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Chemical Reactions in Aqueous Solutions

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  1. Chapter Four Chemical Reactions in Aqueous Solutions

  2. Today… • Turn in: • Nothing • Our Plan: • Notes – Synthesis & Decomposition • Begin Worksheet #1 • Homework (Write in Planner): • WS#1 Due Wednesday

  3. Quick Review • How do we know a chemical reaction has occurred? What do we observe?

  4. Evidence of a Chemical Rxn • Color Change • Precipitate (solid) forms • Gas Evolved • Heat/Light Given Off • Endothermic – heat absorbed (gets cold) • Exothermic – heat released (gets hot) • More about this in Unit 5 - Thermodynamics

  5. Review of Basics • When writing out a balanced equation, it is necessary to indicate what state the substance is in (s, l, g, or aq). • For elements, you can look at the PT to determine their standard state. • Examples: • Mercury • Fluorine • Iron

  6. Review of Basics • There are 7 diatomic elements. • They are only diatomic when they are alone. They are all gases, except Br2,which is a liquid, as indicated on the PT. • Remember the Super 7? • H2, N2, O2, F2, Cl2, Br2, I2

  7. Review of Basics • Soluble means the substance dissolves in solution (water). • Label soluble substances aqueous (aq) • Insoluble means the substance does not dissolve in solution. • Label insoluble substances solid (s) • To determine if a substance is soluble or insoluble, use a solubility table.

  8. Solubility Chart

  9. Other Basic Tips • Acids and bases are aqueous • Water is almost always a liquid • Write water as HOH when it is a reactant • Some common gases that aren’t diatomic are CO2, CO, SO3, SO2, H2S, and NH3

  10. Other Basic Tips • If a reaction indicates that a catalyst was present, it is not part of the actual reaction. Instead, you write the catalyst above the arrow. • Catalysts are not used up in the reaction. They are just there to speed it up. • If a reaction is heated, you draw a triangle above the arrow. Heat is used as a catalyst.

  11. Synthesis Reactions • Occur when two or more reactants combine to form a single product. • In Chem 1 we called this type “Dating” • There are several common types of synthesis reactions.

  12. Synthesis Reactions Type 1: A metal combines with a nonmetal to form a binary salt. Example: A piece of lithium metal is dropped into a container of nitrogen gas. 6Li (s) + N2 (g) → 2Li3N (aq)

  13. Synthesis Reactions Type 2: Metallic oxides and water form bases (metallic hydroxides). Example: Solid sodium oxide is added to water Na2O (s) + HOH (l) → 2NaOH (aq) Example: Solid magnesium oxide is added to water. MgO (s) + HOH (l) → Mg(OH)2 (aq)

  14. Synthesis Reactions Type 3: Nonmetallic oxides and water form acids. The nonmetal retains its oxidation number. Example: Carbon dioxide is bubbled in water. CO2 (g) + H2O (l) → H2CO3 (aq) Example: Dinitrogenpentoxide is bubbled in water. N2O5 (g) + H2O (l) → 2HNO3 (aq)

  15. Synthesis Reactions Type 4: Metallic oxides and nonmetallic oxides form salts. Example: Solid sodium oxide is added to carbon dioxide. Na2O (s) + CO2 (g) → Na2CO3 (aq) Example: Solid calcium oxide is added to sulfur trioxide. CaO (s) + SO3 (g) → CaSO4 (aq)

  16. Try It Out 1 • Solid barium oxide is added to distilled water (p. 142)

  17. Try It Out 2 • Sulfur trioxide gas is added to excess water (p. 146)

  18. Decomposition Reactions • Occur when a single reactant is broken down into two or more products. • In Chem 1 we called this “The Break Up” • There are several common types of decomposition reactions.

  19. Decomposition Reactions Type 1: Metallic carbonates decompose into metallic oxides and carbon dioxide Example: A sample of magnesium carbonate is heated. MgCO3 (s) → MgO (s) + CO2 (g)

  20. Decomposition Reactions Type 2: Metallic chlorates decompose into metallic chlorides and oxygen. Example: A sample of magnesium chlorate is heated. Mg(ClO3)2 (aq) → MgCl2 (aq) + 3O2 (g)

  21. Decomposition Reactions Type 3: Ammonium carbonate decomposes into ammonia, water, and carbon dioxide. Example: A sample of ammonium carbonate is heated. (NH4)2CO3 (aq)→ 2NH3 (g) + H2O (l) + CO2 (g)

  22. Decomposition Reactions Type 4: Sulfurous acid decomposes into sulfur dioxide and water. Example: A sample of ammonium carbonate is heated. H2SO3 (aq) → H2O (l) + SO2 (g)

