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Unit 13: Acids & Bases Chemistry Chapter 19

Unit 13: Acids & Bases Chemistry Chapter 19. Welcome To The GowerHour. Acid. Base. I. Bronsted-Lowry Acids and Bases:. donates. H + , proton. accepts. H + , proton. A. Acid: Substance that _________ a ___________ to another substance.

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Unit 13: Acids & Bases Chemistry Chapter 19

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  1. Unit 13: Acids & BasesChemistryChapter 19 Welcome To The GowerHour

  2. Acid Base I. Bronsted-Lowry Acids and Bases: donates H+, proton accepts H+, proton A. Acid: Substance that _________ a ___________ to another substance. B. Base: Substance that ________ a __________ from another substance. C. Acid – Base reaction: ___ is transferred from the ____ to the _____.   Ex. D. Conjugate acid – base pair: The pair of acid and base that differ by a ___. Ex. ACID BASE H+ acid base HX + Y– ↔ HY + X– Acid Base H+ HF F– Conjugate Acid add H+ HNO2 NO2– Conjugate Base subtract H+ HC2H3O2 C2H3O2– Memorize: NH3 = Base Memorize: NH4+= Acid

  3. Generally: H and OH– H3O+ ★Memorize ★ hydroxide ★ hydronium + H+ – H+ – H+ + H+ H2SO4 SO42– ↔ ↔ CO32– H2CO3 ↔ ↔ (CA) (CB) (CB) (CA) donate accept – H+ + H+ H2O ↔ ↔ HSO4– HCO3– E. Amphoteric substance: Substance that can either _______ or _______ a hydrogen ion.   Ex. H2O F. Example: Label the acids and bases, draw lines to connect the conjugate pairs. (1) NH3 + H2O  NH4+ + OH-(2) HF + H2O  F- + H3O+ G. Categorize each of the following as an Acid, Base, or Amphoteric and below each, write the conjugate acid or base.   H2SO4 H2PO4- NH3 NO2- HSO4- Base Acid Acid Base Acid Ampho Base Base Ampho HSO4– H3PO4 NH4+ HNO2 H2SO4 HPO42– SO42–

  4. II. Water Dissociation Constant (Ion product) aqueous A. The acidic / basic properties of _________ solutions are dependent upon the ____________ that involves the solvent, _______. 1. Reaction: 2. Equilibrium expression:   3. Kw = ___________ @ 25 C. ___________ are favored. equilibrium water H2O (l) ↔ H+ (aq) + OH– (aq) Acid Base Keq = [H+][OH–] = Kw 1.0 x 10−14 Reactants H20  very little conc. of ions

  5. H2O (l) ↔ H+ (aq) + OH– (aq) x x ICE Kw = 1.0 x 10−14 = [H+][OH–] x2 = 1.0 x 10−14M x= 1.0 x 10−7M x= 1.0 x 10−7 M x= 1.0 x 10−7 M B. In pure water:   [H+] = [OH-] = C. In impure water (contains an acidic or basic substance):   If [H+] > 1.0 x 10-7 M, solution is _______. If [H+] < 1.0 x 10-7 M, solution is _________________. If [H+] = 1.0 x 10-7 M, solution is _________. Kw = 1.0 x 10−14 = (1.0 x 10−7 M)(1.0 x 10−7 M) Kw = 1.0 x 10−14 = [H+][OH–]  As [H+]  [OH–]  acidic basic or alkaline neutral 1.0 x 10−7 = pH 7  pH < 7 = acid pH > 7 = base pH = 7 = neutral

  6. p H III. pH and pOH log pH = – log [H+] pOH = – log [OH–] A. Because [H+] and [OH-] are generally very small numbers, a _____ based system of measuring acidity is used. Practice: What is the pH of a solution with: [H+] = 10-5 M? [H+] = 10-11 M?  1. pH and pOH are _________, because we cannot take the logarithm of a unit.  2. For each pH change of 1, the [H+] changes by a factor of ____. Ex. 3. [H+][OH-] = 1.0 x 10-14 Or pH + pOH = 14 5 11 unitless 10 Richter scale (102 = 100) pH of 1 is 100 x’s more concentrated than pH of 3. pH = 1  pH = 5 104 = 10,000

