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Chemical reactions

Chemical reactions. Classifications Reactions in solution Ionic equations. Learning objectives. Distinguish between chemical and physical change Describe concepts of oxidation and reduction Classify reaction according to types of reactants and products

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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

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