Chapter 4, 15 and 19 (Sections 4.1- 4.8) (Sections 15.1, 15.3, 15.4) (Sections 19.1-19.3, 19.8)

1. Reactions in Aqueous Solutions Chapter 4, 15 and 19(Sections 4.1- 4.8) (Sections 15.1, 15.3, 15.4) (Sections 19.1-19.3, 19.8)

2. General Properties • What are the properties of a solution? • Homogeneous mixture • 2 Components: • Solute – is dissolved (smaller amount) • Solvent – does dissolving (larger amount) • Aqueous when the solvent is water

3. Dissociation • What is dissociation? • Within an aqueous solution, the ions in the solute are separated from each other

4. Solvation • What is solvation? • The dissociated ions of the solute spread out and become surrounded by the solvent molecules

5. Three Types • What are the three types of solutions? • Unsaturated • Solvent is still able to dissolve more solute • Saturated • Solvent has dissolved the maximum amount of solute • Supersaturated • Solvent contains more solute than a saturated solution can normally hold

6. Measurement of Volume • What units am I allowed to use for volume? • mL and L are the most common, but IB will not use these • IB will commonly use: • dm3 for L • cm3 for mL • Negative superscript means you invert the unit • dm-3 = 1/dm3 • cm-3 = 1/cm3

7. Concentration • What is concentration? • Concentration is a measure of how much solute is dissolved in the solvent • Tells us how much of a reactant is presentand allows us to do stoichiometry • Use square brackets when expressing concentration • Ex. [H+]

8. Concentration • How can we represent moles in a solution? • Molarity (M) • 16 M, say 16 molar • Formula • High M = concentrated • Low M = dilute

9. Molarity • I dissolved 29.22 g of sodium chloride in 1000 mL of water. • How many moles of NaCl? • (29.22 g)/(58.44 g/mol) = 0.5000 mol • What is the volume? • 1000 mL = 1 L • What is its molarity? • (0.5000 mol)/(1 L) = 0.5 M

10. Molarity Practice • What is the molarity of a barium chloride solution that has 40.0 g of solute dissolved in 5.0 L of water? • Solute  BaCl2 (208.23 g/mol) • Solvent  H2O • Answer  0.038 M

11. Molarity Practice Find the molarity of the following solutions. • 1.2 moles of calcium carbonate in 1.22 liters of solution. • 0.98M • 120 grams of calcium nitrite in 240 mL of solution. • 3.8M • 98 grams of sodium hydroxide in 2.2 liters of solution. • 1.1M • 45 grams of ammonia in 0.75 L of solution. • 3.5M • 734 grams of lithium sulfate are dissolved to make 2500 mL of solution. • 2.7M • 6.7 x 10-2 grams of Pb(C2H3O2)4 are dissolved to make 3.5 mL of solution. • 0.043M

12. Serial Dilution • How can we make a series of solutions starting with a concentrated solution? • Called serial dilution • Start with a concentrated “stock” solution • Use the molarity ratio to figure out your measurements • M1V1 = M2V2 • 1 = initial (have) • 2 = final (wanted) • NOTE: You take the initial volume, transfer it and dilute it to the final volume with solvent

13. Serial Dilution Fructose – C6H12O6

14. Acid-Base Chemistry • Using the Arrhenius definition, what are acids and bases? • Acids produce H+ ions when dissociated in water • Bases produce OH- ions when dissociated in water

15. Acid-Base Chemistry • Using the Arrhenius definition, what happens when you combine an acid and a base? • Examples: HCl + KOH  HNO3 + KOH  3HBr + Al(OH)3 • Neutralization (irreversible) • Produces salt and water • A salt is an ionic (metal/nonmetal) compound that uses ions other than hydrogen and hydroxide KCl + H2O KNO3 + H2O AlBr3 + 3H2O

16. Acid-Base Chemistry • What are the products of the following neutralization reactions? • HCl + NaOH → • HClO3 + NH4OH → • H2SO4 + 2KOH → • NaCl + H2O • NH4ClO3 + H2O • K2SO4 + 2H2O

17. Acid-Base Chemistry • Using the Bronsted-Lowry definition, what are acids and bases? • Acids are any species that can donate a proton (H+) in solution • Proton Donors • 1 – monoprotic acid (HCl) • 2 – diprotic acid (H2SO4) • 3 – triprotic acid (H3PO4) • Bases are any species that accept a proton (H+) in solution • Proton Acceptors

