Acids & Bases Acids: acids are sour tasting Arrhenius acid: Any substance that, when dissolved in water, increases the concentration of hydronium ion (H3O+) Bronsted-Lowry acid: A proton donor Lewis acid: An electron acceptor Bases: bases are bitter tasting and slippery Arrhenius base: Any substance that, when dissolved in water, increases the concentration of hydroxide ion (OH-) Bronsted-Lowery base: A proton acceptor Lewis acid: An electron donor

Lone Hydrogen ions do not exist by themselves in solution Lone Hydrogen ions do not exist by themselves in solution. H+ is always bound to a water molecule to form a hydronium ion

General Equation Brønsted-Lowry Theory of Acids & Bases Conjugate Acid-Base Pairs General Equation

Brønsted-Lowry Theory of Acids & Bases

Brønsted-Lowry Theory of Acids & Bases

Brønsted-Lowry Theory of Acids & Bases Notice that water is both an acid & a base = amphoteric Reversible reaction

ELECTROLYTES Electrolytes are species which conducts electricity when dissolved in water. Acids, Bases, and Salts are all electrolytes. Salts and strong Acids or Bases form Strong Electrolytes. Salt and strong acids (and bases) are fully dissociated therefore all of the ions present are available to conduct electricity. HCl(s) + H2O  H3O+ + Cl- Weak Acids and Weak Bases for Weak Electrolytes. Weaks electrolytes are partially dissociated therefore not all species in solution are ions, some of the molecular form is present. Weak electrolytes have less ions avalible to conduct electricity. NH3 + H2O  NH4+ + OH-

Acids & Bases STRONG vs WEAK H2SO4 NaOH HI KOH HBr Ca(OH)2 HCl Sr(OH)2 _ completely ionized _ partially ionized _ strong electrolyte _ weak electrolyte _ ionic/very polar bonds _ some covalent bonds Strong Acids: Strong Bases: HClO4 LiOH H2SO4 NaOH HI KOH HBr Ca(OH)2 HCl Sr(OH)2 HNO3 Ba(OH)2

Acids & Bases One ionizable proton: HCl → H+ + Cl- Two ionizable protons: H2SO4 → H+ + HSO4- HSO4- → H+ + SO42- Three ionizable protons: H3PO4 → H+ + H2PO4– H2PO4- → H+ + HPO42- HPO42- → H+ + PO4-3 Combined: H2SO4 → 2H+ + SO42- Combined: H3PO4 → 3H+ + PO43-

Acids & Bases a. Al(OH)3 + HCl  b. Ba(OH)2 + HC2H3O2  For the following identify the acid and the base as strong or weak . a. Al(OH)3 + HCl  b. Ba(OH)2 + HC2H3O2  c. KOH + H2SO4  d. NH3 + H2O  Weak base Strong acid Strong base Weak acid Strong acid Strong base Weak acid Weak base

Acids & Bases a. Al(OH)3 + HCl  b. Ba(OH)2 + HC2H3O2  For the following predict the product. To check your answer left click on the mouse. Draw a mechanism detailing the proton movement. a. Al(OH)3 + HCl  b. Ba(OH)2 + HC2H3O2  c. KOH + H2SO4  d. NH3 + H2O  3 AlCl3 + 3 H2O 2 Ba(C2H3O2)2 + 2 H2O K2SO4 + 2 H2O 2 NH4+ + OH-

Conjugate Acid-Base Pairs

Conjugate Acid-Base Pairs

Acids & Bases a. Al(OH)3 + 3 HCl  AlCl3 + 3 H2O For the following Identify the conjugate acid and the conjugate base. The conjugate refers to the acid or base produced in an acid/base reaction. The acid reactant produces its conjugate base (CB). a. Al(OH)3 + 3 HCl  AlCl3 + 3 H2O b. Ba(OH)2 + 2 HC2H3O2  Ba(C2H3O2)2 + 2 H2O c. 2 KOH + H2SO4  K2SO4 + 2 H2O d. NH3 + H2O  NH4+ + OH- CB CA CB CA CB CA CA CB

TITRATION nacid = nbase MAVA = MBVB Titration of a strong acid with a strong base ENDPOINT = POINT OF NEUTRALIZATION = EQUIVALENCE POINT At the end point for the titration of a strong acid with a strong base, the moles of acid (H+) equals the moles of base (OH-) to produce the neutral species water (H2O). If the mole ratio in the balanced chemical equation is 1:1 then the following equation can be used. MOLES OF ACID = MOLES OF BASE nacid = nbase Since M=n/V MAVA = MBVB

Ma Va = Mb Vb rearranges to Ma = Mb Vb / Va TITRATION MAVA = MBVB 1. Suppose 75.00 mL of hydrochloric acid was required to neutralize 22.50 mLof 0.52 M NaOH. What is the molarity of the acid? HCl + NaOH  H2O + NaCl Ma Va = Mb Vb rearranges to Ma = Mb Vb / Va so Ma = (0.52 M) (22.50 mL) / (75.00 mL) = 0.16 M Now you try: 2. If 37.12 mL of 0.843 M HNO3 neutralized 40.50 mL of KOH, what is the molarity of the base? Mb = 0.773 mol/L

