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Acids , Bases , and Salts All are electrolytes

Acids , Bases , and Salts All are electrolytes. Mr. Sharp playing with acids and bases. The Solubility Product K sp. Used for sparingly or slightly soluble salts. Insoluble Salts and Compounds.

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Acids , Bases , and Salts All are electrolytes

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  1. Acids, Bases,andSaltsAll are electrolytes

  2. Mr. Sharp playing with acids and bases

  3. The Solubility Product Ksp Used for sparingly or slightly soluble salts

  4. Insoluble Salts and Compounds • A compound may be said to be “Insoluble” in water, but there are always a few particles or ions that do dissolve. • A compound is said to be: • insoluble if less that 0.1 of a gram of it will dissolve in 100 grams of water • slightly soluble if between 0.1 and 1.0 grams of it will dissolve in 100 grams of water • soluble if more that 1.0 grams of it will dissolve in 100 grams of water

  5. Silver chloride is a sparingly soluble salt • Its equilibrium reaction is: • AgCl(s) Ag+(aq) + Cl-(aq) • Its equilibrium expression is: • Keq = [Ag+] [Cl-] / [AgCl] • The concentration for all pure liquids and solids is constant; therefore [AgCl] becomes part of the Keg which we call Ksp, , the solubility product. • Ksp = [Ag+] [Cl-]

  6. Value of Ksp gives a general idea of how insoluble a substance may be in water • See figure 17-6 on page 566 in Text • In general, the larger the Ksp value the more soluble the substance will be • In general, the smaller the Ksp value the less soluble the substance will be

  7. General Ksp for any compound AbCd • Equilibrium equation: • AbCd bA+d + dC-b • Equilibrium expression(Ksp) • Ksp = [A+d]b [C-b]d • Like practice problems 17-1; 2,6,10

  8. Practice problem 1 Text page565 • Equilibrium equation: • Ag2CrO4 2Ag+1 + CrO4-2 • Ksp Expression: • Ksp = [Ag+1]2 [CrO4-2]

  9. Practice Problem 2 Text page 565 • Equlibrium equation: • PbI2 Pb+2 + 2I-1 • Ksp expression: • Ksp = [Pb+2] [I-1]2

  10. Finding Ksp , given ion concentration Practice problem 3 text page 567 • Like practice problems 17-1; 14,18 • Equilibrium equation: • Cd(OH)2 Cd+2 + 2OH-1 • Ksp expression: Ksp = [Cd+2] [OH-1]2 • Given [Cd+2] = 1.7 x 10-5; therefore, • [OH-1] = 2[Cd+2] = 2(1.7 x 10-5) = 3.4 x 10-5 • Ksp= (1.7 x 10-5) (3.4 x 10-5)2 = 2.0 x 10-14

  11. Practice problem 4 Text page 567 • Equilibrium equation: • Ce(OH)3 Ce+3 + 3OH-1 • Ksp expression: Ksp = [Ce+3] [OH-1]3 • Given: [Ce+3] = 5.2 x 10-6; therefore, • [OH-] = 3[Ce+3] = 3(5.2 x 10-6) = 1.56 x 10-5 • Ksp = (5.2 x 10-6)(1.56 x 10-5 )3 = 2.0 x 10-20

  12. Finding concentration given KspPractice problem 5 Text page 570 • Equilibrium equation: • CaF2 Ca+2 + 2F- • Ksp expression: Ksp = [Ca+2] [F-]2 = 3.9 x 10-11 Let [Ca+2] = x and [F-] = 2x Therefore Ksp = (x)(2x)2 = 3.9 x 10-11 4x3 = 3.9 x 10-11; x3 = (3.9 x 10-11)(4) x3 = 9.75 x 10-12; x = (9.75 x 10-12) [Ca+2] = x = 2.1 x 10 -4 ; [F-] = 2x = 4.2 x 10-4

  13. Practice problem 6 Text page 570 • Like practice problems 17-1; 22,26,30 • Equilibrium equation: • BaCrO4 Ba+2 + CrO4-2 • Ksp = [Ba+2] [CrO4] = 2.0 x 10-10 • Let x = [Ba+2] = [CrO4]; therefore • (x) (x) = x2 = 2.0 x 10-10 • x = (2.0 x 10-10)½ = 1.4 x 10-5

