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Wed. June 4 Announcements:

Wed. June 4 Announcements:. Ch 16 Test Retakes: Take by Friday at latest! Less than 65% See me to review your original test Extra Credit Due Wed June 11 Max 5 points Extra credit cleaning point for Ch 16 test- see me if still need to earn yours! Final Exam Review Packets.

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Wed. June 4 Announcements:

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  1. Wed. June 4 Announcements: • Ch 16 Test Retakes: • Take by Friday at latest! • Less than 65% • See me to review your original test • Extra Credit • Due Wed June 11 • Max 5 points • Extra credit cleaning point for Ch 16 test- see me if still need to earn yours! • Final Exam Review Packets

  2. Chapter 17: Chemistry of Acids & Bases

  3. 17.0 Objectives • 1. To recognize the properties of strong and weak acids and strong and weak bases. • 2. To become familiar with Arrhenius and Bronsted-Lowry theories of acids and bases. • 3. Be able to calculate and understand the following quantities: [H+], [OH-], pH, pOH, Kw, Ka, and Kb. • 4. Calculate the equilibrium concentrations of weak acids and bases. • 5. Classify aqueous salt solutions as acidic, basic, or neutral. • 6. Determine the concentration of unknown acid or base, or the Ka of a weak acid from a titration procedure.

  4. Homework • HW #1 - 11, 13, 15, 17, 19, 21 • Conjugate Acids/Bases; pH • HW #2 - 25, 27, 29, 31, 33, 37 • Equilibrium Constants and pKa • HW#3 - 39, 41, 49 • pKb, Types of Reactions • HW#4 - 51, 53, 55 • pH to calculate K values • HW#5 - 57, 59, 97, 101 • Equilibrium problems

  5. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • A. Definitions • 1. Arrhenius definition of acid • produce H+ in water solution • Ex. HCl; H2SO4 HCl H+ + Cl- • 2. Arrhenius definition of base • produce OH- in water solution • Ex. NaOH, Ca(OH)2 NaOH  Na+ + OH-

  6. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • B. Properties of acids • Ionize when put into water • React with active metals (Group I, II) to produce Hydrogen gas • Neutralize bases to form water and salt • Have a sour taste • Turns blue litmus red

  7. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • C. Properties of bases • Ionize when put into water • Neutralize acids to form water and salt • Have a bitter taste • Feel Slippery • Turn red litmus blue

  8. Definition: Hydronium Ion In aqueous solution, H+ does NOT exist! Note: In problems, [H+] = [H3O+] H+ + H2O  H3O+ (hydronium ion) Dr. Mihelcic AP Chemistry 1

  9. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • D. REVIEW: Strong acids • Strong acids dissociate completely in aqueous solution: There are 6 Strong Acids • HCl, HBr, HI, HNO3, HClO4, H2SO4 • Examples: HCl  H+ + Cl- HNO3 H+ + NO3- Note single arrow!! No Equilibrium Established!

  10. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT Weak acids are much less than 100% ionized in water. (use double arrow; equilibrium established) • Weak acids • Weak acids are partly dissociated in soln. • Treated as an equilibrium • Ex. CH3COOH acetic acid • Species • H+ • H3O+ • [insert H-containing cation here]+

  11. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • F. Example 17.1 Write equations for the aqueous dissociation of the following weak acids: • HPO4-2 • Fe(H2O)63+ • HCN

  12. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • G. Strong Bases: Review • Strong bases completely dissociate in aqueous solution: • Group I metal hydroxides LiOH, NaOH, KOH, RbOH, CsOH • MOST Group II Hydroxides Ca(OH)2, Sr(OH)2, Ba(OH)2, Examples: NaOH - Sr(OH)2

  13. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT Weak bases are less than 100% ionized in water.(use double arrow; equilibrium established) • H. Weak bases • partly dissociated in soln • Treated as an equilibrium • Two types of substances act like weak bases in aqueous solution: • Nitrogen-containing compounds (why?) • Ex. NH3 NH3(aq) + H2O(liq) NH4+(aq) + OH-(aq) • Anions of acids • Ex. F- , HCO3-

  14. Strong Acids and bases are strong electrolytes; weak acids and bases are weak electrolytes!

  15. 17.1 ACIDS, BASES, AND THE EQUILIBRIUM CONCEPT • I. Example 17.2 Write equations for the aqueous dissociation of the following weak bases: NO21- CO32- C2H5NH2

