Plan Ppt 13 Substances in Aqueous Solution PS

Plan, Ppt 13: Substances in Aqueous Solution (PS 5 [end] and PS 6 [beginning] material) 1. Molarity of solutes I (reminder) 2. Reminder: Ionic vs. Molecular (non-acid) vs. Acid 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 Ppt 13 1

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 m. L, 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? Ppt 13 2

NOTE: Once molarity is understood, stoichiometry problems can utilize it! • A new way to get to moles! (mol/L x L) Exp 24, 30! • A new way to go from moles! (mol L = M) Exp 24! • Q 2(d) on Molarity Worksheet: Plan: m. L HCl mol Na 2 O g Na 2 O Plan: m. L HCl mol Na 2 O M Na 2 O Ppt 13 3

NOTE: Once molarity is understood, stoichiometry problems can utilize it! • More example problems of this type will be shown and discussed in Ppt 16 b (as a summary application of ideas—titration as an application of stoichiometry with molarity) Ppt 13 4

Reminder—Ionic Compounds • Metal or “NH 4” 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: Na 3 PO 4 - Made up of Na+ ions and PO 43 - ions - 3 Na+ ions / FU; 1 PO 43 - / FU KNOW YOUR IONS!! (PS 3) • NOTE: Na. OH, KOH, etc. are ionic compounds. Because of the OH- ion (not the cation!), these ionic compounds are also called “bases”. Ppt 13 5

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… Ø Non-acids (e. g. , sulfur trioxide, SO 3), and Ø (binary) acids (e. g. , hydrochloric acid, HCl) • …and one class of ternary molecular compounds Ø (ternary) acids (e. g. , chloric acid, HCl. O 3) • But there are many other kinds Ø 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 Ppt 13 6

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 & PO 43 - H 3 PO 4 (phosphoric acid) - H+’s & S 2 - H 2 S (hydrosulfuric acid) KNOW YOUR (AN)IONS!! (PS 3) • 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]) Ppt 13 7

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. , Na. Cl(aq) • Dissolution is what happens as a solute dissolves in a liquid Ppt 13 8

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) You can assess • Is milk a solution? “dissolution” • Remember the “waviness” in lab visually. – If something “dissolves” in a liquid (to any appreciable extent), it is said to be soluble. • Some substances are soluble in water, some insoluble Ppt 13 9

“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. Ppt 13 10

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 Ppt 13 11

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) Ppt 13 12

Some examples Electrolytes All soluble ionic compounds e. g. , Na. Cl, Na. NO 3, Fe. Cl 3, Ni. SO 4, Ca(OH)2, etc. All Acids and Molecular Bases: e. g. , HCl, HNO 2, HCl. O 4, NH 3, etc. Nonelectrolytes 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) Ppt 13 13

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! Ppt 13 14

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 Ppt 13 15

See Ph. ET: sugar nd salt; salts nd solubility Dissolution only 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 (SO 32 -) remain intact during dissolution, but the ions separate from one another

Reference… • See Ph. ET simulations on solutions and solubility, and salt vs sugar (and maybe solutions and concentration) Ppt 13 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. Ppt 13 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 Ppt 13 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 Ppt 13 20

Analogous to Fig. 5. 12 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: Strong electrolyte Weak electrolyte Nonelectrolyte (a) 100% of FUs are ionized (b) only a small fraction of FUs are ionized (c) NONE are ionized Ppt 13 21

The Six Common Strong Acids (memorize) • HCl • HBr • HI • HNO 3 • H 2 SO 4 • HCl. O 4 All others are weak acids! • HF • HNO 2 • H 3 PO 4 • H 2 CO 3 • HCl. O • HC 2 H 3 O 2 • Etc. Ppt 13 23

Analogous to Figure 5. 8 in Tro. Polar Water Molecules Interact with the Positive and Negative Ions of a Salt Assisting in Dissolution (and Dissociation!!) See Ph. ET: sugar and salt; salts and solubility Ppt 13 24

Na. Cl Dissolves (and Dissociates) (and 100% of FUs separate in solution) Strong electrolyte Na. Cl(s) Na+(aq) + Cl-(aq) Ppt 13 (soluble ionic compound) 25

Analogous to image on p. 177 in Tro. HCl is Completely Ionized (Dissociated) strong electrolyte [one of the six strong acids] Ppt 13 26

Analogous to image on p. 177 in Tro. Acetic Acid (HC 2 H 3 O 2) (weak acid [not one of the six]) Ppt 13 27

An Aqueous Solution of Sodium Hydroxide (strong electrolyte [soluble ionic compound]) Ppt 13 28

Figure 4. 9 (Zumdahl) A solution of NH 3 in Water NOTE: This weak electrolyte does not technically “dissociate” to make ions! It produces ions by reacting (a bit) with water molecules: NH 3(aq) + H 2 O(aq) NH 4+(aq) + OH-(aq) Ppt 13 29

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 Ppt 13 30

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 32 -(aq) • & the concentration of SO 32 - ions in the same solution? Ppt 13 31

Molarity of Electrolytes (remember the meaning of subscripts!) • [Cl-] in 0. 50 M Fe. Cl 3? Consider: Fe. Cl 3(s) Fe 3+(aq) + 3 Cl-(aq) Ppt 13 32

Group Problem (application of demo ideas) • Which solution below will have the greater conductivity (brightest light bulb)? Justify your answer with calculations (e. g, [total ions] = [cation] + [anion]): 1) A 0. 0010 M solution of HCl(aq) , a strong acid, OR 2) A 0. 10 M solution of HNO 2(aq) , a weak acid. Assume that only 7. 0% of the dissolved molecules are dissociated. Answer: 0. 001 M HCl ⇒ 0. 001 M H+ + 0. 001 M Cl- = 0. 002 M total ions 0. 07 x 0. 10 M HNO 2 = 0. 007 M dissociated HNO 2 ⇒ 0. 007 M H+ + 0. 007 M NO 2 - = 0. 014 M total ions Ppt 13 33

Some ionic compounds are soluble, but some are not! • Those that are strong electrolytes • Those that are not, are obviously not strong electrolytes (their formula units never even get into solution!) • Look for patterns: – Na. Cl, Na 2 S, Na. Cl. O 4, Na 3 PO 4, Na. HCO 3, Na 2 C 2 O 4, and Na 2 Cr 2 O 7, Na. OH, Na 2 CO 3 are all soluble • Tentative conclusion? Ppt 13 34

Some ionic compounds are soluble, but some are not! • Look for patterns: – Na. NO 3, Ba(NO 3)2 , Ru(NO 3)2 , Fe(NO 3)3, Pb(NO 3)4, Ni(NO 3)2 , Cu. NO 3 , and Cu(NO 3)2 are all soluble • Tentative conclusion? Ppt 13 35

Common Solubility Rules (You need to memorize only 1 & 2; 3 -6 will be given on Exam 2 b) 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 22+. (i. e. , Ag. Br and Pb. Cl 2 are insoluble. ) 4. Most (common) salts containing sulfate (SO 42 -) as the anion are soluble. However, important exceptions are Ca. SO 4, Sr. SO 4, Ba. SO 4, Pb. SO 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 32 -), phosphate (PO 43 -), chromate (Cr. O 42 -), and sulfide (S 2 -), as the anion are insoluble. However, Ca. S, , Sr. S , and Ba. S are soluble and Rule #2 still follows (e. g. , Na 2 CO 3 is soluble) Ppt 13 36

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Table 4. 1 Partial, “Reformat” Ppt 13 38