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5 Oct. 2010

5 Oct. 2010

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5 Oct. 2010

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  1. 5 Oct. 2010 • Objective: SWBAT describe properties of an aqueous solution, and write an equation for a precipitation reaction. • Do now: What is an electrolyte? Describe why it exhibits the properties of an electrolyte.

  2. Agenda • Do now • Aqueous solutions notes • Precipitation Reaction practice problems Homework: p. 160 #1-6, 8-14 evens (JR) p. 161 #15, 18-24 evens (TTL)

  3. Reactions in Aqueous Solutions

  4. Introduction • Most chemical reactions and virtually all biological processes take place in water! • Three categories of reactions in aqueous solutions: • Precipitation reactions • Acid-Base reactions • Redox reactions

  5. solvent: the part of a solution doing the dissolving, present in larger amount • solute: a substance being dissolved, present in a smaller amount • solution: homogeneous mixture of two or more substances SOLUTE SOLUTION SOLUTION SOLVENT

  6. Examples • KCl in water? • Carbon dioxide in water? • Alcohol in water? • Now: only solutions in which the solvent is water, and the solute is a liquid or a solid.

  7. Properties of an Aqueous Solution • Are either electrolytes or nonelectrolytes • Electrolyte: a substance that, when dissolved in water, results in a solution that can conduct electricity. • ex: NaCl dissolved in water: solid NaCl dissociates into Na+ and Cl- ions • Nonelectrolyte: does not conduct electricity when dissolved in water. • ex: pure water

  8. Strong vs. Weak Electrolytes • Strong: Solute is 100% dissociated in water • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf • Why is water good at hydrating ions?

  9. Acids and bases are electrolytes: • Some are strong: • Some are weak and ionize incompletely: • Double arrow indicates a reversible reaction: reaction can occur in both directions

  10. Reversible Reaction • Molecules ionize and then recombine • Until ionization is occurring as fast as recombination: chemical equilibrium

  11. Precipitation Reactions • formation of an insoluble product (precipitate) which separates from the solution. • This is also an example of a double displacement reaction

  12. Solubility • How do you predict whether a precipitate will form? • Depends on the solubility of the solute • p. 125-126 • Examples: • FeCO3 • KCl • AgCl

  13. Practice Determining Solubility • Ag2SO4 • CaCO3 • Na3PO4 • CuS • Ca(OH)2 • Zn(NO3)2

  14. Writing Equations • We don’t always write the entire chemical equation as if each species existed as a complete molecule • This doesn’t really reflect what’s actually happening!

  15. Molecular Equations • Written as though all species existed as molecules or whole units. • Doesn’t always reflect reality. • What’s actually happening? • Dissolved ionic compounds dissociate into ions!!

  16. Ionic Equation • Shows dissolved species as free ions. • Notice that there are ions that show up on both sides of the equation. • Spectator ions • They can be eliminated.

  17. Net Ionic Equation • To give this net ionic equation showing species that actually take place in the reaction:

  18. Example 1 • Solutions of barium chloride and sodium sulfate react to produce a white solid of barium sulfate and a solution of sodium chloride.

  19. Example 2 • A potassium phosphate solution is mixed with a calcium nitrate solution. Write a net ionic equation.

  20. Example 3 • Solutions of aluminum nitrate and sodium hydroxide are mixed. Write the net ionic equation for the reaction.

  21. 7 Oct. 2010 • Objective: SWBAT define and describe acids and bases as Arrhenius or Bronsted, and as strong or weak. • Do now: Soluble or insoluble? (try first without using your chart!) • Ca3(PO4)2 • Mn(OH)2 • AgClO3 • K2S

  22. Agenda • Homework check (ELS) • Acids and Bases: Definitions, strong and weak • Neutralization Reactions Homework: p. 161 #26, 27, 28, 29, 30, 32, 34 (JMS) Read p. 135-145 and do practice exercises a-d on p. 145

  23. Acid-Base Reactions

  24. Properties of Acids and Bases • Arrhenius definition: • Acids: ionize in water to produce H+ ions • Bases: ionize in water to produce OH- ions

  25. Acids • React with metals like Zn, Mg, Fe to produce hydrogen gas 2HCl(aq) + Mg(s)  MgCl2(aq) + H2(g) • React with carbonates and bicarbonates to produce CO2(g) 2HCl(aq) + CaCO3(s) CaCl2(aq) + H2O(l) + CO2(g) HCl(aq) + NaHCO3(s)  NaCl(aq) + H2O(l) + CO2(g)

  26. Brønsted Definition • Acid: proton donor • Base: proton acceptor • don’t need to be aqueous! • HCl(aq)  H+(aq) + Cl-(aq) proton

  27. But… • HCl(aq)  H+(aq) + Cl-(aq) • H+ is very attracted to the negative pole (O atom) in H2O • HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq) • H3O+ : hydronium ion • Above, a Brønstedacid (HCl) donates a proton to a Brønstedbase (H2O)

  28. Types of Acids • Monoprotic: each one yields one hydrogen ion upon ionization • Ex: HCl, HNO3, CH3COOH, • Diprotic: each gives two H+ ions • Ex: H2SO4 • H2SO4(aq)  H+(aq) + HSO4-(aq) • HSO4-(aq) > H+(aq) + SO42-(aq) • Triprotic: 3 H+

