (one equivalent of OH- added) is called the isoelectric point or isoelectric pH. The isoelectric pH is exactly halfway between the two pKas for glycine. Chemical Environment and the pKa The pKas of the a-carboxyl groups of all amino acids are lower than the pKas of the carboxyl groups in methyl-substituted carboxylic acids such as acetic acid . This is due to the local chemical environment of the a-carboxyl groups in amino acids. Namely, placement near the a-amino group, which is positively charged, makes the a-carboxyl groups of amino acids more acidic than the carboxyl group of acetic acid. Similarly, the chemical environment near a-amino groups makes them more acidic than the amino groups of a methyl-substituted amino compounds such as methylamine. In this case the electron withdrawing properties of the oxygens on the a-carboxyl groups of amino acids make the aamino groups hold onto their protons less tightly than in other environments. Titration of Glutamate The acidic amino acid, glutamate, has a second carboxyl group present in its side-chain. Thus the titration curve for glutamate (and aspartate) has three plateaus, each one corresponding to the dissociation of a proton from the amino acid . Since the R group carboxyl group has a pKR between that of the pK1 and pK2, the second plateau corresponds to the titration of this group. Based on inspection of the ionic forms in solution (top) it is clear that the zwitterionic form