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Chemistry 100(02) Fall 2010

Chemistry 100(02) Fall 2010. Instructor: Dr. Upali Siriwardane e-mail : upali@chem.latech.edu Office : CTH 311 Phone 257-4941 Office Hours : M,W, 8:00-9:00 & 11:00-12:00 a.m Tu,Th,F 9:00 - 10:00 a.m.   Test Dates : March 25, April 26, and May 18; Comprehensive Fina

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Chemistry 100(02) Fall 2010

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  1. Chemistry 100(02) Fall 2010 Instructor: Dr. UpaliSiriwardane e-mail: upali@chem.latech.edu Office: CTH 311 Phone 257-4941 Office Hours: M,W, 8:00-9:00 & 11:00-12:00 a.mTu,Th,F 9:00 - 10:00 a.m.   Test Dates: March 25, April 26, and May 18; Comprehensive Fina Exam: 9:30-10:45 am, CTH 328. October 4, 2010 (Test 1): Chapter 1 & 2 October 27, 2010 (Test 2): Chapter 3 & 4 November 17, 2010 (Test 3): (Chapter 5 & 6) November 18, 2010 (Make-up test) comprehensive: Chapters 1-6 9:30-10:45:15 AM, CTH 328

  2. Chapter 5. Chemical Reactions • 5.1 Exchange Reactions: Precipitation and Net Ionic Equations Page164 • 5.2 Acids, Bases, and Acid-Base Exchange Reactions Page171 • 5.3 Oxidation-Reduction Reactions Page179 • 5.4 Oxidation Numbers and Redox Reactions Page184 • 5.5 Displacement Reactions, Redox, and the Activity Series Page187 • 5.6 Solution Concentration Page191 • 5.7 Molarity and Reactions in Aqueous Solutions Page198 • 5.8 Aqueous Solution Titrations Page201 • CHEMISTRY IN THE NEWS: The Breathalyzer Page189 • CHEMISTRY YOU CAN DO: Pennies, Redox, and the Activity Series of Metals Page192

  3. Chapter 5. KEY CONCEPTSChemical Reactions

  4. Solution Chemistry Collisions of reactants decides the rates of reactions Why Solution Reactions? Why not gaseous Reactions? In Gases: Very fast Reactions Why not solid Reactions? Solids: Very slow Reactions Solution reactions are manageble! Liquids: Fast Reactions

  5. What is a Solution? • A solution: A homogeneous mixture of two or more components. Sugar in water Oxygen in water Air Dental fillings Saline

  6. Solute • substance that is present in smallest quantity • dissolved substance(s) • can be either a gas, a liquid, or a solid • one or more present in a solution Solvent • substance present in largest quantity • only one per solution • water in aqueous solutions

  7. Cola Drinks Solvent • water Solutes • carbon dioxide (gas) • sweetner (solid) • phosphoric acid (liquid) • caramel color (solid)

  8. Reactions in Solutions There are many solvents However, water is most abundant dissolving many chemicals. Water can interact with both cations and anions making it the best solvent for ionic compounds.

  9. Dissolution of (a) Ionic and(b) Molecular Compounds

  10. Types of Chemical Equations Molecular equation: Equation with formula, correct stoichiometric coefficients and physical form written within parenthesis. Ionic equation: All the ionic compounds soluble in water are separated into ions written with their ionic charge and (aq). Net Ionic equation: Ionic equation with all spectator ions removed from both sides.

  11. NaCl(aq)+AgNO3(aq) -->AgCl(s)+ NaNO3(aq) Molecular equation: NaCl (aq) + AgNO3 (aq) --> AgCl (s) + NaNO3 (aq) Ionic Equation: Na+ (aq) + Cl-(aq) + Ag+ (aq) + NO3-(aq) --> AgCl(s) + Na+(aq) + NO3- (aq) Spectator Ions: Na+ (aq) and NO-3 (aq) Net Ionic Equation: Cl- (aq) + Ag+ (aq) --> AgCl (s)

  12. Spectator Ions Ions appearing on both side of an ionic equation. Ionic Equation: Na+ (aq) + Cl-(aq) + Ag+ (aq) + NO3-(aq) --> AgCl(s) + Na+(aq) + NO3- (aq)

  13. More Examples • HCl(aq) + NaOH(aq) ----> NaCl(aq) + H2O(l) • NaOH(aq) + HC2H3O2(aq) ---> NaC2H3O2(aq) + H2O(l)

  14. Types of Chemical Reactions Based on driving force a) Precipitation Reactions b) Acid-base Reactions c)Gas-forming Reactions d)Oxidation-reduction (REDOX)Reactions

  15. Precipitation reactions • They are double displacement reactions of ionic compounds where an insoluble salt is formed when two aqueous salt solutions are mixed.

