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Introduction to Reactions

Introduction to Reactions. Chemical Equation. Reactants  Products Fe + O 2  Fe 2 O 3 A catalyst is a substance that speeds up the reaction but is not changed by it. It is neither a reactant or a product. Signs of a Reaction. Release of a gas

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Introduction to Reactions

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  1. Introduction to Reactions

  2. Chemical Equation • Reactants  Products • Fe + O2  Fe2O3 • A catalyst is a substance that speeds up the reaction but is not changed by it. • It is neither a reactant or a product.

  3. Signs of a Reaction • Release of a gas • CO2 is released when acid is placed in a solution containing CO32- ions • Formation of a solid (precipitate) • A solution containing Ag+ ions mixed with a solution containing Cl- ions • Heat is produced or absorbed • Acid and base are mixed together • Color changes

  4. Common Symbols Symbol Meaning  forms, produces ↔ reversible reaction (s) Solid state (l) Liquid state; water only (g) Gaseous state (aq) aqueous state, all liquids besides water heat/energy is supplied to the reaction Catalyst is used, here platinum

  5. Features of a Chemical Equation Products and reactants must be specified using chemical symbols Reactants – written on the left of arrow Products – written on the right  – energy is needed Physical states are shown in parentheses

  6. Writing Equations • 2H2 (g) + O2(g)  2H2O(g) Identify the substance involved • Coefficients - how many? • Chemical Formula – of what? • Physical State – in what state? • Remember Diatomic Elements • Magic Seven

  7. Example • Two atoms of aluminum react with three units of aqueous copper (II) chloride to produce three atoms of copper and two units of aqueous aluminum chloride? • How many? • Of what? • What physical state?

  8. Example • Two atoms of aluminum react with three units of aqueous copper (II) chloride to produce three atoms of copper and two units of aqueous aluminum chloride? • How many? • Of what? • What physical state? 2 Al(s) + 3 CuCl2(aq)  3 Cu(s) + 2AlCl3(aq)

  9. Describing Equations • Describing Coefficients: • individual atom = “atom” • covalent substance = “molecule” • ionic substance = “unit” 3CO2 2Mg  4MgO 

  10. 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

  11. Describing Equations Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) • How many? • Of what? • In what state?

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

  13. Balancing Reaction

  14. Balancing Reactions • Law of conservation of mass - matter cannot be created or destroyed • mass of the products = mass of the reactants • Coefficient: # of moles of products & reactants • 4Fe + 3O2 2Fe2O3 • Diatomic elements (The Magic 7) H2, N2, O2, F2, Cl2, Br2, I2

  15. Balancing Coefficient- how many of that substance are in the reaction • The equation must be balanced • All the atoms of every reactant must also appear in the products • Number of Hg on left? 2 • on right 2 • Number of O on left? 2 • on right 2

  16. Examine the Equation H2 + O2 H2O • Is the law of conservation of mass obeyed as written? • NO • You never change subscripts • WRONG: H2 + O2 H2O2

  17. Steps in Equation Balancing H2 + O2 H2O The steps to balancing: Step 1. Count the number of moles of atoms of each element on both product and reactant sides ReactantsProducts 2 mol H 2 mol H 2 mol O 1 mol O

  18. Step 2. Determine which elements are not balanced – Oxygen is not balanced Step 3. Balance one element at a time by changing the coefficients H2 + O2 2H2O • This balances oxygen, but is hydrogen still balanced? 2H2 + O2 2H2O Step 4. Make sure the law of conservation of mass is obeyed ReactantsProducts 4 mol H 4 mol H 2 mol O 2 mol O

  19. Practice Equation Balancing Balance the following equations: 1. C2H2 + O2 CO2 + H2O 2. AgNO3 + FeCl3  Fe(NO3)3 + AgCl 3. C2H6 + O2  CO2 + H2O 4. N2 + H2  NH3

  20. Practice Equation Balancing Balance the following equations: 1. 2C2H2 + 5O2 4CO2 + 2H2O 2. 3AgNO3 + FeCl3  Fe(NO3)3 + 3AgCl 3. 2C2H6 + 5O2  4CO2 + 6H2O 4. N2 + 3H2  2NH3

