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Chemistry SM-1232 Week 8 Lesson 1

Chemistry SM-1232 Week 8 Lesson 1. Dr. Jesse Reich Assistant Professor of Chemistry Massachusetts Maritime Academy Spring 2008. Class Today. Hand Back Work Grades Updated Quick Reflection Class on Friday Turn in the quiz today if you haven’t Today Chapter 14: Acids and Bases

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Chemistry SM-1232 Week 8 Lesson 1

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  1. Chemistry SM-1232Week 8 Lesson 1 Dr. Jesse Reich Assistant Professor of Chemistry Massachusetts Maritime Academy Spring 2008

  2. Class Today • Hand Back Work • Grades Updated • Quick Reflection • Class on Friday • Turn in the quiz today if you haven’t • Today Chapter 14: Acids and Bases • Chapter 14 Quiz Monday! • Chapter 13 and 14 Test Wednesday. • Wiki Project next month!

  3. Learning is a frustrating process • Easily learnt things are easily forgotten • Real learning takes time, effort, frustration • Real learning makes you master subjects

  4. Arrhenius acid • HCl  H+ + Cl- • H2SO4  2H+ + SO42- • H3PO4  3H+ + PO43-

  5. Hydronium ion • H+ when in water reacts with water to make an hydronium ion • H+ + H2O  H3O+

  6. Arrhenius Base • A compound that produces OH- ions when dissolved in water. • NaOH  Na+ + OH- • Mg(OH)2  Mg2+ + 2OH-

  7. Typical Bases • NaOH, sodium hydroxide • KOH, potassium hydroxide • NaOCH3, Soidum methoxide • Calcium Carbonate

  8. Bronsted-Lowry Definition • This definition rests on the transfer of H+ ions. • Bronsted acid is a proton H+ donor • Bronsted base is a proton H+ acceptor

  9. Acid Example • HCl + H2O  H3O + Cl- • H2SO4 +2 H2O  2H3O+ + SO42-

  10. Amphoteric • Water is amphoteric because it can act like an acid or base. • HCl + H2O  H3O+ +Cl- • NH3 + H2O NH4+ OH-

  11. Conjugate Acid-Base Pairs • NH3 + H2O  NH4+ + OH- • Base, acid  conjugate acid, conjugate base • On the left NH3 gained it’s a base • H2O gave H+ it’s a base • On the right, now NH4+ has an H+ to give so it’s the conjugate acid • OH- lost the H+ so now it’s the conjugate base

  12. Acid Base Reactions • Neutralization • Acid Reactions • Base Reactions

  13. Neutralization • Most common reaction! • For Arrhenius acid Base Reactions: • Acid + Base = Water + Salt • For Bronstead acid base reactions: • AcidH+ + Base-  conjugate base- + conjugate acid+

  14. Titrations • Use the reaction formula to determine the molar ratio of acid to base. • Use a known amount of acid (or base) and an unknown amount of base (or acid). • Drop in color changing indicator • Add base to acid with indicator until the solution changes color. At that point more base is present than acid.

  15. Titrations • You have 30mL of a 3M solution of NaOH. You perform three titrations using 3M solutions of HCl, H2SO4, and H3PO4. Write three balanced equations. How many mL of each acid solution will it take to make the indicator change color.

  16. Concentration • Another term for molarity is concentration. • You’ve used this before with oranje juice from concentrate. It’s comes in a concentrated form and then you have to dilute it down. • Concentration is written by putting a molecule in brackets like this [HCl], which would mean the concentration of HCl. • A solution that is 1.0M in HCl can be written like [HCl] = 1.0M

  17. Strong Acids • Strong acids fully dissociate in water. That means water tears every molecule of the acid into H+ ions and base – ions. • HCl  H+ + Cl- • Good electrolyte, meaning strong acid solutions conduct electricity since so many charges are in solution.

