types of chemical reactions n.
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  2. AIM: Precipitate Reactions • DO NOW: • 1. Take out the worksheet on predicting the reactions (from before midterm exams) • Go to your note sheet that you printed out on Types of Chemical Reactions and Solution Stoichiometry • Exercise 8 (pg. 8) predict the products of the three double replacement reactions ONLY predict the PRODUCTS

  3. AIM: Precipitate Reactions • Get your 3D- glasses! 3D Video – Physical vs. Chemical Changes Neutralization RedOx

  4. CHEMICAL EQUATIONS • chemical reaction--transforms elements and compounds into new substances • balanced chemical equation--shows the relative amounts of reactants and products • s, l, g, aq--solid, liquid, gas, aqueous solution • NO ENERGY or TIME is given • Antoine Lavoisier (1743-1794)--law of conservation of matter: matter can neither be created • nor destroyed this means “balancing equations”

  5. PRECIPITATION REACTIONS • Precipitate is a driving force for a chemical reaction • A precipitate is an insoluble solid that is formed when two aqueous solutions are mixed • separate the precipitate (ppt) from solution by filtration in what is called a gravimetric analysis.


  7. Practice w/ Solubility Rules • Using the solubility rules, predict what will happen in the examples from your worksheet

  8. Exercise 8 • Using the solubility rules, predict what will happen when the following pairs of solutions are mixed. • KNO3(aq) and BaC12(aq)

  9. Exercise 8 • Na2SO4(aq) and Pb(NO3)2(aq)

  10. Exercise 8 • KOH (aq) and Fe(NO3)3(aq)

  11. DESCRIBING REACTIONS IN SOLUTIONS • COMPLETE balanced equation—gives the overall reaction stoichiometry, but NOT the forms of the reactants & products as they exist in solution • complete ionic equation—represents as IONS all reactants & products that are strong electrolytes • net ionic equations—includes only those solution components undergoing a change. Spectator ions are NOT included • spectator ions--not involved in the reaction process Λ started as an ion AND finished as an ion • THERE IS ALWAYS A CONSERVATION OF CHARGE IN NET IONIC EQN’S.

  12. EXERCISE 9 - For each of the following reactions, write the molecular equation, the complete ionic equation, and the net ionic equation. • Aqueous potassium chloride is added to aqueous silver nitrate to form a silver chloride precipitate plus aqueous potassium nitrate.

  13. EXERCISE 9 • Aqueous potassium hydroxide is mixed with aqueous iron(III) nitrate to form a precipitate of iron(III) hydroxide and aqueous potassium nitrate.

  14. ACID BASE REACTIONS DO NOW: 1. Write out the word equation for what happens when an acid and base react 2. Give examples of strong base/strong acid

  15. ACID – BASE REACTIONS • acids--any cmpd. that, on reaction with water, produces an ion called the hydronium ion, H3O+ [or H+], and an anion (Arrhenius definition) • base--any cmpd. that provides a hydroxide, OH−,and a cation in water (Arrhenius definition) **ammonia, NH3 is an exception, so Bronsted-Lowry defined it as a proton acceptor!! • neutralization—when moles acid = moles base each is neutralized [pH is not necessarily 7.0]. • The products formed are a salt [ask yourself if it is soluble] and water

  16. SOLUBILITY RULES These are strong electrolytes (100% ionized ) and written as ions • 1. Strong Acids: HCl, HBr, HI, H2SO4, HNO3, HClO4, HClO3 • Sulfuric acid (strong acid) can be written as H+ and SO42- or as H+ and HSO4-. If it says CONCENTRATED write it as a molecule***** • 2. Strong Bases: Hydroxides of group IA and IIA(Ba, Sr, Ca are marginal Be and Mg are WEAK)

  17. Rules for Acid – Base Reactions Know SOLUBILITY RULES!!!!! Ionize strong acids and bases Watch out for concentrated sulfuric acid

  18. Example • A solution of sulfuric acid is added to a solution of barium hydroxide until the same number of moles of each compound as been added. • Write out reaction: H2SO4 + Ba(OH)2 BaSO4 +H2O H+ + SO42- + Ba2+ + OH- BaSO4+ H2O

  19. Example • Solutions of ammonia and hydrofluoric acid are mixed

  20. Example • Hydrogen sulfide gas is bubbled through a solution of potassium hydroxide

  21. Example • A solution of sulfuric acid is added to a solution of barium hydroxide until the same number of moles of each compound has been added

