CH 13 Acids and Bases

BrØnsted-Lowry Acid-Base Model In an acid base reaction, a proton (H+) is transferred from an acid to a base. Acids are proton donors Bases are proton acceptors

HB + A-   HA + B- HB, HA are acids A-, B- are bases Conjugate base: the species formed when a proton is removed from an acid. Conjugate acid: the species formed when a proton is added to a base. HB, B- is a conjugate acid-base pair.

A species that can either accept or donate a proton is known as amphiprotic. Ex. H2O Can gain a proton, H3O+ (hydronium ion) Can lose a proton, OH- (hydroxide ion)

What is the conjugate base of HNO2 ? What is the conjugate acid of F- ? HCO3- is amphiprotic, What is the conjugate base?, conjugate acid?

Ion product of water H2O (l) + H2O (l)   H3O+ (aq) + OH- (aq) Write the equilibrium expression for this reaction. Kw = Known as the ion product constant for water. Value for Kw = 1 x 10-14

Neutral solutions The [OH-] = [H+] in a neutral solution. [H+] = 1.0 x 10-7 If the [H+] > [OH-] then the solution is acidic.

pH and pOH pH power of hydrogen ion pH = -log10 [H+] The higher the pH, the less acidic the solution.

Calculate the pH an solution with [H+] = 4.6 x 10-8 M an acid with [H+] = 1.2 x 10-2 M an acid with [H+] = 6.0 M

Strong Acids Strong Acids completely dissociate HCl  H+ + Cl- other strong acids 1.0 M HCl has [H+]=1.0M and [Cl-] = 1.0M Essentially no reverse reaction. Therefore no equilibrium between forward and reverse reactions.

pOH -log10 [OH-] [H+] x [OH-] = 1.0 x 10-14 pH + pOH = 14 The normal pH of blood is 7.4, what is the pOH? What is the [H+] ? The [OH-]?

pH problem (from [OH] ) Dissolve 2.0g of Ba(OH)2 in 1.0 L of water. What is the pH?

Weak Acids and their Ka Weak acids Molecules containing an ionizable Hydrogen atom. HC2H3O2  H+ + C2H3O2- Including Ammonium NH4+  NH3 + H+ Transition metal cations (and Al) Al3+ is really surrounded by 6 water molecules forming a complex ion. Al(H2O)63+

Metal cations Al(H2O)63+ + H2O  Al(H2O)5(OH) 2+ + H3O+ 6 waters are bonded to the central atom. This species transfers an H+ to a solvent (free) water. Forming H3O+

Weak Acids and their Ka Ka is the acid equilibrium constant The larger the Ka the stronger the acid Sulfurous 1.7x10 -2 Acetic 1.8 x 10 -5

% ionization (of a weak acid) HB   H+ + A- The amount of acid that ionizes. Compare the H+ , to the original amount of acid. (HB) Formula: % ionization= [H+]eq / [HB]o (x 100%) Weaker acids ionize less than stronger acids.

% ionization .200M aspirin, has a 0.0025M [H+] at equilibrium. What is the % ionization?

Determine [H+] H+ I C E

Find the pH of .200M acetic acid. The Ka = 1.9 x 10-3

Polyprotic Acids More than 1 ionizable hydrogen Oxalic Acid H2C2O4 the Ka values become successively smaller. H2C2O4   H+ + HC2O4- Ka= 5.9 x 10-2 HC2O4-   H+ + C2O42- Ka= 5.2 x 10-5 The ‘second’ acid is much weaker. 1000x more H+ is put into solution by the first acid.

Weak bases and Kb Solutes acting as weak bases AAA Ammonia: NH3 Amines : R-NH2 Anions derived from a weak acid. F- , C 2H3O2-

Show how each of the following are weak bases NO2- Na2CO3 KHCO3

Kb Write the base equilibrium expression Kb for ammonia. Kb values : the larger the Kb the stronger the base. (just like acids) Use Kb to calculate the [OH-] in a weak base solution.

Calculate the pH of 0.10M NaF The Kb is for F- = 1.4 x 10-11 Write the equation and set up an ICE box. Solve for x to get [OH-]

Ka and Kb The equilibrium constant of a weak acid and its conjugate base multiply to Kw HF and F- K a x Kb = Kw 1 x 10-14

Acid-Base properties of salts Anions F- + H2O CO32- + H2O Most anions behave as weak bases, except those from derived from strong acids. Cl- , NO3- …

Effect of a salt of pH of solution Cation Anion Effect on solution NH4I ZN(NO3)2 KClO4 Na3PO4