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1 INTRODUCTION An acid is a substance that can yield(donate) a hydrogen ion(H+) or hydronium ion when dissolved in water. Eg: HCl, CH3COOH. A base is a substance that can yield hydroxyl ions(OH-). Eg: NaOH. pH: Negative logarithm of H+ concentration. [ pH = -log H+ (mol/L) ] pK: Negative logarithm of dissociation (ionization) constant. Introduction of Acids & Bases They are everywhere.. In your food In your house EVEN IN YOU!!!!! Objectives • • • • Define acids and bases. Know the uses of acids and bases. Understand what is pH and pH scale. Know the use of Henderson-Hasselbalch equation. • Know buffer system of the body. What is an acid? An acid is a substance that can yield(donate) a hydrogen ion(H+) or hydronium ion (H3O+) when dissolved in water. Example: HCl, CH3COOH. It is sour in taste and turns blue litmus paper red The sting of Ant contain formic acid( HCOOH) Uses of Acids • Acetic Acid = Vinegar • Citric Acid = lemons, limes, & oranges. It is in many sour candies such as lemonhead & sour patch. • Ascorbic acid = Vitamin C which your body needs to function. • Sulfuric acid is used in the production of fertilizers, steel, paints, and plastics. • Car batteries What is a base? A base is a substance that can yield hydroxyl ions(OH-). Another word for base is alkali. Example: NaOH, KOH, Mg(OH)2. It is bitter in taste and turns red litmus paper to blue Uses of Bases • Bases give soaps, ammonia, and many other cleaning products some of their useful properties. • The OH- ions interact strongly with certain substances, such as dirt and grease. • Chalk and oven cleaner are examples of familiar products that contain bases. • Your blood is a basic solution. pH 0 pH is generally defined as the negative logarithm of the hydrogen ion activities (concentration) expressed over 14 orders of magnitude. pH = -log10 [H+] 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Acidic Basic Neutral [H+]>[OH-] [H+] = [OH-] [OH-]>[H+] pH • pH is a measure of how acidic or basic a solution is. • The pH scale ranges from 0 to 14. • Acidic solutions have pH values below 7 • A solution with a pH of 0 is very acidic. • A solution with a pH of 7 is neutral. • Pure water has a pH of 7. • Basic solutions have pH values above 7. pH Scale Base: proton acceptor Devised by Sorenson (1902) [H+] can range from 1M and 1 X 10-14M using a log scale simplifies notation pH = -log [H+] Neutral pH = 7.0 Acid: proton donor pH p Scale A pH below the reference range (<7.34 ) is referred to as acidosis, whereas a pH above the reference range (>7.44) is referred to as alkalosis. Dissociation constant - pK • pK indicates the strength of acids and bases and their ability to dissociate in water. • Strong acids have pK values < 3, whereas strong bases have pK > 9. • pH = pK when the protonated and unpronated (dissociated) forms are in equal concentration. Henderson Hesselbach equation • The Henderson Hesselbach equation expresses acid-base relationships in a mathematical formula: cA pH = pK + log------cHA • WhereA- = proton acceptor(eg. HCO3-), • HA = proton donor or weak acid (eg.H2CO3), • pK = pH at which there is an equal concentration of protonated and unprotonated species. SOURCES OF HYDROGEN IONS Hydrogen ions[H+] are produced in the body as a result of metabolism, particularly from the oxidation of sulphur containing amino acids of proteins ingested as food. The total amount of H+ produced each day in this way is of the order of 60mmol.If all this was to be diluted in the extracellular fluid, [H+] would be 100, 000 times more than normal! This just does not happen. because of effective buffer systems present in the body and the normal ventilatory and excretory function of lungs and kidneys. Regulatory Mechanisms That Maintain pH Buffer systems of the blood. React very rapidly (less than a second) Kidney buffer system. Concentration of bicarbonate controlled by the kidneys. Reacts slowly (minutes to hours) Can reabsorb bicarbonate Can synthesize bicarbonate Respiratory buffer system. Concentration of carbonic acid controlled by the lungs. Reacts rapidly (seconds to minutes) Buffers • Buffers are solutions in which the pH remains relatively constant, even when small amounts of acid or base are added. – Contain a weak acid (CH3COOH) and its salt(CH3COONa); H2CO3 & NaHCO3 How buffers work • Equilibrium between acid and base. • Example: Acetate buffer CH3COOH CH3COO- + H+ • If more H+ is added to this solution, it simply shifts the equilibrium to the left, absorbing H+, so the [H+] remains unchanged. • If H+ is removed (e.g. by adding OH-) then the equilibrium shifts to the right, releasing H+ to keep the pH constant. Buffer systems of the body Blood Buffer System Kidney Buffer System Respiratory Buffer System Blood Buffer Systems • Plasma Buffer Systems-bicarbonate buffer system -phosphate buffer system -protein buffer system • RBC buffer systems-haemoglobin buffer system -bicarbonate buffer system Buffer Systems of body (contd) Kidney buffer systems• Phosphate buffer systems- most important • Ammonia buffer system • Bicarbonate buffer system Respiratory buffer system• Bicarbonate buffer system