  7. B. Example calculations 1. Calculate the pH of the following solutions, and indicate if the solution is acidic or basic: (a) [H+] = 1.0 x 10-11 M (b) [H+] = 2.11 x 10-2 M (c) [OH-] = 3.98 x 10-7 M pH = 11 basic or alkaline pH = – log [H+] pH = – log (2.11 x 10-2) pH = 1.68 acidic pOH = – log (3.98 x 10-7) pOH = 6.40 pH = 7.60 basic

  8. 10x = antilog (opposite of log) pH = – log [H+] 10–9.35 = [H+] 10x Key 2. Calculate the [H+] and [OH-] of the following and indicate if they are acidic or basic: • pH = 9.35 • pH = 1.10 • pOH = 2.98 basic 9.35 = – log [H+] [H+] = 4.47 x 10– 10 M –9.35 = log [H+] [H+][OH-] = 1.0 x 10-14 Divide both side by log [OH–] = 2.24 x 10– 5 M [H+] = 7.94 x 10– 2 M acidic [OH–] = 1.26 x 10– 13 M [OH–] = 1.05 x 10– 3 M [H+] = 9.55 x 10– 12 M

  9. NaOH Lye H2SO4 High quality H2O Acid Rain C. pH Scale: Neutral Alkaline/Basic Acidic 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

  10. Properties of Acids and Bases A. Properties of Acids 1. _________ taste (ex. _________) 2. Reacts with some metals to form ___ (ex.___________) 3. __________electricity (__________ in solution) 4. React with _________ B. Properties of Bases 1. _______taste (ex.________) 2. _____________ (ex._____________) 3. ______________electricity (___________ in solution) Sour Vinegar H2 Magnesium Conducts electrolytes bases Bitter Soap Slippery Soap, NaOH Conducts electrolytes

  11. H2O ionize reversible + Cl(aq) H+(aq) HCl V. Strong vs. Weak A. Strong Acids: Completely ______ in solution (Not _________reaction). Ex. HCl (1) (2) *****MEMORIZE***** HCl H2SO4 HNO3 HBr HI HClO4 (H3O+ = Hydronium ion) + Cl H3O+ HCl + H2O Chemists use H+ & H3O+ interchangeably. H+ is often used for simplicity, but H3O+ more closely represents reality. hydrochloric sulfuric nitric hydrobromic hydroiodic perchloric

  12. H2O H2O NaOH +H2O Na+ + OH + H2O ionize + OH Na+ B. Strong Bases: Completely _______ in solution. (Group I and II hydroxides) Ex. NaOH Practice: Ca(OH)2 *****Memorize***** LiOH NaOH KOH Ca(OH)2 Ba(OH)2 Sr(OH)2 Ca2+ + OH 2

  13. H2O  partially reversible 1 % H+ + HCO3 H2CO3 C. Weak Acids: Ionize __________ (_________reaction). Generally less than _____ of the molecules ionize. Ex. H2CO3 (1) (2) Example of weak acids: HC2H3O2, H2CO3 H3O+ + HCO3 H2CO3 + H2O  If it is not a strong acid then it is a weak acid!!

  14. partially reversible NH4+ + OH NH3 + H2O  D. Weak Bases: Ionize __________ (_________reaction). Partial dissociation of an weak base in water: Ex. NH3 (1) Example of weak bases: NH3, NH2CH2CH3 (aminoethane) If it is not a strong base then it is a weak base!!