  23. Decomposition Reactions Type 5: Carbonic acid decomposes into carbon dioxide and water. Example: A sample of carbonic acid is heated. H2CO3 (aq) → H2O (l) + CO2 (g)

  24. Decomposition Reactions Type 6: A binary compound may break down to produce two elements. Example: Molten sodium chloride is electrolyzed. 2NaCl(l) → 2Na (s) + Cl2 (g)

  25. Decomposition Reactions Type 7: Hydrogen peroxide decomposes into water and oxygen. Example: 2H2O2(aq) → 2H2O (l) + O2 (g)

  26. Decomposition Reactions Type 8: Ammonium hydroxide decomposes into ammonia and water. Example: NH4OH(aq) → NH3 (g) + HOH (l)

  27. Try It Out 1 • Solid calcium chlorate is heated in the presence of a magnesium dioxide catalyst

  28. Try It Out 2 • A sample of lithium carbonate is heated strongly

  29. STOP • Complete Worksheet #1 by next class.

  30. Today… • Turn in: • Get Out WS#1 • Our Plan: • Questions on WS#1 • Quick Review • Notes – Single & Double Replacement • Begin WS#2 • Homework (Write in Planner): • WS#2 Due Friday

  31. Quick Review • Write formulas for the reactants and predicted products for the chemical reactions that follow: • Solid calcium carbonate is strongly heated in a test tube. • Solid lithium chlorate is heated in the presence of a manganese dioxide catalyst. • Excess chlorine gas is passed over hot iron filings.

  32. Single Replacement • Reactions that involve an element replacing one part of a compound. • The products include the displaced element and a new compound. • In Chem 1 we called this “Cheating” • An element can only replace another element that is less active than itself.

  33. Single Replacement General Activity Series for Metals: Most Active Li Ca Na Mg Al Zn Fe Pb H2 Cu Ag Pt Au Least Active

  34. Single Replacement General Activity Series for Nonmetals: Most Active F2 Cl2 Br2 I2 Least Active

  35. Remember the Trend? • Or you can just look at the Periodic Table • The most reactive metals are on the bottom left (Francium) • The most reactive nonmetals are on the top right (Fluorine) • Hydrogen is the tricky one. • Use the pink sheet for electronegativity potentials otherwise.

  36. For metals, the more negative the reduction potential, the more reactive. For nonmetals, the more positive the reduction potential, the more reactive.

  37. Single Replacement Type 1: Active metals replace less active metals from their compounds in aqueous solution. Example: Magnesium turnings are added to a solution of iron (III) chloride. 3Mg (s) + 2FeCl3(aq)→ 2Fe (s) + 3MgCl2 (aq)

  38. Single Replacement Type 2: Active metals replace hydrogen in water. Example: Sodium is added to water. 2Na (s) + 2HOH(l)→ 2NaOH (aq) + H2 (g)

  39. Single Replacement Type 3: Active metals replace hydrogen in acids. Example: Lithium is added to hydrochloric acid. 2Li (s) + 2HCl(aq) → 2LiCl (aq) + H2 (g)

  40. Single Replacement Type 4: Active nonmetals replace less active nonmetals from their compounds in aqueous solution. Example: Chlorine gas is bubbled into a solution of potassium iodide. Cl2 (g) + 2KI(aq) → I2 (g) + 2KCl (aq)

  41. Single Replacement Type 5: If a less reactive element is combined with a more reactive element in compound form, their will be no resulting reaction. Example: Chlorine gas is bubbled into a solution of potassium iodide. Cl2 (g) + KF (aq) → No Reaction Example: Zinc is added to a solution of sodium chloride. Zn (s) + NaCl (aq) → No Reaction

  42. Try It Out 1 • A piece of aluminum metal is added to a solution of gold nitrate. (p. 146)

  43. Try It Out 2 • Liquid bromine is shaken with a potassium iodide solution. (p. 152)

  44. Try It Out 3 • Small chunks of potassium are added to water. (p. 164)

  45. Try It Out 4 • Small strips of platinum are placed in hydrochloric acid.

  46. Double Replacement • Reactions between two compounds in aqueous solution where the cations and anions appear to “switch partners”. AX + BY → AY + BX • In Chem 1 we called this “Swapping”

  47. Double Replacement • All double replacement reactions (aka metathesis) must have a “driving force” or reason why the reaction will occur or “go to completion”.

  48. Double Replacement • “Driving Force” for reactions: • Formation of a precipitate • Formation of a gas • Formation of primarily molecular species (nonelectrolytes, water, weak acids)

  49. Double Replacement • If one of these “driving forces” is NOT present, then the reaction does not go to completion. • This type of reaction is indicated by a double arrow