Reactions in Aqueous Solution

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1 Reactions in Aqueous Solution AP Chemistry Chapter 4

2 Reactions in Aqueous Solution Fundamental factor that causes reaction involving ions to occur is the starting ions are removed from solution. Ways this is accomplished: 1) Formation of a precipitate 2) Formation of a gas 3) Formation of a primarily molecular species 4) Form original ions into different ions. (oxidation - reduction)

3 1) Precipitation Reactions Use solubility Rules ---> (table 4.1 pg. 144) (desk tops or handout) 1. Most nitrates (NO 3- ) are soluble. 2. Most salts containing the alkali metal ions (Li +, Na +, K +, Cs +, Rb + ) and the ammonium ion (NH 4+ ) are soluble. 3. Most chloride, bromide, and iodide salts are soluble. Exceptions: Ag +, Pb 2+, and Hg Most sulfate salts are soluble. Exceptions: BaSO 4, PbSO 4, HgSO 4, and CaSO Most hydroxide are only slightly soluble. Exceptions: Group I & Calcium down in Group II. 6. Most sulfide (S 2- ), carbonate (CO 2-3 ), chromate (CrO 2-4 ), and phosphate (PO 3-4 ) salts are only slightly soluble. S 2- exceptions: Group I, Group II, NH + 4 CO 2-3, CrO 2-4, PO 3-4 exceptions: Group I, NH + 4

4 Solubility Rules Put these in a special spot. I recommend starting a small (1/2 inch) 3-ring binder for only predicting reactions papers. MEMORIZE! We ll try to quiz on these rules daily this unit.

5 1) Precipitation Reactions 1st example - Mix solutions of Ba(NO 3 ) 2 and Na 2 CO 3. What happens?

6 1) Precipitation Reactions 2nd example - Mix solutions of CuCl 2 and NaOH. What happens?

7 1) Precipitation Reactions 3rd example - Mix solutions of MgCl 2 and KNO 3. What happens?

8 Precipitate Colors PbI 2 HgCl 2 HgS AgCl (white)

9 Precipitate Colors Cu 2+ (aq) + 2NH 3(aq) + 3 H 2 O (l) <=> Cu(OH) 2(s) + 2NH 4 + (aq) Cu(OH) 2 (s) + 4 NH 3 (aq) <=> [Cu(NH 3 ) 4 ] 2+ (aq) + 2OH- (aq) Cu 2+ (aq) + 2OH- (aq) <=> Cu(OH) 2 (s)

10 2) Gas formation Reactions (H 2 S, CO 2, SO 2, NH 3 ) H 2 CO 3 H 2 O + CO 2 H 2 O + SO 2 H 2 O + NH 3 H 2 S (g)

11 2) Gas formation Reactions (H 2 S, CO 2, SO 2, NH 3 ) 4th example - Mix solutions of Na 2 CO 3 and HCl. What happens?

12 3) Acid-Base Reactions Acids form H + ions in solution Difference between strong and weak acids: * Must know 6 strong acids - any other acid is weak

13 3) Acid-Base Reactions Strong Acid Reaction: Weak Acid Reaction: Ex M HF: [HF] = 0.97 M, [H+] = [F-] = 0.03 M 3% ionized

14 3) Acid-Base Reactions EXAMPLES - (Species present at highest concentration is written as the reactant.) A) Strong Acid - Strong Base HCl + Ca(OH) 2 Net ionic equation:

15 3) Acid-Base Reactions EXAMPLES - (Species present at highest concentration is written as the reactant.) B) Strong Acid - Weak Base HCl + NH 3 Net ionic equation:

16 3) Acid-Base Reactions EXAMPLES - (Species present at highest concentration is written as the reactant.) C) Weak Acid - Strong Base HF + NaOH Net ionic equation

17 REVIEW OF ELECTROLYTES & NET IONIC REACTIONS Strong Electrolytes exist 100% as ions in solution Weak Electrolytes exists partially as ions in solution Non-electrolytes do NOT exist as ions

18 REVIEW OF ELECTROLYTES

19 Problem Set 4.1 Summary Problem Set 4.1 is all about writing net ionic reactions. When writing an ionic equation write everything that exists as ions as ions. Write as Ions Use Solubility Rules on pg. 141 (or rules below Aqueous (soluble) ionic compound the overhead) to determine solubility. Strong acids Write as molecules Insoluble ionic compounds (or any ionic cpd not in solution) Weak acids Molecular compounds

20 Problem Set 4.1 Summary See back page of your note packet! NOTE: The strong bases are also the soluble hydroxides. Ex. A solution of sodium hydroxide is mixed with hydrochloric acid Total ionic: Na + (aq) + OH - (aq) + H + (aq) + Cl - (aq) H 2 O (l) + Na + (aq) + Cl - (aq) Net ionic: OH - (aq) + H + (aq) H 2 O (l) Ex. Magnesium hydroxide is mixed with hydrochloric acid Total ionic: Mg(OH) 2(s) + 2 H + (aq) + 2 Cl - (aq) 2 H 2 O (l) + Mg 2+ (aq) + 2 Cl - (aq) Net ionic: Mg(OH) 2(s) + 2 H + (aq) 2 H 2 O (l) + Mg 2+ (aq)

21 Problem Set 4.1 Summary NOTE: THE CALCULATIONS ARE SIMPLY 3-STEP MOLE PROBLEMS, BUT THEY ARE MUCH HARDER THAN MOST YOU VE DONE BEFORE. (YOU LL NEED TO WRITE BALANCED EQUATIONS FOR MOST OF THEM.)

22 Stoichiometry Flow Chart Go back to pg. 3 of note packet 1. Identify Reaction Type 2. Write a balanced Equation 3. Calculate moles of reactants Given V & M: n = M V Mass: n = g/mm 4. Determine Limiting Reagent 5. Calculate moles of products

23 Quantitative Analysis Quantitative analysis the determination of how much of a given component is present in a sample Titration a process in which one reagent is added to another (usually acids and bases) with which it reacts. An indicator is used to determine the which equivalent quantities of the two reagents have been added. Equivalence point which just enough titrant has been added to fully react with B. end point the point during a titration where an indicator turns color ex. Phenolphthalein

24 Quantitative Analysis Ex ml of fruit juice is titrated with ml of an NaOH solution which contains 1.82 mg/ml. Determine the mg of vitamin C, C 6 H 8 O 6, per milliliter of juice.

25 Solution Stoichiometry Examples: M H+ V H+ = M OH- V OH- 1. When aqueous solutions of sodium hydroxide and Iron (III) nitrate are mixed, a red gelatinous precipitate forms. Calculate the mass of precipitate formed when ml of M NaOH and ml of M Fe(NO 3 ) 3 are mixed.

26 Solution Stoichiometry Examples: 2. Barium Hydroxide reacts with acetic acid, HC 2 H 3 O 2. Calculate the concentration of all ions and molecules except H 2 O after the reaction if ml of M Ba(OH) 2 is added to ml of M HC 2 H 3 O 2. (Assume the final is the sum of the initial volumes.)

27 Solution Stoichiometry Examples: 3. What volume of M AlCl 3 is needed to react fully with 25.0 ml of M Pb(NO 3 ) 2?

28 Solution Stoichiometry Examples: ml of acetic acid, CH 3 COOH, is titrated with ml of LiOH. The LiOH solution is 1.42 mg/ml. Determine the mg of the acid per ml of solution.

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