  15. H2O ↔ H+ weaker HA (aq) H+(aq) + A(aq) Ka = acid-dissociation constant VI. Calculations with weak acids and weak bases. • Weak Acids (Ka): As Ka decreases, less ____ is formed in solution, so acid is _______. B. Weak Bases (Kb): As Kb decreases, less _____ is formed in solution, so base is _______. OH weaker + H2O(l) B (aq) BH+(aq) + OH(aq) Kb = base-dissociation constant ↔

  16. x C. Examples 1. Example: Aspirin is a weak organic acid whose molecular formula is HC9H7O4. An aqueous solution of aspirin is prepared by dissolving 3.60 g/ L. The pH of this solution is found to be 2.60. Calculate Ka for aspirin. HC9H7O4↔ H+ + C9H7O4- 2.00 x 10-2 0 0 – x + x + x 0.0175 2.51 x 10-3 2.51 x 10-3 3.60 g 1 mol = 2.00 x 10-2 M = I pH = 2.60 L 180.15 g [H+] = 2.51 x 10-3

  17. HCN H+ + CN ↔  HCl H+ + Cl I C E Assume x is small [H+] = 0.100 M 0.100 0 0 • Compare the pH of 0.100 M HCl and 0.100 M HCN (Ka = 4.90 x 10-10) pH = 1.00 - x + x + x 0.100 - x x x x = 7.0 x 10–6 M pH = 5.15

  18. Kw = [H+] [OH] K1 • K2 = Knet H+ + C2H3O2 1.8 x 10 – 5 ↔ VII. Multiple Equilibria: When two reactions are added to give a third (net reaction) the equilibrium constants are multiplied (______________).   Ex. HC2H3O2 Ka = C2H3O2- Kb =   Kw = _______________ • ___________ Equation: 5.6 x 10 – 10 HC2H3O2 + OH + H2O ↔ H2O ↔H+ + OH– Kw = 1.0 x 10−14 [H+] [C2H3O2] [HC2H3O2] [OH] [HC2H3O2] [C2H3O2] Ka • Kb = Kw

  19. decrease inverse A. As Ka increases, Kb ________ (_______ relationship). B. The stronger the acid, the ________ its conjugate base. Ex. If Ka = 1.0 x 1030, determine Kb. weaker Kb = 1.0 x 10 – 44 Ka = strong acid; Kb = weak base

  20. H2O H2O H2O H2O H2O ↔ ↔ ↔ ↔ ↔ , H2PO4 HCO3 H+ 2 Ka1 = 4.2 x 10 – 7 H+ + HCO3 H2CO3 Ka2 = 4.8 x 10 – 11 HCO3 H+ + CO32- VIII. Polyprotic Acids: Certain weak acids contain more than one ionizable ___. These acids dissociate in multiple steps.  A. Diprotic acid: Dissociates to form ___ hydrogen ions.   Ex. B. Triprotic acid: Dissociates to form ___ hydrogen ions.   Ex.  C. The _________ formed in one step (e.g. _______________) dissociates in the next step.  D. The dissociation constant (___________________) becomes smaller with each successive step: Ka1 >  E. The acids formed in successive steps become progressively ________. 3 + H2PO4 Ka1 = 7.5 x 10 – 3 H3PO4 H+ H2PO4 H+ + HPO42- Ka2 = 6.2 x 10 – 8 HPO42- H+ + PO43- Ka3 = 4.8 x 10 – 13 substance equilibrium constant Ka2 > Ka3 weaker

  21. H2O ↔ H2O H2O H2O H2O ↔ ↔ ↔ ↔ H+ + HSO3 H2SO3 HSO3 H+ + SO32- F. Example: Write the dissociation reactions of sulfurous acid, H2SO3.  G. Example: Write the dissociation reactions of citric acid, H3C5H5O7. H+ + H2C5H5O7  H3C5H5O7 H+ + HC5H5O7 2- H2C5H5O7  + C5H5O7 3- H+ HC5H5O7 2-

  22. H2O H2O ionic H+ OH ion IX. Acid / Base Properties of salt solutions: A. A salt is an ____ compound not containing ___ or____. B. When a salt dissolves in water, the ___ are formed. Ex. NaCl(s) Ex. K2CO3(s) C. Cations: Weak acids or “spectator” ions? 1. Cations derived from ___________ are _________ ions. (Do not react with water, therefore have _______ on pH.) 2. Other cations are slightly _____. (Lewis acid = ______acceptor). D. Anions: Weak bases or “spectator” ions? 1. Anions derived from __________are ________ ions. 2. Other anions are slightly ____. + Cl Na+ 2 K+ + CO32- strong bases spectator no effect e - pair acidic strong acids spectator basic no effect

