1. Unit 4 (Chapter 4): Aqueous Reactions & Solution Stoichiometry John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten

2. Solutions: homogeneous mixtures. ( _____ throughout) solvent is present in greatest abundance. solute dissolved in/by solvent + same

3. 0.250 L Molarity Molarity (M) is a measure of the concentration of a solution. moles of solute (mol) liters of solution (L) Molarity (M) = units:mol/L or mol ∙ L –1 What is the molarity of a solution with 29.2 g of sodium chloride in 250. mL of water? 29.2 g NaCl x 1 mol NaCl 58.44 g NaCl = 0.500 mol NaCl = 2.00 M NaCl

4. Preparing a Solution 1-mass solute 2-add solvent, swirl to dissolve 3-add solvent to mark WS #1-2 Conc. Calc’s HW p.160 #60, 67

5. WS Concentration & Dilutions 0.100 mol NaHCO 3 5.00 g NaHCO 3 x 1 mol NaHCO 3 84.01 g NaHCO 3 x 1 L NaHCO 3 = 0.595 L NaHCO 3 #1 1 mol CuSO 4 0.275 L CuSO 4 x 1.20 mol CuSO 4 1 L CuSO 4 x 159.62 g CuSO 4 = 52.7 g CuSO 4 #2

6. Dilution M 1 V 1 = M 2 V 2 1-calc M 1 V 1 =M 2 V 2 2-pipet V 1 from concentrated 3-fill to mark w DI

7. Dilution M 1 V 1 = M 2 V 2 What volume of water must be added to prepare 2.0 L of 3.0 M CuSO 4 from a stock solution of concentration 8.0 M ? WS #3-4 Dilutions

8. WS Concentration & Dilutions #3 #4 M 1 V 1 = M 2 V 2 (12.0 M)V 1 = (1.25 M)(500 mL) V 1 = 52.1 mL (0.0521 L) M 1 V 1 = M 2 V 2 (2.50 M)V 1 = (0.200 M)(250 mL) V 1 = 20.0 mL (0.0200 L)

9. HCl + H 2 O  H 3 O + + Cl – H H H H H H O O Cl + – NH 3 + H 2 O  NH 4 + + OH – H H H H H H H H H H N O O N + –  Acid: proton (H + ) donor  Base: proton (H + ) acceptor

10. Strength of Acids and Bases STRONG (complete dissociation) HA (aq) + H 2 O (l) H 3 O + (aq) + A – (aq) B (aq) + H 2 O (l) BH + (aq) + OH – (aq) WEAK (partial dissociation) HA (aq) + H 2 O (l) H 3 O + (aq) + A – (aq) B (aq) + H 2 O (l) BH + (aq) + OH – (aq)

11. Strong Acids: Only 6 strong acids: Nitric (HNO 3 ) Sulfuric (H 2 SO 4 ) Hydrochloric (HCl) Hydrobromic (HBr) Hydroiodic (HI) Perchloric (HClO 4 ) proton (H + ) donors HI + H 2 O  H 3 O + + I –

12. Strong Bases: The strong bases are soluble hydroxides (OH – ) of… Group 1 (Li,Na,K) CBS (Ca, Ba, Sr) Mg(OH) 2 not as soluble proton (H + ) acceptors OH – + H 3 O +  H 2 O + H 2 O ase

13. Salts: Ionic Solids: (metal-nonmetal) dissociate (dissolve) by separation into ions Electrolytes: ions in solution that conduct electricity

14. StrongWeak C 11 H 22 O 11 CH 3 OH H 2 O Non CH 3 COOH HNO 2 NH 3 NaOH HNO 3 KCl completely dissociate partially ionize only molecules NO ions ALL ions SOME ions HW p.159 #33

15. Electrolytes: Strong, Weak, or Non? Compound Ionic STRONG Molecular Acid (H____) STRONG (6) WEAK (& NH 3 ) Not Acid NON HW p.157-159 #1,2,4,5,38 nonmetals (Covalent)metal-nonmetal C 11 H 22 O 11 C 2 H 5 OH H 2 O CH 3 COOH HNO 2 HF KBr CaI 2 FeCl 3 NaOH Ca(OH) 2 (strong bases) HCl, HBr, HI HNO 3 H 2 SO 4 HClO 4 (ions conduct electricity)

