Acids and Bases – A Summary

ACIDS AND BASES – A SUMMARY There are three separate definitions of what makes an acid and a base. The Arrhenius definition is the simplest: An acid is a compound that produces hydrogen ions on solution in water. A base is a compound that produces hydroxide ions on solution in water. HCl  H+ + Cl- hydrochloric acid, an Arrhenius acid NaOH  Na+ + OH- sodium hydroxide, an Arrhenius base The Bronsted-Lowry definition goes a little bit further: An acid is a substance that is a proton (hydrogen ion) donor. A base is a substance that is a proton (hydrogen ion) acceptor. HCl + H2O == H3O+ + ClIn the above reaction, HCl donates a proton to water, so it is an acid. Water accepts a proton from HCl, so it is a base. In the reverse reaction, H3O+ donates a proton to Cl-, so it is an acid. Cl- accepts a proton, so it is a base. The Bronsted-Lowry definition introduces the idea of conjugate acid-base pairs. In the above reaction, the conjugate base of HCl (an acid) is Cl-. Since HCl is a very strong acid, Cl- would be a very weak base. The third definition of acids and bases is the Lewis definition. An acid is an electron pair acceptor, and a base is an electron pair donor. This definition extends the idea of acids and bases to solvents other than water, as well. Pure water undergoes dissociation to a very slight extent. H2O == H+ + OHIn this dissociation, the concentrations of hydrogen ion and hydroxide ion are equal to 10-7 M each. The equilibrium constant, called the ion product constant for water would be: Kw = [H+][OH-] = [10-7][10-7] = 10-14 This is a very important concept: In any water solution, the Kw = 10-14. In other words, the product of the concentrations of hydrogen ion and hydroxide ion will always equal 10-14. Note that in this equilibrium constant, the concentration of water is not included, as water is a pure liquid. Now, using this idea, we introduce a scale, called the pH scale, to quantify how acidic or basic a solution might be. This scale is based on the following definition: pH = - log [H+] In pure water, [H+] = 10-7, so pH = -log (10-7) = - (-7) = 7 We call this condition neutral, as it is neither acidic or basic. The hydrogen ion and hydroxide ion concentrations are equal. The above chart shows the relationship between pH, [H+], [OH-], and pOH. pH = -log [H+] pOH = -log [OH-] [H-] = antilog (-pH) [OH-] = antilog (-pOH) pH + pOH = 14 [H+][OH-] = 10-14 Note: the pH scale is a logarithmic scale. Solutions differing by a pH of 1 have a 10 fold difference in hydrogen ion concentration. Acids and bases are electrolytes. This means that when an acid or a base is dissolved in water, the solution will conduct electricity. There are weak acids and strong acids, just as there are weak bases and strong bases. A strong acid or a strong base would be a strong electrolyte. That is, it would be completely dissociated into its respective ions in solution. Examples of a strong acid and a strong base are: HCl  H+ + Cl- hydrochloric acid, a strong acid NaOH  Na+ + OH- sodium hydroxide, a strong base Weak acids and weak bases only partially dissociate into their respective ions on solution. Examples of a weak acid and a weak base are (along with their equilibrium constants): CH3COOH == CH3COO- + H+ Ka = [CH3COO-][H+] / [CH3COOH] NH4OH == NH4+ + OH- Kb = [NH4+][OH-] / [NH4OH] The equilibrium constants for the above dissociations are both 1.8 x 10-5, which indicates that only a small dissociation takes place. The size of the dissociation constants for weak acids and weak bases can be used as a measure of the strength of the acid or base. The smaller the dissociation constant the weaker the acid or base. A neutralization reaction is a reaction between an acid and a base, and the products are water and a salt. Examples are: HCl + NaOH  H2O + NaCl 2HCl + Ca(OH)2  2H2O + CaCl2 A titration is an experiment where the known concentration of one solution is used to determine the concentration of an unknown solution. In acid-base chemistry, the known concentration of an acid could be used to determine the unknown concentration of a base of vice versa using a neutralization reaction. You would use the following setup for a titration: Let’s assume (for illustration) that we have an unknown solution of HCl whose concentration we want to determine. We could accurately measure a volume of the HCl solution into an Erlenmeyer flask using a pipet (say 25 ml). We could take a solution of 0.1 M NaOH and fill a buret. Then, we could add 1 milliliter increments of the base to the flask and measure the pH of the solution after each addition. If we plot pH versus volume of NaOH added, the curve would look something like: The equivalence point is the pH at which all of the acid has been reacted with the base. In the case of a strong acid reacting with a strong base, the equivalence point would be at pH 7. If you know the volume of base added, and you know its concentration, you can calculate the number of moles used, as we know if we multiply concentration in molarity (moles per liter) by volume in liters, we get # moles. If our volume is in milliliters, we get millimoles. A millimole is 1/1000 of a mole, as a milliliter is 1/1000 of a liter. Assume in the example that it took 20 milliliters of 0.1 M NaOH to neutralize the HCl in 25 ml of the unknown solution. # millimoles NaOH used = 20 x 0.1 = 2.0 millimoles conc. of HCl = 2.0 millimoles divided by 25 milliliters = 0.08 M Salts of weak acids and weak bases. Salts of weak acids will produce basic solutions because of hydrolysis. Consider sodium acetate, NaCH3CO2, or for simplicity, we’ll write it NaAc: NaAc  Na+ + Ac- The salt will dissociate completely. However, the acetate ion will undergo hydrolysis as Ac- + H2O == HAc + OHAnother way to think of it would be: Ac- + H+ == HAc Which is the reverse reaction in the dissociation of HAc. In the case of the salt of a weak base, such as NH4Cl, NH4Cl  NH4+ + Cl- The salt will dissociate completely. However, the ammonium ion would under go hydrolysis as: NH4+ + H2O == NH4OH + H+ giving an acidic solution Thought question: how would all of this affect the pH of the equivalence points in titrations of weak acids and weak bases?