BCH 201: General Biochemistry I (2 Units) WATER, ACID, BASE, pH and BUFFERS PDF

BCH 201: GENERAL BIOCHEMISTRY I (2 UNITS) WATER, ACID, BASE, pH and BUFFERS BY ABDUL-MALIK ABDUL-QADIR INTRODUCTION Water is an inorganic, transparent, tasteless, odorless, and nearly colorless chemical substance,...

BCH 201: GENERAL BIOCHEMISTRY I (2 UNITS) WATER, ACID, BASE, pH and BUFFERS BY ABDUL-MALIK ABDUL-QADIR INTRODUCTION Water is an inorganic, transparent, tasteless, odorless, and nearly colorless chemical substance, which is the main constituent of Earth's hydrosphere and the fluids of most living organisms. It is vital for all known forms of life, even though it provides no calories or organic nutrients. Its chemical formula is H2O, meaning that each of its molecules contains one oxygen and two hydrogen atoms, connected by covalent bonds. Water is the name of the liquid state of H2O at standard ambient temperature and pressure. It forms precipitation in the form of rain and aerosols in the form of fog. Clouds are formed from suspended droplets of water and ice, its solid state. When finely divided, crystalline ice may precipitate in the form of snow. The gaseous state of water is steam or water vapor. Water moves continually through the water cycle of evaporation, transportation (evapotranspiration), condensation, precipitation, and runoff, usually reaching the sea. Water is widely distributed on Earth as freshwater and salt water in the oceans. The Earth is often referred to as the "blue planet" because when viewed from space it appears blue. This blue color is caused by reflection from the oceans which cover roughly 70% of the area of the Earth. Water plays an important role in the world economy. Approximately 70% of the freshwater used by humans goes to agriculture. Fishing in salt and fresh water bodies is a major source of food for many parts of the world. Much of long-distance trade of commodities (such as oil and natural gas) and manufactured products is transported by boats through seas, rivers, lakes, and canals. Large quantities of water, ice, and steam are used for cooling and heating, in industry and homes. Water is an excellent solvent for a wide variety of substances both mineral and organic; as such it is widely used in industrial processes, and in cooking and washing. Water, ice and snow are also central to many sports and other forms of entertainment, such as swimming, pleasure boating, boat racing, surfing, sport fishing, diving, ice skating and skiing. PROPERTIES OF WATER The major chemical and physical properties of water are;  Water is a liquid at standard temperature and pressure.  It is tasteless and odorless.  The intrinsic color of water and ice is a very slight blue hue, although both appear colorless in small quantities.  Water vapor is essentially invisible as a gas.  Water is transparent in the visible electromagnetic spectrum. Thus aquatic plants can live in water because sunlight can reach them. Ultra-violet and infrared light is strongly absorbed. BIOCHEMISTRY OF WATER WATER is the solvent of choice for biological systems. It constitutes about 70-85% of cell weight, it serves as an important solvent and a reactant in biochemical reactions. It helps regulate temperature since it is able to absorb large amounts of heat. It regulate intracellular pH and it also used for transport. Water is dipole because of its geometry and the difference in electronegativity between hydrogen and oxygen. Oxygen is more electronegative than hydrogen - Oxygen is sp3 hybridized; tetrahedral electron geometry; BENT molecular geometry - H-bonding is especially strong in water because:  The O—H bond is very polar  There are 2 lone pairs on the O atom  Each H2O molecule can form four H bonds to other molecules, resulting in a tetrahedral arrangement. Dissociation of Water Amphipathic compounds Amphipathic compounds are the molecules which contain both hydrophobic groups (large nonpolar hydrocarbon chains) and polar or ionic groups (hydrophilic groups). They don’t dissolve in water as individual molecules. When they reach at a definite concentration (critic micelle concentration) in water, they associate with each other in submicroscopic aggregations of molecules called micelles. Micelles have hydrophilic groups on their exterior (bonding with solvent water), and hydrophobic groups clustered in their interior. They occur in spherical shapes. Micelle structures are stabilized by hydrogen bonding with water by van der Waals attractive forces between hydrocarbon groups in the interior, and by energy of hydrophobic interactions. Osmotic pressure Osmotic pressure is a measure of the tendency of water molecules to migrate from a diluted solution to a concentrated solution through a semipermeable membrane. This migration of water molecules is termed osmosis. A solution containing 1 mol of solute particles in 1 kg of water is a one-osmolal solution. Non-covalent Interactions: Relatively weak and reversible  Hydrogen Bonds: Special dipole-dipole interaction (electronegative atom (e.g. O or N) interacts with H atom that is partially positive (i.e. attached to N, O, F) Very important for protein and DNA structure.  