  14. Properties ofAcids • Taste sour (we test not taste) • Turns blue litmus red • Neutralizes bases and basic(metallic)oxides • React with metals like Zn to produce hydrogen gas • Solution in water are electrolytes • Examples:HCl, HNO3, H2SO4, • HCH3COO (CH3COOH) (HC2H3O2)

  15. Propeties of Bases • Taste bitter(we test not taste) • Feel slippery • Turn red litmus blue • Neutralize acids and acidic(nonmetallic) oxides • Solutions in water are electrolytes • Corrosive • Examples: NaOH, KOH, Ca(OH)2, NH3

  16. AcidandBaseTheories • Arrhenius - 1887 • Bronsted-Lowery - 1923 • Lewis - 1923

  17. Arrhenius Theory • Acids - substances that ionize in water to produce hydrogen ions(H+) • Bases - substances that ionize in water to produce hydroxide ions(OH-)

  18. Arrhenius Examples • Arrhenius Acids: • HCl  H+ + Cl- • HNO3  H+ + NO3- • Arrhenius Bases: • NaOH  Na+ + OH- • Ca(OH)2  Ca2+ + 2OH-

  19. Bronsted-Lowry Theory • Acids - substancesthat donate protons(Proton donors) • Bases - substances that accept protons(Proton acceptors) • Hydrogen ions are really the same as protons

  20. Hydronium Ion(hydrogen ion riding piggy-back on a water molecule)

  21. Bronsted-Lowry Example: Water acting as an acid • Ammonia + water yield ammonium ion plus hydroxide ion • NH3 + H2O  NH4+ + OH- base acid conjugate conjugate acid base

  22. Bronsted Lowry Example:Water acting as a base Acid base Conjugate acid conjugate base

  23. Amphiteric = being ableto act as anacidor base • acid = proton donor • base = proton acceptor

  24. Lewis Theory • Acids - substances that accept electron-pairs • Bases - substances that donate electron-pairs • Must draw electron dot formula to determine

  25. Lewis example: Lewis base Lewis acid

  26. Lewis Example

  27. Lewis Example: • Lewis Lewis • acid base

  28. Lewis Theory is most inclusive, but we will use Bronsted-Lowry, mostly

  29. Naming BinaryAcids • Made up of Hydrogen and a nonmetal (two elements) • Use prefix of hydro • Use the nonmetal as the root of the name • Add suffix of ic • Example: HCl = hydrochloric

  30. HBr HI HF H2O H2S H2Te H2Se hyrdobromic acid hydroiodic acid hydrofluoric acid hydroxyic acid hydrosulfuric acid hydrotelluric acid hydroselenic acid Practice NamingAcids

  31. Naming Oxyacids • Name most common acid in family with the root of nonmetal other than oxygen then add suffix of ic • Acid of family with one more oxygen than most common add prefix of per and suffix of ic • Acid with one less oxygen than most common use suffix of ous • Acid with two less oxygens than most common add prefix of hypo and keep suffix of ous

  32. HClO4 (one more O) per means more than HClO3(most common) HClO2 (one less O) HClO (two less O’s) hypo means less than perchloric acid chloric acid chlorous acid hypochlorous acid Naming tertiary(oxyacids) Usechlorineoxyacidfamily as a guide:

  33. H2SO4(most common) H2SO3 HNO3(most common) HNO2 HBrO4 HBrO3(most common) HBrO2 HBrO Sulfuric acid sulfurous acid nitric acid nitrous acid perbromic acid bromic acid bromous acid hypobromous acid Namingoxyacids

  34. NamingArrhenius bases • Name positive ion first. Then add name hydroxide last • NH4OH = ammoniumhydroxide • KOH = potassiumhydroxide • Ca(OH)2 = calcium hydroxide • Mn(OH)7 = manganese(VII) hydroxide

  35. LiOH Ba(OH)2 Al(OH)3 Sn(OH)4 Lithium hydroxide barium hydroxide aluminum hydroxide tin(IV) hydroxide Practice Naming bases

