Chapter 15: Acids and Bases Bronsted Lowry Acids HCl  H + (aq) + Cl - (aq) HNO 3  H + (aq) + NO 3 - (aq) HC 2 H 3 O 2  H + (aq) + C 2 H 3 O 2 - (aq)

Chapter 15: Acids and Bases Bronsted-Lowry Acids Are Active in Water Acids are Electrolytes. HCl + H 2 O  H 3 O + (aq) + Cl - (aq) HC 2 H 3 O 2 + H 2 O  H 3 O + (aq) + C 2 H 3 O 2 - (aq)

Bronsted-Lowry Acids and Bases: Role of H 2 O HCl + H 2 O  H 3 O + (aq) + Cl - (aq) H 3 O + is the Hydronium Ion. In an Acidic Solution : –H 3 O + is the H + donor.

Chapter 15: Bronsted-Lowry Bases Water is involved in the activity of all bases! A Bronsted Lowry Base Accept H + through interaction with water. NaOH  Na + (aq) + OH - (aq) NH 3 + H 2 O  NH 4 + (aq) + OH - (aq)

Chapter 15: Bronsted-Lowry Conjugate Acid Base Pairs HCl + H 2 O  H 3 O + (aq) + Cl - (aq) HC 2 H 3 O 2 + H 2 O  H 3 O + (aq) + C 2 H 3 O 2 - (aq) NH 3 + H 2 O  NH 4 + (aq) + OH - (aq)

Bronsted-Lowry Conjugate Acid Base Pairs Conjugate base form of an acid has a “-” charge. –Cl - C 2 H 3 O 2 - (aq) Conjugate acid form of a base has a “+’ charge. –NH 4 +

Bronsted-Lowry Conjugate Acid Base Pairs Weak acid has strong conjugate base form. Strong Acid has weak conjugate base form. Weak base has strong Conjugate Acid form.

Problem Write the formula for the conjugate acid of each of the following (a) F - (b) N 2 H 4 (c) C 5 H 5 N (d) O 2 -

Problem Write the formula for the conjugate base of each of the following: (a) NH 2 OH (b) HSO 3 - (d) H 5 IO 6

Chapter 15 Acids and Bases… Bronsted Lowry Strong Acids –HCl  H + (aq) + Cl - (aq) –Use ratios in Stoichiometric Considerations –Can be Binary Compounds of Hydrogen and Nonmetals HCl, HBr, HI Strength increases from top to bottom in same group.

Chapter 15: Acids and Bases Bronsted Lowry Strong Acids –Can Be Oxyacids –HNO 3, H 2 SO 4, HClO 4 –Strength depends on electronegativity of central atom and number of oxygen atoms which are also very electronegative.

Chapter 15: Acids and Bases… Bronsted Lowry Weak Acids HC 2 H 3 O 2 + H 2 O (l)  H 3 O + (aq) + C 2 H 3 O 2 - (aq) HC 2 H 3 O 2  H + (aq) + C 2 H 3 O 2 - (aq) Stoichiometric Calculations Use Equilibrium Constant Expression Ka = [H + ] [C 2 H 3 O 2 - ]/[HC 2 H 3 O 2 ]

Chapter 15: Bronsted Lowry Bases Strong Bases NaOH  Na + (aq) + OH - (aq) Ca(OH) 2  Ca 2+ (aq) + 2OH - (aq) Stoichiometric Ratios Used in Calculations

Chapter 15: Bronsted Lowry Bases Weak Bases NH 3 + H 2 O (l) NH 4 + (aq) + OH - (aq) Usually Contain Nitrogen: Amines, Amides Stoichiometric Calculations Use Equilibrium Constant Expression K b = [NH 4 + ] [OH - ]/[NH 3 ]

Autoionization of Water: pH Scale In pure water at 25 ◦ C, H 2 O (l)  H + (aq) + OH - (aq) [H + ] [OH - ] = 1.0E-14

Autoionization of Water Calculate the [OH - ] when the [H + ] = 1.8E - 3 M Calculate the [H + ] when the [OH - ] is 6.4E-2M

Auto Ionization of Water Acid has more H + than OH - Base has more OH - than H + Neutral Solution has equal amounts of H + and OH -

From PowerPoint #16 (a) Calculate the [OH - ] when the [H + ] = 1.8E - 3 M [OH - ] = 5.6E-12 (b) Calculate the [H + ] when the [OH - ] is 6.4E-2M [H + ] = 1.6E-13

pH Equation pH = -log [H + ] (a) Calculate pH of solution when ([H + ] = 1.8E - 3 M) (b) Calculate pH when [H + ] = 1.6E-13

pH Scale Remember that [H + ] [OH - ] = 1.0E-14 Therefore [H + ] has limited range from approximately 1.0E-1 M to 1.0E-13 M.

pH Scale Acidic pH More H + than OH - pH range of Neutral pH [H + ] = [OH - ] pH = 7.0

pH Scale Basic pH More OH - than H + pH range of 7.1 to 13.99

Other pH = -log [H + ] pOH = - log [OH - ] pH + pOH = 14

Problem A solution is labeled as M Sr(OH) 2. What is the pH of the solution. Sr(OH) 2  Sr 2+ (aq) + 2 OH - (aq)

Problem Calculate pH when 10.0 ml of M HBr is mixed with 20.0 ml of M HCl.

Problem 15.24, page 703 Calculate the number of moles of KOH in 5.50 ml of M KOH solution. What is the pOH of the solution.

Problem A: Page 703 Calculate the Hydrogen ion concentration in mole/Liter for: (a) a solution whose pH is 5.20

Lewis Acids and Bases (1) BRONSTED-LOWRY NH 3(aq) + H 2 O  NH 4 + (aq) + OH - (aq) (2) Lewis Acids:Bases NH 3(aq) + BF 3  H 3 N - BF 3

Lewis Acids and Bases Lewis base is electron pair donor. Lewis Acid is electron pair acceptor. Neutralization involves formation of coordinate covalent bond. –Electrons in shared pair come from one of the two atoms participating in the bond.

Lewis Acids and Bases Fe 2+ (aq) +H 2 O (l)  Fe[H 2 O)(OH)] + (aq) + H + (aq) : NH 3(aq) + H 2 O  NH 4 + (aq) + OH - (aq)

Lewis Bases Molecules or Ions with Unshared Pair Electrons. Molecules or Ions with Complete Valence Shells O 2 - :NH 3

Lewis Acids Molecules or Ions with Incomplete Valence Shells –BF 3 H + Molecules or Ions with Complete Valence Shells but With Multiple Bonds that can be shifted to make room for more electrons :O=C=O: Transition Metal Cations –Fe 2+, Zn 2+, Cu 2+

Lewis Acids Molecules or Ions that Have Central Atoms Capable of Holding Additional Electrons –Usually Atoms of Elements in Period 3 and Below. SO 2