  16. Solubility rules for ionic compounds • All acids are soluble. • All Na+, K+ and NH4+ salts are soluble. • All nitrate and acetate salts are soluble. • All chlorides except AgCl and Hg2Cl2 PbCl2 are soluble. • All sulfates are soluble except PbSO4, Hg2SO4, SrSO4 and BaSO4. • All sulfides are insoluble except those of the Group IA (1), IIA (2) and ammonium sulfide. • All hydroxides are insoluble except those of the group IA(1) and IIA Ba(OH)2. Sr(OH)2 and Ca(OH)2

  17. Illustration of SomeSolubility Rules

  18. Precipitation of Silver Chloride AgNO3 + NaCl  AgCl + NaNO3 precipitate

  19. H2O H2O Ionic equations • When ionic substances dissolve in water, they dissociate into ions. • AgNO3 Ag++ NO3- • KClK+ + Cl- • When a reaction occurs, only some of the ions are actually involved in the reaction. Ag++ NO3- +K+ + Cl-AgCl(s) + K+ + NO3-

  20. Ionic equations • To help make the reaction easier to see, we commonly list only the species actually involved in the reaction. • Full ionic equation • Ag++ NO3- +K+ + Cl- AgCl(s) + K+ + NO3- • Net ionic equation • Ag++Cl-AgCl(s) • NO3- and K+ are referred to as spectator ions.

  21. Precipitation of Barium Sulfate Double Displacement: BaCl2(aq) + Na2SO4(aq)  2NaCl(aq) + BaSO4(s) precipitate

  22. Ionic Equations Molecular Equation: BaCl2(aq) + Na2SO4(aq)  2NaCl(aq) + BaSO4(s) precipitate Total Ionic Equation: Ba+2 + 2Cl-1 + 2Na+ + SO4-2  2Na+ +Cl-1 + BaSO4(s) Net Ionic Equation: Ba+2 + SO4-2BaSO4

  23. Precipitation or Not • MgI2 + NaNO3= 2 NaI + Mg(NO3)2 • Ba(NO3)2+Na2SO4= BaSO4 + 2 NaNO3 • AgCl +NaNO3 = AgNO3 + NaCl

  24. Acid/base Reactions Acid • substance that donates H+ ions to solution • sour-tasting substances • substances whose aqueous solutions are capable of turning blue litmus indicators red • dissolves certain metals to form salts • react with bases or alkalis to form salts • substance that donates a OH-1 ion to solution • hydroxides and oxides of metals • bitter tasting, slippery solutions • turn litmus blue • react with acids to form salts Base

  25. Neutralization Reactions Formation of water is the diving force acid + base  “salt” + water HCl + NaOHNaCl + H2O H2SO4 + 2KOH  K2SO4 + 2H2O Salt • Substances produced by the reaction of an acid with a base • Characterized by ionic bonds and high melting points • Electrical conductivity when melted or when in solution • Has a crystalline structure when in the solid state

  26. How do find precursor Acid and base of a Salt • Acid (A) + Base(B) = Salt + water (H2O) • HA + BOH = BA + H2O • E.g. LiNO3 • B (Li) A (NO3) • OH H • BOH(LiOH) HA(HNO3)

  27. Some acids, bases and their salts Acid Sodium salt Name Formula Name Formula Acetic acid HC2H3O2 Sodium acetate NaC2H3O2 Hydrogen chloride HCl Sodium chloride NaCl Nitric acid HNO3 Sodium nitrate NaNO3 Phosphoric acid H3PO4 Sodium phosphate Na3PO4 Sulfuric acid H2SO4 Sodium sulfate Na2SO4 Base Chloride salt Name Formula Name Formula Sodium hydroxide NaOH Sodium chloride NaCl Barium oxide BaO Barium chloride BaCl2 Sodium oxide Na2O Sodium chloride NaCl Ammonia NH3 Ammonium chloride NH4Cl

  28. Ionization of Acids in Water

  29. Common Acids and Bases

  30. Ionic Equations StrongAcid/base Molecular Equation: HCl(aq) + NaOH(aq) ----> NaCl(aq) + H2O(l) Total Ionic Equation: H+ + Cl-1 + Na+ + OH-1  Na+ +Cl-1 + H2O Net Ionic Equation: H+ + OH-1  H2O NaOH(aq) + HC2H3O2(aq) -----> NaC2H3O2(aq) + H2O(l) Na+(aq) + OH -(aq) + HC2H3O2(aq) -----> Na+(aq) +C2H3O2-(aq) + H2O(l) HC2H3O2(aq) + OH-1 C2H3O2-(aq) + H2O

  31. Ionic Equations WeakAcid/base Molecular Equation: NaOH(aq) + HC2H3O2(aq) -----> NaC2H3O2 (aq) + H2O(l) Total Ionic Equation: Na+(aq) + OH - (aq) + HC2H3O2aq) -----> Na+(aq) +C2H3O2 -(aq) + H2O(l) Net Ionic Equation: HC2H3O2 (aq) + OH-C2H3O2 -(aq) + H2O

  32. Acid/base Ionic Equations H2SO4 + 2KOH  K2SO4 + 2H2O Total Ionic Equation: 2H+ + SO4-2 + 2Na+ + 2OH-1  2Na+ +2Cl-1 + 2H2O Net Ionic Equation: 2H+ + 2OH-1  2H2O

  33. Strong Acids • strong – completely ionized • weak –partially ionized • Strong Acids(strong electrolytes) • HCl, HBr, HI, H2SO4, HNO3, HClO4 (all others are weak) • Weak acids( weak electrolytes) • CH3COOH, HF, HCN, H3PO4, HCOOH, H2CO3 • Table 5.2 page 171.

