Acids Bases pH PDF

WATER, pH, ACIDS, BASES AND BUFFER BY Oluwadare AGUNBIADE  WATER  Water as a reactant.  Water is a direct participant in many biochemical reactions. When a biomolecule is split apart by water, the reaction is called a hydrolysis reaction...

WATER, pH, ACIDS, BASES AND BUFFER BY Oluwadare AGUNBIADE  WATER  Water as a reactant.  Water is a direct participant in many biochemical reactions. When a biomolecule is split apart by water, the reaction is called a hydrolysis reaction.  When a biomolecule is formed from two components with the elimination of water, the reaction is called a condensation reaction.  Water molecules have a very slight tendency to ionize to a hydrogen ion and a hydroxide ion  H2O  H+ + OH-.  Actually, free protons (H+) do not exist in solution, and instead, hydrogen ions formed in water are immediately hydrated to form hydronium ions (H3O+). Another important characteristic of water… Water can form acids and bases Dissociation of Water Neutral water has equal amounts of H+ and OH - Acids: Excess of H+ in aqueous solution Bases: Excess of OH- in aqueous solution Acids & bases neutralize each other.  The properties of Water Water is the predominant chemical component of all living organisms. Most chemical reactions in the cell take place in aqueous environment. Hydrogen bonds hold the oxygen and hydrogen atoms together in a water molecule. The oxygen of water is very electronegative, while hydrogen is electropositive, as a result water is dipolar and exhibit slight tendency to dissociate. Biological importance of water A molecule with electrical charge distributed unequally about its structure is referred to as a dipole. H 3O = H+ + OH- The strong dipole and high dielectric constant of water enables it to dissolve large quantities of charged compounds. Presence of hydrogen bond also enables water to dissolve many organic molecules that contain functional groups. Water provide environment for macromolecules to achieve stable structure in solution  ACIDS An acid is any ionic compound that releases hydrogen ions (H+) in solution. Weak acids have a sour taste. Strong acids are highly corrosive (So don’t go around taste-testing acids.) Examples:  Ascorbic acid (C H O , Vitamin C) 6 8 6  Citric acid (C6H8O7, a weak organic acid in citrus fruits)  Phosphoric acid (H3PO4, in pop…this stuff is also used to remove rust… hmmm) Acid is a compound that dissociates in aqueous solution to produce proton (H+) and a conjugate base (A-). HA = H + A Types of Acids Strong and Weak Acids A strong acid dissociate completely while a weak acid dissociate partially in solution e.g. HCl H+ + Cl- Example of strong acid H3PO4 H+ + H2PO4- Example of Weak acids Based on the number of ionizable proton Monoprotic e.g. HCl Diprotic e.g. H2SO4 Polyprotic e.g. H3PO4  BASE A base is an ionic compound that releases hydroxyl ions (OH-) in solution. Bases are also called alkaline substances. Some general properties of bases include: Taste: Bitter taste (opposed to sour taste of acids and sweetness of aldehydes and ketones). Touch: Slimy or soapy feel on fingers. Reactivity: Strong bases are caustic on organic matter, react violently with acidic substances. Examples:  Sodium hydroxide, NaOH, of lye or caustic soda used in oven cleaners.  Magnesium hydroxide, Mg(OH)2, also known as milk of magnesia, a weak base used in antacids and laxatives.  Dissociation Constants of Weak Acids and Bases (I) Unlike strong acids (HCl, H2SO4) and strong bases (NaOH, KOH), weak acids and bases (e.g., acetic acid (CH3COOH) and amino acids) are not completely ionized when dissolved in water. Instead, depending upon the pH, both the proton donor species (conjugate acid) and proton acceptor species (the conjugate base) occur together in solution. Each weak acid has a characteristic tendency to lose its proton in aqueous solution. The stronger the acid, the greater the tendency to ionize. The tendency of any conjugate acid (HA) to lose a proton and form its conjugate base (A -) is defined by the equilibrium constant (Keq) for the reversible reaction HA  H+ + A- For which Keq = [H+][A-]/[HA] = Ka. Equilibrium constants for ionization reactions are usually called acid dissociation constants, Ka. (Continued on the next slide) Buffers Defination: Buffer is a solution that resists change in pH when acid or base is added. When either acid or base is added, there will be temporary change in pH, but the pH is quickly restored. Components of Buffer A buffer contains a weak acid and its conjugate base. Examples of buffer solutions are: Acetate buffer (acetic acid and acetate salt), bicarbonate buffer (carbonic acid and bicarbonate salt). Buffers Definition: a solution that resists change in pH  Typically a mixture of the acid and base form of a chemical  Can be adjusted to a particular pH value Why use them?  Enzyme reactions and cell functions have optimum pH’s for performance  Important anytime the structure and/or activity of a biological material must be maintained Factors in choosing a buffer Be sure it covers the pH range you need Generally: pKa of acid ± 1 pH unit Consult tables for ranges or pKa values Be sure it is not toxic to the cells or organisms you are working with. Be sure it would not confound the experiment (e.g. avoid phosphate buffers in experiments on plant mineral nutrition). Buffers of Physiological Importance PHOSPHATE BUFFER BICARBONATE BUFFER PROTEIN BUFFER How buffers regulate the pH of solution If hydrogen ions are added to a buffer solution, the conjugate base react with the excess hydrogen ions to form the acid. If OH- ions are added, they react with the acid present in the buffer to produce water and conjugate base. Biological importance of buffer Body fluids such as blood, cerebrospinal fluid, saliva etc. have constant pH under normal physiological conditions. This is possible due to the presence of buffer in these fluids. Hydrogen and hydroxyl ions are constantly added to the body fluids as products of metabolism. Approximate pH of some body fluids are : blood - 7.4, cerebrospinal fluid - 7.30 to 7.50, saliva – 6.5 to 7.5. PREPARATION OF BUFFER A buffer can resist pH changes if the pH is at or near a weak acid pK value. Buffer range: the pH range where maximum resistance to pH change occurs when adding acid or base. It is = +1 pH from the weak acid pK  If pK is 4.8 the buffering range is 3.8 5.8 Why?  The Henderson-Hasselbalch Equation (I) The shape of the titration curves for all weak acids is described by the Henderson-Hasselbalch equation. This equation simply restates the expression for the ionization of a weak acid, HA, as shown in the following derivation: Ka = [H+] [A-] / [HA] After solving for [H+], the equation becomes [H+] = Ka [HA] /[A-]. Then take the negative logarithm of both sides -log [H+] = -log Ka - log [HA] / [A-] After substituting pH and pKa into the equation, it becomes pH = pKa - log [HA] / [A-] Finally, the Henderson-Hasselbach equation results after mathematical manipulation of the log term on the right, namely pH = pKa + log [A-] / [HA] Worked Example THANK YOU