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Chapter 8-Chemical Equations & Reactions

Chapter 8-Chemical Equations & Reactions

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Chapter 8-Chemical Equations & Reactions

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  1. Chapter 8-Chemical Equations & Reactions 8.1-Describing Chemical Reactions 8.2-Types of Chemical Reactions 8.3-Activity Series of the Elements

  2. Describing Chemical Reactions

  3. Signs of a Chemical Reaction • Evolution of heat and light • Formation of a gas • Formation of a precipitate • Color change

  4. Characteristics of Chemical Equations • The equation must represent known facts. • The equation must contain the correct formulas for the reactants and products. • You must remember & be able to apply the knowledge from Ch. 7 regarding writing names & formulas for ionic & covalent compounds. • The Law of Conservation of Mass must be satisfied.

  5. Law of Conservation of Mass • Mass is neither created nor destroyed in a chemical reaction • Total mass stays the same • Atoms can only rearrange 4 H 2 O 4 H 2 O 36 g 4 g 32 g

  6. Chemical Equations A+B  C+D PRODUCTS REACTANTS

  7. Chemical Equations

  8. Word vs. Formula/Chemical Eqns. • Word equations are equations in which the reactants & products are represented by words. • Methane + oxygen  carbon dioxide + water • Formula equations are equations in which the reactants & products are represented by their symbols or formulas. • CH4(g) + O2(g) CO2(g) + H2O(g)

  9. Diatomic Elements • Exist as two atoms bonded together. • Remember brinclhof! • Bromine • Iodine • Nitrogen • Chlorine • Hydrogen • Oxygen • Fluorine

  10. Writing Word & Formula Eqns. • Write word and chemical equations for each of the following: • Hydrogen peroxide in an aqueous solution decomposes to produce oxygen and water. • Hydrogen peroxide  oxygen + water • H2O2(aq) O2(g) + H2O(l)

  11. Writing Word & Formula Eqns. • Solid copper metal reacts with aqueous silver nitrate to produce solid silver metal and aqueous copper(II) nitrate. • Copper + silver nitrate  silver + copper(II) nitrate • Cu(s) + AgNO3(aq)  Ag(s) + Cu(NO3)2(aq) • Solid zinc metal reacts with aqueous copper(II) sulfate to produce solid copper metal and aqueous zinc sulfate. • Zinc + copper(II) sulfate  copper + zinc sulfate • Zn(s) + CuSO4(aq)  Cu(s) + ZnSO4(aq)

  12. Balancing Equations • In a chemical equation, the law of conservation of mass is satisfied by “balancing” the number of atoms of each element in the reactants and products.

  13. Balancing Steps 1. Write the unbalanced equation. 2. Count atoms on each side. • Add coefficients to make #s equal. *Coefficient-small whole number that appears in front of a formula in a chemical eqn. Coefficient  subscript = # of atoms 4. Reduce coefficients to lowest possible ratio, if necessary. 5. Double check atom balance!!!

  14. Helpful Tips • Balance one element at a time. • Update ALL atom counts after adding a coefficient. • If an element appears more than once per side, balance it last. • Balance polyatomic ions as single units. • “1 SO4” instead of “1 S” and “4 O”

  15. Al + CuCl2 Cu + AlCl3 Al Cu Cl Balancing Example Aluminum and copper(II) chloride react to form copper and aluminum chloride. 2 3 3 2  2  6 1 1 1 1 2 3 2  3  6   3

  16. Balancing Examples • H2O2(aq) O2(g) + H2O(l) • 2H2O2(aq) O2(g) + 2H2O(l) • Cu(s) + AgNO3(aq)  Ag(s) + Cu(NO3)2(aq) • Cu(s) + 2AgNO3(aq)  2Ag(s) + Cu(NO3)2(aq) • Zn(s) + CuSO4(aq)  Cu(s) + ZnSO4(aq) • Zn(s) + CuSO4(aq)  Cu(s) + ZnSO4(aq)

  17. Writing Equations • Identify the substances involved. • Use symbols to show: • How many? – coefficient • Of what? – chemical formula • In what state? – physical state • Remember the diatomic elements.

  18. Describing Equations • Describing Coefficients: • individual atom = “atom” • covalent substance = “molecule” • ionic substance = “unit” 3CO2 2Mg  4MgO  3 molecules of carbon dioxide 2 atoms of magnesium 4 units of magnesium oxide

  19. Describing Equations Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g) to produce • How many? • Of what? • In what state? One atom of solid zinc reacts with two molecules of aqueous hydrochloric acid one unit and one of aqueous zinc chloride molecule of hydrogen gas.

  20. Describing Equations • Describe the following equation. Remember: how many? of what? in what state? • One unit of aqueous lead(II) chloride reacts with one unit of aqueous sodium chromate to produce one unit of solid lead(II) chromate and two units of aqueous sodium chloride. PbCl2(aq) + Na2CrO4(aq) PbCrO4(s) + 2NaCl(aq)

  21. Review • Write balanced chemical equations for each of the following reactions: • Solid sodium combines with chlorine gas to produce solid sodium chloride. • When solid copper reacts with aqueous silver nitrate, the products are aqueous copper(II) nitrate & solid silver. • In a blast furnace, the reaction between solid iron(III) oxide & carbon monoxide gas produces solid iron & carbon dioxide gas. • 2Na(s) + Cl2(g) 2NaCl(s) • Cu(s) + 2AgNO3(aq) Cu(NO3)2(aq) + 2Ag(s) • Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g)

  22. Types of Chemical Reactions Pages 262-269

  23. Classifying reactions • Enables us to predict the products of reactions • 5 basic types: • Synthesis • Decomposition • Single-replacement • Double-replacement • Combustion

  24. Synthesis • AKA composition reaction • the combination of 2 or more substances to form a compound • only one product A + B  AB

  25. Synthesis H2(g) + Cl2(g) 2HCl(g)

  26. Reactions of Metals with Oxygen, Sulfur, & Halogens • When a metal reacts with oxygen, sulfur, or a halogen (gp 17), an ionic compound will form. • Use the criss cross method to determine the formula of the product. • Nonmetals can also react, but their products are harder to predict. • Carbon can react with oxygen to form both CO2 and CO.