  34. Strong Bases • Strong bases Metal Hydroxides • (Group IA):Li, Na K. (Group IIA): Ca, Sr, Ba (all others metal hydroxides are weak) • Weak bases • NH3, amines -CH3NH2 • Table 5.2 page 171.

  35. Gas-Forming Exchange Reaction:CO2, SO2, H2S

  36. Reaction of Metal Carbonates with Acids Molecular Equation: CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2CO3(aq) H2CO3(aq) H2O +CO2(g) Total Ionic Equation: CaCO3(s) + 2H+ + 2Cl-1 Ca+2 + 2Cl-1 + H2O + CO2(g) Net Ionic Equation: CaCO3(s) + 2H+ Ca+2 + H2O + CO2(g)

  37. Reaction of Metal Carbonates with Acids Alka-Seltzer NaHCO3(aq) + HCl(aq)NaCl(aq) + H2O +CO2(g) Net Ionic Equation: HCO3-1 + 2H+ H2O + CO2(g) Tums CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O +CO2(g) Net Ionic Equation: CO3-2 + 2H+ H2O + CO2(g)

  38. Reaction of Metal Sulfites with Acids Molecular Equation: CaSO3(s) + 2HCl(aq) CaCl2(aq) + H2SO3(aq) H2SO3(aq) H2O +SO2(g) Total Ionic Equation: CaSO3(s) + 2H+ + 2Cl-1 Ca+2 + 2Cl-1 + H2O + SO2(g) Net Ionic Equation: CaSO3(s) + 2H+ Ca+2 + H2O + SO2(g)

  39. Reaction of Metal Sulfides with Acids Molecular Equation: Na2S(aq) + 2HCl(aq) 2NaCl(aq) + H2S(g) Total Ionic Equation: 2Na++ S-2 + 2H+ + 2Cl-1 2Na+ + 2Cl-1 + H2S(g) Net Ionic Equation: S-2 + 2H+ H2S(g)

  40. Oxidation-Reduction Reactions

  41. Recognizing Redox Reactions

  42. Oxidation-Reduction(Redox) Reactions Electrons are transferred from one compound to the other resulting in a chemical change. E.g. Zn(s) + 2HCl(aq) --->ZnCl2(aq) + H2(g) includes Single Replacement Reactions Oxidation–loss of electrons Reduction– gain of electrons oxidizing agent– substance that causes oxidation reducing agent– substance that cause reduction

  43. Oxidation number or State • A number assigned to a atom in compounds, ions and polyatomic ions to show the number of electrons relative to an atom in the element.

  44. Rules for Assigning Ox # a) Oxidation number of atoms in an element is zero (0). e.g. O2 b)Monoatomic ions: Ox # equal to charge. E.g. Na+, Ox # = +1 c) Sum of the oxidation numbers in an element, compound is equal to zero. Sum of the oxidation numbers in an ion, cation or anion is equal to the ionic charge d) As a rule ONs of H =+1, and O=-2 almost most of the time. The group number in the periodic table could be used for main group elements (p and s blocks). d and f block elements show variable ONs E.g. Fe shows either +3 or +2.

  45. Oxidation State What is the oxidation state of Cl in HClO4? • H  +1 • O  -2 • neutral compound, thus sum equals zero • 4O  4  -2 = -8 • H  1  +1 = +1 • 0 = +1 + (y) + (-8) y = +7

  46. Oxidation State What is the oxidation state of S in H2SO4? • H  +1 • O  -2 • neutral compound, thus sum equals zero • 4O  4  -2 = -8 • 2H  2  +1 = +2 • 0 = +2 + (x) + (-8) x = +6

  47. Assigning the Oxidation State • Assign the oxidation states to each atom in • NaCl • O2 • CBr4 • S8 • MnO2 • KMnO4 • K2Cr2O7

  48. Which of the following reactions are redox? a) NaCl + AgNO3 ----> AgCl + NaNO3 b) NaOH + HCl ----> NaCl + H2O c) Zn + 2HCl ----> ZnCl2 + H2 d) 2Cr + 6HCl ----> 2CrCl3 + 3H2 e) MnO2 + 4HBr ----> Br2 + MnBr2 + 2H2O

  49. Half Redox Reactions two half-reactions in a redox reaction: one where the oxidation is talking place and the other where reduction is taking place.

  50. Half Reactions E.g. 2 Na + Cl2 ----> 2NaCl ON 0 0 +1 -1 Oxidation Na ----> Na+ + e- ; Na increase ON, 0 ----> +1 Reduction Cl2 + 2e- ----> 2Cl- ; Cl decrease ON, 0 ----> -1

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