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CHEM 120: Introduction to Inorganic Chemistry

CHEM 120: Introduction to Inorganic Chemistry. Instructor: Upali Siriwardane (Ph.D., Ohio State University) CTH 311, Tele: 257-4941, e-mail: upali@chem.latech.edu Office hours: 10:00 to 12:00 Tu & Th ; 8:00-9:00 and 11:00-12:00 M,W,& F. Chapters Covered and Test dates.

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CHEM 120: Introduction to Inorganic Chemistry

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  1. CHEM 120: Introduction to Inorganic Chemistry Instructor: Upali Siriwardane (Ph.D., Ohio State University) CTH 311, Tele: 257-4941, e-mail: upali@chem.latech.edu Office hours: 10:00 to 12:00 Tu & Th ; 8:00-9:00 and 11:00-12:00 M,W,& F

  2. Chapters Covered and Test dates • Tests will be given in regular class periods  from  9:30-10:45 a.m. on the following days: September 22,     2004 (Test 1): Chapters 1 & 2 • October 8,         2004(Test 2):  Chapters  3, & 4 • October 20,         2004 (Test 3): Chapter  5 & 6 • November 3,        2004 (Test 4): Chapter  7 & 8 • November 15,      2004 (Test 5): Chapter  9 & 10 • November 17,      2004 MAKE-UP: Comprehensive test (Covers all chapters • Grading: • [( Test 1 + Test 2 + Test3 + Test4 + Test5)] x.70 + [ Homework + quiz average] x 0.30 = Final Average •                               5

  3. Chapter 7: Reactions and Solutions 1. Types of chemical reactions : combination, decomposition, or replacement. 2. Classes of chemical reactions: precipitation, reactions with oxygen, acid–base, and oxidation–reduction. 3. Solution terms: solution, solute, and solvent. 4. Kinds of solutions: gas/liquid, liquid/ liquid, solid/liquid 5. Solubility and equilibrium. 6. Solution concentration: weight/volume percent and weight/weight percent. 7. Solution concentration: molarity. 8. Dilution: Preparing solutions. 9. Interconvert molar concentration of ions and illiequivalents/liter. 10. Concentration-dependent solution properties: Collegative properties. 11. Unique chemical and physical properties of water. 12. Role of electrolytes in blood and their relationship to the process of dialysis.

  4. Writing chemical reactions • We can classify some chemical reactions according to certain patterns that are observed. This helps us to predict the products of reactions.

  5. Combination reactions • Combination Reaction

  6. Decomposition reactions

  7. Replacement reactions: single replacement

  8. Replacement reactions: double replacement

  9. Classify as to type of reaction • 2Al(OH)3(s) g Al2O3(s) + 3H2O(g) • Fe2S3(s) g 2Fe(s) + 3S(s) • Na2CO3(aq) + BaCl2(aq) g BaCO3(s) + NaCl(aq) • C(s) + O2(g) g CO2(g)

  10. Types of chemical reactions • Precipitation reactions: mix reactants together and get an insoluble precipitate (not soluble in water). • How do you know what cmpds are insoluble? • Learn table 7.1

  11. 1A salts, NO3-’s, acetates (CH3COO-), NH4+’s are very soluble. • Cl-’s, Br-’s and I-’s are soluble except when combined with Ag+, Pb2+, Hg22+. • CO32-’s, PO43-’s, S 2-’s generally insoluble (except for 1A’s and NH4+) • OH-’s: Ba2+, 1A’s are soluble; others generally insoluble • **The SO42-’s of Ba2+, Ag+, Ca2+, Sr2+, Pb2+are insoluble. Others are soluble.

