1. Plan, Ppt 13: Substances in Aqueous Solution (PS5, 19-22 and PS6, 1-7 material) 1.Reminder: Ionic vs. Molecular (non-acid) vs. Acid 2.Molarity of solutes I 3.Qualitative Issues (types of solutes/solutions) Electrolytes (Operational) Electrolytes (Conceptual) Strong vs. Weak vs Non Electrolyes 4.Molarity Reprise (electrolyte issues) 5.Solubility rules for ionic compounds 1 Ppt13

2. Reminder—Ionic Compounds Metal or “NH 4 ” listed first in formula -Write the formulas of the cation and anion -State the number of each ion present in one FU of the compound Made up of ions (cation and anion type) You need to be able to: Example: Na 3 PO 4 -Made up of Na + ions and PO 4 3- ions -3 Na + ions / FU; 1 PO 4 3- / FU NOTE: NaOH, KOH, etc. are ionic compounds. Because of the OH - ion (not the cation!), these ionic compounds are also called “bases”. KNOW YOUR IONS!! (PS3) 2 Ppt13

3. Reminder—Molecular Compounds In this class, a compound that is not ionic! NO metal first; a nonmetal (unless “NH 4 ”) Learned to name binary molecular compounds…  Non-acids (e.g., sulfur trioxide, SO 3 ), and  (binary) acids (e.g., hydrochloric acid, HCl) …and one class of ternary molecular compounds  (ternary) acids (e.g., chloric acid, HClO 3 ) But there are many other kinds  Organic compounds: o C 6 H 12 O 6 (glucose); CH 3 COCH 3 (acetone); C 8 H 8 O 4 (aspirin); most drugs, actually 3 Ppt13

4. Reminder—Acids (a subset of molecular compounds) H is listed first in formula (for acids in this class [simple]) -H + ’s & PO 4 3-  H 3 PO 4 (phosphoric acid) -H + ’s & S 2-  H 2 S (hydrosulfuric acid) 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: KNOW YOUR (AN)IONS!! (PS3) -Determine the anion from which the acid is derived You need to be able to: 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]) 4 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 5 Ppt13

6. Operational Definition, Dissolution 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. Some substances are soluble in water, some insoluble You can assess “dissolution” visually. 6 Ppt13

7. “Theory” note: Warning: Dissolution is not the same thing as Dissociation!! 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: You cannot assess “dissociation” visually. 7 Ppt13

8. Molarity (Worksheet) Try these now! a)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) b)How many moles of sugar (C 6 H 12 O 6 ) are in 5.4 L of a 0.15 M solution? c)How many liters of a 2.5 M solution of KBr will contain 3.6 moles of KBr? PS5, problems19-21 reflect this material 8 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 Na 2 O  g Na 2 O Plan: mL HCl  L HCl  mol HCl  mol Na 2 O  M Na 2 O 9 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) 10 Ppt13

11. Back to the qualitative discussion of solutions 11 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 with increased 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 12 Ppt13

13. ElectrolytesNonelectrolytes All soluble ionic compounds e.g., NaCl, NaNO 3, FeCl 3, NiSO 4, Ca(OH) 2, etc. Most soluble molecular substances that are not acids or bases CH 3 COCH 3 (acetone) CH 3 OH (methanol) O 2 C 12 H 22 O 11 (sucrose) All Acids and Molecular Bases: e.g., HCl, HNO 2, HClO 4, NH 3, etc. Some examples 13 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! 14 Ppt13

15. What happens during 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 15 Ppt13

16. **Note that each individual cation (Na + ) and anion (SO 3 2- ) remain intact during dissolution, but the ions separate from one another (i.e., the FUs dissociate)! **Note that molecules remain intact during dissolution; they just separate from one another (without dissociating)! Dissolution only Dissolution & dissociation PS6, problem 3 reflects this material 16

17. Dissolution is not the same as Dissociation (reprise) 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. Recall: You cannot assess “dissociation” visually. Need “light bulb” or conductivity meter. 17 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 18 Ppt13

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

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

21. 21 Ppt13

22. The Six Common Strong Acids (memorize) HCl HBr HI HNO 3 H 2 SO 4 HClO 4 All others are weak acids! HF HNO 2 H3PO4 H 2 CO 3 HClO HC 2 H 3 O 2 Etc. PS5, problem 22 reflects this material 22 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!!) 23 Ppt13

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

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

26. Analogous to image on p. 148 in Tro. Acetic Acid (HC 2 H 3 O 2 ) (weak acid [not one of the six]) 26 Ppt13

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

28. Figure 4.9 (Zumdahl) A solution of NH 3 in Water 28 Ppt13 NH 3 (aq) + H 2 O (aq) NH 4 + (aq) + OH - (aq) NOTE: This weak electrolyte does not technically “dissociate” to make ions! It produces ions by reacting (a bit) with water molecules:

29. Demonstration (if not already done)— Dissolution is not the same as Dissociation! To see if a substance dissolves (e.g., in H 2 O): – 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) 29 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 Na 2 SO 3 means: 0.15 moles of Na 2 SO 3 (FUs) per liter of solution, regardless of the fact that Na 2 SO 3 is a strong electrolyte – 1.5 M acetone means: 1.5 moles of acetone (molecules) per liter of solution 30 Ppt13

31. Molarity of Electrolytes (remember the meaning of subscripts!) What is the concentration of Na + ions in a 2.0 M solution of Na 2 SO 3 ? Consider: Na 2 SO 3 (s)  2 Na + (aq) + SO 3 2- (aq) 31 Ppt13 & the concentration of SO 3 2- ions in the same solution?

32. Molarity of Electrolytes (remember the meaning of subscripts!) [Cl - ] in 0.50 M FeCl 3 ? Consider: FeCl 3 (s)   Fe 3+ (aq) + 3 Cl - (aq) PS6, problems 4-5 reflect this material 32 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, Na 2 S, NaClO 4, Na 3 PO 4, NaHCO 3, Na 2 C 2 O 4, and Na 2 Cr 2 O 7, NaOH, Na 2 CO 3 are all soluble Tentative conclusion? 33 Ppt13

34. Some ionic compounds are soluble, but some are not! Look for patterns: – NaNO 3, Ba(NO 3 ) 2, Ru(NO 3 ) 2, Fe(NO 3 ) 3, Pb(NO 3 ) 4, Ni(NO 3 ) 2, CuNO 3, and Cu(NO 3 ) 2 are all soluble Tentative conclusion? 34 Ppt13

35. Common Solubility Rules (You need to memorize only 1 & 2; 3-6 will be given on Exam 2b) 1.All (common) nitrate salts (salts that contain NO 3 - as the anion) are soluble. 2.All (common) salts containing an alkali metal ion (Li +, Na +, K +, etc.) or ammonium ion (NH 4 + ) as the cation are soluble. 3.Most (common) salts containing Cl -, Br -, or I - as the anion are soluble. However, important exceptions are those containing Ag +, Pb 2+, or Hg 2 2+. (i.e., AgBr and PbCl 2 are insoluble.) 4.Most (common) salts containing sulfate (SO 4 2- ) as the anion are soluble. However, important exceptions are CaSO 4, SrSO 4, BaSO 4, PbSO 4, and Ag 2 SO 4. ------- 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 (CO 3 2- ), phosphate (PO 4 3- ), chromate (CrO 4 2- ), and sulfide (S 2- ), as the anion are insoluble. However, CaS,, SrS, and BaS are soluble and Rule #2 still follows (e.g., Na 2 CO 3 is soluble) 35 Ppt13

36. 36 Ppt13

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