Van Der Waals interactions a. Dispersion Forces (London Forces) (induced dipoles in non-polar molecules) b. Dipole-dipole  Ionic Bonds: Electrostatic interaction between two oppositely charged ions. Hydrophobic Interactions/Hydrophobic Effect: Relations between water and hydrophobic molecules (low water-soluble molecules). Nonpolar substances tend to aggregate in aqueous solution and exclude water molecules. Thermodynamically unfavorable to dissolve hydrophobic substance in water entropy-driven process. ΔG = ΔH – TΔS Water molecules align themselves around non-polar molecule and lose freedom to form hydrogen bonds. Entropy lost in system results in thermodynamic barrier. Multiple molecules aggregate – increase the entropy of the system because fewer water molecules needed to surround the aggregate than to hydrate each dispersed molecule.Water soluble compounds are those in which the interactions between the solute and water are greater than those between solute molecules. i.e, salts, biological molecules that have polar or ionic groups (e.g. glucose, ethanol) - Reversible self-dissociation = ionization that generates H+ and OH and Can be described by the following equilibrium Note hydrogen atoms do not exist as free H+ in solution. Actually are hydronium ions (H3O+). For simplicity, we just write H+. Express extent of ionization quantitatively: Use law of mass action to define the equilibrium point of the dissociation reaction: Keq is defined as the ratio of the concentrations of the products and reactants. Units used to define concentration are Molarity (M) = moles/L pH The concept of p[H] was first introduced by Danish chemist Søren Peder, Lauritz Sørensen at the Carlsberg Laboratory in 1909 revised to the modern pH in 1924 after it became apparent that electromotive force in cells depends on activity rather than concentration of hydrogen ions. In the first papers, the notation had the H as a subscript to the lowercase p, like so: pH. It is unknown what the exact definition of 'p' in pH is. A common definition often used in schools is "percentage". However some references suggest the p stands for “Power”, others refer to the German word “Potenz” (meaning power in German), and still others refer to “potential”. Jens Norby published a paper in 2000 arguing that p is a constant and stands for “negative logarithm”; H then stands for Hydrogen. According to the Carlsberg Foundation pH stands for "power of hydrogen". Other suggestions that have surfaced over the years are that the p stands for puissance (also meaning power, but, then, the Carlsberg Laboratory was French-speaking) or that pH stands for the Latin terms pondus Hydrogenii or potentia hydrogenii. It is also suggested that Sørensen used the letters p and q (commonly paired letters in mathematics) simply to label the test solution (p) and the reference solution (q). Pure (neutral) water has a pH around 7 at 25 °C (77 °F); this value varies with temperature. When an acid is dissolved in water, the pH will be less than 7 (if at 25 °C (77 °F)). When a base, or alkali, is dissolved in water, the pH will be greater than 7 (if at 25 °C (77 °F)). A solution of a strong acid, such as hydrochloric acid, at concentration 1 mol dm−3 has a pH of 0. A solution of a strong alkali, such as sodium hydroxide, at concentration 1 mol dm−3, has a pH of 14. Thus, measured pH values will lie mostly in the range 0 to 14. Since pH is a logarithmic scale, a difference of one pH unit is equivalent to a tenfold difference in hydrogen ion concentration. An approximate measure of pH may be obtained by using a pH indicator. A pH indicator is a substance that changes color around a particular pH value. It is a weak acid or weak base and the color change occurs around 1 pH unit either side of its acid dissociation constant, or pKa, value. For example, the naturally occurring indicator litmus is red in acidic solutions (pH7 at 25 °C (77 °F)) solutions. Universal indicator consists of a mixture of indicators such that there is a continuous color change from about pH 2 to pH 10. Universal indicator paper is simple paper that has been impregnated with universal indicator. ACIDS, BASE AND BUFFERS An Acid is a molecule or ion capable of donating a proton (hydrogen ion H+) (a Brønsted–Lowry acid), or, alternatively, capable of forming a covalent bond with an electron pair (a Lewis acid). The first category of acids are the proton donors, or Brønsted–Lowry acids. In the special case of aqueous solutions, proton donors form the hydronium ion H3O+ and are known as Arrhenius acids. Brønsted and Lowry generalized the Arrhenius theory to include non-aqueous solvents. A Brønsted or Arrhenius acid usually contains a hydrogen atom bonded to a chemical structure that is still energetically favorable after loss of H+. Aqueous Arrhenius acids have characteristic properties which provide a practical description of an acid. Acids form aqueous solutions with a sour taste, can turn blue litmus red, and react with bases and certain metals (like calcium) to form salts. The word acid is derived from the Latin acidus/acēre meaning sour. An aqueous solution of an acid has a pH less than 7 and is colloquially also referred to as 'acid' (as in 'dissolved in acid'), while the strict definition refers only to the solute. A lower pH means a higher acidity, and thus a higher concentration of positive hydrogen ions in the solution. Chemicals or substances having the property of an acid are said to be acidic. Common aqueous acids include hydrochloric acid (a solution of hydrogen chloride which is found in gastric acid in the stomach and activates digestive enzymes), acetic acid (vinegar is a dilute aqueous solution of this liquid), sulfuric acid (used in car batteries), and citric acid (found in citrus fruits). As these examples show, acids (in the colloquial sense) can be solutions or pure substances, and can be derived from acids (in the strict sense) that are solids, liquids, or gases. Strong acids and some concentrated weak acids are corrosive, but there are exceptions such as