  23. H2O H2O H2O H2O Ex. HCl  H+ + Cl– (All others weak acids) Li+ ; Na+ ; K+ ; Ca2+ ; Sr2+ ; Ba2+ ; SO42- ; ClO4 ; Cl ; Br ; I (All others weak bases) NO3 E. Spectator ions:  1. Cations:  2. Anions: F. Example: Consider water solutions of these four salts: (a) NH4I, (b) Zn(NO3)2, (c) KClO4, (d) Na3PO4 Classify each salt solution as acidic, basic, or neutral. (Show the dissociation reactions of each.) a) NH4I NH4+ + I¯ Acidic (w. acid) (Spec) b) Zn(NO3)2 Zn2+ + NO3¯ Acidic 2 (w. acid) (Spec) c) KClO4 + ClO4¯ Neutral K+ (Spec) (Spec) d) Na3PO4 3 Na+ + PO43- Basic (Spec) (w. base) Which will undergo hydrolysis? The salts that react w/water to change the pH!

  24. X. Acid – Base Reactions H2O salt A. Strong acid + Strong base --> ____ + ____ Example: HCl(aq) + NaOH(aq) --> Example: HNO3(aq) + Ca(OH)2(aq) --> Net reaction: NaCl (aq) + H2O (l) 2 Ca(NO3)2 (aq) + H2O (l) 2 2 H+ + 2 NO3¯ + Ca2+ + 2 OH– Ca2+ + 2 NO3¯ + 2 H2O (l) H+ + OH– H2O (neutralization)

  25. water weak base  weak base (Soln. will still be basic) H2O + C2H3O2 H2O + F • Weak acid + Strong base ______+__________ Note: Strong Base Example: HC2H3O2 + OH-   Example: HF + OH- • Strong acid + Weak base __________Note: Strong acid Example: H+ + NH3   Example: H+ + ClO-  D. Weak acid + Weak base _________+__________ Example: HC2H3O2 + NH3   Example: HF + ClO- weak acid  weak acid (Soln. will still be acidic) NH4+ HClO weak acid weak base ↔ NH4+ + C2H3O2 ↔ HClO + F ↔

  26. indicator equal neutralization XI. Titration: A process in which one reagent is added to another with which it reacts; an ________ is used to determine the point at which _____ quantities of the two reagents have been added. A. Equivalence point: the point at which the _____________ reaction is complete. B. End point: The point at which the ________ changes color. 1. Indicators change at different ______. 2. In doing a titration, one must choose an _________ where the equivalence point and the _________ coincide. 3. Example Indicators: IndicatorColor changepH at end point Methyl Orange ___________ ___________ Bromothymol Blue ___________ ___________ Phenolphthalein ___________ ___________ nA = nB indicator pH’s indicator end point orange - red 5 yellow - blue 7 clear - pink 9

  27. color C. Acid – Base Indicators:Produce a _____ change in an acid-base reaction. 1. Example: Phenolphthalein, bromothymol blue 2. Natural indicators: purple cabbage, hydrangeas • Reversible reactions: (Phenolphthalein) H-In  H+ + In- CLEAR PINK Add acid, rxn shifts ____, solution turns _____. Add base, rxn shifts _____, solution turns _____. left clear right pink

  28. ~ pH 7 D. Titration Curves: 1. Strong acid – Strong Base (Equivalence point = _______) 0.100 M HCl is titrated with 0.100 M NaOH

  29. ~pH 8.8 2. Weak Acid – Strong Base (Equivalence point = _______)   0.100 M CH3COOH is titrated with 0.100 M NaOH

  30. ~ pH 5.3 3. Strong Acid – Weak Base (Equivalence point = _______)   0.100 M NH3 is titrated with 0.100 M HCl

  31. coefficients nA = nB (mol acid = mol base) E. Calculations with titrations: 1. STOICHIOMETRY   2. At equivalence point:

