Chemistry of Acids and Bases

CHEMISTRY 122 CHAPTER 17 NOTES

CHAPTER 17 EXERCISES (THESE ARE NOT TO BE TURNED IN, BUT QUIZ QUESTIONS WILL BE TAKEN FROM THESE EXERCISES SO IT IS TO YOUR BENEFIT TO DO THEM): 2, 3, 5, 6, 8, 10, 13, 15, 16, 18, 19, 21, 23, 25, 27, 28, 29, 31, 33, 37, 39, 41, 43, 44, 47, 48, 50, 53, 57, 60, 63, 64, 67, 70, 71, 73, 75, 78, 80, 87, 95, 102, 112

Chemistry of Acids and Bases

Svante Arrheius (1884) developed a theory of acid-base reactions

Arrhenius acid = substance that contains H and produces H+ ions in aqueous solutions

Arrhenius base = substance that contains OH and produces OH- ions in aqueous solutions

Neutralization = H+(aq) + OH-(aq) --„Ñ- H2O(l)

Hydrated hydrogen ion

H+ does not exist independently in aqueous solutions. It is associated with one or more water molecules.

Therefore H+ is said to be hydrated and exists as H+(H2O)n where n is a small integer. The H+ ions are attracted to the Ñ"Ф- part of the water molecule centered on the O atom.

The exact value of n is not known for most solutions but one way to represent it is as H3O+ which is called the hydronium ion.

For our purposes H+(aq) and H3O+ are the same.

Bronsted-Lowry theory expanded on the Arrhenius theory of acids and bases.

Bronsted-Lowry acid = a proton (H+) donor

Bronsted-Lowry base = a proton (H+) acceptor

This means that an acid-base reaction is due to the transfer of H+ from an acid to a base.

Example: HCl(aq) + NaOH(aq) --„Ñ- NaCl(aq) + H2O(l)

Arrhenius description:

Bronsted-Lowry description:

Bronsted-Lowry acid -base reactions can also be described in terms of conjugate acid-base pairs.

The substance left when an acid loses H+ is known as its conjugate base.

Acid Conjugate Base

HCl Cl-

HNO3 NO3-

CH3COOH CH3COO-

H2SO4 HSO4-

HSO4- SO42-

H3O+ H2O

H2O OH-

The substance produced when a base gains H+ is known as its conjugate acid.

Base Conjugate acid

OH- H2O

H2O H3O+

CO32- HCO3-

HCO3- H2CO3

NH3 NH4+

Reactions can be described in terms of what the conjugate acid-base pairs are.

HCl(aq) + NaOH(aq) --„Ñ- NaCl(aq) + H2O(l)

The stronger the acid the weaker its conjugate base and the stronger the base the weaker its conjugate acid.

Water can act like an acid in the presence of a base or like a base in the presence of an acid.

NH3(aq) + H2O(l) „І„Ñ- NH4+(aq) + OH-(aq)

CH3COOH(aq) + H2O(l) „І„Ñ- H3O+(aq) + CH3COO-(aq)

Autoionization of water

Autoionization = reaction in which water slightly ionizes to produce hydrated hydrogen and hydroxide ions

H2O + H2O „І„Ñ- H3O+(aq) + OH-(aq)

Or H2O „І„Ñ- H+(aq) + OH-(aq)

One water molecule donates a H+ ion to the other water molecule

Water is said to be amphiprotic because it can both accept and donate protons.

Amphoterism is a property of a substance to act as either an acid or a base.

Amphiprotism is a more specific type of amphoterism in which transfer of H+ is involved.

Some of the insoluble metal hydroxides are amphoteric because they dissolve in both acids and bases.

Example: Al(OH)3

Al(OH)3 + 3HCl--„Ñ- AlCl3 + 3H2O

Al(OH)3 + NaOH --„Ñ- NaAl(OH)4(aq)

Autoionization of Water

Pure water ionizes to a very slight extent:

H2O(l) + H2O(l) „І--„Ñ- H3O+(aq) + OH-(aq)

Kw = ion product for water = [H3O+][OH-] = 1.0x10-14 at 250C

For pure water [H3O+] = [OH-] = 1.0x10-7

The Kw expression [H3O+][OH-] = 1.0x10-14 at 250C holds for dilute aqueous solutions as well as for pure water

The value of Kw does vary with temperature but unless otherwise stated assume that the temperature is 250C

Example: Calculate the concentrations of OH- and H3O+ in a solution of 0.00050 M Ca(OH)2(aq)

Acidic solution

Basic solution

Neutral solution

pH and pH scales

pH = -log[H3O+] or [H3O+] = 10-pH

pOH = -log[OH-] or [OH-] = 10-pOH

pKw = -log Kw = pH + pOH = 14.00 at 250C

Example: Determine the [H3O+], [OH-], pH, and pOH for an aqueous solution of 0.00035 M HCl