Acid–Base Reactions for AP Chemistry (page 2)
Practice problems for these concepts can be found at:
- Reactions and Periodicity Review Questions for AP Chemistry
- Reactions and Periodicity Free-Response Questions for AP Chemistry
Acids and bases are extremely common, as are the reactions between acids and bases. The driving force is often the hydronium ion reacting with the hydroxide ion to form water. The chapter on Equilibrium describes the equilibrium reactions of acids and bases, as well as some information concerning acid–base titration. After you finish this section, you may want to review the acid–base part of the Equilibrium chapter.
Properties of Acids, Bases, and Salts
At the macroscopic level, acids taste sour, may be damaging to the skin, and react with bases to yield salts. Bases taste bitter, feel slippery, and react with acids to form salts.
At the microscopic level, acids are defined as proton (H+) donors (Brønsted–Lowry theory) or electron-pair acceptors (Lewis theory). Bases are defined as proton (H+) acceptors (Brønsted–Lowry theory) or electron-pair donors (Lewis theory). Consider the gas-phase reaction between hydrogen chloride and ammonia:
- HCl + :NH3(g) → HNH3+Cl– ( or NH4+Cl–)
HCl is the acid, because it is donating an H+ and the H+ will accept an electron pair from ammonia. Ammonia is the base, accepting the H+ and furnishing an electron pair that the H+ will bond with via coordinate covalent bonding. Coordinate covalent bonds are covalent bonds in which one of the atoms furnishes both of the electrons for the bond. After the bond is formed, it is identical to a covalent bond formed by donation of one electron by both of the bonding atoms.
Acids and bases may be strong, dissociating completely, or weak, partially dissociating and forming an equilibrium system.(See Chapter 15 for the details on weak acids and bases.) Strong acids include:
- Hydrochloric, HCl
- Hydrobromic, HBr
- Hydroiodic, HI
- Nitric, HNO3
- Chloric, HClO3
- Perchloric, HClO4
- Sulfuric, H2SO4
The strong acids above are all compounds that ionize completely in aqueous solution, yielding hydrogen ions and the anions from the acid.
Strong bases include:
- Alkali metal (Group IA) hydroxides (LiOH, NaOH, KOH, RbOH, CsOH)
- Ca(OH)2, Sr(OH)2, and Ba(OH)2
The strong bases listed above are all compounds that dissociate completely, yielding the hydroxide ion (which is really the base, not the compound).
Unless told otherwise, assume that acids and bases not on the lists above are weak and will establish an equilibrium system when placed into water.
Some salts have acid–base properties. For example, ammonium chloride, NH4Cl, when dissolved in water will dissociate and the ammonium ion will act as a weak acid, donating a proton. We will examine these acid–base properties in more detail in the next section.
Certain oxides can have acidic or basic properties. These properties often become evident when the oxides are dissolved in water. In most case, reactions of this type are not redox reactions.
Many oxides of metals that have a +1 or +2 charge are called basic oxides (basic anhydrides), because they will react with acids.
Many times they react with water to form a basic solution:
Many nonmetal oxides are called acidic oxides (acidic anhydrides), because they react with water to form an acidic solution:
- CO2(g) + H2O(I) → H2CO3(aq)
H2CO3(aq) is named carbonic acid and is the reason that most carbonated beverages are slightly acidic. It is also the reason that soft drinks have fizz, because carbonic acid will decompose to form carbon dioxide and water.
In general, acids react with bases to form a salt and, usually, water. The salt will depend upon which acid and base are used:
Reactions of this type are called neutralization reactions.
The first two neutralization equations are represented by the same net ionic equation:
H+(aq) + OH–(aq) → H2O (I)
In the third case, the net ionic equation is different:
H+(aq) + NH3(aq) → NH4+ (aq)
As mentioned previously, certain salts have acid–base properties. In general, salts containing cations of strong bases and anions of strong acids are neither acidic nor basic. They are neutral, reacting with neither acids nor bases. An example would be potassium nitrate, KNO3. The potassium comes from the strong base KOH and the nitrate from the strong acid HNO3.
Salts containing cations not of strong bases but with anions of strong acids behave as acidic salts. An example would be ammonium chloride, NH4Cl.
Cations of strong bases and anions not of strong acids are basic salts. An example would be sodium carbonate, Na2CO3. It reacts with an acid to form carbonic acid, which would then decompose to carbon dioxide and water:
The same type of reaction would be true for acid carbonates, such as sodium bicarbonate, NaHCO3.
Another group of compounds that have acid–base properties are the hydrides of the alkali metals and of calcium, strontium, and barium. These hydrides will react with water to form the hydroxide ion and hydrogen gas:
Note that in this case, water is behaving as H+OH–.
A common laboratory application of acid–base reactions is a titration. A titration is a laboratory procedure in which a solution of known concentration is used to determine the concentration of an unknown solution. For strong acid/strong base titration systems, the net ionic equation is:
- H+(aq) + OH–(aq) → H2O (I)
For example, suppose you wanted to determine the molarity of an HCl solution. You would pipet a known volume of the acid into a flask and add a couple drops of a suitable acid–base indicator. An indicator that is commonly used is phenolphthalein, which is colorless in an acidic solution and pink in a basic solution. You would then fill a buret with a strong base solution (NaOH is commonly used) of known concentration. The buret allows you to add small amounts of the base solution to the acid solution in the flask. The course of the titration can also be followed by the use of a pH meter. Initially the pH of the solution will be low, since it is an acid solution. As the base is added and neutralization of the acid takes place, the pH will slowly rise. Small amounts of the base are added until one reaches the equivalence point. The equivalence point is that point in the titration where the number of moles of H+ in the acid solution has been exactly neutralized with the same number of moles of OH–:
- molesH+ = molesOH– at the equivalence point
For the titration of a strong acid with a strong base, the pH rapidly rises in the vicinity of the equivalence point. Then, as the tiniest amount of base is added in excess, the indicator turns pink. This is called the endpoint of the titration. In an accurate titration the endpoint will be as close to the equivalence point as possible. For simple titrations that do not use a pH meter, it is assumed that the endpoint and the equivalence point are the same, so that:
- molesH+ = molesOH– at the endpoint
After the equivalence point has been passed, the pH is greater than 7 (basic solution) and begins to level out somewhat. Figure 6.1 shows the shape of the curve for this titration.
Reaction stoichiometry can then be used to solve for the molarity of the acid solution. See the Stoichiometry chapter for a discussion of solution stoichiometry.
An unknown base can be titrated with an acid solution of known concentration. One major difference is that the pH will be greater than 7 initially and will decrease as the titration proceeds. The other major difference is that the indicator will start off pink, and the color will vanish at the endpoint.
Practice problems for these concepts can be found at:
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