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Background. A ( g ) + 2 B ( g ) 3 C ( g ) + D ( g ).

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  1. Background A(g) + 2 B(g) 3 C(g) + D(g) A florence flask was getting dressed for the opera. All of a sudden she screamed: "Erlenmeyer, my joules! Somebody has stolen my joules!". The husband replied: "Take it easy honey, do not overreact. We'll find a solution". [Products] Equilibrium constant (Keq) = [Reactants] 3 [C] [D] Keq = LeChatelier’s Principle (lu-SHAT-el-YAY’s) 2 [A][B]

  2. Acids, Bases, and Salts Ch. 19

  3. Acids • Properties • Taste sour or tart • Change the color of an acid-base indicator • Can be strong or weak electrolytes in aqueous solution

  4. Bases • Properties • Taste bitter • Feel slippery • Will change the color of an acid-base indicator • Can be strong or weak electrolytes in aqueous solution Q: Why do chemistry professors like to teach about ammonia? A: Because it's basic stuff.

  5. During class, the chemistry professor was demonstrating the properties of various acids. “Now I’m going to drop this silver coin into this glass of acid. Will it dissolve?” “No sir,” one student called out. “No?” queried the professor. “Perhaps you can explain why the silver won’t dissolve in this particular acid.” “Because if it would, you wouldn’t have dropped it in!”

  6. Acid vs. Base Alike Different Different Affects pH and litmus paper pH < 7 pH > 7 Topic Topic sour taste Related to H+ (proton) concentration bitter taste Acid Base react with metals pH + pOH = 14 does not react with metals

  7. electrolytes Properties ACIDS BASES • electrolytes • bitter taste • sour taste • turn litmus red • turn litmus blue • react with metals to form H2 gas • slippery feel • vinegar, milk, soda, apples, citrus fruits • ammonia, lye, antacid, baking soda ChemASAP

  8. Strong Bases (strong electrolytes) NaOH sodium hydroxide KOH potassium hydroxide Ca(OH)2 calcium hydroxide Weak Acids (weak electrolytes) Weak Base (weak electrolyte) CH3COOH acetic acid H2CO3 carbonic NH3 ammonia Common Acids and Bases Strong Acids (strong electrolytes) HCl hydrochloric acid HNO3 nitric acid HClO4 perchloric acid H2SO4 sulfuric acid NH3 + H2O  NH4OH

  9. Common Acids FormulaName of AcidName of Negative Ion of Salt HF hydrofluoric fluoride HBr hydrobromic bromide HI hydroiodic iodide HCl hydrochloric chloride HClO hypochlorous hypochlorite HClO2 chlorous chlorite HClO3 chloric chlorate HClO4 perchloric perchlorate H2S hydrosulfuric sulfide H2SO3 sulfurous sulfite H2SO4 sulfuric sulfate HNO2 nitrous nitrite HNO3 nitric nitrate H2CO3 carbonic carbonate H3PO3 phosphorous phosphite H3PO4 phosphoric phosphate

  10. Common Bases Sodium hydroxide NaOH lye or caustic soda Potassium hydroxide KOH lye or caustic potash Magnesium hydroxide Mg(OH)2 milk of magnesia Calcium hydroxide Ca(OH) 2 slaked lime Ammonia water NH3 H2O household ammonia Have you heard the one about a chemist who was reading a book about helium and just couldn't put it down?

  11. Arrhenius Acids and Bases • Arrhenius said that acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution • Bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solution • Monoprotic acids = acids that contain 1 ionizable hydrogen like nitric acid (HNO3) • Diprotic acids = acids that contain 2 ionizablehydrogens like sulfuric acid (H2SO4) • Triprotic acids = acids that contain 3 ionizablehydrogens like phosphoric acid (H3PO4) • Hydroxides of group I metals are very soluble in water and caustic to skin, hydroxide of group II metals are not very soluble in water and very dilute even when saturated (can be taken internally) Q: if both a bear in Yosemite and one in Alaska fall into the water which one disolves faster? A: The one in Alaska because it is Polar.

  12. Bronsted-Lowry Acids and Bases • But what about bases like sodium carbonate (Na2CO3) and ammonia (NH3)??? • The Bronsted-Lowry theory defines an acid as a hydrogen-ion donor, and a base as a hydrogen-ion acceptor • More complete definition

  13. Conjugate Acids and Bases • Conjugate acid = the particle formed when a base gains a hydrogen ion • Conjugate base = the particle that remains when an acid has donated a hydrogen ion • Conjugate acids and bases are always paired w/ a base or an acid, respectively • Conjugate acid-base pair = consists of 2 substances related by the loss or gain of a single hydrogen ion.