  29. Strong vs. Weak Acids Strong Acids Dissociate completely Weak Acids Dissociate Incompletely • HCl hydrochloric • HBrhydrobromic • HI hyroiodic • HNO3 nitric • H2SO4 sulfuric • HClO4perchloric • HF hydrofluoric • HNO2 nitrous • H3PO4 phosphoric • CH3COOH acetic

  30. Brønsted Bases • H+(aq) + OH-(aq)  H2O(l) • Here, the hydroxide ion accepts a proton to form water. • OH- is a Brønstedbase. • NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)

  31. Brønsted Acid or Base? • HBr • NO2- • HCO3- • SO42- • HI

  32. Acid-Base Neutralization • reaction between an acid and a base • produce water and a salt • salt: ionic compound (not including H+ or OH- or O2-) • acid + base  water + salt • Strong acid + Strong base example • HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) • Write the ionic and net ionic equations! • Which are spectator ions?

  33. Weak acid + Strong base example: • HCN(aq) + NaOH(aq)  NaCN(aq) + H2O(l) • HCN does not ionize completely • HCN(aq) + Na+(aq) + OH-(aq)  Na+(aq) + CN-(aq) + H2O(l) • Write the net ionic equation

  34. Acid-Base Reaction: Gas Formation • Some salts (with CO32-, SO32-, S2-, HCO3-) react with acids to form gaseous products Na2CO3(aq) + 2HCl(aq)  2NaCl(aq) + H2CO3(aq) Then the carbonic acid breaks down: H2CO3(aq)  H2O(l) + CO2(g)

  35. Homework • p. 161 #26, 27, 28, 29, 30, 32, 34 (JMS) • Bring your book to class!

  36. 13 October 2010 • Objective: SWBAT model the transfer of electrons between reactants in redox reactions by correctly writing oxidation and reduction half reactions and overall reactions; determine oxidation numbers. • Do now: Write balanced molecular, ionic and net ionic equations for this reaction between a weak acid and a strong base: H3PO4(aq) + Ba(OH)2(aq)

  37. Agenda • Do now • Read objective, debrief do now, review agenda • Homework presentation • Notes: Writing redox half reactions and assigning oxidation numbers • Practice problems • Discussion of types of redox reactions (p. 139-145) Homework: p. 162 #37, 40, 43, 44, 45, 47, 50, 54, 55 (PD), read p. 139-145

  38. 14 October 2010 • Objective: SWBAT model the transfer of electrons between reactants in redox reactions by correctly writing oxidation and reduction half reactions and overall reactions; determine oxidation numbers. • Do now: Write and label (ox. or red.) the two half-reactions: Cu(s) + 2AgNO3(aq)  Cu(NO3)2 (aq) + 2Ag(s)

  39. Agenda • Do now • Read objective, debrief do now, review agenda • Homework presentation • Notes: Writing redox half reactions and assigning oxidation numbers • Practice problems • Discussion of types of redox reactions (p. 139-145) Homework: p. 162 #37, 40, 43, 44, 45, 47, 50, 54, 55 (PD), read p. 139-145

  40. Oxidation-Reduction Reactions • What was being transferred in acid-base reactions? • Protons! • Redox reactions: electron transfer!

  41. 2Mg(s) + O2(g)  2MgO(s) • Mg2+ bonds with O2- • What’s happening with electrons? • Two steps, 2 half reactions: 2Mg  2Mg2+ + 4e- O2 + 4e-  2O2- 2Mg + O2 + 4e-  2Mg2+ + 202- + 4e- 2Mg + O2  2Mg2+ + 2O2- 2Mg2+ + 2O2-  2MgO

  42. Oxidation: Half reaction that refers to the LOSS of electrons • Reduction: Half reaction that refers to the GAIN of electrons 2Mg  2Mg2+ + 4e- O2 + 4e-  2O2- • Reducing agent: donates electrons • Oxidizing agent: accepts electrons

  43. Another Example • Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s) • (What type of reaction?) • For which elements is the charge different as a reactant and a product?

  44. Oxidation Numbers • Keeps track of electrons in redox reactions • The number of charges the atom would have in a molecule (or ionic compound) if electrons were transferred completely.

  45. Assigning Oxidation Numbers • Free elements = 0 (ex: H2, Na, K, O2) • Monotomic ions = charge of ion (ex: Li+ = +1, Fe3+ = +3) • Oxygen = -2 (peroxide O22- = -1) • Hydrogen = +1, except with metals in binary compounds (ex: LiH) then = -1 • Fluorine = -1 • In a neutral molecule, sum must = 0 • Not always integers

  46. Examples • Li2O • HNO3 • Cr2O72- • PF3 • SO2 • MnO4-

  47. 4 Types of Redox Reactions • Combination • S(s) + O2(g)  SO2(g) • Decomposition • 2HgO(s)  2Hg(l) + O2(g) • Combustion • C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l) • Displacement • Three types…

  48. Three types of displacement • Hydrogen displacement • With alkali metals and some alkaline earth metals and cold water or HCl • 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g) • Metal displacement (use activity series) • TiCl4(g) + 2Mg(l)  Ti(s) + 2MgCl2(l) • Halogen displacement • F2>Cl2>Br2>I2 (moves down group 17) • Cl2(g) + 2KBr(aq)  2KCl(aq) + Br2(l)