Chemical reactions Chemical reactions are classified into: 1 1- Irreversible reaction, which proceeds in one direction e.g. reaction of acids with alkalies to form salts and water Acid + alkali salt + water 2- Reversible reaction, which proceeds in both directions. Most of the reactions in biological systems are of this type. In reversible reaction, the reactants (A and B) react together to give the products (C and D) and vise versa i.e. the reversible reaction does not proceed to completion in either direction. A + B C + D Law of mass action The law of mass action states that “In reversible reactions the velocity of the reaction at a certain temperature is directly proportional to the molecular concentration of the reacting substances”. If we have a reversible reaction in which A and B are the reacting substances, C and D are the products, V1 is the velocity of the reaction in the forward direction and V2 is the velocity of the reaction in the reverse direction. V1 A + B C + D V2 According to the law of mass action, the velocity of the reaction in the forward direction (V1) is directly proportional to the molecular concentration of A and B. V1 ∝ [A] × [B] ∴ V1 = [A] × [B] × K1 where K1 is a constant. Also, the velocity of the reaction in the reverse direction (V2) is directly proportional to the molecular concentration of C and D. V2 ∝ [C] × [D] ∴ V2 = [C] × [D] × K2 where K2 is a constant. At equilibrium, the velocity of the reaction in the forward direction (V1) is equal to the velocity of the reaction in the reverse direction (V2). V1 = V2 ∴ [A] × [B] × K1 = [C] × [D] × K2 [C] × [D] K1 ∴ = = Keq [A] × [B] K2 2 Where Keq is the equilibrium constant. So, in reversible reactions at equilibrium the product of molecular concentrations of the resulting substances divided by the product of molecular concentrations of the reacting substances is a constant called equilibrium constant (Keq). Water Water is the liquid in which all known life forms exist. By far the most common molecule in the living organisms is water. Since most of the molecular interactions in the body take place in an aqueous environment, we must always consider the effects of water on biomolecules. Water molecule (H2O) is formed of one oxygen atom covalently linked to 2 hydrogen atoms. Since O has a much higher electron affinity (electronegativity) than H, it attracts the electrons away from the hydrogen atoms, resulting in a partial negative charge on the oxygen and a partial positive charge on each of the hydrogen atoms. This creates a positive pole near the hydrogen atoms and a negative one on the opposite side. So, water is best described as a polar − compound. One end, or pole, of the 2δmolecule has a partial positive charge delta+ (δ+) on the hydrogen atoms, and the other end has a partial negative charge (2δ-) on the oxygen atom. + + δ δ water molecule Water's polarity gives it many of the life-critical properties. Its polarity is responsible for hydrogen bonding which makes water molecules cohesive (sticking together), creates surface tension and keeps water liquid. Likewise its polarity makes it a wonderful solvent for ionic molecules. Biomolecules interacting with water fall into three categories: 1. Hydrophilic (water loving) molecules which readily form hydrogen bonds with water and therefore dissolve in water. These are substances that are ionic or can engage in hydrogen bonding with water molecules e.g. sodium chloride and glucose. 2. Hydrophobic (water hating)molecules that can not form hydrogen bond with water and can not form ionic bonds. These are nonionic and nonpolar substances that are unable to engage in attractive interactions with water molecules e.g. neutral fat and oils. 3. Amphipathic molecules which have both a hydrophobic and a 3 hydrophilic portion. In aqueous solutions these compounds spontaneously aggregate with the hydrophilic portion exposed to the water and the hydrophobic portion in the interior where it is shielded from the water molecules e.g. lipoproteins and phospholipids. Dissociation (Ionization) of water H2O can act as both a proton donor (acid) and proton acceptor (base) for itself. A proton can be transferred from one water molecule to another, resulting in the formation of one hydroxyl ion (OH-) and one + hydronium ion (H3O ). + - H2O + H2O H3O + OH acid base acid base This is called the autoionization or dissociation of water. This equilibrium can also be expressed as: + - H2O H + OH although this is misleading because free protons (H+) are rarely present in pure water. At equilibrium : + - [H ] X [OH ] Keq = [H2O] Where: -16 Keq is the equilibrium constant. It is a constant that equals 1.8 x 10 [H+] is hydrogen ion concentration in mole/liter (mol/L) - [OH ] is the hydroxyl ion concentration in mol/L. [H2O] is the concentration of water in mol/L. Each liter of pure water weighs 1000 gram and the molecular weight of water is 18. So, in pure water [H2O] = 1000/18 =55.5 M By substitution: + - [H ] X [OH ] Keq = 55.5 Rearranging gives: + - Keq X 55.5 = [H ][OH ] Keq X 55.5 is another constant called ion product of water (Kw) + - ∴ Kw = Keq X 55.5 = [H ][OH ] -16 + - ∴ Kw = 1.8 x 10 X 55.5 = [H ][OH ] -14 + - ∴ Kw = 1 x 10 = [H ][OH ] Pure water which is neutral has the same concentration of H+ as OH- ions + - -7 ∴ [H ] = [OH ] = 10 4 As the ion-product constant of water (Kw) always remains constant at equilibrium, consequently: • If the concentration of either H+ or OH- increases in solution, then the other must decrease to compensate. • An acid causes an increase in the concentration of H+ ions and an alkali causes an increase in the concentration of OH- ions. • So, in acidic solutions [H+] is more than [OH-], while in alkaline solutions [OH-] is more than [H+]. Hydrogen ion concentration is used to express acidity or alkalinity in aqueous solution as follows: + - -7 • In neutral solution [H ] = [OH ] = 10 + - + -7 • In acidic solution [H ] > [OH ] and [H ] > 10 while [OH-] < 10-7 - + + -7 • In alkaline solution [OH ] > [H ] and [H ] < 10 while [OH-] > 10-7 pH pH is a measure of the concentration of hydrogen ions. It is defined as the negative logarithm of hydrogen ions concentration in mol/L. pH = -log[H+] As mentioned before, the [H+] in pure water at 25 οC is 10-7; therefore the pH of pure water is: pH = -log(10-7) pH = - (-7) pH = 7 pH ranges from 0 to 14 pH = 0, the most acidic (1 molar HCl) pH < 7, solution is acidic pH = 7, solution is neutral pH > 7, solution is alkaline pH = 14, the most alkaline (1 molar NaOH) A lower pH always means a higher concentration of H+. The most acidic fluid in the body is gastric juice, which has a pH range 1-2. While the most alkaline juice in the body is the pancreatic juice with pH about 8. Milk (pH = 6.6 - 6.9) and urine (pH=6) are slightly acidic while blood, saliva and aqueous humor of the eye are slightly alkaline (pH = 7.4). Blood pH is kept within very narrow range (7.37 - 7.43) by blood buffers. Also, the negative logarithmic scale is useful for measuring hydroxyl 5 ion concentration [OH-]: pOH = - log [OH-] As discussed in dissociation of water: + - -14 Kw = [H ][OH ] = 1 x 10 [H+][OH-] = 1 x 10-14 By logarithmic transformation, we obtain the following useful expressions: -14 pH + pOH = - log 10 = 14 This means that the sum of pH and pOH is always 14 and if either of them increases then the other must decrease to compensate. The Relationship Between pH and pOH pH pOH [H+] mol/L [OH-] mol/L 0 14 1.0 10-14 2 12 0.01 10-12 4 10 0.0001 10-10 6 8 10-6 10-8 8 6 10-8 10-6 10 4 10-10 0.0001 12 2 10-12 0.01 14 0 10-14 1.0 N.B.: The pH scale is logarithmic scale representing the concentration of H+ ions in a solution. The pH scale is logarithmic, so pH difference of 1 means a tenfold difference in the relative concentration of hydrogen ions. Measurement of pH pH can be measured by colorimetric and electrometric methods 1- Colorimetric method It depends on indicators, which are weak organic acids that change their colours on ionization i.e. they have certain colour in the non- ionized state and another colour in the ionized state. For example the use of phenolphthalein solution or litmus paper to measure pH. 2- Electrometric method pH is measured by pH meter, which uses hydrogen sensitive electrode that measures hydrogen ion concentration in the solution. 6 Buffers They are solutions that resist appreciable changes in pH when an acid or alkali is added to it. Biological systems use buffers to control pH. Buffers may be formed of one of the following mixtures: 1- Weak acid and its salt with strong base e.g. carbonic acid and sodium bicarbonate. 2- Weak base and its salt with strong acid e.g. ammonium hydroxide and ammonium chloride. Mechanism of buffer action In the case of bicarbonate buffer (NaHCO3 / H2CO3) If strong alkali as sodium hydroxide (NaOH) is added it will react with the carbonic acid forming sodium bicarbonate, which is a weak alkali. So there is no appreciable change in the pH. NaOH + H2CO3 NaHCO3 + H2O If strong acid as hydrochloric acid (HCl) is added it will react with the sodium bicarbonate forming carbonic acid, which is a weak acid.