Experiment 24Page 257

Dr. Scott Buzby Ph.D.

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

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

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 carborane superacid 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

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

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

Strong BaseA 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 BaseA 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

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

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

Salt of a strong acid and a strong baseNeither 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 baseThe 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

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

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

AOHOHHA 32

OHOHH 2

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

][

][log

acidweak

baseconjugatepKpH a

Hydrolysis of Salts – Page 265

pH of Buffer Solutions – Page 266Preparation of a buffer solutionOperation of the pH meterEffect of acid and base on the buffer pH

Report Sheet – Pages 269-273

Questions – Page 272

Pre-Lab Experiment 25 – Page 281