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

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

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+

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

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

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

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

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

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

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

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

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