1. Acids & Bases Acids:  acids are sour tasting  Arrhenius acid  Arrhenius acid: Any substance that, when dissolved in water, increases the concentration of hydronium ion (H 3 O + )  Bronsted-Lowry acid  Bronsted-Lowry acid: A proton donor  Lewis acid  Lewis acid: An electron acceptorBases:  bases are bitter tasting and slippery  Arrhenius base  Arrhenius base: Any substance that, when dissolved in water, increases the concentration of hydroxide ion (OH - )  Bronsted-Lowery base  Bronsted-Lowery base: A proton acceptor  Lewis acid  Lewis acid: An electron donor

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

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

4. Brønsted-Lowry Theory of Acids & Bases

6. Notice that water is both an acid & a base = amphoteric Reversible reaction

7. Conjugate Acid-Base Pairs

9. Acids & Bases CB 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  AlCl 3 + 3 H 2 O b. Ba(OH) 2 + 2 HC 2 H 3 O 2  Ba(C 2 H 3 O 2 ) 2 + 2 H 2 O c. 2 KOH + H 2 SO 4  K 2 SO 4 + 2 H 2 O d. NH 3 + H 2 O  NH 4 + + OH - CB CA CA CB

10. 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) + H 2 O  H 3 O + + 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. NH 3 + H 2 O  NH 4 + + OH -

12. Acids & Bases STRONG vs WEAK STRONG vs WEAK _ completely ionized_ partially ionized _ strong electrolyte_ weak electrolyte _ ionic/very polar bonds_ some covalent bonds Strong Acids : Strong Bases: HClO 4 LiOH H 2 SO 4 NaOH HIKOH HBrCa(OH) 2 HClSr(OH) 2 HNO 3 Ba(OH) 2

13. Acids & Bases One ionizable proton: HCl → H + + Cl - Two ionizable protons: H 2 SO 4 → H + + HSO 4 - HSO 4 - → H + + SO 4 2- Three ionizable protons: H 3 PO 4 → H + + H 2 PO 4 – H 2 PO 4 - → H + + HPO 4 2- HPO 4 2- → H + + PO 4 -3 Combined: H 2 SO 4 → 2H + + SO 4 2- Combined: H 3 PO 4 → 3H + + PO 4 3-

14. Acids & Bases For the following identify the acid and the base as strong or weak. a. Al(OH) 3 + HCl  b. Ba(OH) 2 + HC 2 H 3 O 2  c. KOH + H 2 SO 4  d. NH 3 + H 2 O  Weak baseStrong acid Weak acid Strong acid Strong base Weak base Weak acid

15. Acids & Bases 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 + HC 2 H 3 O 2  c. KOH + H 2 SO 4  d. NH 3 + H 2 O  AlCl 3 + 3 H2OH2O Ba(C 2 H 3 O 2 ) 2 + 2 H2OH2O K 2 SO 4 + 2 H2OH2O NH 4 + + OH - 2 2 3

16. Water Equilibrium

17. K w = [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

18. pH A measure of the hydronium ion The scale for measuring the hydronium ion concentration [H 3 O + ] 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 [H 3 O + ] of 0.1 mol/L or 10 -1 A solution with a pH of 3 has [H 3 O + ] of 0.001 mol/L or 10 -3 A solution with a pH of 7 has [H 3 O + ] of 0.0000001 mol/L or 10 -7 pH = - log [H 3 O + ]

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

20. Manipulating pH Algebraic manipulation of: pH = - log [H 3 O + ] allows for: [H 3 O + ] = 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!

21. TITRATION 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 (H 2 O). 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 n acid = n base Since M=n/V M A V A = M B V B

22. TITRATION 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  H 2 O + NaCl M a V a = M b V b rearranges to M a = M b V b / V a so M a = (0.52 M) (22.50 mL) / (75.00 mL) = 0.16 M = 0.16 M Now you try: 2. If 37.12 mL of 0.843 M HNO 3 neutralized 40.50 mL of KOH, what is the molarity of the base? M b = 0.773 mol/L

23. Molarity and Titration

24. TITRATION 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 (H 2 O). 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 n acid = n base

25. TITRATION 1. If 37.12 mL of 0.543 M LiOH neutralized 40.50 mL of H 2 SO 4, what is the molarity of the acid? 2 LiOH + H 2 SO 4  Li 2 SO 4 + 2 H 2 O 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 H 2 SO 4 / 2 mol LiOH)= 0.0101 mol H 2 SO 4 : Last calculate the Molarity: M a = n/V = 0.010 mol H 2 SO 4 / 0.4050 L = 0.248 M M a = n/V = 0.010 mol H 2 SO 4 / 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? M b = 0.308 mol/L

26. 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 NaHCO 3 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: H 2 SO 4 + 2 NaHCO 3 → Na 2 SO 4 + 2 H 2 O + 2 CO 2

27. PRACTICE PROBLEM #25 1. How many milliliters of 1.25 M LiOH must be added to neutralize 34.7 mL of 0.389 M HNO 3 ? 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 H 2 SO 4 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

28. 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 H 2 SO 4 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 ?