  16. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • A. Arrhenius definitions are Narrow Problems: • Only Includes aqueous systems • Limits acids and bases to contain H+ and OH-

  17. 17.2 BRONSTED CONCEPT OF ACIDS AND BASESA more general definition than Arrehenius theory… • B. Bronsted-Lowry Acid • ACIDS DONATE H+ IONS • Acid = Proton Donor (ADP: Acids Donate Protons) • Ex. HNO3 • C. Bronsted-Lowry Base • BASES ACCEPT H+ IONS • Base= Proton acceptor (BAP: Bases Accept Protons)

  18. The Brønsted definition means NH3 is a BASEin water, and water is itself an ACID Bases Accept Protons Acids Donate Protons

  19. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • D. Conjugate acids and bases • Conjugate acid: Bronsted Base with an extra proton (H+) attached • Every base has a conjugate acid • Conjugate base: Bronsted acid minus a proton (H+) • Every acid has a conjugate base

  20. Conjugate Pairs • NH3 / NH4+ is a conjugate pair — related by the gain or loss of H+ • Every acid has a conjugate base, formed when H+ is removed from the acid. • Every base has a conjugate acid, formed when H+ is added to the base.

  21. Conjugate Pairs Generalized equation: HB (aq) + A (aq)  HA (aq) + B- (aq)

  22. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES Some substances can function as both an ACID OR a BASE, depending on what they are reacted with. (can donate OR accept H+) • E. Amphiprotic substances • Substances that can ionize as either an acid or a base depending on the properties of the other species in soln. • Ex. H2O • can act as a conjugate acid or a conjugate base H2O + H2O  H3O+ + OH-

  23. H2O + H2O  H3O+ + OH- • Ex. H2O • can act as a conjugate acid or a conjugate base

  24. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • F. Example 17.3 Identify the Bronsted Lowry acids and bases and their conjugates in the following equations: • NH4+(aq) + OH1-(aq)  NH3(aq) + HOH(l) • N2H4(aq) + HOH(l)  N2H51+(aq) + OH1-(aq)

  25. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • G. Example 17.4 Write equations to show that HSO41- can act like an amphiprotic substance.

  26. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES Extra Example: Example 17.4 b Write the Bronsted Lowry equations for the weak acid HCO3 - and the weak base NH3, identifying the conjugate acid-base pairs in each equilibrium.

  27. Extra Example: Example 17.5 Write two equations for each substance below showing amphiprotic behavior. The first equation should show proton donation, and the second proton acceptance. • a. HSO4- • b. HCO3-

  28. 17.2 BRONSTED CONCEPT OF ACIDS AND BASES • H. Polyprotic acids • Acids that give off more than 1 H+ when put into water • 1st Proton usually given off rapidly • Subsequent protons are given off with increasing difficulty (stronger bases at each step- why?) • Ex. H2SO4 HSO4- SO42- Another example: Oxalic acid (diprotic)

  29. Oxalic acid is diprotic, having two protons to donate Another example: phosphoric acid (Triprotic) Donates one H+ at a time: H3PO4 H+ + H2PO4- H2PO4-  H+ + HPO4-2 HPO4-2  H+ + PO4-3

  30. 17.3 WATER AND THE pH SCALE • A. Water -self-ionization, and the water dissociation constant, Kw Self-Ionization: (H2O ↔ H+ + OH-) H2O + H2O ↔H3O+ + OH- Acid(1) Base(2) ↔ Acid(2) Base(1) • For any sample of water molecules: 2 H2O (l)↔ H3O+ (aq) + OH- (aq) • Kw= [H3O+] [OH-] = 1.00 x 10-14(at 25 oC)

  31. 17.3 WATER AND THE pH SCALE • B. Relationships for all aqueous solutions • In a neutral soln: • [H3O+] = [OH-] = 1.0 x 10-7 M • In an acidic soln: [H3O+] > [OH-] • [H3O+] > 1.0 x 10-7 M • [OH-] < 1.0 x 10-7 M • In a basic soln: [H3O+] < [OH-] • [H3O+] < 1.0 x 10-7 M • [OH-] > 1.0 x 10-7 M Remember: Kw = [H3O+] [OH-] = 1.00 x 10-14