  21. Types of Reaction

  22. Combination Reactions • Synthesis reactions • The joining of two or more elements or compounds, producing a product of different composition Examples: • metal + nonmetal  salt: 2Na(s) + Cl2(g)  2NaCl(s) • H + Cl HCl • MgO(s) + CO2(g)  MgCO3(s) A + B  AB

  23. Decomposition Reactions • Produce two or more products from a single reactant • Reverse of a combination reaction Examples: • 2HgO(s)  2Hg(l) + O2(g) • CaCO3(s)  CaO(s) + CO2(g) • Removal of water from a hydrated material AB  A + B

  24. Replacement Reactions Single-replacement • One atom replaces another in the compound producing a new compound Examples: • Cu(s)+2AgNO3(aq)  2Ag(s)+Cu(NO3)2(aq) • 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g) A + BC  B + AC

  25. Single Replacement Rxn. • Activity Series – lists metals in order of decreasing reactivity (p.333) • Reactive metals will replace any metal listed below it in the activity series • If the metal is below, no reaction occurs • Halogen(7A) can replace other halogens that are below it in the periodic table

  26. Activity Series

  27. Single Replacement Rxn. • 2K(s) + 2H2O(l)  • Zn(s) + Cu(NO3)2(aq)  • Cu(s) + Al2O3(aq)  • Br2(aq) + 2NaI(aq)  • Br2(aq) + NaCl 

  28. Single Replacement Rxn. • 2K(s) + 2H2O(l)  2KOH(aq) + H2(g) • Zn(s) + Cu(NO3)2(aq)  Cu(s) + Zn(NO3)2(aq) • Cu(s) + Al2O3(aq)  No reaction • Br2(aq) + 2NaI(aq)  2NaBr(aq) + I2(aq) • Br2(aq) + NaCl  No reaction

  29. Double Replacement • Two compounds undergo a “change of partners” • Two compounds react by exchanging atoms to produce two new compounds AB + CD  AD + CB

  30. Double Replacement Rxn. • Double-displacement reaction • Exchange of positive ions • Occur in aqueous solution To occur: • One of the products is slightly soluble and a precipitates forms • One product is a gas • One of the products is a molecular compound, like water

  31. Types of Double-Replacement • Acid + base  water and salt HCl(aq)+NaOH(aq) NaCl(aq)+H2O(l) • Formation of solid lead chloride from lead nitrate and sodium chloride Pb(NO3)2(aq) + 2NaCl(aq)  PbCl2(s) + 2NaNO3(aq) AB + CD  AD + CB

  32. Precipitation Reactions • Chemical change in a solution that results in one or more insoluble products Solubility Rules (p.344) 1. salts of alkali metals and ammonia  soluble 2. nitrate salts and chlorate salts  soluble 3. sulfate salts, except compounds with Pb, Ag, Hg, Ba, Sr, and Ca  soluble 4. Chloride salts, except with Ag, Pb, and Hg soluble 5. carbonates, phosphates, chromates, sulfides, and hydroxides  most are insoluble

  33. Predicting Whether Precipitation Will Occur • Recombine the ionic compounds to have them exchange partners • Examine the new compounds formed and determine if any are insoluble • Any insoluble salt will be the precipitate • Pb(NO3)2(aq) + NaCl(aq)  (s) PbCl2 (?) + NaNO3 ( ?) (aq)

  34. Precipitates Predict Whether These Reactions Form Precipitates • Potassium chloride and silver nitrate • Potassium acetate and silver nitrate

  35. Precipitates Predict Whether These Reactions Form Precipitates • Potassium chloride and silver nitrate • KCl(aq) + AgNO3(aq) KNO3(aq) + AgCl(s) • Potassium acetate and silver nitrate • KC2H3O2+ AgNO3(aq) KNO3(aq) + AgC2H3O2(s)

  36. Reactions with Oxygen • Reactions with oxygen generally release energy in the form of light or heat Combustion • Reactants: Oxygen and a hydrocarbons • Products: CO2 and H2O • Combustion of natural gas • CH4+2O2CO2+2H2O • Rusting or corrosion of iron • 4Fe + 3O2 2Fe2O3

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