Plan, Ppt 13: Substances in Aqueous Solution (PS5, 19-22 and PS6, 1-7 material)

1. Plan, Ppt 13: Substances in Aqueous Solution(PS5, 19-22 and PS6, 1-7 material) • Reminder: Ionic vs. Molecular (non-acid) vs. Acid • Molarity of solutes I • Qualitative Issues (types of solutes/solutions) • Electrolytes (Operational) • Electrolytes (Conceptual) • Strong vs. Weak vs Non Electrolyes • Molarity Reprise (electrolyte issues) • Solubility rules for ionic compounds Ppt13

2. Reminder—Ionic Compounds • Metal or “NH4” listed first in formula • Made up of ions (cation and anion type) • You need to be able to: • Write the formulas of the cation and anion • State the number of each ion present in one FU of the compound • Example: Na3PO4 • Made up of Na+ ions and PO43- ions • 3 Na+ ions / FU; 1 PO43- / FU KNOW YOUR IONS!! (PS3) • NOTE: NaOH, KOH, etc. are ionic compounds. Because of the OH- ion(not the cation!), these ionic compounds are also called “bases”. Ppt13

3. Reminder—Molecular Compounds • In this class, a compound that is not ionic! • NO metal first; a nonmetal (unless “NH4”) • Learned to name binary molecular compounds… • Non-acids (e.g., sulfur trioxide, SO3), and • (binary) acids (e.g., hydrochloric acid, HCl) • …and one class of ternary molecular compounds • (ternary) acids (e.g., chloric acid, HClO3) • But there are many other kinds • Organic compounds: • C6H12O6 (glucose); CH3COCH3 (acetone); C8H8O4 (aspirin); most drugs, actually Ppt13

4. Reminder—Acids (a subset of molecular compounds) • H is listed first in formula (for acids in this class [simple]) • Made up of molecules, BUT…. • …can be thought of as being made by taking any anion and adding H+’s until a neutral FU results: • H+’s & PO43- H3PO4 (phosphoric acid) • H+’s & S2- H2S (hydrosulfuric acid) KNOW YOUR (AN)IONS!! (PS3) • You need to be able to: • Determine the anion from which the acid is derived • H+ is often called a “proton” (since an H atom has 1 p and 1 e-; if you remove the e-, you get H+ [1 p only]) Ppt13

5. A solution is a homogeneous mixture • Solvent is often a liquid • Solute is “what’s dissolved in” the liquid • Solute can be solid, liquid, or gas at room temp. • Solute can be a molecular compound (non-acid or acid), or an ionic compound • If solvent is water, solution is called “aqueous” [indicated by (aq) after solute] • E.g., NaCl(aq) • Dissolution is what happens as a solute dissolves in a liquid Ppt13

6. Operational Definition, Dissolution You can assess “dissolution” visually. • Some substances are soluble in water, some insoluble • Dissolution: Formation of a homogeneous mixture (a “solution”) • one substance “dissolves” in another • A “solution” is something you can see through (it is clear, not cloudy) • Is milk a solution? • Remember the “waviness” in lab • If something “dissolves” in a liquid (to any appreciable extent), it is said to be soluble. Ppt13

7. “Theory” note:Warning: Dissolution is not the same thing as Dissociation!! You cannot assess “dissociation” visually. • Discussion of “dissociation” will come shortly. • For now just note that dissociation refers to a process that occurs after dissolution. • As such, in contrast to dissolution: Ppt13

8. Molarity (Worksheet)Try these now! PS5, problems19-21 reflect this material • If 3.12 g of KF is dissolved in some water to make the final volume equal to 50.0 mL, what is the molarity of this solution? • (using P.T. info: MM of KF = 58.10 g/mol) • How many moles of sugar (C6H12O6) are in 5.4 L of a 0.15 M solution? • How many liters of a 2.5 M solution of KBr will contain 3.6 moles of KBr? Ppt13

