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Chemical reactions

Chemical reactions

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Chemical reactions

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  1. Chemical reactions ClassificationsReactions in solution Ionic equations

  2. Learning objectives • Distinguish between chemical and physical change • Describe concepts of oxidation and reduction • Classify reaction according to types of reactants and products • Distinguish among strong, weak and non-electrolytes • Identify common acids and bases by from chemical formula • Predict formation of precipitates by application of solubility rules • Write total and net ionic equations from balanced molecular equations

  3. One approach to classification

  4. Oxidation – reduction: focusing on electrons • Oxidation is loss of electrons • Reduction is gain of electrons • Oxidation is always accompanied by reduction • The total number of electrons is kept constant • Oxidizing agents oxidize and are themselves reduced • Reducing agents reduce and are themselves oxidized

  5. Redox in chemistry • All reactions involve rearrangement of atoms and molecules • Some reactions involve rearrangement of atoms and molecules and electrons • Photosynthesis, respiration, combustion... • These are called redox reactions • Any reaction involving elements must be redox

  6. Combination reactions • Element + element  compound (redox) • S + O2→ SO2 • Metal + nonmetal  binary ionic compound • Nonmetal + nonmetal  binary covalent compound • Compound + element  compound (redox) • CO + O2→ CO2 • Compound + compound  compound • SO2 + H2O →H2SO3

  7. Decomposition reactions • Compound  element + element (redox) • HgO → Hg + O2 • Compound  element + compound (redox) • PCl5→ PCl3 + Cl2 • Compound  compound + compound • CaCO3→ CaO + CO2

  8. Single replacement (displacement) • Element displaces another element from compound (redox) • Zn + CuSO4→ ZnSO4 + Cu

  9. Double replacement (displacement) • Compounds exchanging partners • Usually ionic compounds in solution • Identify ions and swap them • KCl + AgNO3 → KNO3 + AgCl(s) • Very often a solid is produced

  10. Acid – base neutralization:special case of double replacement • KOH(aq) + HNO3(aq) = KNO3(aq) + H2O(l) • Product is liquid water not a solid BASE ACID SALT WATER

  11. Combustion • Element or compound reacting with oxygen (redox) • CH4 + O2→ CO2 + H2O • Associated with production of heat and light • Often involves hydrocarbons (CxHy) • CO2 and H2O are products

  12. Sorting solution reactions: dissolved species • Electrolytes: • Ionic compounds produce ions in solution (NaCl, NH4NO3 etc.) • Non-electrolytes: • Covalent compounds do not produce ions in solution (CH3OH, C6H12O6 etc.)

  13. Electrolytes: distinguishing by strength • Strong electrolytes are characterized by complete dissociation in water • Weak electrolytes dissociate to a much smaller extent.

  14. Strong, weak or non electrolyte? • All soluble salts are strong electrolytes • Strong acids and bases are strong electrolytes • Weak acids and bases are weak electrolytes • Insoluble compounds are non-electrolytes • Molecular compounds are non-electrolytes

  15. Classification of electrolytes

  16. Four classes of substance with solution reactions Yes No Yes No cov weak strong ionic

  17. Recognizing acids and bases • Acids usually have H at the beginning of the formula – HCl • Bases usually have OH in the formula – NaOH • Not in organic compounds though - CH3OH

  18. Focus on double replacement • Driven by removal of ions from solution • Formation of an insoluble solid (precipitate) • Formation of non-ionized molecules (acid – base) • Formation of a gas

  19. 1. Predicting precipitation reactions • Does one of the new cation-anion combinations produce insoluble salt? • How do I know? • Initial combinations are all soluble • Use solubility rules to investigate • If yes, a precipitate is produced

  20. Solubility rools • Group I and ammonium compounds are generally soluble • Nitrates and acetates are generally soluble • Chlorides, bromides and iodides are generally soluble {except Pb(II), Ag(I) and Hg(I)} • Carbonates and phosphates are generally insoluble (except group I) • Hydroxides and sulphides are generally insoluble(except groups I and II)

  21. 3. Production of a gas • If product is a gas that has a low solubility in water, reaction in solution is driven to produce the gas • Tums relief • Any carbonate with an acid NaHCO3(s) + HCl(aq) = NaCl(aq) + H2O(l) + CO2(g)

  22. Writing balanced molecular equations for double replacement reactions • Use correct formulae • Metal ion charge predicted from group number • Use table for correct formula and charge for polyatomic ions • Identify as solid (s), gas (g), liquid (l) or dissolved (aq) • Balance: atoms (groups) on left = atoms (groups) on right

  23. Balancing double replacement equations • It’s very much a matter of states – show them! Pb(NO3)2(aq) + 2KI(aq) = 2KNO3(aq) + PbI2(s) • Balance polyatomic ions as units: • Pb2+, K+, I-, NO3-

  24. Molecular equation for reaction of Na2SO4 + Ba(NO3)2

  25. Total ionic equations • Dissolved substances: • Strong electrolytes show as ions • Weak or non- electrolytes show as molecular formula • All others show as molecular formula Pb(NO3)2(aq) + 2KI(aq) = 2KNO3(aq) + PbI2(s) Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) = 2K+(aq) + 2NO3-(aq) + PbI2(s)

  26. Net ionic equations • Spectator ions are those ions that do not undergo a change; they do not participate in the chemical change and are the same on both sides of the equation • Remove all spectator ions from the equation Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) = 2K+(aq) + 2NO3-(aq) + PbI2(s)

  27. Net ionic equations Pb2+(aq) + 2I-(aq) = PbI2(s) • Mass and charge must still balance, although overall charge may not be neutral in a net ionic equation

  28. Net ionic equation for reaction of Na2SO4 + Pb(NO3)2