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Announcements

Announcements. HAVE A GREAT BREAK! REMINDER: NO LAB. Types of Chemical Reactions. Chapter 4 Goals: To be able to predict chemical reactivity. To know how to synthesize specific compounds. Types of Reactions. Acid-Base Oxidation-Reduction Precipitation Gas Forming Organic: Substitution

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Announcements

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  1. Announcements HAVE A GREAT BREAK! REMINDER: NO LAB

  2. Types of Chemical Reactions Chapter 4 Goals: • To be able to predict chemical reactivity. • To know how to synthesize specific compounds.

  3. Types of Reactions • Acid-Base • Oxidation-Reduction • Precipitation • Gas Forming • Organic: • Substitution • Addition • Elimination

  4. Reactions in Aqueous Solution • Aqueous- solvent is water • Reactions we’ll discuss today/next week are in aqueous solution, unless otherwise noted • Acid-Base • Redox • Precipitation

  5. Electrolytes • Strong: All of the solute comes apart to yield ions in solution • Dissolution of KMnO4 • Weak: Some of the solute comes apart to yield ions • Nonelectrolytes: No ions formed • Let’s compare

  6. Electrolytes in the Human Body • Most important: Na+, Cl-, K+, Ca2+, Mg2+ ,HCO3-, and PO43-,SO42- • Elevated K+ cardiac arrythmia • Decreased extracellular K+ paralysis • Excess extracellular Na+ fluid retention • Decreased plasma Ca2+ and Mg2+  muscle spasms

  7. Acids and Bases • Theories- there’s lots of them • Ones we’ll use in this course • Lewis (later) • Brønsted-Lowry (now) • An acid is a substance that donates a proton (H+) to a base • The hydronium ion

  8. Brønsted-Lowry Definitions • Acid= donates a proton (H+) to a base • Base= accepts a proton (H+) from an acid • Acid base reactions are reversible (almost always)

  9. Brønsted-Lowry Definitions • Acid= donates a proton (H+) to a base • Base= accepts a proton (H+) from an acid • Acid base reactions are reversible (almost always)

  10. Important Acids and Bases

  11. Strong Acids • 100% of acid molecules produce ions in water • Dissociation vs. ionization

  12. Weak Acids • Only a few acid molecules produce ions (≤5%) • Strong vs. Weak acid ionization

  13. Polyprotic Acids • Polyprotic acids can donate more than one H+ • Sulfuric acid • Citric acid (C6H8O7) : Not all H’s are acidic H2SO4 H+ + HSO4- HSO4-  H+ + SO42-

  14. Bases Strong bases are hydroxide salts For now, the only important weak base is NH3.

  15. If H3PO4 reacts as an acid, which of the following can it not make? • H4PO4+ • H2PO4- • HPO42- • PO43- 20

  16. If C2O42- (oxalate ion) reacts in an acid-base reaction, which of the following can it not make? • H2C2O4 • HC2O4- • 2 CO2 20

  17. Acid Base Reactions

  18. Acid Base Reactions • Strong Acid + Strong Base HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) acidbase“salt”water • What do we get if we mix: HBr + LiOH 

  19. Acid Base Reactions • Diprotic acids or bases H2SO4(aq) + NaOH(aq)  H2SO4(aq) + Ba(OH)2(aq)  HCl(aq) + Ba(OH)2(aq) 

  20. Acid-Base Reactions Diprotic Acids or Bases H2SO4(aq) + NaOH(aq)  H2SO4(aq) + Ba(OH)2(aq)  HCl(aq) + Ba(OH)2(aq) 

  21. Acid-Base Reactions Strong Acid + Weak Base HCl(aq) + NH3(aq)  NH4Cl(aq)

  22. Acid-Base Reactions Weak Acid + Strong Base HCN(aq) + NaOH(aq)  NaCN(aq) + H2O(l) acidbase “salt” water

  23. Net Ionic Equations HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) What really happens: H+(aq) + OH-(aq)  H2O(l) Sodium ion and chloride ion are “spectator ions”

  24. Reactions involving weak bases HCl(aq) + NH3(aq)  NH4+(aq) + Cl-(aq) Net-Ionic Equation: NH3(aq) + H+(aq)  NH4+(aq)

  25. CH3CO2H(aq) + NaOH(aq)  • 1. CH3CO2H2+(aq) + NaO(aq) • 2. CH3CO2-(aq) + H2O(l) + Na+(aq) • 3. CH4(g) + CO2(g) + H2O(l)

  26. HCN(aq) + NH3(aq)  • 1. NH4+(aq) + CN-(aq) • 2. H2CN+(aq) + NH2-(aq) • 3. C2N2(s) + 3 H2(g)

  27. Solution Concentration: Molarity • Molarity = moles solute per liter of solution • 0.30 mol NH3 dissolved in 0.500 LConcentration = • Written like: [NH3] = 0.60 M

  28. pH Scale • In pure water, a few molecules ionize to form H3O+ and OH–H2O + H2O  OH– + H3O+ • In acidic and basic solutions, these concentrations are not equalacidic: [H3O+] > [OH–]basic: [OH–] > [H3O+]neutral: [H3O+] = [OH–]

  29. pH Scale • Measure how much H3O+ is in a solution using pH • pH < 7.0 = acidic • pH > 7.0 = basic • pH = 7.0 = neutral • Measure of H3O+ and OH–concentration (moles per liter) in a solution • As acidity increases, pH decreases

  30. pH Scale • The pH scale is logarithmic:100 102 log(102) = 210 101 log(101) = 11 100 log(100) = 00.1 10–1 log(10–1) = –10.01 10–2 log(10–2) = –2 • pH = –log [H3O+] • pH if [H3O+] = 10–5? 10–9? Acidic or basic? • pH if [H3O+] = 0.000057 M?

  31. Finding [H3O+] from pH [H3O+] = 10-pH What is [H3O+] if pH = 8.9?

  32. pH: Quantitative Measure of Acidity • Acidity is related to concentration of H+ (or H3O+) • pH = -log[H3O+]

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