carboranes and boric acid. The second category of acids are Lewis acids, which form a covalent bond with an electron pair. An example is boron trifluoride (BF3), whose boron atom has a vacant orbital which can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom in ammonia (NH3). Lewis considered this as a generalization of the Brønsted definition, so that an acid is a chemical species that accepts electron pairs either directly or by releasing protons (H+) into the solution, which then accept electron pairs. However, hydrogen chloride, acetic acid, and most other Brønsted-Lowry acids cannot form a covalent bond with an electron pair and are therefore not Lewis acids. Conversely, many Lewis acids are not Arrhenius or Brønsted-Lowry acids. In modern terminology, an acid is implicitly a Brønsted acid and not a Lewis acid, since chemists almost always refer to a Lewis acid explicitly as a Lewis acid. CHEMICAL CHARACTERISTICS Monoprotic acids Monoprotic acids, also known as monobasic acids, are those acids that are able to donate one proton per molecule during the process of dissociation (sometimes called ionization) as shown below (symbolized by HA): HA(aq) + H2O(l) ⇌ H3O+(aq) + A−(aq) Ka Common examples of monoprotic acids in mineral acids include hydrochloric acid (HCl) and nitric acid (HNO3). On the other hand, for organic acids the term mainly indicates the presence of one carboxylic acid group and sometimes these acids are known as monocarboxylic acid. Examples in organic acids include formic acid (HCOOH), acetic acid(CH3COOH) and benzoic acid (C6H5COOH). Polyprotic acids Polyprotic acids, also known as polybasic acids, are able to donate more than one proton per acid molecule, in contrast to monoprotic acids that only donate one proton per molecule. Specific types of polyprotic acids have more specific names, such as diprotic (or dibasic) acid (two potential protons to donate), and triprotic (or tribasic) acid (three potential protons to donate). A diprotic acid (here symbolized by H2A) can undergo one or two dissociations depending on the pH. Each dissociation has its own dissociation constant, Ka1 and Ka2. H2A(aq) + H2O(l) ⇌ H3O+(aq) + HA−(aq) Ka1 HA−(aq) + H2O(l) ⇌ H3O+(aq) + A2−(aq) Ka2 The first dissociation constant is typically greater than the second; i.e., Ka1 > Ka2. For example, sulfuric acid (H2SO4) can donate one proton to form the bisulfate anion (HSO−4), for which Ka1 is very large; then it can donate a second proton to form the sulfate anion (SO2−4), wherein the Ka2 is intermediate strength. The large Ka1 for the first dissociation makes sulfuric a strong acid. In a similar manner, the weak unstable carbonic acid (H2CO3) can lose one proton to form bicarbonate anion (HCO−3) and lose a second to form carbonateanion (CO2−3). Both Ka values are small, but Ka1 > Ka2. BASE In chemistry, bases are substances that, in aqueous solution, release hydroxide (OH−) ions, are slippery to the touch, can taste bitter if an alkali, change the color of indicators (e.g., turn red litmus paper blue), react with acids to form salts, promote certain chemical reactions (base catalysis), accept protons from any proton donor or contain completely or partially displaceable OH− ions. Examples of bases are the hydroxides of the alkali metals and the alkaline earth metals (NaOH, Ca(OH)2, etc In water, by altering the autoionization equilibrium, bases yield solutions in which the hydrogen ion activity is lower than it is in pure water, i.e., the water has a pH higher than 7.0 at standard conditions. A soluble base is called an alkali if it contains and releases OH− ions quantitatively. However, it is important to realize that basicity is not the same as alkalinity. Metal oxides, hydroxides, and especially alkoxidesare basic, and conjugate bases of weak acids are weak bases. PROPERTIES OF BASE General properties of bases include:  Concentrated or strong bases are caustic on organic matter and react violently with acidic substances.  Aqueous solutions or molten bases dissociate in ions and conduct electricity.  Reactions with indicators: bases turn red litmus paper blue, phenolphthalein pink, keep bromothymol blue in its natural colour of blue, and turn methyl orange-yellow.  The pH of a basic solution at standard conditions is greater than seven.  Bases are bitter. Acid/Base Balance pH balance is important to homeostasis of organisms. Homeostasis is the tendency of the body to maintain a balanced internal environment, even when faced with external changes. Such as the body's ability to maintain an internal temperature around 98.6 degrees F, whatever the temperature outside. Examples: Digestion needs acidic environment (pH 2-3) Urine is slightly acidic Blood must stay in neutral range near 7.35 to 7.45 Compounds that dissociate in water and produce cations other than H+ and anions other than OH- are called salts. BUFFERS - Buffers are extremely important in biological systems - Buffers are solutions that resist changes in pH upon addition of acid or base – Examples: Maintaining blood pH. Maintaining physiological pH inside cells Role of buffer - Certain salts, called buffers, can combine with excess hydrogen (H+) or hydroxide (OH-) ions. - Produce substances less acidic or alkaline. - Act like a chemical sponge to soak up excess acid or base, keep pH constant. - Buffers can be “used up”. Once used up, no longer help regulate pH. - - Buffers are vital to maintaining pH in organisms Henderson-Hasselbalch Equation Equation that describes the behavior of weak acids in solution. It allows us to calculate the concentration of an acid and conjugate base