SAT Chemistry Acids, Bases, and Salts - Definitions and Properties
SAT Chemistry Acids, Bases, and Salts - Definitions and Properties
AcidsThere are some characteristic properties by which an acid may be defined. The most important are:
1. Water (aqueous) solutions of acids conduct electricity. The degree of conduction depends on the acid’s degree of ionization. A few acids ionize almost completely, while others ionize to only a slight degree. Table 10 indicates some common acids and their degrees of ionization.
2. Acids will react with metals that are more active than hydrogen ions (see page 206) to liberate hydrogen. (Some acids are also strong oxidizing agents and will not release hydrogen. Somewhat concentrated nitric acid is such an acid.)
3. Acids have the ability to change the color of indicators. Some common indicators are litmus and phenolphthalein. Litmus is a dyestuff obtained from plant life. When litmus is added to an acidic solution, or paper impregnated with litmus is dipped into an acid, the neutral purple color changes to pink-red. Phenolphthalein is pink in a basic solution and becomes colorless in a neutral or acid solution.
4. Acids react with bases so that the properties of both are lost to form water and a salt. This is called neutralization. The general equation is:
Acid + Base —> Salt + Water
An example is:
Mg(OH)2(aq) + H2SO4(aq) → MgSO4(aq) + 2H2O(l)
5. Acids react with carbonates to release carbon dioxide. An example:
The most common theory used in first-year chemistry is the Arrhenius Theory, which states that an acid is a substance that yields hydrogen ions in an aqueous solution. Although we speak of these hydrogen ions in the solution, they are really not separate ions but become attached to the oxygen of the polar water molecule to form the H3O+ ion (the hydronium ion). Thus, it is really this hydronium ion we are concerned with in an acid solution.
The general reaction for the dissociation of an acid, HX, is commonly written as
HX ⇌ H+ + X-
To show the formation of the hydronium ion, H30+, the complete equation is:
HX + H2O ⇌— H3O+ + X-
A list of common acids and their formulas is given in Chapter 4, Table 7; an explanation of the naming procedures for acids precedes Table 7.
Bases
Bases may also be defined by some operational definitions that are based on experimental observations. Some of the important ones are as follows:
1. Bases are conductors of electricity in an aqueous solution. Their degrees of conduction depend on their degrees of ionization. The degrees of ionization of some common bases are shown in Table 11.
2. Bases cause a color change in indicators. Litmus changes from red to blue in a basic solution, and phenolphthalein turns pink from its colorless state.
3. Bases react with acids to neutralize each other and form a salt and water.
4. Bases react with fats to form a class of compounds called soaps. Earlier generations used this method to make their own soap.
5. Aqueous solutions of bases feel slippery, and the stronger bases are very caustic to the skin.
The Arrhenius Theory defines a base as a substance that yields hydroxide ions (OH) in an aqueous solution.
Some common bases have familiar names, for example:
Sodium hydroxide = lye, caustic soda
Potassium hydroxide = caustic potash
Calcium hydroxide = slaked lime, hydrated lime, limewater
Ammonium hydroxide = ammonia water, household ammonia
Much of the sodium hydroxide produced today comes from the Hooker cell electrolysis apparatus. The electrolysis process for the decomposition of water was discussed in Chapter 7. When an electric current is passed through a saltwater solution, hydrogen, chlorine, and sodium hydroxide are the products. The formula for this equation is:
Broader Acid-Base Theories
Besides the common Arrhenius Theory of acids and bases discussed for aqueous solutions, two other theories, the Bronsted-Lowry Theory and the Lewis Theory, are widely used.
The Bronsted-Lowry Theory (1923) defines acids as proton donors and bases as proton acceptors. This definition agrees with the aqueous solution definition of an acid giving up hydrogen ions in solution, but goes beyond to other cases as well.
An example is the reaction of dry HCl gas with ammonia gas to form the white solid NH4Cl.
HCl(g) + NH3(g) → NH4C1(s)
The HCl is the proton donor or acid, and the ammonia is a Bronsted-Lowry base that accepts the proton.
Conjugate Acids and Bases
In an acid-base reaction, the original acid gives up its proton to become a conjugate base. In other words, after losing its proton, the remaining ion is capable of gaining a proton, thus qualifying as a base. The original base accepts a proton, so it now is classified as a conjugate acid since it can release this newly acquired proton and thus behave like an acid.
Some examples are given below:
Strength of Conjugate Acids and Bases
The extent of the reaction between a Bronsted-Lowry acid and base depends on the relative strengths of the acids and bases involved. Consider the following example. Hydrochloric is a strong acid. It gives up protons readily. It follows that the Cl-ion has little tendency to attract and retain a proton. Consequently, the Cl-ion is an extremely weak base.
This observation leads to an important conclusion: the stronger an acid is, the weaker its conjugate base; the stronger a base is, the weaker its conjugate acid. This concept allows strengths of different acids and bases to be compared to predict the outcome of a reaction. As an example, consider the reaction of perchloric acid, HClO4, and water.
Another important general conclusion is that proton-transfer reactions favor the production of the weaker acid and the weaker base. For a reaction to approach completion, the reactants must be much stronger as an acid and as a base than the products.
The Lewis Theory (1916) defines acids and bases in terms of the electron-pair concept, which is probably the most generally useful concept of acids and bases. According to the Lewis definition, an acid is an electron-pair acceptor; and a base is an electron-pair donor. An example is the formation of ammonium ions from ammonia gas and hydrogen ions.
Notice that the hydrogen ion is in fact accepting the electron pair of the ammonia, so it is a Lewis acid. The ammonia is donating its electron pair, so it is a Lewis base.
Another example is boron trifluoride. It is an excellent Lewis acid. It forms a fourth covalent bond with many molecules and ions. Its reaction with a fluoride ion is shown below.
The acid-base systems are summarized below.
Acid Concentration Expressed as pH
100 = 102 so logarithm of 100, base 10 = 2
10,000 = 104 so logarithm of 10,000, base 10 = 4
0.01 = 10-2 so logarithm of 0.01, base 10 = -2
The logarithms of more complex numbers can be found in a logarithm table. An example of a pH problem is:
Find the pH of a 0.1 molar solution of HCl.
1st step. Because HC1 ionizes almost completely into H+ and Cl-, [H+] = 0.1 mole/liter.
2nd step. By definition
pH = -log [H+] so
pH = -log [10-1]
3rd step. The logarithm of 10-1 is -1 so
pH = -(-1)
4th step. The pH then is = 1.
Because water has a normal H+ concentration of 10-7 mole/liter because of the slight ionization of water molecules, the water pH is 7 when the water is neither acid nor base. The normal pH range is from 0 to 14.
The pOH is the negative logarithm of the hydroxide ion concentration:
pOH = -log [OH-]
If the concentration of the hydroxide ion is 10-9 M, then the pOH of the solution is +9. From the equation
[H+][OH-] = 1.0 x 10-14 at 298 K
the following relationship can be derived:
pH + pOH = 14.00
In other words, the sum of the pH and pOH of an aqueous solution at 298 K must always equal 14.00. For example, if the pOH of a solution is 9.00, then its pH must be 5.00.
Sample Problem ______________________________________________
What is the pOH of a solution whose pH is 3.0?
Substituting 3.0 for pH in the expression
pH + pOH = 14.0 gives
3.0+ pOH = 14.0
pOH = 14.0 - 3.0
pOH = 11.0
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