  22. Example • A solution of sodium hydroxide is added to a solution of sodium dihydrogen phosphate until the same number of moles of each compound has been added

  23. Example • Dilute nitric acid is added to crystals of pure calcium oxide

  24. Example • Equal volumes of 0.1 molar sulfuric acid and 0.1 molar potassium hydroxide are mixed

  25. Example • A solution of ammonia is added to a dilute solution of acetic acid

  26. RedOxREACTIONS DO NOW: 1. What occurs in oxidation-reduction reactions? 2. What do the following mean? OiLRiG LEO says GER 3. Using the Rules for Assigning oxidation number on pg. 13 of Types of Chemical Reactions and Solution Stoichiometry do Exercise 16

  27. Terms to Know • OIL RIG – oxidation is loss, reduction is gain (of electrons) • Oxidation – the loss of electrons, increase in charge • Reduction – the gain of electrons, reduction of charge • Oxidation number – the assigned charge on an atom • Oxidizing agent (OA) – the species that is reduced and thus causes oxidation Reducing agent (RA) – the species that is oxidized and thus causes reduction

  28. Rules for Assigning Oxidation Numbers

  29. Exercise 17

  30. Exercise 18

  31. Balancing RedOx Reactions by Half Reaction Method • Divide the equation into oxidation and reduction half reactions. [OILRIG] • Balance all elements besides hydrogen and oxygen. • Balance O’s by adding H2O’s to the appropriate side of each equation. • Balance H’s by adding H+ • Balance the charge by adding electrons. [OILRIG again] • Multiply the half reactions to make electrons equal for both half-reactions. • Cancel out any common terms and recombine the two half reactions. • IF BASIC, neutralize any H+ by adding the SAME NUMBER of OH- to EACH side of the • balanced equation. [This creates some waters that will cancel!] • CHECK!!

  32. Sample Problem • Assign oxidation states to all atoms in the following equation, identify the oxidation and reduction half reactions, and the OA and RA. MnO4−(aq) + Fe2+(aq)  Mn+2(aq) + Fe3+(aq)

  33. Sample Problem • Balance the following equation using the half-reaction method. (acidic) MnO4−(aq) + I−(aq)  Mn+2(aq) + I2(aq)

  34. Sample Problem • (basic) Ag(s) + CN− + O2Ag(CN)2−(aq)

  35. Exercise 19

  36. Exercise 20

  37. Oxidation Reduction Reactions • Redox reactions involve the transfer of electrons. The oxidation numbers of at least two elements must change. Single replacement, some combination and some decomposition reactions are redox reactions. • To predict the products of a redox reaction, look at the reagents given to see if there is both an oxidizing agent and a reducing agent. When a problem mentions an acidic or basic solution, it is probably redox.

  38. Oxidation Reduction Reactions

  39. Oxidation Reduction Reactions

  40. EXAMPLE • Iron (III) ions are reduced by iodide ions

  41. EXAMPLE • Solution of tin (II) sulfate is added to a solution of iron (II) sulfate

  42. EXAMPLE • Metallic copper is heated with concentrated sulfuric acid

  43. EXAMPLE • Manganese (IV) is added to warm concentrated hydrobromic acid

  44. EXAMPLE • Chlorine gas is bubbled into cold dilute sodium hydroxide

  45. EXAMPLE • Solid iron (III) oxide is heated in excess carbon monoxide

  46. EXAMPLE • Hydrogen peroxide solution is added to acidified potassium iodide solution

  47. EXAMPLE • Potassium permanganate solution is added to concentrated hydrochloric acid

  48. Decomposition Reactions • Reactions where a compound breaks down into two or more elements or compounds. Heat, electrolysis, or a catalyst is usually necessary. • A compound may break down to produce two elements. • A compound may break down to produce an element and a compound. • A compound may break down to produce two compounds.

  49. Decomposition Reactions • Metallic carbonates break down to yield metallic oxides and carbon dioxide, • Metallic chlorates break down to yield metallic chlorides and oxygen. • Hydrogen peroxide decomposes into water and oxygen. • Ammonium carbonate decomposes into ammonia, water and carbon dioxide. • Sulfurous acid decomposes into water and sulfur dioxide. • Carbonic acid decomposes into water and carbon dioxide.

  50. Example • A solution of hydrogen peroxide is heated