  32. x x x x x Problems: 1. Example titration problem: 35.00 mL of 0.2500 M sodium hydroxide is titrated with 0.4375 M HCl. (a) Write the balanced chemical equation. (b) Determine the volume of HCl added at the equivalence point (i) using stoichiometry and (ii) using the titration equation. (a) HCl + NaOH NaCl + H2O (b) (i) 1 ml 0.2500 mol NaOH 1 mol HCl 35.00 ml NaOH 10-3 L 1 L HCl 0.4375 mol HCl 1 ml 1 L NaOH 1 mol NaOH 10-3 L = 20.00 ml HCl (b) (ii) VA = 20.00 mL HCl

  33. 2. Example: A 15.0 mL sample of a solution of H2SO4 with an unknown molarity is titrated with 32.4 mL of 0.145 M NaOH to the bromothymol endpoint. What is the molarity of the sulfuric acid solution? H2SO4 + NaOH 2 Na2SO4 + H2O 2 MA = 0.157 M H2SO4

  34. 3. A Ca(OH)2 solution was used to titrate 15.0 mL of a 0.125 M H3PO4 solution. If 12.4 mL of Ca(OH)2 are used to reach the endpoint, what is the concentration of the Ca(OH)2? 2 H3PO4 + Ca(OH)2 3 Ca3(PO4)2 + H2O 6 MB = 0.227 M Ca(OH)2

  35. H2O H2O IQ 2 a) Ca(ClO3)2 Ca2+ + ClO3¯ 2 Basic • Show the dissociation rxn of each of the following salts and determine if each solution is acidic, basic, or neutral. • Predict the products and balance the following: a. HNO3 + Sr(OH)2 b. HClO2 + OH– (Spec) (w. base) a) FeCl2 Fe2+ + Cl¯ 2 Acidic (w. acid) (spec) 2 Sr(NO3)2 + H2O 2 H2O + ClO2¯

  36. IQ 1 Li2CO3 CaCO3 • Write formulas for two salts that: (a) contain CO32- and are basic: (b) contain Li+ and are neutral: 2. Predict the products and balance the following. Write the net rxn for each. a. HNO3 + Ca(OH)2 b. HClO2 + LiOH Net Rxn: LiBr Li2SO4 Ca(NO3)2 + H2O 2 2 NO3- Ca2+ Ca2+ NO3- H+ OH- H2O Net Rxn: H+ + OH- H2O + LiClO2 ClO2- Li+ Li+ ClO2- H+ OH- HClO2 + OH- H2O + ClO2-

  37. Titration of a Strong Acid w/ a Strong Base • Clean burette w/ NaOH; Drain into “waste beaker”. • Fill burette with NaOH to 0.00 mL (hundredths)(meniscus). Record initial vol. • Measure out 10.00 mL (hundredths)(meniscus) of acid using the graduated cylinder (Use the pipet). Record volume of acid; add to the Erlenmeyer flask. • Add 2 drops of phenolphthalein to the acid. IMPORTANT!!!! • Add a little water to Erlenmeyer flask (Rinse acid off sides). • Start adding NaOH slowly; swirl flask as you add. • At the endpoint (when soln. remains light pink) stop adding base and record volume of base added. (Check: rinse flask) • Dispose titrated solution in sink; rinse flask and repeat steps 3-8. • The final vol. of the base is your initial vol. for the next trial.

  38. Reminders • Wear safety glasses/goggles. • Only use the pipette for acid! • Base in burette • Clean up after 3rd trial. • Must get a clean up stamp before leaving. • Leave excess acid and base in appropriate beakers.

