Sodium hydroxide, a base, produces hydroxide ions in solution:

Transcription

1 Assistant Lecture Aayad Amaar Acids and Bases The properties of acids and bases are related to their chemical structure. All acids have common characteristics that enable them to increase the hydrogen ion concentration in water. All bases lower the hydrogen ion concentration in water. Two theories, one developed from the other, help us to understand the unique chemistry of acids and bases. Arrhenius Theory of Acids and Bases One of the earliest definitions of acids and bases is the Arrhenius theory. According to this theory, an acid, dissolved in water, dissociates to form hydrogen ions or protons (H + ), and a base, dissolved in water, dissociates to form hydroxide ions (OH - ). For example, hydrochloric acid dissociates in solution according to the reaction Sodium hydroxide, a base, produces hydroxide ions in solution: The Arrhenius theory satisfactorily explains the behavior of many acids and bases. However, a substance such as ammonia, NH3, has basic properties but cannot be an Arrhenius base, because it contains no OH -. The Br? nsted-lowry theory explains this mystery and gives us a broader view of acid-base theory by considering the central role of the solvent in the dissociation process. Br?nsted-Lowry Theory of Acids and Bases The Br?nsted-Lowry theory defines an acid as a proton (H + ) donor and a base as a proton acceptor. Hydrochloric acid in solution donates a proton to the solvent water thus behaving as a Br?nsted-Lowry acid: H3O + is referred to as the hydrated proton or hydronium ion.

2 The basic properties of ammonia are clearly accounted for by the Br?nsted-Lowry theory. Ammonia accepts a proton from the solvent water, producing OH -. An equilibrium mixture of NH3, H2O, NH4 +, and OH - results. For aqueous solutions, the Br?nsted-Lowry theory adequately describes the behavior of acids and bases. We shall limit our discussion of acid-base chemistry to aqueous solutions and use the following definitions: An acid is a proton donor. A base is a proton acceptor. Conjugate Acids and Bases The Br?nsted-Lowry theory contributed several fundamental ideas that broadened our understanding of solution chemistry. First of all, an acid-base reaction is a charge-transfer process. Second, the transfer process usually involves the solvent. Water may, in fact, accept or donate a proton. Last, and perhaps most important, the acid-base reaction is seen as a reversible process. Consequently, any acid-base reaction can be represented by the general equation In the forward reaction, the acid (HA) donates a proton (H + ) to the base (B) leading to the formation of BH + and A -. However, in the reverse reaction, it is the BH + that behaves as an acid; it donates its extra proton to A -. A - is therefore a base in its own right because it accepts the proton. These product acids and bases are termed conjugate acids and bases. A conjugate acid is the species formed when a base accepts a proton. A conjugate base is the species formed when an acid donates a proton.

3 The acid and base on the opposite sides of the equation are collectively termed a conjugate acid-base pair. In the above equation: BH + is the conjugate acid of the base B. A - is the conjugate base of the acid HA. B and BH + constitute a conjugate acid-base pair. HA and A - constitute a conjugate acid-base pair. Rewriting our model equation: Although we show the forward and reverse arrows to indicate the reversibility of the reaction, seldom are the two processes equal but opposite. One reaction, either forward or reverse, is usually favored. Consider the reaction of hydrochloric acid in water: HCl is a much better proton donor than H 3 O +. Consequently the forward reaction predominates, the reverse reaction is inconsequential, and hydrochloric acid is termed a strong acid. The dissociation of hydrochloric acid is so favorable that we describe it as 100% dissociated and use only a single forward arrow to represent its behavior in water: The degree of dissociation, or strength, of acids and bases has a profound influence on their aqueous chemistry. For example, vinegar (a 5% [w/v] solution of acetic acid in water) is a consumable product; aqueous hydrochloric acid in water is not. Why? Acetic acid is a weak acid and, as a result, a dilute solution does no damage to the mouth and esophagus. The following section looks at the strength of acids and bases in solution in more detail.