  18. Mono or multiprotic • Some acids only release 1 proton (H+ ion). Others release more than one. • HCl is monoprotic • Sulfuric acid H2SO4 is diprotic • Phosphoric acid H3PO4 is triprotic

  19. Weak Acids • Weak acids do not completely ionize • HF is a weak acid • In solution it become H+ and F- when it ionizes, but there are strong electrostatic attractions F- and H+ so they come back together. • HF + H2O  H3O+ + F- • As a consequence weak acids are poor electrolytes and electricity is not conducted through their solutions well.

  20. Tug of War • The solvent pulls the charges apart and dissolves them, but the oppositely charged particles are attracted to each other. • Generically speaking • HA + H2O  H3O + A- • If the acid is strong the products are favored and the reactants are barely present. If the acid is weak the reactants are favored and the products are barely present.

  21. Rule of thumb • HA + H2O  H3O + A- • HA is the acid • A- is the conjugate base • The stronger the acid the weaker the conjugate base.

  22. Strong Bases • A strong base completely dissociates in solution into the ions that make it up. • NaOH  Na+ + OH- • A 1M solution of NaOH will have [Na+]= 1M and [OH-]=1M. • Strong bases make good electrolytes. • For the purpose of this class any hydroxide is a strong base.

  23. Weak Bases • Weak bases do not fully dissociate in water. Most weak bases do not have a hydroxide ion as part of them. • In order to act in a basic manner the weak base reacts with water to steal a proton, and that interaction creates the OH- ion. • NH3 + H2O  NH4+ + OH-

  24. Weak Bases • Generically speaking • B + H2O  BH+ + OH- • B= base • BH+= conjugate acid • The stronger the base the weaker the conjugate acid. • The book has a list of common weak bases on page 505

  25. Time to hurt you! • Water is amophoteric. It can act like a base or an acid. • HCl + H2O  H3O + Cl- • NH3 + H2O  NH4+ + OH- • Water can act like it’s own acid and base • H2O + H2O  H3O+ + OH-

  26. Water’s Ion Product Constant • H2O + H2O  H3O+ + OH- • Scientists have measured how much this happens at 25C in pure water. • The concentration (aka molarity) of each ion is 1.0e-7M. Aka [H3O+]=[OH-]=1.0e-7 • If you multiply the concentration of [H3O]X[OH-] you get 1.0e-14. • This is considered the ion product constant for water (Kw).

  27. Kw • Kw= [H3O+]X[OH-] • Kw= 1.0e-7 X 1.0e-7 • Kw= 1.0e-14

  28. Quick Math Break • Answer the following • 1.0e6 X 1.0e7= • 1.0e 6 / 1.0e12= • 1.0e6 X 4.5= • 1.0e6 / 3.2= • 4.2e-13/ 3.3 e-4 =

  29. How does this help us? • It becomes a 3 variable problem! • Kw=[H3O+] X [OH-] • We know Kw and if we know H3O or OH- we can solve for the other one! • Work this one out. If [H3O+]=1e-3 then what must the [OH-] be equal to?

  30. Solution • Kw=[H3O+] X [OH-] • Kw/[H3O+] = [OH-] • 1e-14/1e-3 = [OH-] • 1e-11 = [OH-]

  31. Acidic or basic? • In the previous example you found that the acid concentration was 1e-3 and the base concentration was 1e-11. • Is there more acid or base in solution? So, is the solution acidic or basic? • It’s acidic! If you don’t see why write out the number 1e-3 and write out the number 1e-11.

  32. Handy Reference • In acidic solutions [H3O+] > [OH-] • In neutral solutions [H3O+] = [OH-] • In basic solutions [OH-] > [H3O+]

  33. Work • Read 509 through the end of the chapter. • Copy over the example problems • Chapter 14 HW due Monday. • Quiz on Chapter 14 Monday. • Test on Chapter 13 and 14 Wednesday April 29th. We will have class May 1st to start chapter 15.

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