Molarity and Titration

TITRATION nacid = nbase Titration of a strong acid with a strong base ENDPOINT = POINT OF NEUTRALIZATION = EQUIVALENCE POINT At the end point for the titration of a strong acid with a strong base, the moles of acid (H+) equals the moles of base (OH-) to produce the neutral species water (H2O). If the mole ratio in the balanced chemical equation is NOT 1:1 then you must rely on the mole relationship and handle the problem like any other stoichiometry problem. MOLES OF ACID = MOLES OF BASE nacid = nbase

TITRATION 1. If 37.12 mL of 0.543 M LiOH neutralized 40.50 mL of H2SO4, what is the molarity of the acid? 2 LiOH + H2SO4  Li2SO4 + 2 H2O First calculate the moles of base: 0.03712 L LiOH (0.543 mol/1 L) = 0.0202 mol LiOH Next calculate the moles of acid: 0.0202 mol LiOH (1 mol H2SO4 / 2 mol LiOH)= 0.0101 mol H2SO4 Last calculate the Molarity: Ma = n/V = 0.010 mol H2SO4 / 0.4050 L = 0.248 M 2. If 20.42 mL of Ba(OH)2 solution was used to titrate29.26 mL of 0.430 M HCl, what is the molarity of the barium hydroxide solution? Mb = 0.308 mol/L

Molarity and Titration A student finds that 23.54 mL of a 0.122 M NaOH solution is required to titrate a 30.00-mL sample of hydr acid solution. What is the molarity of the acid? A student finds that 37.80 mL of a 0.4052 M NaHCO3 solution is required to titrate a 20.00-mL sample of sulfuric acid solution. What is the molarity of the acid? The reaction equation is: H2SO4 + 2 NaHCO3 → Na2SO4 + 2 H2O + 2 CO2

Water Equilibrium

Equilibrium constant for water Water Equilibrium Kw = [H+] [OH-] = 1.0 x 10-14 Equilibrium constant for water Water or water solutions in which [H+] = [OH-] = 10-7 M are neutral solutions. A solution in which [H+] > [OH-] is acidic A solution in which [H+] < [OH-] is basic

pH A measure of the hydronium ion The scale for measuring the hydronium ion concentration [H3O+] in any solution must be able to cover a large range. A logarithmic scale covers factors of 10. The “p” in pH stands for log. A solution with a pH of 1 has [H3O+] of 0.1 mol/L or 10-1 A solution with a pH of 3 has [H3O+] of 0.001 mol/L or 10-3 A solution with a pH of 7 has [H3O+] of 0.0000001 mol/L or 10-7 pH = - log [H3O+]

The pH scale pH = - log [H3O+] 1 2 3 4 5 6 7 8 9 10 11 12 13 14 The pH scale ranges from 1 to 10-14 mol/L or from 1 to 14. pH = - log [H3O+] 1 2 3 4 5 6 7 8 9 10 11 12 13 14 acid neutral base

pH + pOH = 14 ; the entire pH range! Manipulating pH Algebraic manipulation of: pH = - log [H3O+] allows for: [H3O+] = 10-pH If pH is a measure of the hydronium ion concentration then the same equations could be used to describe the hydroxide (base) concentration. [OH-] = 10-pOH pOH = - log [OH-] thus: pH + pOH = 14 ; the entire pH range!

PRACTICE PROBLEM #25 10.8 mL 0.0101 g 29.6 mL 5.623 x 10-9 M pH = 4.0 1. How many milliliters of 1.25 M LiOH must be added to neutralize 34.7 mL of 0.389 M HNO3? 2. What mass of Sr(OH)2 will be required to neutralize 19.54 mL of 0.00850 M HBr solution? 3. How many mL of 0.998 M H2SO4 must be added to neutralize 47.9 mL of 1.233 M KOH? 4. What is the molar concentration of hydronium ion in a solution of pH 8.25? 5. What is the pH of a solution that has a molar concentration of hydronium ion of 9.15 x 10-5? 6. What is the pOH of a solution that has a molar concentration of hydronium ion of 8.55 x 10-10? 10.8 mL 0.0101 g 29.6 mL 5.623 x 10-9 M pH = 4.0 pOH = 4.9

GROUP STUDY PROBLEM #25 ______1. How many milliliters of 0.75 M KOH must be added to neutralize 50.0 mL of 2.50 M HCl? ______2. What mass of Ca(OH)2 will be required to neutralize 100 mL of 0.170 M HCl solution? ______3. How many mL of 0.554 M H2SO4 must be added to neutralize 25.0 mL of 0.9855 M NaOH? ______ 4. What is the molar concentration of hydronium ion in a solution of pH 2.45? ______ 5. What is the pH of a solution that has a molar concentration of hydronium ion of 3.75 x 10-9? ______ 6. What is the pOH of a solution that has a molar concentration of hydronium ion of 4.99 x 10-4?