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Reactions in solution

Reactions in solution. A subset of chemical reactions. Learning objectives. Define solution and its components Distinguish among strong, weak and non-electrolyte Identify strong acids and strong bases Apply solubility rules to prediction of precipitate formation

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Reactions in solution

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  1. Reactions in solution A subset of chemical reactions

  2. Learning objectives • Define solution and its components • Distinguish among strong, weak and non-electrolyte • Identify strong acids and strong bases • Apply solubility rules to prediction of precipitate formation • Classify types of chemical reaction • Predict course of reaction based on activity series • Define oxidation and reduction • Identify oxidizing and reducing agent in reactions • Determine oxidation numbers in ions and compounds

  3. Solution • A homogeneous mixture of two or more substances • Not just limited to liquid state • Solutions may or may not contain electrolytes • Electrolytes are substances that conduct electricity when dissolved

  4. Electrolytes and ionic compounds • All ionic compounds are electrolytes when dissolved in water • Not all ionic compounds are soluble • How do we tell? • Rules to predict solubility • Covalent molecular compounds* are non-electrolytes – no ions produced • *Except acids and bases

  5. Dissociation and ionization: same or different? • Ionic compounds dissociate in water • Ions already exist in the solid • Acids or bases* ionize in water • A pure acid or base contains no ions • *Except strong bases like NaOH, Ca(OH)2 are ionic

  6. When the weak are made strong • Strong electrolytes are characterized by their nearly complete dissociation in water • Weak electrolytes dissociate to a much smaller extent.

  7. Strong, weak or non electrolyte? • All soluble (ionic) 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 (except acids/bases)

  8. Know your acids • The six strong acids • HCl, HBr, HI (but not HF) • HNO3 (but not HNO2) • H2SO4 (but not H2SO3) • HClO4 (maybe HClO3) • All other acids are weak

  9. Recognizing acids • Mineral acids: HCl, HNO3 etc. • Conventionally H appears first in the formula • All strong acids are mineral • May be strong or weak • Organic acids: CH3COOH etc • Harder to spot • Sometimes written with H in front – HCH3CO2 • Always weak • Presence of –OH (-SH): necessary but not sufficient • Not all –OH are acidic (CH3OH is not an acid)

  10. Recognizing bases • Mineral bases usually distinguished by OH groups – all strong • NaOH, Ca(OH)2 • Ammonia, NH3, is an exception – is weak • Organic bases do not contain –(OH) – all weak

  11. Classifying chemical reactions • Acid-base reactions • Oxidation-reduction reactions • Combination reactions • Decomposition reactions • Single displacement reactions • Double displacement (metathesis)/(partner exchange)reactions(in solution)

  12. Neutralization • Combine acid with base: ACID + BASE = SALT + WATER HCl(aq) + NaOH(aq) = H2O(l) + NaCl(aq) Mg(OH)2(s) + 2HCl(aq) = MgCl2(aq) + 2H2O(l) • Salt contains anion of acid and cation of base: HCl + NaOH = NaCl + H2O HCl + KOH = KCl + H2O HNO3 + KOH = KNO3 + H2O 2HCl + Ca(OH)2 = CaCl2 + 2H2O HCN + NaOH = NaCN + H2O

  13. Acid-base reaction with gas formation • Tums... HCl(aq) + NaHCO3(aq) = NaCl(aq) + H2CO3(aq) • H2CO3 is unstable: H2CO3(aq) = H2O(l) + CO2(g) • Bad egg gas: 2HCl + Na2S = H2S(g) + 2NaCl(aq)

  14. Oxidation - reduction • 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

  15. Oxidation number is the number of electrons gained or lost by the element in making a compound Oxidation numbers • Metals are typically considered more 'cation-like' and would possess positive oxidation numbers, while nonmetals are considered more 'anion-like' and would possess negative oxidation numbers.

  16. Predicting oxidation numbers • Oxidation number of atoms in element is zero • Oxidation number of element in monatomic ion equals charge • Sum of oxidation numbers in compound is zero • Sum of oxidation numbers in polyatomic ion equals charge • F has ON –1 • H has ON +1; except in metal hydrides where it is –1 • Oxygen is usually –2. Exceptions: • O is –1 in hydrogen peroxide, and other peroxides • O is –1/2 in superoxides KO2 • In OF2 O is +2

  17. Position of element in periodic table determines oxidation number • G1A is +1 • G2A is +2 • G3A is +3 (some rare exceptions) • G5A are –3 in compounds with metals, H or with NH4+ • Exceptions are compounds with elements to right (e.g. NO2, PF5); in which case use rules 3 and 4. • G6A below O (S, Se etc.) are –2 in binary compounds with metals, H or NH4+ • When combined with O or lighter halogen (e.g. SeO2, SF6) use rules 3 and 4. • G7A elements are –1 in binary compounds with metals, H or NH4+ or with a heavier halogen (e.g. Cl in BrCl3) • When combined with O or a lighter halogen, use rules 3 and 4 (e.g. Br in BrCl3 or Cl in ClO4-).

  18. Identifying reagents • Those elements that tend to give up electrons (metals) are typically categorized as reducing agents and those that tend to accept electrons (nonmetals) are referred to as oxidizing agents.

  19. Identify redox by change in oxidation numbers • Reducing agent increases its oxidation number (Na) • Oxidizing agent decreases its oxidation number (H in H2O)

  20. Nuggets of redox processes • Where there is oxidation there is always reduction

  21. Iron reduces Cu2+ to Cu • Iron reduces Cu2+ ions to Cu • Cu does not reduce Fe2+

  22. Applying activity series to metals in acids • Mg is higher than H in activity series – forms H2 • Cu is lower than H in activity series – no H2 produced

  23. Element can be oxidizer and reducer depending on relative positions in activity series • Fe reduces Cu2+ • Cu reduces Ag+ (lower activity) • Fe2+ is reduced by Zn (higher activity)

  24. Combination reactions • Element + element  compound (redox) • Metal + nonmetal  binary ionic compound • Nonmetal + nonmetal  binary covalent compound • Compound + element  compound (redox) • Compound + compound  compound

  25. Decomposition reactions • Compound  element + element (redox) • Compound  element + compound (redox) • Compound  compound + compound

  26. Single replacement (displacement) • Element displaces another element from compound (redox)

  27. Metathesis (double displacement) reactions involve changing partners • AX + BY = AY + BX • Driven by removal of ions from solution • Formation of an insoluble solid (precipitate) • Formation of nonionized molecules (eg H2O) • Acid-base neutralization • Formation of a gas (eg CO2)

  28. Precipitation reactions • Does one of the possible cation-anion combinations produce an insoluble salt? • Initial compounds are all soluble • Use solubility rules to investigate • If yes, a precipitate is produced

  29. Solubility rools OKApplied not remembered

  30. Production of a gas • If product is a gas that has low solubility in water, reaction produces gas • Any carbonate with an acid for example

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