By the Arrhenius definitions, acids are substances that produce H3O+ ions in water and bases are substances that produce OH- ions in water. A strong acid or base produces high concentrations of these ions. The Bronsted-Lowry definitions expand this concept beyond water: An acid is a proton donor, and a base is a proton acceptor. Every acid has a conjugate base. An amphoteric substance, such as water, can act as either an acid or a base.
The strengths of weak acids are expressed by Ka values; the higher the Ka, the stronger the acid. Strong acids and bases are corrosive. Acids neutralize metals, metal hydroxides, and oxides to give salts, which are made up of positive and negative ions. Acids also react with carbonates, bicarbonates, ammonia, and amines.
In pure water, a small percentage of molecules undergo the reaction 2H2O---H3O++OH-, enough to produce a concentration of 10-7 M of each of H3O+ and OH-. The ion product, Ka=[H3O+][OH-], is equal to 10-14 in any aqueous solution. Hydronium ion concentrations are generally expressed in pH units, with pH=-log[H3O+]. Solutions with pH less than 7 are acidic; those with pH higher than 7 are basic. Neutral solutions have a pH of 7. The pH is measured with indicators or with a pH meter.
Many salts hydrolyze in water. Salts of strong acids and weak bases are acidic; salts of weak acids and strong bases are basic; salts of strong acids and strong bases are neutral.
A buffer solution does not change its pH very much when a strong acid or base is added. Buffers are made up of approximately equal concentrations of a weak acid and its conjugate base. Every buffer solution has a pH and a capacity. The most important buffers for blood are carbonate and phosphate.
The concentration of aqueous solutions of acids and bases can be measured by titration. For instance, a base of known concentration is added gradually to an acid of unknown concentration (or vice versa) until an end point is reached, at which point the solution is completely neutralized. The unknown acid concentration can then be calculated by knowing the volume of a known base concentration needed for neutralization.
The pH Scale
The pH scale correlates the hydronium ion concentration with a number, the pH, which serves as a useful indicator of the degree of acidity or basicity of a solution. The pH scale specifies "how acidic" or "how basic" a solution is.
1. Addition of an acid (proton donor) to water increases the [H3O+) and decreases the [OH-].
2. Addition of a base (proton acceptor) to water decreases the [H3O+) and increases the [OH-].
3. [H3O+] = [OH-] when equal amounts of acid and base are present.
4. In all of the above cases, [H3O+] [OH-] = 1.0 x 10-14 = ion product for water. The pH of a solution is defined as the negative logarithm of the molar concentration of the hydronium ion.
pH = -log [H30+]
Neutralization
The reaction of and acid with a base to produce a salt and water is referred to as neutralization. In the strictest sense, neutralization requires equal numbers of moles of H3O+ and OH- to produce a neutral solution (no excess acid or base).
A neutralization reaction may be used to determine the concentration of an unknown acid or base solution. The technique of titration involves the addition of measured amounts of a standard solution (one whose concentration is known) to neutralize the second, unknown solution. From the volumes of the two solutions and the concentration of the standard solution, the concentration of the unknown solution may be determined.
Acid-base reactions may occur in various combining ratios if they involve a polyprotic substance. Polyprotic substances accept or donate more than one proton per formula unit. Examples include H2SO4, H3PO4, (acids) and Ca(OH)2 (base).
Acids-Bases Buffers
A buffer solution contains components that enable the solution to resist large changes in pH when acids or bases are added.
The Buffer Process
The basis of buffer action is the establishment of an equilibrium between either a weak acid and its salt, or a weak base and its salt. A buffer solution functions in accordance with LeChatelier's Principle, which states that an equilibrium system, when stressed, will shift its equilibrium to alleviate that stress.
Addition of Base (OH-) to a Buffer Solution
For the acetic acid/sodium acetate system:
CH3COOH + H2O = H3O+ + CH3COO-
OH- added, equilibrium shifts to the right
Addition of Acid (H3O+) to a Buffer Solution
For the acetic acid/ sodium acetate system:
CH3COOH + H2O+ + CH3COO-
H3O+ added, equilibrium shifts to the left
Higher-than-normal CO2 levels shift the above equilibrium to the right increasing [H3O+] and lowering the pH. A situation of high blood CO2 levels and low pH is acidosis.
Lower-than-normal Co2 levels shift the equilibrium to the left, decreasing [H3O+] and making the pH more basic; this condition is termed alkalosis (from alkali, implying basic in nature).