Chapter 4: Chemical Reactions & Solution Stoichiometry

1. Chapter 4: Chemical Reactions & Solution Stoichiometry OWL Deadline: 15-October-12 (7:50 AM) Exam #2 on 15-Oct with Dr. Vining. Types of Reactions Solubility Rules Aqueous Reactions and Stoichiometry Oxidation-Reduction Reactions

2. Adding lead(II) nitrate to potassium iodide gives a yellow precipitate of lead(II) iodide. Precipitation reactions are important in biological organisms— for example, in the production of bone (calcium phosphate) and seashell (calcium carbonate).

3. Combination Reaction A reaction in which two or more substances combine to form a third substance. For example: 2Na(s) + Cl2(g)  2NaCl(s)

4. Decomposition Reaction A reaction in which a single compound reacts to give two or more substances. For example: 2HgO(s)  2Hg(l) + O2(g)

5. Let’s look at some reactions: Sodium metal reacts with oxygen gas to produce sodium oxide. -What kind of reaction ? -Write the reaction (equation) -Balance the equation Aluminum Carbonate is heated and forms aluminum oxide and carbon dioxide gas. -What kind of reaction ? -Write the reaction (equation) -Balance the equation

6. Displacement Reaction A reaction in which an element reacts with a compound, displacing another element from it. For example: Zn(s) + 2HCl(aq)  H2(g) + ZnCl2(aq)

7. Acid-Base Reaction (usually a type of Double Displacement Reaction that happens in solution) An acid that reacts with a base will usually for a salt and water. For example: CH3COOH (aq) + NaOH (aq)CH3COONa (aq) + HOH (l)

8. Precipitation Reaction (usually a type of Double Displacement Reaction in Solution) A reaction in which an or ion reacts with a compound or ion, displacing another element from it. A solid precipitate forms from the reactants in solution. For example: MgCl2 (aq) + 2 AgNO3 (aq)  Mg(NO3)2 (aq) + 2 AgCl(s)

9. Gas Forming Reaction (usually a type of Double Displacement Reaction in Solution) A reaction in which an or ion reacts with a compound or ion, displacing another element from it. A solid precipitate forms from the reactants in solution. For example: Na2CO3(aq) + 2HCl(aq)  2NaCl(aq) + H2O(l) + CO2(g)

10. Combustion Reaction A reaction in which a substance reacts with oxygen, usually with the rapid release of heat to produce a flame. For example: 4Fe(s) + 3O2(g)  2Fe2O3(s)

11. Let’s look at some reactions: Calcium carbonate reacts with hydrochloric acid to produce water, carbon dioxide and calcium chloride . -What kind of reaction ? -Write the reaction (equation) -Balance the equation A solution of barium chloride reacts with a solution of sodium sulfate. Barium sulfate precipitates. Sodium chloride remains in solution. -What kind of reaction ? -Write the reaction (equation) -Balance the equation

12. What happens in aqueous solution? • Molecules and ions become solvated. • The solvent is the larger component, usually a liquid • The solute is the minor component, usually starting as a solid or a liquid (but it can be a gas) • When water is the solvent, then the solute becomes hydrated

13. Nonelectrolytes can dissolve in water, but do not conduct charge • They don’t form ions • CH3OH (l)  CH3OH (aq) methanol dissolves in water • Where is this reaction important for your car? • Weak electrolytes dissolve, but only partially dissociate into their component ions. The solutions are weak conductors • CH3COOH (l) + H2O (l)  CH3COO- (aq) + H3O+ (aq) • Strong electrolytes dissolve and completely dissociate into their component ions. The solutions are good conductors • NaCl (s)  Na+ (aq) + Cl- (aq)

14. Non-, Strong- or Weak-Electrolytes?

16. Solubility of Ionic Compounds • Soluble: to dissolve. A solute that is soluble, dissolves in the solvent. • Say that 10 times really fast.  • There are specific rules for predicting the solubility of ionic compounds. • When a solid forms from a solution, it is called a precipitate

17. Solubility Rules (know them) • 1. Group IA and ammonium compounds are soluble. • 2. Acetates and nitrates are soluble. • 3. Most chlorides, bromides, and iodides are soluble. • Exceptions: AgCl, Hg2Cl2, PbCl2;AgBr, Hg2Br2, HgBr2, PbBr2;AgI, Hg2I2, HgI2, PbI2 • 4. Most sulfates are soluble. • Exceptions: CaSO4, SrSO4, BaSO4, Ag2SO4, Hg2SO4, PbSO4

20. Reactions in Aqueous Solution • Acid-Base Reactions • An acid + base produces a salt + water • “Salt” refers to an ionic compound. • Acids produce H+ ions when dissolved in water. Bases increase the amount of OH- ions when dissolved in water. • Simplest definitions • What else do we know? • Taste • pH • Slippery? • Strong acids are strong electrolytes • Weak acids are weak electrolyotes • Acids can also be: • Monoprotic (one proton produced) • Diprotic (two protons produced) • Triprotic….

