1. Hydrolysis of Salts and pH of Buffer Solutions Experiment 24 Page 257 Dr. Scott Buzby Ph.D.

2. Objectives • Learn about the concept of hydrolysis • Acids • Bases • Hydrolysis • Gain a familiarity with acid-base indicators • Learn about the behavior of buffer solutions • Henderson-Hasselbalch equation • Safety Note!!! • Strong acids and bases attack living tissue and cause serious burns wear proper PPE

3. Acids • An acid (from the Latin acidus meaning sour) is traditionally considered any chemical compound that, when dissolved in water, gives a solution with a pH less than 7.0 • Arrhenius acids • a substance that increases the concentration of hydronium ions, H3O+, when dissolved in water • Brønsted-Lowry acids • A Brønsted-Lowry acid (or simply Brønsted acid) is a species that donates a proton from a Brønsted-Lowry base • Lewis acids • A Lewis acid is a species that accepts a pair of electrons from another species; in other words, it is an electron pair acceptor

4. Strength of Acids • Strong Acids • A strong acid is an acid that dissociates completely in an aqueous solution by losing one proton • Weak Acids • A weak acid is an acid that dissociates incompletely and does not release all of its hydrogens in a solution (i.e. it does not completely donate all of its protons) • While strong acids are generally assumed to be the most corrosive, this is not always true. The carboranesuperacid which is one million times stronger than sulfuric acid, is entirely non-corrosive, whereas the weak acid hydrofluoric acid (HF) is extremely corrosive and can dissolve glass and all metals except iridium

5. Examples of Acids • Strong Acids • HCl - hydrochloric acid • HNO3 - nitric acid • H2SO4 - sulfuric acid • HBr - hydrobromic acid • HI - hydroiodic acid • HClO4 - perchloric acid • Weak Acids • HCH3O2 – acetic acid • HF – hydrofluoric acid • Most organic acids • NH4+ - ammonium

6. Bases • A base is traditionally considered any chemical compound that, when dissolved in water, gives a solution with a pH greater than 7.0 • Arrhenius bases • a substance that increases the concentration of hydroxide ions, OH-, when dissolved in water • Brønsted-Lowry bases • A Brønsted-Lowry base (or simply Brønsted base) is a species that accepts a proton from a Brønsted-Lowry acid • Lewis bases • A Lewis base is a species that donates a pair of electrons to another species; in other words, it is an electron pair donor

7. Strength of Bases • Strong Base • A strong base is a base which hydrolyzes completely and is able to deprotonate very weak acids in an acid-base reaction, common examples of strong bases are the hydroxides of alkali metals and alkaline earth metals like NaOH and Ca(OH)2 • Weak Base • A weak base is a chemical base that does not ionize fully in an aqueous solution or as Brønsted-Lowry bases are proton acceptors, a weak base may also be defined as a chemical base in which protonation is incomplete

8. Examples of Bases • Strong Bases • Potassium hydroxide (KOH) • Barium hydroxide (Ba(OH)2) • Caesium hydroxide (CsOH) • Sodium hydroxide (NaOH) • Strontium hydroxide (Sr(OH)2) • Calcium hydroxide (Ca(OH)2) • Lithium hydroxide (LiOH) • Rubidium hydroxide (RbOH) • Magnesium hydroxide (Mg(OH)2) • Weak Bases • Alanine, C3H5O2NH2 • Ammonia, NH3 • Methylamine, CH3NH2 • Pyridine, C5H5N • Other weak bases are essentially any bases not on the list of strong bases

9. Hydrolysis of Salts • A hydrolysis reaction is the reaction of a ion with water • Anions of weak acids (C2H3O2-) react with water to form OH-, raising the pH of the solution • Cations of weak bases (NH4+) react with water to form H+, lowering the pH of the solution

10. Summary of Hydrolysis Behavior • Salt of a strong acid and a strong base • Neither ion hydrolyzes, and the solution has a pH of 7 (Neutral) • Salt of a strong acid and a weak base • The cation hydrolyzes, forming H+ and the solution has a pH < 7 (Acidic) • Salt of a weak acid and a strong base • The anion hydrolyzes, forming OH- and the solution has a pH > 7 (Basic) • Salt of a weak acid and a weak base • Both ions hydrolyze, and the pH of the solution is determined by the relative extent to which each ion hydrolyzes

11. Buffer Solutions • A buffer solution is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid • It has the property that the pH of the solution changes very little when a small amount of strong acid or base is added to it • For example the addition of 0.036g of HCl to 1L of water causes the pH to drop from 7.0 to 3.0 • Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical/biological applications • Human blood has a pH ≈ 7.4, changes of as little as ± 0.05 can be dangerous or even fatal

12. Theory of Buffer Solutions • In a solution there is an equilibrium between a weak acid, HA, and its conjugate base, A- • When hydrogen ions (H+) are added to the solution, equilibrium moves to the left, as there are hydrogen ions (H+ or H3O+) on the right-hand side of the equilibrium expression • When hydroxide ions (OH-) are added to the solution, equilibrium moves to the right, as hydrogen ions are removed in the reaction • Thus, in both cases, some of the added reagent is consumed in shifting the equilibrium in accordance with Le Chatelier's principle and the pH changes by less than it would if the solution were not buffered

13. pH of a Buffer Solution • Henderson-Hasselbalch equation is used to determine the pH of a buffer solution • This equation is convenient for preparing buffer solutions because you can neglect the amounts of the acid and base that ionize and use the initial concentrations of the acid and the conjugate base

14. Procedure • Hydrolysis of Salts – Page 265 • pH of Buffer Solutions – Page 266 • Preparation of a buffer solution • Operation of the pH meter • Effect of acid and base on the buffer pH

15. Due Next Week • Report Sheet – Pages 269-273 • Questions – Page 272 • Pre-Lab Experiment 25 – Page 281