18. For Example, Ammonia • Why is ammonia considered to be a base? • NH3 • It cannot produce a hydroxide ion (OH-) • It can, however, accept a proton (H+) from another acid like HCl or even H2O • After accepting the proton, it is now ammonium (NH4+)

19. Amphoteric • What does amphoteric mean? • A substance can act as an acid or a base • Ex. Water • OH- (hydroxide ion) • H3O+ (hydronium ion)

20. Acid-Base Chemistry • Using the Bronsted-Lowry definition, what happens when you combine an acid and a base? • Proton transfer (reversible) • Produces two conjugates • Conjugate base – acid after proton lost • Conjugate acid – base after proton is gained

21. Acid-Base Chemistry • What are conjugate acid/base pairs? • Pairing up the original acid or base with it’s conjugate partner • Acid/Conjugate Base • CH3COOH/CH3COO- • Base/Conjugate Acid • H2O/H3O+

22. Acid-Base Chemistry • What are the conjugate pairs? NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(aq) H2O/ OH- NH3/ NH4+

23. Conjugate Acids and Bases • What are the conjugate bases of these acids? • HNO3 • H2O • H3O+ • H2SO4 • HBr • HCO3- • What are the conjugate acids of these bases? • OH- • H2O • HCO3- • SO42- • ClO4-

24. Electrolytes • What is the difference between an electrolyte and a non-electrolyte? • An electrolyte conducts electricity when dissolved in water • Strong electrolytes conduct electricity better than weak ones • All acids and bases are electrolytes

25. Acid-Base Chemistry • What is the difference between weak and strong acids/bases? • Weak acids and bases are weak electrolytes • Do not dissociate completely, reversible reaction, so not many ions in solution • Ex. CH3COOH ↔ CH3COO- + H+ • Strong acids and bases are strong electrolytes • Dissociate completely, so more ions are in the solution • Ex. HCl → Cl- + H+

26. Table 4.1 on Page 111

27. pH • What does pH represent? • Stands for the power of hydrogen • Represents the concentration of hydrogen ions, [H+], in a solution

28. pH Scale • What is the pH scale? • A scale that indicates how acidic or basic a solution is • Ranges from 0-14 • <7 = acidic • 7 = neutral • >7 = basic/alkaline

29. pH Scale Common Substances

30. pH • How do you calculate pH? • Formula • pH is unitless • A negative log means as [H+] increases, pH decreases • Each time pH decreases by 1 the [H+] is 10 times more • ie. pH 2 has 10 times the [H+] as pH 3

31. pOH • How is pOH related to pH? • Indirectly • Formula • pOH is unitless • As pH increases, pOH decreases • As [H+] decreases, [OH-] is increases

32. Indicators • What is an acid/base indicator? • Indicates whether a solution is acidic or basic based on color changes • Common Indicators • Litmus Paper • Bromothymol Blue • Methyl Red • Phenolphthalein • Universal Indicator • Data Booklet – Table 16

33. Titrations • What is a titration? • A volumetric analysis to determine the unknown concentration of an acid with a base of known concentration • A buret is used to determine the volume of base that was used

34. Equivalence Point • What is an equivalence point? • The exact point when the volume of base added completely neutralizes the acid • Can then stoichiometrically determine the concentration of the acid using molar ratios • MacidVacid = MbaseVbase

35. Types of Reactions • What are the five general types of reactions? • Decomposition (breakdown) • C → A + B • Synthesis (combination) • A + B → C • Combustion • A + O2 → B + Water • Single Displacement • A + BC → AC + B • Double Displacement • AB + CD → AC + BD

36. Precipitation Reactions • What is a precipitation reaction? • Occurs when aqueous anions and cations combine to form an ionic solid that is insoluble • Almost always a double displacement reaction

37. Precipitation Reactions • What is the difference between soluble and insoluble? • What is a precipitate? • Soluble will dissolve in a particular solvent and insoluble will not • The insoluble solid product formed in an aqueous chemical reaction • Process of creating a precipitate is called nucleation