  12. Are these soluble? • Ag2SO4 • Li2S • Pb(NO3)2 • AgCl • BaSO4 • (NH4)2SO4

  13. Predict whether get precipitation when mix • Zinc sulfate and barium chloride • Sodium sulfate and potassium chloride) • Ammonium carbonate and calcium chloride • Strontium nitrate and potassium phosphate

  14. Reactions with oxygen • Combustion Reactions: When compounds containing C and H react with O2 (burning or combustion) get CO2 and H2O. (Greenhouse effect) • Corrosion: 4Fe(s) + 3O2(g) g 2Fe2O3(s) rust

  15. Acid-base reactions • Acid-base reactions involve transfer of a H+ from the acid (starts with H) to the base (hydroxide) to form a salt and water • Oxidation-reduction will be covered in Chapter 9

  16. Properties of solutions • A solution is a homogeneous mixture of two or more substances and is made up of a solvent and one or more solutes. • The solutes are the species that are being dissolved in the solvent. The solvent is usually present in the greater amt. • An aqueous solution has ________ as the the solvent.

  17. Types of solutions • gas in gas --air • gas in liquid--soda • gas in solid--gas on solid catalyst • liquid in liquid • liquid in solid--mercury amalgam • solid in liquid • solid in solid--14-karat gold

  18. Ionic compounds (electrolytes) dissociate into ions when dissolved in water. The solutions conduct electricity. • Molecular compounds in general do not dissociate into ions in aqueous solution. The solutions do not conduct electricity and are ________________.

  19. Properties of solutions • 1. A solution is a ___________ mixture. Each species in the solution • 2. retains its ________ identity. • 3. __________composition • 4.clear and transparent (but can have color) • 5. remains uniform throughout for all time • 6. can be separated into its components by ___________ means.

  20. Solute size is _________________. • The solute cannot be separated from the solvent by paper filtration.

  21. Colloids (colloidal suspension) • Colloids: have similar properties as solutions but the particle size is • Colloids behave differently when exposed to light. • ________ effect helps distinguish between solutions and colloids. • Homogenized milk is a _____________.

  22. Emulsions and Suspensions • Emulsions, suspensions; larger particle size than solutions. Solute separates on standing. • Particle size is > 200nm. • Filterable.

  23. Degree of solubility • Can dissolve different amounts of solute in solvent. • Maximum amount of solute that can be dissolved in a given amt of solvent at a given temp:____________

  24. Say a solution is __________ if the maximum amount of solute is dissolved in a given solvent. Solid solution • Dynamic equilibrium set up between dissolved and undissolved solute • Maximum amout of solute that can be dissolved in a given amt of solvent at a given temp: solubility

  25. If a solution has less than the max amt dissolved: • If a solution has more than the max amount dissolved: ______________: unstable--excess solute will fall to bottom and form a precipitate.

  26. Factors that affect solubility • Like dissolves like--smaller difference in polarity between solute and solvent, more soluble. • Polar solutes in polar solvents • Non-polar solutes in non-polar solvents • Ionic solids in polar solvents

  27. Factors that affect solubility I. Temperature • A. ionic compound in water: • B. Gas in water:

  28. Factors that affect solubility II. Pressure • A. Pressure changes have little effect on the solubility of a solid or liquid in another liquid.

  29. Pressure effects continued • B. The solubility of a gas in a liquid

  30. Under 1 atm total pressure, the partial pressure of O2 is 159 mmHg and the solubility of O2 in blood is 44g/100mL. • In Denver (mile high city) PO2 = 132mmHg and the solubility of O2 in the blood is 37g/100mL • Mt Whitney (2.5miles high) PO2 = 98 mmHg and the solubility of O2 in the blood is 27g/100mL • Mt Everest (5.8 miles) PO2 = 52 mmHg and the solubility of O2 in the blood is 14g/100mL

  31. Problem • The solubility of N2 in blood at 37oC and at a partial pressure of 0.80 atm is 5.6 x 10-4 mol/L. • A deep sea diver breathes compressed air with the partial pressure of N2 equal to 4.0 atm. How much N2 is dissolved in the blood at this pressure?

  32. Concentration of solutions: • Concentration gives us the amount of solute dissolved in a given amt of solvent or in a given amt of solution. • There are different ways of expressing concentration.