  32. A common way to express acidity and basicity is with pH (the “power of hydrogen”) 17.3 WATER AND THE pH SCALE • C. The pH scale (Ch. 5 review) • Ranges from 0 to 14 • 0~7 = ACID • 7~14 = BASE • 7 = Neutral

  33. 17.3 WATER AND THE pH SCALE • D. Definition of pH and pOH • pH = - log [H+] • [H3O+] = 10-pH • pOH = - log [OH-] • [OH-] = 10-pOH • pH + pOH = 14 Remember, pure water (neutral) [H3O+] = [OH-] = 1.0 x 10-7 M

  34. pH 0: [H+] = 0.1 M pH 14: [H+] =0.00000000000001 M pOH 14: [OH-]=0.1 M pOH 0: [OH-]=0.00000000000001 M

  35. pH to [H+] Calculations If the pH of Coke is 3.12, it is acidic. Because pH = - log [H3O+] then log [H3O+] = - pH Take antilog and get [H3O+] = 10-pH [H3O+] = 10-3.12 = 7.6 x 10-4 M

  36. 17.3 WATER AND THE pH SCALE [H+]=5.0 x 10-3M pH = 7.4 • E. Example 17.5 Calculate the [H+], [OH-], pH, and pOH for a. lemon juice with [H+]=5.0 x 10-3M, and b. blood with a pH of 7.4.

  37. 17.3 WATER AND THE pH SCALE [H+]=5.0 x 10-3M pH = 7.4 • E. Example 17.5 Calculate the [H+], [OH-], pH, and pOH for a. lemon juice with [H+]=5.0 x 10-3M, and b. blood with a pH of 7.4.

  38. Note: Strong Acids and Bases Since strong Acids dissociate completely in aqueous solution: E.g. HCl  H+ + Cl- [H+] can be calculated from the molarity of the acid. 2.0 M 2.0 M 2.0 M Likewise, Strong Bases dissociate completely in aqueous solution: E.g. Sr(OH)2 Sr+2 + 2OH- So [OH-] can be calculated from the molarity of the base. 0.5 M 0.5 M 1.0 M

  39. Note: Weak Acids and Weak Bases- exixist in solution in equilibrium with conjugate Since srong Acids/bases do ionize less than 100% , we Must use ICE tables for all calculations involving weak acids!

  40. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • A. Weak acids, Ka • Equilibrium of an acid with conjugate base • **REMEMBER** H2O (l) is not included! HA + H2O <-----> H3O+ + A- Ka=

  41. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • B. Value • Less than 1 because it is a weak acid (more reactant compared to product) • Ka increases as the strength of acid increases • Kb increases as strength of base increases Weak acid has Ka < 1 Leads to small [H3O+] Weak base has Kb < 1 Leads to small [OH-]

  42. Acids Conjugate Bases Increase strength Increase strength Dr. Mihelcic AP Chemistry 1

  43. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • C. pKa and value • pKa = - log Ka • Value of pKa is smaller as the strength of the acid increases Higher Ka means stronger acid Higher Ka means lower pKa

  44. 17.4 EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES • D. Example 17.6 In a solution prepared by dissolving 0.100 mole of lactic acid per liter, [H+] = 3.7 x 10-3M. Calculate the Ka for lactic acid. Strategy: • Write the equation showing the disassociation of the acid. • Set up an IRE table and substitute [H+] onto the “E” line. • Calculate the other “E” concentrations. • Solve for the Ka.

  45. Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc- and the pH. Note that we neglect [H3O+] from H2O. Step 1. Set up IRE table. [HOAc] [H3O+] [OAc-] initial reacted equilib

  46. Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 3. Write Ka expression This is a quadratic. How can we simplify it??

  47. Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 4. Solve Ka expression First assume x is very small because Ka is so small. Therefore, Now we can more easily solve this approximate expression. Dr. Mihelcic AP Chemistry 1

  48. Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 5. Solve Kaapproximate expression x = [H3O+] = [OAc-] = [Ka • 1.00]1/2 x = [H3O+] = [OAc-] = 4.2 x 10-3 M pH = - log [H3O+] = -log (4.2 x 10-3) = 2.37 Dr. Mihelcic AP Chemistry 1

  49. Approximations for Weak Acids Consider the approximated expression : For many weak acids, If Ka <= 1.0 x 10-4, can simplify the math as follows: Initial conc. – x ~ Initial concentration Dr. Mihelcic AP Chemistry 1

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