14.1 The Nature of Acids and Bases

• Acids were first recognized as a class of substances that taste sour

  • The first person to recognize the essential nature of acids and bases was Svante Arrhenius.

  • Arrhenius postulated that acids produce hydrogen ions in an aqueous solution, while bases produce hydroxide ions.

  • At the time, the Arrhenius concept of acids and bases was a major step forward in quantifying acid-base chemistry, but this concept is limited because it applies only to aqueous solutions and allows for only one kind of base

• Conjugate base: Everything that remains of the acid molecule after a proton is lost

• The conjugate acid is formed when the proton is transferred to the base

14.2 Acid Strength

• ****A strong acid yields a weak conjugate base—one that has a low affinity for a proton

• A weak acid is one for which the equilibrium lies far to the left

  • The weaker the acid, the stronger its conjugate base.

  • Dilution of a weak acid increases its percent dissociation

• The common strong acids are sulfuric acid, hydrochloric acid, nitric acid, and perchloric acid

  • Perchloric acid can explode if handled improperly

  • Most acids are oxyacids, in which the acidic proton is attached to an oxygen atom

• Water is the most common amphoteric substance.

• H2O is never included because it is assumed to be constant

• Kw is the ion-product constant for water

14.3 The pH Scale

• The pH scale is a compact way to represent solution acidity

• The rule is that the number of decimal places in the log is equal to the number of significant figures in the original number

• The main reason that acid-base problems sometimes seem difficult is that a typical aqueous solution contains many components

• Since pH is a log scale, the pH changes by 1 for every 10-fold change in H+

14.4 Calculating the pH of Strong Acid Solutions

• Container labels indicate the substance(s) used to make up the solution but do not necessarily describe the solution components after dissolution.

• Always write the major species present in the solution

14.5 Calculating the pH of Weak Acid Solutions

• First, always write the major species present in the solution

• Typically, the Ka values for acids are known to have an accuracy of only about _+5%

• A mixture of three acids might lead to a very complicated problem

• However, the situation is greatly simplified by the fact that even though HNO2 is a weak acid, it is much stronger than the other two acids present

• To avoid clutter we do not show the units of concentration in the ICE tables. All terms have units of mol/L

• It is often useful to specify the amount of weak acid that has dissociated in achieving equilibrium in an aqueous solution

• For a given weak acid, the percent dissociation increases as the acid become more dilute

• The more dilute the weak acid solution, the greater is the percent dissociation

14.6 Bases

• Strong bases are hydroxide salts, such as NaOH and KOH

• The alkaline earth hydroxides are also strong bases

• The alkaline earth hydroxides are not very soluble and are used only when the solubility factor is not important

• Calcium hydroxide, Ca(OH)2, often called slaked lime**,** is widely used in industry because it is inexpensive and plentiful

  • Slaked lime is also widely used in water treatment plants for softening hard water, which involves the removal of ions

• A base does not have to contain hydroxide ions

• Bases such as ammonia typically have at least one unshared pair of electrons that is capable of forming a bond with a proton

14.7 Poly protic acids

• **** A polyprotic acid has more than one acidic proton

• Polyp protic acids dissociate one proton at a time

• For a typical polyprotic acid in the water, only the first dissociation step is important in determining the pH

  • Each step has a characteristic Ka value

  • Typically for weak polyprotic acid, Ka1 7 Ka2 7 Ka3

• Sulfuric acid is unique

  • It is a strong acid in the first dissociation step

  • It is a weak acid in the second step

14.8 Acid-Base Properties of Salts

• ****It can produce acidic, basic, or neutral solutions.

• For any salt whose cation has neutral properties and whose anion is the conjugate base of a weak acid, the aqueous solution will be basic

• Salts that contain: Cations of strong bases and anions of strong acids produce neutral solutions

  • Cations of strong bases and anions of weak acids produce basic solutions

  • Cations of weak bases and anions of strong acids produce acidic solutions

14.9 The Effect on Structure on Acid-Base Properties

• Many substances that function as acids or bases contain the H¬O¬X grouping

  • Molecules in which the O¬X bond is strong and covalent tend to behave as acids

  • As X becomes more electronegative, the acid becomes stronger

• The net effect is to both polarize and weaken the O---H bond; this effect becomes more important as the number of attached oxygen atoms increases.

• There is an excellent correlation between the electronegativity of X and the acid strength for oxyacids

14.10 Acid- Base Properties of Oxides

• A compound containing the H¬O¬X group will produce an acidic solution in water if the O¬X bond is strong and covalent

  • If the O¬X bond is ionic, the compound will produce a basic solution in water

• Other common covalent oxides that react with water to form acidic solutions are sulfur dioxide, carbon dioxide, and nitrogen dioxide

• Thus, when a covalent oxide dissolves in water, an acidic solution forms (acidic oxides)

• Most ionic oxides produce basic solutions when they are dissolved in water, which is called basic oxides

14.11 The Lewis Acid--Base Model

• **** A Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor

• Another way of saying this is that a Lewis acid has an empty atomic orbital that it can use to accept (share) an electron pair from a molecule that has a lone pair of electrons

• The Lewis model encompasses the Bronsted–Lowry model, but the reverse is not true

• The reaction between a covalent oxide and water to form a Bronsted–Lowry acid can be defined as a Lewis acid-base reaction

  • An example is a reaction between sulfur trioxide and water

14.12 Strategy for Solving Acid-Base Problems: A Summary

• ****When analyzing an acid-base equilibrium problem, do not ask yourself how a memorized solution can be used to solve the problem

• Solving Acid-Base Problems: List the major species in solution

  • Determine the concentration of the products

  • Write down the major species in solution after the reaction

  • Look at each major component of the solution and decide if it is an acid or a base

  • Pick the equilibrium that will control the pH