  33. Percent (W/V %) • Weight/volume percent weight/volume% (W/V%)= (mass of solute in g vol of soln in mL) x 100% • Note that the volume of a solution does not equal the volume of solute and solvent.

  34. Percent (W/W%) • Weight/weight percent • Weight/weight % (W/W%) = (mass of solute mass of solution) x 100% • The mass of the solution =

  35. Problems • Calc the composition of the soln in W/V%: 20.0g acetic acid in 2.50L sollution • Calc the W/W % of 31.0 g of KCl in 152 g of water. • Calc the W/W% of 50.0 g KCl in 5.00 x 102 mL solution (d = 1.14g/mL)

  36. How many grams of solute are needed to prepare: • A. 2.50 x 102 g of 5.00% (W/W) NH4Cl • B. 2.50 x 102mL of 3.50% (W/V) Na2CO3 • Calc the amount of water that must be added to 5.00 g of urea (NH2)2CO in the preparation of a 16.2 W/W % by mass solution.

  37. Molarity M • Molarity (M) = moles of solute vol of soln in L • units of molarity: mol of solute/L solution • M = moles solute(M )/V and #moles = M x V • Molarity and W/V% are temperature dependent.

  38. Molarity problems • Calc the molarity of 20.0g acetic acid in 2.50L solution. • Calc the no of grams of solute needed to make 2.50 x 102 mL of 0.200M KOH. • Calc the volume, in ml, needed to provide 2.14g of NaCl from a 0.270M solution.

  39. Dilution • When a solution is diluted you add more solvent. The no. of moles of solute does not change. (#moles= M x V) • So initial vol x molarity (mol/L) = # mol solute and final vol x new molarity = same # moles # moli = # molf • So MiVi = MfVf

  40. Dilution problems • 50.0mL of a 0.250M sucrose soln was diluted to 5.00 x 102 mL. What is the molar conc of the resulting solution? • A 6.00 mL portion of an 8.00M stock solution is to be diluted to 0.400M. What will be the final volume after dilution?

  41. Molality (m) • molality (m) = moles of solute/mass in kg of solvent • A 2.5m (molal) NaCl solution has 2.5moles of NaCl dissolved in 1000g or 1kg of solvent • Molality is temperature independent.

  42. Concentration-dependent solution properties • Colligative properties are properties

  43. So NaCl(aq) g Na +(aq) + Cl-(aq) • K2SO4(aq) g 2K+(aq) + SO42-(aq) • C12H22O11(aq) gC12H22O11(aq)

  44. We are going to examine the effect of adding a solute to a solvent on • 1. vapor pressure • 2. freezing point • 3. boiling point • 4. osmosis • Remember that colligative properties depend only on the number of particles in solution and not on their identity.

  45. Vapor pressure • When a solute is added to a solvent the vapor pressure (equilibrium) of the solution is lower than that of the pure solvent. (explain) • Raoult observed the relationship between the amount of the solvent and the vapor pressure of the solution

  46. Effects on freezing and boiling points • What effect does vp lowering have on the freezing pt and boiling pt of a solution? • Since the vapor pressure of the solution is lower than the vapor pressure of the pure solvent,

  47. When a nonvolatile solute is added to a solvent the freezing point of the solution is lowered. (explain)

  48. Math relationship for b.pt. elevation and f. pt depression Dtf = I mkf kf is the freezing pt constant i = no of particles in solution per formula unit For molecular species i =1 For NaCl i = These i values For K2SO4 i = assume 100% For Al2(SO4)3 i = ionization. For water kf = 1.86oC/m

  49. Dtb = imkb kb is the b. pt. constant; m is the molality of the solution and i is the no. of particles in solution. • For water kb =0.52oC/m

  50. What are the normal freezing and boiling pts of • a. 58.5g NaCl in 100. g of water • b. 60.0 of urea [(NH2)2CO] in 100. g of water.

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