4 Acid-Base Properties of Water The role that the solvent, water, plays in acid-base reactions is noteworthy. In the example above, the water molecule accepts a proton from the HCl molecule. The water is behaving as a proton acceptor, a base. However, when water is a solvent for ammonia (NH 3 ), a base, the water molecule donates a proton to the ammonia molecule. The water, in this situation, is acting as a proton donor, an acid. Water, owing to the fact that it possesses both acid and base properties, is termed amphiprotic. The solvent properties of water are a consequence of this ability to either accept or donate protons. Water is the most commonly used solvent for acids and bases. These interactions promote solubility and dissociation of acids and bases Acid and Base Strength The terms acid or base strength and acid or base concentration are easily confused. Strength is a measure of the degree of dissociation of an acid or base in solution, independent of its concentration. Concentration, as we have learned, refers to the amount of solute (in this case, the amount of acid or base) per quantity of solution. The strength of acids and bases in water depends on the extent to which they react with the solvent, water. Acids and bases are classified as strong when the reaction with water is virtually 100% complete and as weak when the reaction with water is much less than 100% complete. Important strong acids include: Note that the equation for the dissociation of each of these acids is written with a single arrow. This indicates that the reaction has little or no tendency to proceed in the reverse direction to establish equilibrium. All of the acid molecules are dissociated to form ions. All common strong bases are metal hydroxides. Strong bases completely dissociate in aqueous solution to produce hydroxide

5 ions and metal cations. Of the common metal hydroxides, only NaOH and KOH are soluble in water and are readily usable strong bases: Weak acids and weak bases dissolve in water principally in the molecular form. Only a small percentage of the molecules dissociate to form the hydronium or hydroxide ion. Two important weak acids are: The double arrow implies an equilibrium between dissociated and undissociated species. We have already mentioned the most common weak base, ammonia. Many organic compounds function as weak bases. Several examples of weak bases follow: The fundamental chemical difference between strong and weak acids or bases is their equilibrium ion concentration. A strong acid, such as HCl, does not, in aqueous solution, exist to any measurable degree in equilibrium with its ions, H 3 O + and Cl -. On the other hand, a weak acid, such as acetic acid, establishes a dynamic equilibrium with its ions, H 3 O + and CH3COO -. The relative strength of an acid or base is determined by the ease with which it donates or accepts a proton. Acids with the greatest proton-donating capability (strongest acids) have the weakest conjugate bases. Good proton acceptors (strong bases) have weak conjugate acids. Solutions of acids and bases used in the laboratory must be handled with care. Acids burn because of their exothermic reaction with water present on and in the skin. Bases react with proteins, which are principal components of the skin and eyes. Such solutions are more hazardous if they are strong or concentrated. A strong acid or base produces more H3O + or OH - than does the corresponding weak acid or base. More-

6 concentrated acids or bases contain more H 3 O + or OH - than do less concentrated solutions of the same strength. The Dissociation of Water Aqueous solutions of acids and bases are electrolytes. The dissociation of the acid or base produces ions that can conduct an electrical current. As a result of the differences in the degree of dissociation, strong acids and bases are strong electrolytes; weak acids and bases are weak electrolytes. The conductivity of these solutions is principally dependent on the solute and not the solvent (water). Although pure water is virtually 100% molecular, a small number of water molecules do ionize. This process occurs by the transfer of a proton from one water molecule to another, producing a hydronium ion and a hydroxide ion: This process is the autoionization, or self-ionization, of water. Water is therefore a very weak electrolyte and a very poor conductor of electricity. Water has both acid and base properties; dissociation produces both the hydronium and hydroxide ion. Pure water at room temperature has a hydronium ion concentration of 1.0x10-7 M. One hydroxide ion is produced for each hydronium ion. Therefore, the hydroxide ion concentration is also 1.0x10-7 M. Molar equilibrium concentration is conveniently indicated by brackets around the species whose concentration is represented: The product of hydronium and hydroxide ion concentration in pure water is referred to as the ion product for water. The ion product is constant because its value does not depend on the nature or concentration of the solute, as long as the temperature does not change. The ion product is a temperature-

7 dependent quantity. The nature and concentration of the solutes added to water do alter the relative concentrations of H 3 O + and OH - present, but the product, [H 3 O + ][OH - ], always equals 1.0x10-14 at 25 C. This relationship is the basis for a scale that is useful in the measurement of the level of acidity or basicity of solutions. This scale, the ph scale.

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