22. Let’s look at some reactions…. Autoprotolysis of water Nitric acid plus sodium hydroxide (SA + SB) Oxalic acid plus sodium hydroxide (WA + SB) Acetic acid plus water (WA + WB) Acetic acid plus lithium hydroxide (WA + SB) Ammonia plus water (WB + WA)

23. Precipitation Reactions • You MUST know solubility rules • Ordinarily double displacement reactions 3 CaCl2 (aq) + 2 Na3PO4 (aq)  Ca3(PO4)2 (s) + 6 NaCl (aq) If the polyatomic ions are constant, you can balance them as “groups” • Ammonium chloride + lead (II) nitrate solutions are mixed. • Does a reaction occur? • If so, is there a precipitate? • Write and balance the chemical equation • Lithium phosphate and potassium iodide are mixed.

24. Net Ionic Equations (NIE) • In many reactions, there are important compounds/ions and then there are spectators • Spectator ions are present in the solution, but don’t actually react or change what they are associated with • They are sometimes an ion in a strong electrolyte • Net Ionic Equations eliminate spectators AgNO3(aq) + NaCl(aq)AgCl(s) + NaNO3(aq) Ag+(aq) + NO3-(aq) + Na+(aq)+ Cl-(aq) AgCl(s) + Na+(aq) + NO3(aq) Ag+(aq) + Cl-(aq)AgCl(s)

26. Potassium iodide reacts with lead (II) nitrate. Does a reaction take place? Write the balanced equation. Identify strong electrolytes, soluble compounds. Identify spectator ions. Write Net Ionic Equation.

27. Net Ionic Equations & Acid/Base Reactions • Acids and bases are strong or weak electrolytes.

28. Lithium hydroxide reacts with phosphoric acid to produce lithium phosphate and water.

29. Net Ionic Equations & Gas Forming Reactions • Even though non-electrolytes are a big part of the reaction (the gas), there are still spectators.

30. Calcium carbonate reacts with hydrochloric acid to produce carbon dioxide, water and calcium chloride.

31. Oxidation Reduction (REDOX) Reactions • Now we start looking at when electrons move from one atom to another • Charges Change • Oxidation: when electrons are lost • Reduction: when electrons are gained • Because the charge is reduced

33. Half-Reactions • Half-reactions are like mini net ionic equations. • Oxidation half-reaction • Reduction half-reaction • They are used to identify the two processes and balance the overall equation. • Plus, lots of times a chemist is only worried about one half of the reaction.

34. Oxidation Numbers (charges) • The charge on some ion or species • Cations (+) or anions (-) • Now we are looking at: • Elements becoming ions • Ions changing their charge 2 HCl(aq) + Zn(s)  ZnCl2(aq) + H2(g) Write the half-reactions Identify oxidizing and reducing agents

36. Chromate, dichromate, sulfite, sulfate, nitrite, nitrate. What are the oxidation numbers in each?

37. What is the oxidation number of Cr in dichromate, Cr2O72-? Cr O 2(oxidation number of Cr) + 7(-2) = -2 2(oxidation number of Cr) + (-14) = -2 2(oxidation number of Cr) = -12 Oxidation number of Cr = +6

38. Solution Stoichiometry…… In the titration above, the indicator changes color to indicate when the reaction is just complete.

39. Molarity (M): moles solute/liter of solution • The solution volume is the FINAL volume (includes everything. • Molarity changes with temperature, which can be a problem because solutions expand and contract. • One of the more common units we will use • Molality (m): moles of solute / kg of solvent • The mass of a solvent used is constant, regardless of temperature • A significant unit in geochemistry, especially when dealing with various melted rocks (hot boiling magma)….

40. How much sodium chloride do I need to make a 1.0 liter solution that is 0.35 M NaCl? A solution is made from 45.0 g of Calcium Nitrate dissolved in 2.0 liters of water. What is the concentration (M) of each ion in the solution? Describe how you would make a 2.0 molar solution of phosphoric acid (500 mL).

41. Stoichiometry Calculations

42. 250 mL of a 1.0 M calcium chloride solution is mixed with a solution of sodium phosphate (in excess). After the reaction is complete, the resulting volume is 1.0 liters. • Write the balanced equation • Identify any precipitates • Calculate the amount of any precipitate that is formed • What is the resulting concentration of calcium ion in solution?

43. Acid Base Titrations

44. The titrant is added to the buret. • The equivalence point is when the moles of titrant are exactly the amount needed to react completely with the analyte • An indicator is a substance that is used to show the equivalence point (a dye of some sort) • The indicator changes color at the end point-this is sometimes a little different from the equivalence point.

45. In the titration above, the indicator changes color to indicate when the reaction is just complete.

46. 2.00 grams of oxalic acid dihydrate are titrated to the endpoint with 25.0 mL of sodium hydroxide. What is the molarity of the NaOH solution? • The standardized NaOH solution above is used to titrate 10.0 mL of sulfuric acid. • What is the molarity of the sulfuric acid solution • How many grams of sulfuric acid were in the solution?