  14. Cont… • A water molecule that gains a hydrogen ion becomes a positively charged Hydronium ion (H3O+) • Amphoteric = a substance that can act as both an acid and a base – EX: Water

  15. Lewis Acids and Bases • Lewis proposed that an acid accepts a pair of electrons during a reaction while a base donates a pair of electrons • More general than either of the other 2 theories • Lewis acid = a substance that can accept a pair of electrons to form a covalent bond • Lewis base = a substance that can donate a pair of electrons to form a covalent bond

  16. Acid – Base Systems

  17. Many Lewis acids are also Bronsted-Lowry acids and vice versa but not all • *PP 1-2, 19.1 sect. assessment #8 pg. 593 Copper leaves Copper Sulfate and says see you: he answers CuSO4!!!!!

  18. Ion Product Constant for Water • Self-ionization of water = the reaction in which water molecules produce ions • For aqueous solutions, the product of the hydrogen-ion concentration and the hydroxide-ion concentration equals 1.0 x 10-14 • [H+] x [OH-] = 1.0 x 10-14 • ion-product constant for water (KW) = the product of the concentrations of the hydrogen ions and hydroxide ions in water • Acidic solution = one in which [H+] is greater than [OH-] • *The [H+] is greater than 1 x 10-7M* • Basic solution = one in which [H+] is less than [OH-] • *The [H+] is less than 1 x 10-7M* • *SP 19.1, PP 9-10 pg. 596

  19. [H+] pH 10-14 14 10-13 13 10-12 12 10-11 11 10-10 10 10-9 9 10-8 8 10-7 7 10-6 6 10-5 5 10-4 4 10-3 3 10-2 2 10-1 1 100 0 1 M NaOH Ammonia (household cleaner) 7 Acid Base 0 14 Blood Pure water Milk Acidic Neutral Basic Vinegar Lemon juice Stomach acid 1 M HCl pH Scale

  20. pH of Common Substances water (pure) 7.0 vinegar 2.8 soil 5.5 gastric juice 1.6 carbonated beverage 3.0 drinking water 7.2 bread 5.5 1.0 M NaOH (lye) 14.0 blood 7.4 potato 5.8 orange 3.5 1.0 M HCl 0 milk of magnesia 10.5 urine 6.0 apple juice 3.8 detergents 8.0 - 9.0 bile 8.0 lemon juice 2.2 milk 6.4 tomato 4.2 ammonia 11.0 seawater 8.5 bleach 12.0 coffee 5.0 13 11 12 14 1 6 9 10 0 2 3 4 7 5 8 basic neutral acidic [H+] = [OH-]

  21. pH of Common Substance pH [H1+] [OH1-] pOH 14 1 x 10-14 1 x 10-0 0 13 1 x 10-13 1 x 10-1 1 12 1 x 10-12 1 x 10-2 2 11 1 x 10-11 1 x 10-3 3 10 1 x 10-10 1 x 10-4 4 9 1 x 10-9 1 x 10-5 5 8 1 x 10-8 1 x 10-6 6 6 1 x 10-6 1 x 10-8 8 5 1 x 10-5 1 x 10-9 9 4 1 x 10-4 1 x 10-10 10 3 1 x 10-3 1 x 10-11 11 2 1 x 10-2 1 x 10-12 12 1 1 x 10-1 1 x 10-13 13 0 1 x 100 1 x 10-14 14 NaOH, 0.1 M Household bleach Household ammonia Lime water Milk of magnesia Borax Baking soda Egg white, seawater Human blood, tears Milk Saliva Rain Black coffee Banana Tomatoes Wine Cola, vinegar Lemon juice Gastric juice More basic 7 1 x 10-7 1 x 10-7 7 More acidic

  22. Søren Sorensen (1868 - 1939) pH Concept • pH = the negative logarithm of the hydrogen-ion concentration of a solution • A solution in which [H+] if greater than 1 x 10-7 M has a pH less than 7.0 and is acidic. The pH of pure water or a neutral aqueous solution is 7.0. A solution with a pH greater than 7 is basic and has a [H+] of less than 1 x 10-7M.

  23. Cont… • The pOH of a solution equals the negative logarithm of the hydroxide-ion concentration • A solution w/ a pOHless than 7 is basic, greater than 7 is acidic • For pH calculations, you should express the hydrogen-ion concentration in scientific notation • *SP 19.2, PP 11 pg. 599 • *Given pH = 4.6 determine the hydronium ion • *SP 19.3-19.4, PP 13-16 pg. 600-601

  24. pH Calculations pH [H3O+] pH = -log[H3O+] [H3O+] = 10-pH pH + pOH = 14 [H3O+] [OH-] = 1 x10-14 pOH [OH-] pOH = -log[OH-] [OH-] = 10-pOH