  27. Example • Mg(s) + O2(g)? • Type-synthesis • Form ionic compound (metal & non)criss cross charges • Mg+2 O-2  MgO • Mg(s) + O2(g) MgO (unbalanced!) • 2Mg(s) + O2(g)  2MgO

  28. Decomposition • a compound breaks down into 2 or more simpler substances • only one reactant AB  A + B

  29. Decomposition 2H2O(l) 2H2(g) + O2(g)

  30. Decomposition of Binary Compounds • Binary compounds will decompose into their component elements. • 2H2O(l)  2H2(g) + O2(g)

  31. Decomposition of Metal Carbonates • Ionic compounds formed from a metal cation and the carbonate anion (CO32-) will decompose to form a metal oxide and carbon dioxide. • CaCO3(s)  CaO(s) + CO2(g)

  32. Decomposition of Metal Hydroxides • Ionic compounds formed from a metal cation and the hydroxide anion (OH-) will decompose to form a metal oxide and water. • Ca(OH)2(s)  CaO(s) + H2O(g)

  33. Decomposition of Metal Chlorates • Ionic compounds formed from a metal cation and the chlorate anion (ClO3-) will decompose to form a metal chloride and oxygen. • 2KClO3(s)  2KCl(s) + 3O2(g)

  34. Decomposition of Acids • Certain acids decompose to form nonmetal oxides and water. • Carbonic acid (H2CO3) will decompose to form H2O and CO2. • Sulfuric acid (H2SO4) will decompose into water and sulfur trioxide.

  35. Single Replacement • one element replaces another in a compound • metal replaces metal (+) • nonmetal replaces nonmetal (-) A + BC  B + AC

  36. Single Replacement Cu(s) + 2AgNO3(aq) Cu(NO3)2(aq) + 2Ag(s)

  37. Double Replacement • ions in two compounds “change partners” • cation of one compound combines with anion of the other AB + CD  AD + CB

  38. Double Replacement Pb(NO3)2(aq) + K2CrO4(aq) PbCrO4(s) + 2KNO3(aq)

  39. Combustion • the burning of any substance in O2 to produce heat A + O2 B CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

  40. Combustion • Predicting Products: • contain oxygen • hydrocarbons form CO2 + H2O Na(s)+ O2(g) 4 2 Na2O(s) 5 3 4 C3H8(g)+ O2(g) CO2(g)+ H2O(g)

  41. Classify each rxn: • N2(g) + 3H2(g) 2NH3(g) • 2C6H14(l) + 19O2(g)  12CO2(g) + 14H2O(g) • AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) • Synthesis • Combustion • Double replacement

  42. Example • HgO(s) ? • Type-decomposition • Binarysplit into component elements • HgO(s)  Hg(l) + O2(g) (unbalanced) • 2HgO(s)  2Hg(l) + O2(g)

  43. Predicting products:double replacement rxns • AB + CD  AD + CB • For this type of rxn to occur, at least one of the products must be insoluble in water (precipitate or be a solid) • Check solubility rules/table • Rules pg. R64 • Table B-12

  44. Example • AgNO3(aq) + NaCl(aq) ? • Type-double replacement • Expected products: • AgCl and NaNO3 • Check the solubility of these compounds • Insoluble, soluble • Rxn will occur b/c a product is insoluble • AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)

  45. Example • KOH(aq) + Ca(NO3)2(aq) ? • Type-double replacement • Expected products: • KNO3 and Ca(OH)2 • Check the solubility of these compounds • Soluble, soluble • Rxn will NOT occur b/c both products are soluble!! • KOH(aq) + Ca(NO3)2(aq)  no rxn

  46. Activity Series of the Elements

  47. Predicting products: single replacement rxns • A + BC  B + AC • To determine if A will replace B in compound BC, use the activity series • Activity series-list of elements organized according to the ease with which the elements undergo certain chemical reactions. • An element can replace any element below it in the series. • Pg. 266

  48. Example • Al(s) + ZnCl2(aq)? • Type-single replacement • What is replacing what? • Al replacing Zn • Will this happen? Check to see if Al is more active than Zn. • Al is above Zn, so it is more active • Al(s) + ZnCl2(aq) Zn(s) + AlCl3(aq) (unbalanced) • 2Al(s) + 3ZnCl2(aq) 3Zn(s) + 2AlCl3(aq)

  49. Example • Co(s) + 2NaCl(aq)? • Type-single replacement • What is replacing what? • Co replacing Na • Will this happen? Check to see if Co is more active than Na. • Co is NOT above Na in the activity series, so it will NOT replace Na. • Co(s) + 2NaCl(aq) no rxn (NR)

  50. Review • Using the activity series, predict whether each of the possible reactions will occur. For reactions that will occur, write the products & balance the eqn. • MgCl2(aq) + Zn(s) • Al(s) + H2O(g)  • Cd(s) + O2(g)  • I2(s) + KF(g)  NR 2Al(s) + 3H2O(g) Al2O3(s) + 3H2(g) 2Cd(s) + O2(g) 2CdO(s) NR