Chapters 14 and 15

Aqueous solutions have a sour taste. Acids change the color of acid-base indicators. Some acids react with active metals to release hydrogen gas, H 2. Acids react with bases to produce salts and water (i.e. a neutralization reaction) Some acids conduct electric current.

Aqueous solutions taste bitter. Bases change the color of acid-base indicators. Dilute aqueous solutions of bases feel slippery (ex. soap). Bases react with acids to produce salts and water. (i.e. neutralization reactions) Bases conduct electric current.

Theories of acids and bases

Strong acids and bases are those that ionize completely in aqueous solution. They are also typically strong electrolytes. Weak acids and bases are those that do not ionize completely in aqueous solution. They are also typically weak electrolytes.

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Conjugate acid-base reactions HF (aq) + H 2 O (l)  F - (aq) + H 3 O + (aq) c.a.1 c.b.2 c.b.1 c.a.2

Amphoteric substances are those that can react as either an acid or a base.

Molecules containing – OH groups (hydroxyl groups) can be acidic or amphoteric. The more oxygen atoms, more polar  more acidic

Oxyacids of chlorine

Acid-dissociation equilibrium constant (K a ) - A measure of the relative strength of an acid. For the generic acid dissociation reaction with water, HA (aq) + H 2 O (l)  H 3 O + (aq) + A - (aq) K a = [H 3 O + ][A - ] [HA] As the K a value of an acid increases, so does the strength of the acid. By definition: strong acid: K a > 1 weak acid: K a < 1

strong acid + water  weaker acid + weaker base (water acts as a strong base) weak acid + water ATTEMPTS TO CONVERT TO  stronger acid + stronger base (water acts as a weak base) BUT the reaction cannot naturally proceed in this direction. This is why strong acids dissociate nearly completely whereas weak acids dissociate only slightly.

The larger the value of K a, the stronger is the acid. K a is a better measure of the strength of an acid than pH because adding more water to the acid solution will not change the value of the equilibrium constant K a, but it will change the H + ion concentration on which pH depends.

Base dissociation constant or equilibrium constant, K b For the reaction in which the Arrhenius base, BOH, dissociates to form the ions OH - and B + : BOH  OH - + B + For a Brønsted-Lowry base: B + H 2 O  BH + + OH - K b = [BH + ][OH - ] [B]

K b provides a measure of the strength of a base 1.if K b is large, the base is largely dissociated so the base is strong 2.if K b is small, very little of the base is dissociated so the base is weak.

H 2 O(l) + H 2 O(l)  H 3 O + (aq) + OH - (aq) K w = [H 3 O + ] [OH - ] = 1 x Measurements at 25˚C show that the [H 3 O + ] and [OH - ] are each 1 x M

When something is acidic [H 3 O + ] > [OH - ]; [H 3 O + ] >1 x M and [OH - ] < 1 x M When something is basic [H 3 O + ] < [OH - ]; [H 3 O + ] 1 x M [H 3 O + ] and [OH - ] are inversely proportional

pH means proportion of H + ion scale from 0 – 14 pH = -log[H 3 O + ]pOH = -log[OH - ] As K w = 1 x we can conclude that pH + pOH = 14

pH[H 3 O + ]pOH[OH - ] x M 11 x M131 x M 21 x M121 x M 31 x M111 x M 41 x M101 x M 51 x M91 x M 61 x M81 x M 71 x M7 81 x M61 x M 91 x M51 x M 101 x M41 x M 111 x M31 x M 121 x M21 x M 131 x M11 x M 141 x M01 Acidic pH 0-7 [H 3 O + ] > [OH - ]; [H 3 O + ] > 1 x M [OH - ] < 1 x M pH = 7 NEUTRAL [H 3 O + ] = [OH - ] 1 x M Basic pH 7-14 [H 3 O + ] < [OH - ]; [H 3 O + ] < 1 x M [OH - ] > 1 x M

pH of common items

compounds whose colors are sensitive to pH. Indicators change colors because they are either weak acids or weak bases. In basic solution HIn  H + + In - In acidic solution

Indicators come in many different colors. The pH range over which an indicator changes color is called its transition interval.

Blue litmus paper turns red under acidic conditions and red litmus paper turns blue under basic (i.e. alkaline) conditions

A pH meter determines the pH of a solution by measuring the voltage between the two electrodes that are placed in the solution.

DR between strong acids and bases always form water and a salt (an ionic compound formed from the cation of the base and the anion of the acid). The net ionic equation is: H 3 O + (aq) + OH - (aq)  2H 2 O (l)

Titration is the controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concentration.

The point at which the two solutions used in a titration are present in chemically equivalent amounts is the equivalence point. The point in a titration at which an indicator changes color is called the end point.

The solution that contains the precisely known concentration of a solute is known as a standard solution. A primary standard is a highly purified solid compound used to check the concentration of the known solution in a titration. Burettes for acid-base titrations

Titration curve strong acid + strong base

Titration curve weak acid + strong base

Titration curves chemguide.co.uk

solutions composed of a weak acid or base and its salt (ex. CH 3 COOH and NaCH 3 COO) and can withstand large changes in pH. Buffer capacity is the amount of acid or base a buffer solution can absorb without a significant change in pH. This has very important biological implications.

Buffer examples Ex. Blood maintains a pH of 7.4. The conjugate acid-base pair that acts as a buffer system is H 2 CO 3 and HCO 3 -. When there is excess acid present H 3 O + + HCO 3 -  H 2 O + H 2 CO 3 When there is excess base present OH - + H 2 CO 3  H 2 O + HCO 3 -