GJ

Chapter 14 - Acids and Bases

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 - Polyprotic acids

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