  25. Strength of Acids and Bases • Strong Acids = completely ionized in aqueous solution • HCl and H2SO4 • Weak Acids = ionize only slightly in aqueous solution • Ethanoic acid (acetic acid)

  26. Comparison of Strong and Weak Acids Reversibility of reaction Ions existing when acid, HA, dissociates in H2O Type of acid, HA Ka value Not reversible H+ and A-, only. No HA present. Ka value very large Strong Weak reversible Ka is small H+, A-, and HA HA(aq) + H2O(l) H3O+(aq) + A-(aq) The equilibrium expression for the reaction is [H3O+] [A-] Note: H3O+ = H+ Ka = [HA]

  27. Relative Strengths of Acids and Bases Acid Formula Conjugate base Formula perchlorate ion ClO4- chloride ion Cl- nitrate ion NO3- hydrogen sulfate ion HSO4- water H2O sulfate ion SO42- dihydrogen phosphate ion H2PO4- acetate ion C2H3O2- hydrogen carbonate ion HCO3- hydro sulfide ion HS- ammonia NH3 carbonate ion CO32- hydroxide ion OH- amide ion NH2- hydride ion H- perchloric HClO4 hydrogen chloride HCl nitric HNO3 sulfuric H2SO4 hydronium ion H3O+ hydrogen sulfate ion HSO4- phosphoric H3PO4 acetic HC2H3O2 carbonic H2CO3 hydrogen sulfide H2S ammonium ion NH4+ hydrogen carbonate ion HCO3- water H2O ammonia NH3 hydrogen H2 Decreasing Acid Strength Decreasing Base Strength acid conjugate base + H+

  28. Acid Dissociation Constant • For dilute solutions, the conc. of water is a constant. It can be combined w/ Keq to give the acid dissocation constant. • Acid dissociation constant = the ratio of the concentration of the dissociated (or ionized) form of an acid to the concentration of the undissociated (nonionized) form. • Weak acids have small Ka values. The stronger an acid is, the larger is its Ka value. • Nitrous acid (HNO2) has a Kaof 4.4 x 10-4, acetic acid has a Ka of 1.8 x 10-5 so nitrous acid is more ionized and has a higher [H3O+] or [H+] thus is a stronger acid • Di and triprotic acids lose each H separately so they have multiple dissociation constants

  29. [A-] [H3O+] 1 x 10-14 Ka = [HA] = [OH-] 1 x 10-14 = [H3O+] Equilibrium and pH Calculations Weak acid Strong acid HA H+ + A- HA + H2O H3O+ + A- HA + H2O H3O+ + A- acid-dissociation constant calculations [HA] = [H3O+] [H3O+] + antilog(-pH) [OH-] -log [H3O+] pH - 1 x 10-14 = [H3O+][OH-] 0 7 14 Kw = [H3O+][OH-]

  30. Base Dissociation Constant • Strong bases = dissociate completely into metal ions and hydroxide ions in aqueous solution • Ex: Ca(OH)2 • Weak bases = react w/ water to form the hydroxide ion and the conjugate acid of the base • Ex: ammonia NH3 • Base dissociation constant (Kb) = the ratio of the concentration of the conjugate acid times the conc. of the hydroxide ion to the conc. of the base

  31. Calculating Dissociation Constants • To find the Ka of a weak acid or the Kb of a weak base, substitute the measured concentration of all the substances present at equilibrium into the expression for Ka or Kb. • *SP 19.5, PP 22-23 pg. 610

  32. Weak Acids (pKa) Weak Acids – dissociate incompletely (~20%) Strong Acids – dissociate completely (~100%) A(g) + 2 B(g) 3 C(g) + D(g) [Products] Equilibrium constant (Keq) = [Reactants] 3 [C] [D] Keq = LeChatelier’s Principle (lu-SHAT-el-YAY’s) 2 [A][B]

  33. Values of Ka for Some Common Monoprotic Acids Formula Name Value of Ka* HSO4- hydrogen sulfate ion 1.2 x 10-2 HClO2chlorous acid 1.2 x 10-2 HC2H2ClO2monochloracetic acid 1.35 x 10-3 HF hydrofluoric acid 7.2 x 10-4 HNO2 nitrous acid 4.0 x 10-4 HC2H3O2 acetic acid 1.8 x 10-5 HOClhypochlorous acid 3.5 x 10-8 HCN hydrocyanic acid 6.2 x 10-10 NH4+ ammonium ion 5.6 x 10-10 HOC6H5 phenol 1.6 x 10-10 *The units of Ka are mol/L but are customarily omitted. Increasing acid strength

