1 / 82

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

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). General Properties. What are the properties of a solution?. Homogeneous mixture 2 Components: Solute – is dissolved (smaller amount)

mead
Download Presentation

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

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  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)

More Related