9. NOTE: Once molarity is understood, stoichiometry problems can utilize it! • “Another way to get to moles” (mol/L x L) • “Another way to go from moles” (mol  L = M) • PS6, Q2; Question (d) on Molarity Worksheet: Plan: mL HCl  L HCl  mol HCl  mol Na2O  g Na2O Plan: mL HCl  L HCl  mol HCl  mol Na2O  M Na2O Ppt13

10. NOTE: Once molarity is understood, stoichiometry problems can utilize it! • More example problems of this type will be shown and discussed in Ppt 16b (as a summary application of ideas—titration as an application of stoichiometry with molarity) Ppt13

11. Back to the qualitative discussion of solutions Ppt13

12. Some solutes form solutions that conduct electricity and some do not! Operational Definition of “Electrolyte” • Electrolyte: a substance that is soluble in water and produces a solution withincreased electrical conductivity • **Can test with a light bulb apparatus or conductivity meter** • By definition, electrolytes are a subset of all soluble substances. • All electrolytes are soluble, but not all soluble substances are electrolytes!! • Soluble substances that do not produce a solution with increased electrical conductivity are nonelectrolytes Ppt13

13. Some examples Ppt13

14. Conceptual Definition of Electrolyte(theory) • Electrical current is “moving charge” • Need to have “mobile charges” to increase electrical conductivity • Electrolyte: a substance for which at least some of its dissolved formula units end up turning into ions • Process is called dissociation(or “ionization”) • Ions in solution are free to move around, hence increasing electrical conductivity! Ppt13

15. What happensduring dissolution?(theory) • See next slide for pics that show the following: • Molecules will separate from one another, but generally remain intact (if substance dissolves) • Too small to scatter light, that’s why solution is “clear” • Ions will separate from one another once dissolved (if substance dissolves) • Too small to scatter light, that’s why solution is “clear” • Ionic compounds that dissolve will also undergo dissociation Ppt13

16. Dissolution only PS6, problem 3 reflects this material Dissolution & dissociation **Note that molecules remain intact during dissolution; they just separate from one another (without dissociating)! **Note that each individual cation (Na+) and anion (SO32-) remain intact during dissolution, but the ions separate from one another (i.e., the FUs dissociate)!

17. Dissolution is not the same as Dissociation (reprise) Recall: You cannot assess “dissociation” visually. Need “light bulb” or conductivity meter. Dissolution applies to any substance that dissolves, whether molecular or ionic Dissociation(in aqueous solution) applies only to electrolytes; it refers to the production of ions in solution, after FUs dissolve. Ppt13

18. Dissolution ≠ Disscociation (continued) • Most molecular substances’ FUs do not dissociate after they have dissolved • I.e., Most molecular substances are not electrolytes • But FUs of ionic compounds that dissolve DO also dissociate • I.e., Soluble ionic compounds are electrolytes Ppt13

19. Electrolytes come in two “types” • NOTE: “nonelectrolyte” refers to a soluble substance that is not an electrolyte • Thus an insoluble substance is technically neither a strong, weak, nor nonelectrolyte “strong”: ~100% of dissolved FUs are ionized “weak”: significantly less than 100% are ionized Ppt13

20. Analogous to Fig. 4.14 in Tro. Comparison of Strong, Weak, and Nonelectrolyte solutions (operationally and conceptually) All three substances are SOLUBLE, but they differ in the fraction of FUs that exist as separated ions: • 100% of FUs are ionized • only a small fraction of FUs are ionized • NONE are ionized Strong electrolyte Weak electrolyte Non- electrolyte Ppt13

21. Ppt13

22. The Six Common Strong Acids (memorize) • HCl • HBr • HI • HNO3 • H2SO4 • HClO4 • All others are weak acids! • HF • HNO2 • H3PO4 • H2CO3 • HClO • HC2H3O2 • Etc. PS5, problem 22 reflects this material Ppt13

23. Analogous to Figure 4.10 in Tro. Polar Water Molecules Interact with the Positive and Negative Ions of a Salt Assisting in Dissolution (and Dissociation!!) Ppt13

24. NaCl Dissolves (and Dissociates)(and 100% of FUs separate in solution) Strong electrolyte (soluble ionic compound) NaCl(s)  Na+(aq) + Cl-(aq) Ppt13