38. Johnny was diligently working in the lab trying to generate a little nucleation. But things weren't going so well and he's just not having any luck. Then all of a sudden his lab partner (Dittmore) fumbles in, accidently knocking Johnny's beaker of silver nitrate into some potassium chloride which spills all over Johnny. 'Heavens to Betsy!' Johnny gleefully proclaims as a beautiful white solid of silver chloride materializes. And that's why, the legend goes…

39. General Solubility Rules for Ionic Compounds in Water (pg. 113)

40. Solubility Practice • Please determine whether the following are soluble or insoluble in water and why. • AgNO3 • NaOH • RbClO3 • AgI • CaSO4 • Soluble (nitrate) • Soluble (alkali metal ex) • Soluble (chlorate) • Insoluble (silver ion ex) • Insoluble (calcium ion ex)

41. Precipitation Reactions • How is an ionic equation different from a molecular equation? • An ionic equation shows all aqueous chemicals completely dissociated into cations and anions • Remember! – Insoluble products do not dissociate

42. Ionic Equation Example #1 • Molecular: • Pb(NO3)2 (aq)+ 2NaI (aq) → PbI2 (s)+ 2NaNO3 (aq) • What is the precipitate? • Lead (II) Iodide – only solid product • Ionic: • Pb2+ + 2NO31- + 2Na1+ + 2I1- → PbI2 + 2Na1+ + 2NO31- • What is a spectator ion? • An ion that has nothing to do with the overall reaction • What is the net ionic equation? • An ionic equation that removes the spectator ions • Ex. Pb2+ + 2I1- → PbI2

43. Ionic Equation Example #2 • AlCl3 (aq) + NaOH (aq) → • What is the balanced reaction? • AlCl3 (aq) + 3NaOH (aq) → Al(OH)3 (s) + 3NaCl (aq) • Which product is the precipitate? • Al(OH)3 • What is the ionic equation? • Al3+ + 3Cl1- + 3Na1+ + 3OH1- → Al(OH)3 + 3Na1+ + 3Cl1- • What is the net ionic equation? • Al3+ + 3OH1- → Al(OH)3

44. Precipitation Reactions • What would be the net ionic equation if all products were soluble? • There would not be a net ionic equation • All ions would be spectator ions and cancel each other out in the equation CoCl2(aq) + Na2SO4(aq) → CoCl2 (aq) + Na2SO4 (aq) → CoSO4 (aq) + 2NaCl (aq)

45. Ionic Equation Practice #1 • Determine the products of the reaction and then determine the ionic AND net equation. • AgNO3 (aq) + NaCl (aq)  • AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq)

46. Ionic Equation Practice #2 • Determine the products of the reaction and then determine the ionic AND net equation. • AgNO3 (aq) + K2CrO4 (aq)  • 2AgNO3 (aq) + K2CrO4 (aq)  Ag2CrO4 (s) + 2KNO3 (aq)

47. Ionic Equation Practice #3 • Determine the products of the reaction and then determine the ionic AND net equation. • Mg(NO3)2 (aq) + Na2CO3 (aq)  • Mg(NO3)2 (aq) + Na2CO3 (aq)  MgCO3 (s) + 2NaNO3 (aq)

48. Ionic Equation Practice #4 • Determine the products of the reaction and then determine the ionic AND net equation. • Pb(NO3)2 (aq) + KI (aq) → • Pb(NO3)2 (aq) + 2KI (aq) → PbI2 (s) + 2KNO3 (aq)

49. Ionic Equation Practice #5 • Determine the products of the reaction and then determine the ionic AND net equation. • Pb(NO3)2 (aq) + NaOH (aq) → • Pb(NO3)2 (aq) + 2NaOH (aq) → Pb(OH)2 (s) + 2NaNO3 (aq)

50. 5 Precipitation Reactions • AgNO3(aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) • 2AgNO3(aq) + K2CrO4 (aq)  Ag2CrO4 (s) + 2KNO3 (aq) • Mg(NO3)2(aq) + Na2CO3 (aq)  MgCO3 (s) + 2NaNO3 (aq) • Pb(NO3)2(aq) + 2KI (aq) → PbI2 (s) + 2KNO3 (aq) • Pb(NO3)2(aq) + 2NaOH (aq) → Pb(OH)2 (s) + 2NaNO3 (aq)