16. Electrolytes: Strong, Weak, or Non? Compound Ionic STRONG Molecular Acid (H____) WEAK (& NH 3 ) Not Acid NON nonmetals (Covalent)metal-nonmetal STRONG (6) QUIZ!!! (at the bell)

17. Cl Na H O strong acid (H + A – ) strong base (M + OH – ) HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O (l) ACID + BASE SALT + WATER HH H O Cl +– ionic compound (M + A – ) water H 2 O (HOH) Acid-Base Neutralization Reactions HW p.158 #40a

18. Double Replacement:(precipitate) precipitate: insoluble ionic compound (as predicted by solubility rules) Precipitation Reactions Pb 2+ I–I–

19. ALWAYS Soluble Ions: Na +, K +, etc. group I (alkali metals) NH 4 + ammonium NO 3 – nitrate HCO 3 – bicarbonate C 2 H 3 O 2 – (CH 3 COO – )acetate (ethanoate) ClO 3 –, ClO 4 – chlorate, perchlorate Solubility Rules Common Precipitates form with:examples Ag +, Pb 2+, Hg 2+ (AP/H)AgCl, PbI 2 OH – (hydroxide)Cu(OH) 2 CO 3 2 – (carbonate)CaCO 3 ****** WS Solubility & NIE’s #1

20. Molecular Equation The molecular equation lists the reactants and products in their molecular form. AgNO 3 (aq) + KCl (aq)  AgCl (s) + KNO 3 (aq)

21. Ionic Equation In the ionic equation all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions. This more accurately reflects the species that are found in the reaction mixture. Ag + (aq) + NO 3 – (aq) + K + (aq) + Cl – (aq)  AgCl (s) + K + (aq) + NO 3 – (aq) AgNO 3(aq) + KCl (aq)  AgCl (s) + KNO 3(aq)

22. Net Ionic Equation Ag + (aq) + NO 3 – (aq) + K + (aq) + Cl – (aq)  AgCl (s) + K + (aq) + NO 3 – (aq) NIE:Ag + (aq) + Cl – (aq)  AgCl (s) Cross out Spectator Ions that do not change (same state & same charge) from the left side of the equation to the right. The only species left are those things that react (change) during the course of the reaction.

23. Balanced Net Ionic Equations 1. Write a complete molecular equation. 2. Dissociate all strong electrolytes (aq). 3. Cross out spectators (same charge & state) 4. Write the net ionic equation with the species that remain and balance it. comp – diss – cross – net – bal (solubility rules)

24. BaSO 4 + NH 4 NO 3 Balanced Net Ionic Equations 1)(NH 4 ) 2 SO 4 + Ba(NO 3 ) 2 → 2)NaOH + MgBr 2 → comp – diss – cross – net – bal +2–2+––+ Ba 2+ + SO 4 2– → BaSO 4 NaBr + Mg(OH) 2 +–2+––+ Mg 2+ + 2 OH – → Mg(OH) 2 (s) HW p.158 #21

25. HCl (aq) + NaOH (aq)  NaCl (aq) + H 2 O (l) H + + Cl – + Na + + OH –  Na + + Cl – + H 2 O Neutralization Reactions H + + OH –  H 2 O When a Strong Acid reacts with a Strong Base, the net ionic equation is…

26. HF (aq) + KOH (aq)  KF (aq) + H 2 O (l) HF + Na + + OH –  Na + + F – + H 2 O Neutralization Reactions HX + OH –  X – + H 2 O When a Weak reacts with a Strong, the net ionic equation is… HW p.159 #40 (finish)

27. BaSO 4 + NH 4 NO 3 Balanced Net Ionic Equations (NH 4 ) 2 SO 4 + Ba(NO 3 ) 2 → comp – diss – cross – net – bal +2–2+––+ Ba 2+ + SO 4 2– → BaSO 4 (s) HF (aq) + KOH (aq)  KF (aq) + H 2 O (l) WS Solubility & NIE’s #2 HF + OH –  F – + H 2 O –+–+

