Acids, Bases, & Salts

I. Properties of Acids & Bases A.Properties of Acids 1.Aqueous solutions have a sour taste 2.Acids change the color of acid-base indicators 3.Some acids react with active metals to release hydrogen Zn(s) + H 2 SO 4 (aq)  ZnSO 4 (aq) + H 2 (g) 4.Acids react with bases to produce salts and water HCl(aq) + NaOH(aq)  NaCl(aq) + H 2 O(l) 5.Acids conduct electric current

B.Properties of Bases 1.Aqueous solutions of bases have a bitter taste 2.Bases change the color of acid-base indicators 3.Dilute aqueous solutions of bases feel slippery 4.Bases react with acids to produce salts and water 5.Bases conduct electric current

II.Arrhenius Acids and Bases - Svante Arrhenius, Swedish chemist ( ) A.Arrhenius Acid – A chemical compound that increases the concentration of hydrogen ions, H +, in aqueous solution [H + ] > [OH - ] B.Arrhenius Base – A substance that increases the concentration of hydroxide ions, OH-, in aqueous solution [H + ] < [OH - ]

C.Aqueous Solutions of Acids 1.Acids are molecular compounds that ionize in solution HNO 3 + H 2 O  H 3 O + + NO 3 - H 2 SO 4 + H 2 O  H 3 O + + HSO 4 - H 2 O + HCl  H 3 O + + Cl -

D.Strength of Acids 1.Strong acids ionize completely in solution 2.Weak acids ionize only slightly and are weak electrolytes Strong AcidsWeak Acids H 2 SO 4 HSO 4 - HClO 4 H 3 PO 4 HClHF HNO 3 CH 3 COOH HBrH 2 CO 3 HIH2SH2S HCN HCO 3 -

E.Aqueous Solutions of Bases 1.Ionic bases dissociate to some extent when placed in water NaOH (s) Na + (aq) + OH − (aq) 1.Basic solutions are referred to as “alkaline” 2.Molecular bases react with water to produce hydroxide ions NH 3 (g) + H 2 O (l) ↔ NH 4 + (aq) + OH - (aq)

F.Strength of Bases 1.Strength of ionic bases is linked to solubility a)High solubility = strong base b)Low solubility = weak base 2.Molecular bases tend to be weak regardless of solubility

III.Bronsted-Lowry Acids & Bases A.Bronsted-Lowry Acid – molecule or ion that donates a proton (H + ) B.Bronsted-Lowry Base – molecule or ion that accepts protons 1.Hydroxide ions (OH - ) are acceptor of ionic bases

IV.Conjugate Acids & Bases A.Conjugate Base 1.The species that remains after an acid has given up a proton H 3 PO 4 (aq) + H 2 O (l) ↔ H 3 O + (aq) + H 2 PO 4 - (aq) acid conjugate base 2.The stronger an acid, the weaker its conjugate base

B.Conjugate Acid 1.The species that is formed when a base gains a proton H 3 PO 4 (aq) + H 2 O (l) ↔ H 3 O + (aq) + H 2 PO 4 - (aq) base conjugate acid 2.The stronger a base, the weaker its conjugate acid

IV. Ionization Constants A.Ionization constant = K a B.Compares to relative strength of acids (high K a = stronger acid)

B.pH and pOH 1.H 2 O (l) + H 2 O (l) ↔ H 3 O + + OH - 2.At 25 o C, [H 3 O + ] and [OH - ] = 1.0 x mol/L, and remains constant in pure water and dilute aqueous solutions. 3.This constant, K w, is called the ionization constant of water. K w = [H 3 O + ][OH - ] = (1.0 x )(1.0 x ) = 1.0 x mol/L

V. pH Scale & [H 3 O + ] [OH - ] A.pH Scale 1.pH [OH - ] 2.pH = 7 : neutral [H 3 O + ] = [OH - ] 3.pH > 7 : basic/alkaline [H 3 O + ] < [OH - ]

B.Formulas for pH, pOH, [H 3 O + ], & [OH - ] pH = - log [H 3 O + ]pOH = - log [OH - ]pH + pOH = 14 Calculating pH: [H 3 O + ] must be in scientific notation! Calculate the log of the [H 3 O + ] Multiply by -1 If [H 3 O + ] = , then pH = = 2.61 x (log 2.61 x ) x -1 = 1.58 (In calculators type: 2.61, EE, 2, +/-, log, +/-)

Calculating pOH [OH - ] must be in scientific notation! Calculate the log of the [OH - ] Multiply by -1 If [OH - ] = , then pOH = = 4.8 x log 4.8 x x-1 = 8.3 (In calculators type: 4.8, EE, 9, +/-, log, +/-)

Calculating [H3O+] and [OH-] From pH & [OH-] from [H3O+] (and vice versa) [H3O+] = antilog (-pH) If pH = 7.52, what are the [H3O+] and [OH-] concentrations? a) [H3O+] = antilog (-7.52) (In calculators type: 7.52, +/-, 2nd, 10x or 7.52, +/-, 2nd, log) b) [H3O+] = c) [H3O+] [OH-] = 1.0 x d) [OH-] = 1.0 x = 1.0 x = 3.3 x 10-7 M OH- [H3O+]