  34. H2SO4 H+ + HSO41- & HSO41- H+ + SO42- Sample 1) One gram of concentrated sulfuric acid (H2SO4) is diluted to a 1.0 dm3 volume with water. What is the molar concentration of the hydrogen ion in this solution? What is the pH? Solution) First determine the number of moles of H2SO4 1 mol H2SO4 x mol H2SO4 = 1 g H2SO4 = 0.010 mol H2SO4 98 g H2SO4 OVERALL: H2SO4 2 H+ + SO42- in dilute solutions...occurs ~100% 0.010 M 0.020 M pH = - log [H+] substitute into equation pH = - log [0.020 M] pH = 1.69

  35. A volume of 5.71 cm3 of pure acetic acid, HC2H3O2, is diluted with water at 25 oC to form a solution with a volume of 1.0 dm3. What is the molar concentration of the hydrogen ion, H+, in this solution? (The density of pure acetic acid is 1.05 g/cm3.) Step 1) Find the mass of the acid Mass of acid = density of acid x volume of acid = 1.05 g/cm3 x 5.71 cm3 = 6.00 g (From the formula of acetic acid, you can calculate that the molar mass of acetic acid is 60 g / mol). Step 2) Find the number of moles of acid. 1 mol HC2H3O2 x mol acetic acid = 6.00 g HC2H3O2 = 0.10 mol acetic acid (in 1 L) 60 g HC2H3O2 Molarity: M = mol / L Substitute into equation M = 0.10 mol / 1 L M = 0.1 molar HC2H3O2 Step 3) Find the [H+] Ka =

  36. Ka = 1.8 x 10-5 = HC2H3O2 H+ + C2H3O21- weak acid ? 0.1 M Step 3) Find the [H+] 0.1 M Ka = 1.8 x 10-5 @ 25 oC for acetic acid How do the concentrations of H+ and C2H3O21- compare? Substitute into equation: pH = - log[H+] x2 = 1.8 x 10-6M pH = - log [1.3 x10-3M] x = 1.3 x 10-3 molar = [H+] pH = 2.9

  37. Practice Problems: 1a) What is the molar hydrogen ion concentration in a 2.00 dm3 solution of hydrogen chloride in which 3.65 g of HCl is dissolved? 1b) pH 2a) What is the molar concentration of hydrogen ions in a solution containing 3.20 g of HNO3 in 250 cm3 of solution? 2b) pH 3a) An acetic acid solution is 0.25 M. What is its molar concentration of hydrogen ions? 3b) pH 4) A solution of acetic acid contains 12.0 g of HC2H3O2 in 500 cm3 of solution. What is the molar concentration of hydrogen ions? 1a) 0.0500 M 2a) 0.203 M 3a) 2.1 x 10-3M 4) 2.7 x 10-3M 1b) pH = 1.3 2b) pH = 0.7 3b) pH = 2.7

  38. [Products] Ka = [H3O+] [CN1-] [Reactants] Ka = [HCN] Weak Acids Cyanic acid is a weak monoprotic acid. If the pH of 0.150 M cyanic acid is 2.32. calculate Ka for cyanic acid. H3O+(aq) + CN1-(aq) H+(aq) HCN(aq) 4.8 x 10-3M 4.8 x 10-3M 0.150 M pH = -log[H3O+] 10-pH = [H3O+] [4.8 x 10-3M] 10-2.32 = [H3O+] [4.8 x 10-3M] [CN1-] Ka = [0.150 M] 4.8 x10-3 M = [H3O+] Ka = 1.54 x 10-4

  39. Titration Q: How did the chemist survive the famine? A: By subsisting on titrations. • Neutralization Reaction = a reaction b/w an acid and a base in aqueous solution to produce salt and water • Equivalence point = when the # of moles of hydrogen ions equals the number of moles of hydroxide ions • Titration = the process of adding a known amt. of solution of known conc. to determine the conc. of another solution • Standard solution = the solution of known conc. • End point = the point at which the indicator changes color • The point of neutralization is the end pt. of the titration • *SP 19.7, PP 32-33 pg. 616 Q: How do you get lean molecules?A: Feed them titrations. Q: What did one titration say to the other? A: Let's meet at the endpoint!

  40. Buffers • Buffer = a solution in which the pH remains relatively constant when small amts. of acid or base are added • A buffer is a solution of weak acids and one of its salts or a solution of a weak base and one of its salts • A buffer solution is better able to resist drastic changes in pH than is pure water • Buffer capacity = the amt. of acid or base that can be added to a buffer solution before a significant change in pH occurs

  41. Naming Acids Anion Acid _________ ide (chloride, Cl1-) Hydro____ ic acid (hydrochloric acid, HCl) add H+ ions _________ic acid (chloric acid, HClO3) (perchloric acid, HClO4) _________ ate (chlorate, ClO3-) (perchlorate, ClO4-) add H+ ions _________ite (chlorite, ClO2-) (hypochlorite, ClO-) ______ous acid (chlorous acid, HClO2) (hypochlorous acid, HClO) add H+ ions

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