28. Gas-Forming Reactions Single Rep: Metal + Acid Metal Ion + H 2 Ex: Zn (s) + H 2 SO 4 (aq) ZnSO 4 (aq) + H 2 (g) NIE: Zn (s) + 2 H + (aq) Zn 2+ (aq) + H 2 (g) H 2 Demo (M 0 )(H + )(M + )(gas) +2– 2+2– Double Rep: Acid + Carbonate Salt + Ex: HCl (aq) + CaCO 3 (s) CaCl 2 (aq) + H 2 O (l) + CO 2 (g) NIE: 2 H + (aq) + CaCO 3 (s) Ca 2+ (aq) + H 2 O (l) + CO 2 (g) (or Bicarbonate) (HCO 3 – ) (CO 3 2– ) (H + ) CH 3 COOH + NaHCO 3  CH 3 COONa +H 2 O + CO 2 HW p. 159 #43 H 2 CO 3 (aq) H 2 O (l) + CO 2 (g) (decomposes immediately) (gas) CO 2 Demo

29. g A L of A g B mol BL of B g A 1 mol A mol A 1 L g B 1 mol B mol B 1 L molar mass A molar mass B molarity A (M) molarity B (M) mol-to-mol ratio mol A Rxn: A (aq) + 2 B (aq)  C + 2 D Solution Stoichiometry HW p. 161 #81

30. Oxidation-Reduction Reactions (REDOX) video clip (One cannot occur without the other) LEO says GER

31. Oxidation Numbers Is it a redox reaction? To find out… 1)assign oxidation numbers* (or oxidation states) to each element in a reaction. 2)check if any oxidation states changed ( ↓ reduced, ↑ oxidized) of elements describe electrons that would be lost or gained IF the compound was 100% ionic. of ions show electrons transferred IN an ionic compound *oxidation numbers *charges

32.  F is always −1.  other halogens are −1, but can be positive, like in oxyanions. Ex. ClO 3 – or NO 3 – or SO 4 2– Assigning Oxidation Numbers 1.All elements are 0. (all compounds are 0) 2.Monatomic ion is its charge. (Ex. Na + ion) 3.Most nonmetals tend to be negative, but some are positive in certain compounds or ions. (Ex. SO 3 )  O is −2 always, but in peroxide ion is −1 (O 2 –2 ).  H is +1 with nonmetals, −1 with a metals.

33. Oxidation Numbers The sum of the ox. #’s in a neutral compound is 0. The sum of the ox. #’s in a polyatomic ion is the charge on the ion. Calculate the oxidation number of each: Sulfur in… SO 3 Chromium in… K 2 Cr 2 O 7 Nitrogen in… NH 4 + Cobalt in… [CoCl 6 ] 3–

34. Classifying REDOX Reactions All rxns (but…NOT double replacement) Decomposition AB → A + B1 → 2 (+/– → 0 0) Combustion C x H y + O 2 → CO 2 + H 2 O (–/+ 0 → +/– +/–) Single Replacement AB + C → A + CB (+/– 0 → 0 +/–) Synthesis A + B → AB2 → 1 (0 0 → +/–)

36. Single Replacement (REDOX) Cu 2+ (aq) + 2 Ag (s)  Cu (s) + 2 Ag + (aq) X Cu (s) + 2 Ag + (aq)  Cu 2+ (aq) + 2 Ag (s) silver ions oxidize copper metal

37. Activity Series of Metals increasing ease of oxidation Cannot displace H + from acid to make H 2 (g)

38. Write an overall equation for the synthesis of calcium sulfide from its elements. Then write the RED and OX half- equations to identify the REDOX process. Writing REDOX Reactions Ca + S  CaS Ca 0  Ca +2 0 +2–20 OX: RED: S 0  S –2 + 2 e – 2 e – + Ca + S  CaS

39. 4 Al +3  4 Al 0 12 e – + O –2  3 O 2 0 + 12 e – 4 ( ) 3 ( ) Writing REDOX Reactions Write an overall equation for the decomposition of aluminum oxide into its elements. Then write the RED and OX half- equations to identify the REDOX process. Al 2 O 3  Al + O 2 O –2  O 2 0 0 +3–2 0 OX: RED: Al +3  Al 0 + 4 e – 3 e – + 2 Al 2 O 3  4 Al + 3 O 2 2 6

40. Write the net ionic equation for the reaction of solid zinc in a solution of hydrochloric acid. Writing REDOX Reactions Mg (s) + HCl (aq)  MgCl 2 (aq) + H 2 (g) comp – diss – cross – net – bal 0+1–1+2–10 Mg + H +  Mg +2 + H 2 2 Classify the reaction in two ways. Single-Replacement and Redox

41. Mg + 2 H +  Mg 2+ + H 2 (g) WS 5c #1-2 What is red & what is ox? red ox