  39. Titration of a Strong Acid w/ a Strong Base HCl + NaOH NaCl + H2O MA = ? MB = 0.0915 M VA = 10.00 mL VB = ? mL nB = 1 nA = 1

  40. Titration of Vinegar Lab • Fill burette with NaOH to 0.00 mL (hundredths)(meniscus). Record initial vol. 2. Measure out 1.00 mL (hundredths)(meniscus) of vinegar using the graduated cylinder (Use the pipet). Record volume of vinegar; add to the Erlenmeyer flask. 3. Add 2 drops of phenolphthalein to the vinegar. IMPORTANT!!!! 4. Add a little water to Erlenmeyer flask (Rinse acid off sides). 5. Start adding NaOH slowly; swirl flask as you add. 6. At the endpoint (when soln. remains light pink) stop adding base and record volume of base added. 7. Dispose titrated solution in sink; rinse flask and repeat steps 2-7. 8. The final vol. of the base is your initial vol. for the next trial.

  41. Reminders • Wear safety glasses/goggles. • Only use the pipette for acid! • Base in burette • Clean up after 3rd trial. • Must get a clean up stamp before leaving. • Leave excess acid and base in appropriate beakers.

  42. Average vol. of vinegar (i.e. 1 mL) Density of vinegar = 1.02 g/mL Vinegar Calculations HC2H3O2 + NaOH H2O + C2H3O2– • Vinegar = Acetic acid (solute) in water (solvent). • Average number of moles of acetic acid: • Mass (g) of acetic acid (HC2H3O2): 4. % mass of acetic acid: Average [ ] of acid

  43. Average vol. of vinegar (i.e. 1 mL) Density of vinegar = 1.02 g/mL Vinegar Calculations HC2H3O2 + NaOH H2O + NaC2H3O2 • Vinegar = Acetic acid (solute) in water (solvent). • Average number of moles of acetic acid: • Mass (g) of acetic acid (HC2H3O2): 4. % mass of acetic acid: Average [ ] of acid

  44. Neutralization Capacity of an Antacid • Fill burette with 0.0953 M NaOH to 0.00 mL (hundredths)(meniscus). Record initial vol. • Weigh ~ 0.10 g of Tums (CaCO3). Record and add to Erlenmeyer flask. • Add water to E. flask and swirl to dissolve powder. • Measure out 5.00 mL (hundredths)(meniscus) of 0.336 M HCl using the graduated cylinder (Use the pipet). Record volume of acid; add to the E. flask containing antacid and swirl. • Add a little water to E. flask (Rinse acid off sides). • Add 2 drops of phenolphthalein to the acid. IMPORTANT!!!! • Start titrating; swirl E. flask as you add NaOH. • At the endpoint (when soln. remains light pink) stop adding base and record volume of base added. • Dispose titrated solution in sink; rinse flask and repeat steps 2-7. • The final vol. of the base is your initial vol. for the next trial.

  45. Reminders • Wear safety glasses/goggles. • Only use the pipette for acid! • Base in burette • Clean up after 3rd trial. • Must get a clean up stamp before leaving. • Leave excess acid and base in appropriate beakers.

  46. Antacid Calculations Acid (HCl) (VA) Antacid (CaCO3)(VA1) NaOH (VA2) VA1 = Volume of acid neutralized by the antacid. VA2 = Volume of acid neutralized by the NaOH. VA = Total volume of acid (5 mL). VA = VA1 +VA2 Step 1: Solve for VA2 Step 2: Solve for VA1 VA = VA1 +VA2

  47. Antacid Calculations Step 3: Solve for VA1/ g. Mass of antacid used in the trial (i.e. 0.10) Step 4: Solve for VA1/ Tablet Antacid : 1 Tablet = 1.30 g Step 5: Solve for VA1/ cent Antacid : 72 Tablet = $5.49 Labs Due: Tuesday 05/15

  48. Acid Nomenclature Review • Acetic acid (w) 11. Carbonic acid (w) • Oxalic acid (w) 12. Perchloric acid (s) • Hydrocyanic acid (w) 13. Hypobromous acid (w) • Cyanic acid (w) 14. Nitric acid (s) • Sulfurous acid (w) 15. Chloric acid (w) • Sulfuric acid (s) 16. Hydrofluoric acid (w) • Hypochlorous acid (w) 17. Phosphorous acid (w) • Bromous acid (w) 18. Hydroiodic acid (s) • Periodic acid (w) 19. Nitrous acid (w) • Phosphoric acid (w) 20. Hydrochloric acid (s)

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