25. Analogous to image on p. 148 in Tro.HCl is Completely Ionized(strong electrolyte [one of the six strong acids]) Ppt13

26. Analogous to image on p. 148 in Tro. Acetic Acid (HC2H3O2)(weak acid [not one of the six]) Ppt13

27. An Aqueous Solution of Sodium Hydroxide(strong electrolyte [soluble ionic compound]) Ppt13

28. Figure 4.9 The Reaction of NH3 in Water(NH3 is a weak base—it produces a bit of OH- by reacting with H2O) Ppt13

29. Demonstration (if not already done)—Dissolution is not the same as Dissociation! • To see if a substance dissolves (e.g., in H2O): • Put a small amount of solid in a large test tube • Add water until about 1/3 full • Swirl for awhile: waviness indicates dissolution, and… • …if tube ends up CLEAR (not cloudy or opaque), then all of the substance dissolved, and the substance is said to be “soluble” • If a solution is clear, you cannot tell whether or not is contains an electrolyte just by looking at it! • You must test for electrical conductivity (light bulb lights) Ppt13

30. Connecting nanoscopic pictures to “molarity” • Molar concentration is like “number density” • Number of FU’s(regardless of mass) in a given amount of space • Molarity = moles of solute per liter of solution • 0.15 M Na2S means: • 0.15 moles of Na2S (FUs) per liter of solution, regardless of the fact that Na2S is a strong electrolyte • 1.5 M acetone means: • 1.5 moles of acetone (molecules) per liter of solution Ppt13

31. Molarity of Electrolytes (remember the meaning of subscripts!) Consider: Na2S(s)   2 Na+(aq) + S2-(aq) What is the concentration of Na+ ions in a 2.0 M solution of Na2S? What about the concentration of S2- ions in the same solution? Ppt13

32. Molarity of Electrolytes (remember the meaning of subscripts!) Consider: FeCl3(s)   Fe3+(aq) + 3 Cl-(aq) PS6, problems 4-5 reflect this material [Cl-] in 0.50 M FeCl3? Ppt13

33. Some ionic compounds are soluble, but some are not! • Those that are are strong electrolytes • Those that are not, are obviously not strong electrolytes (their formula units never even get into solution!) • Look for patterns: • NaCl, Na2S, NaClO4, Na3PO4, NaHCO3, Na2C2O4, and Na2Cr2O7, NaOH, Na2CO3 are all soluble • Tentative conclusion? Ppt13

34. Some ionic compounds are soluble, but some are not! • Look for patterns: • NaNO3, Ba(NO3)2 , Ru(NO3)2 , Fe(NO3)3, Pb(NO3)4, Ni(NO3)2 , CuNO3 , and Cu(NO3)2 are all soluble • Tentative conclusion? Ppt13

35. Common Solubility Rules(You need to memorize only 1 & 2; 3-6 will be given on Exam 2b) • All (common) nitrate salts (salts that contain NO3- as the anion) are soluble. • All (common) salts containing an alkali metal ion (Li+, Na+, K+, etc.) or ammonium ion (NH4+) as the cation are soluble. • Most (common) salts containing Cl-, Br-, or I- as the anion are soluble. However, important exceptions are those containing Ag+, Pb2+, or Hg22+. (i.e., AgBr and PbCl2 are insoluble.) • Most (common) salts containing sulfate (SO42-) as the anion are soluble. However, important exceptions are CaSO4, SrSO4, BaSO4, PbSO4, and Ag2SO4. • ------- • 5. Most (common) salts containing hydroxide(OH-) as the anion are insoluble. However, Ca(OH)2, Sr(OH)2 , and Ba(OH)2 are slightly soluble. • 6. Most (common) salts containing carbonate (CO32-), phosphate (PO43-), chromate (CrO42-), and sulfide (S2-), as the anion are insoluble. However, CaS,, SrS, and BaS are soluble and Rule #2 still follows (e.g., Na2CO3 is soluble) Ppt13

36. Ppt13

37. Table 4.1 Partial, “Reformat” PS6, problem 7 reflects this material (but used later as well) Ppt13