  36. Anydrides means without water • Acid anhydride = an acid without water(nonmetallic oxides • SO2 + H2O  H2SO3 • N2O5 + H2O  2HNO3 • Basic anydride = base without water(metallic oxides) • BaO + H2O  Ba(OH)2 • Na2O + H2O  2NaOH

  37. Determining anhydride by subtracting water • H2SO4 - H2O = SO3 • 2HNO3 - H2O = N2O5 • 2H3AsO4 - 3H2O = As2O5 • Ba(OH)2 - H2O = BaO • 2NaOH - H2O = Na2O • 2Al(OH)3 - 3H2O = Al2O3 • All H atoms must add out

  38. Acidsand bases neutralize each other • acid + base salt + water parent parent  child + water acid base salt 1. HCl + NaOH  NaCl + H2O 2. HClO4+NaOH NaClO4 + H2O 3. HClO3+NaOH NaClO3 + H2O 4. HClO2+ NaOH  NaClO2+ H2O 5. HClO+NaOH NaClO + H2O

  39. Rules for namingsalts • Named from parent acid and base • First name comes from parent base • Second name comes from parent acid • Salts from binary acids end in ide • Salts from oxyacids: • Salt from most common, use suffix of ate • Salt from one more oxygen use per prefix and ate suffix • Salt from one less oxygen use ite suffix • Salt from two less oxygens use hypo prefix and ite suffix

  40. NaCl NaClO4 NaClO3 NaClO2 NaClO sodium chloride sodium perchlorate sodium chlorate sodium chlorite sodium hypochlorite Practice naming salts from slide #38

  41. KBrO3 Mn2(SO4)7 CuSO3 Mg(BrO4)2 BaI2 LiBrO NaNO3 KNO2 Potassium bromate manganese(VII) sulfate copper(II) sulfite magnesium perbromate barium iodide lithium hypobromite sodium nitrate potassium nitrite More Practice namingsalts: Go to slide #39 to answer the following

  42. Strongacidsand bases • Strong acids and bases are nearly 100% ionized • HCl + H2O H3O+ + Cl- • If 0.10 mole of HCl are placed in water, we get 0.10 mole of H3O+ • NaOH  Na+ + OH- • If 0.02 mole of NaOH are place inwaterwe get 0.02 mole of OH-

  43. Weakacidsand bases • Weak acids and bases are only slightly ionized • HC2H3O2 + H2O H3O+ + C2H3O2- • Weak acid produces very few H3O+, [H3O+] must be calculated • NH3 + H2O  NH4+ + OH- • Weak base produces very few OH-, •  [OH-] must be calculated

  44. Self Ionization ofwater • H2O + H2O H3O++ OH- • Keq = [ H3O+] [ OH-] / [ H2O]2 • multiply each side by [H2O]2 and let Keq[H2O]2 = Kw • Kw = [H3O+] [ OH-] = 1.00 x 10-14

  45. Concentration of ions in pure water • In pure water [H3O+] = [OH-] = 10-7 M • In an acid solution, [H3O+] > [OH-] • In a basic solution, [H3O+] < [OH-] • In a neutral solution [H3O+] = [OH-]

  46. Calculating [OH-] or [H3O+] • Given that [OH-] = 4.78 x 10-12, determine the [H3O+] and if the solution is acidic or basic. • [[H3O+] [OH-] = 1 x 10-14 • [H3O+] = (1 x 10-14)  [OH-] • [H3O+] = (1 x 10-14)/ (4.78 x 10-12) = • 2.09 x 10-3 • [H3O+] > [OH-] therefore solution is acidic

  47. Solving # 2, 19-1 Practice problems • [H3O+] [OH-] = 1 x 10-14 • [OH-] = 1 x 10-8 • [H3O+] = (1 x 10-14)  (1 x 10-8) • [H3O+] = 1 x 10-6 M • [H3O+] > [OH-], therefore solution is acid

  48. pH is a method ofexpressing theacidity of a water solution • pH = -log[H3O+] • [H3O+] = 10-pH • In pure water, pH = 7; neutral • In acid solution, pH < 7 • In basic solution, pH > 7

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