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 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? How do we tell? Rules to predict solubility Rules to predict solubility  Covalent molecular compounds* are non- electrolytes – no ions produced *Except acids and bases *Except acids and bases

5. Dissociation and ionization: same or different?  Ionic compounds dissociate in water Ions already exist in the solid Ions already exist in the solid  Acids or bases* ionize in water A pure acid or base contains no ions A pure acid or base contains no ions *Except strong bases like NaOH, Ca(OH) 2 are ionic *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) HCl, HBr, HI (but not HF) HNO 3 (but not HNO 2 ) HNO 3 (but not HNO 2 ) H 2 SO 4 (but not H 2 SO 3 ) H 2 SO 4 (but not H 2 SO 3 ) HClO 4 (maybe HClO 3 ) HClO 4 (maybe HClO 3 )  All other acids are weak

9. Recognizing acids  Mineral acids: HCl, HNO 3 etc. Conventionally H appears first in the formula Conventionally H appears first in the formula All strong acids are mineral All strong acids are mineral May be strong or weak May be strong or weak  Organic acids: CH 3 COOH etc Harder to spot Harder to spot Sometimes written with H in front – HCH 3 CO 2 Sometimes written with H in front – HCH 3 CO 2 Always weak Always weak Presence of –OH (-SH): necessary but not sufficient Presence of –OH (-SH): necessary but not sufficient Not all –OH are acidic (CH 3 OH is not an acid) Not all –OH are acidic (CH 3 OH is not an acid)

10. Recognizing bases  Mineral bases usually distinguished by OH groups – all strong NaOH, Ca(OH) 2 NaOH, Ca(OH) 2  Ammonia, NH 3, 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) = H 2 O(l) + NaCl(aq) Mg(OH) 2 (s) + 2HCl(aq) = MgCl 2 (aq) + 2H 2 O(l)  Salt contains anion of acid and cation of base: HCl + NaOH = NaCl + H 2 O HCl + KOH = KCl + H 2 O HNO 3 + KOH = KNO 3 + H 2 O 2HCl + Ca(OH) 2 = CaCl 2 + 2H 2 O HCN + NaOH = NaCN + H 2 O

13. Acid-base reaction with gas formation  Tums... HCl(aq) + NaHCO 3 (aq) = NaCl(aq) + H 2 CO 3 (aq)  H 2 CO 3 is unstable: H 2 CO 3 (aq) = H 2 O(l) + CO 2 (g)  Bad egg gas: 2HCl + Na 2 S = H 2 S(g) + 2NaCl(aq)

14. Oxidation - reduction  Oxidation is loss of electrons  Reduction is gain of electrons Oxidation is always accompanied by reduction Oxidation is always accompanied by reduction The total number of electrons is kept constant The total number of electrons is kept constant  Oxidizing agents oxidize and are themselves reduced  Reducing agents reduce and are themselves oxidized

16. 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.  Oxidation number is the number of electrons gained or lost by the element in making a compound

17. Predicting oxidation numbers 1.Oxidation number of atoms in element is zero 2.Oxidation number of element in monatomic ion equals charge 3.Sum of oxidation numbers in compound is zero 4.Sum of oxidation numbers in polyatomic ion equals charge F has ON –1 F has ON –1 H has ON +1; except in metal hydrides where it is –1 H has ON +1; except in metal hydrides where it is –1 Oxygen is usually –2. Exceptions: Oxygen is usually –2. Exceptions: O is –1 in hydrogen peroxide, and other peroxides O is –1 in hydrogen peroxide, and other peroxides O is –1/2 in superoxides KO 2 O is –1/2 in superoxides KO 2 In OF 2 O is +2 In OF 2 O is +2

18. Position of element in periodic table determines oxidation number G1A is +1 G1A is +1 G2A is +2 G2A is +2 G3A is +3 (some rare exceptions) G3A is +3 (some rare exceptions) G5A are –3 in compounds with metals, H or with NH 4 + G5A are –3 in compounds with metals, H or with NH 4 + Exceptions are compounds with elements to right (e.g. NO 2, PF 5 ); in which case use rules 3 and 4. Exceptions are compounds with elements to right (e.g. NO 2, PF 5 ); in which case use rules 3 and 4. G6A below O (S, Se etc.) are –2 in binary compounds with metals, H or NH 4 + G6A below O (S, Se etc.) are –2 in binary compounds with metals, H or NH 4 + When combined with O or lighter halogen (e.g. SeO 2, SF 6 ) use rules 3 and 4. When combined with O or lighter halogen (e.g. SeO 2, SF 6 ) use rules 3 and 4. G7A elements are –1 in binary compounds with metals, H or NH 4 + or with a heavier halogen (e.g. Cl in BrCl 3 ) G7A elements are –1 in binary compounds with metals, H or NH 4 + or with a heavier halogen (e.g. Cl in BrCl 3 ) When combined with O or a lighter halogen, use rules 3 and 4 (e.g. Br in BrCl 3 or Cl in ClO 4 - ). When combined with O or a lighter halogen, use rules 3 and 4 (e.g. Br in BrCl 3 or Cl in ClO 4 - ).

19. 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.

20. Identify redox by change in oxidation numbers  Reducing agent increases its oxidation number (Na)  Oxidizing agent decreases its oxidation number (H in H 2 O)

21. Nuggets of redox processes  Where there is oxidation there is always reduction Oxidizing agent Reducing agent Is itself reduced Is itself oxidized Gains electrons Loses electrons Causes oxidation Causes reduction

22. Iron reduces Cu 2+ to Cu  Iron reduces Cu 2+ ions to Cu  Cu does not reduce Fe 2+

23. Applying activity series to metals in acids  Mg is higher than H in activity series – forms H 2  Cu is lower than H in activity series – no H 2 produced

24. Element can be oxidizer and reducer depending on relative positions in activity series  Fe reduces Cu 2+  Cu reduces Ag + (lower activity)  Fe 2+ is reduced by Zn (higher activity)

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

26. Decomposition reactions  Compound  element + element (redox)  Compound  element + compound (redox)  Compound  compound + compound

27. Single replacement (displacement)  Element displaces another element from compound ( redox)

28. 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 an insoluble solid (precipitate) Formation of nonionized molecules (eg H 2 O) Formation of nonionized molecules (eg H 2 O) Acid-base neutralization Acid-base neutralization Formation of a gas (eg CO 2 ) Formation of a gas (eg CO 2 )

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

30. Solubility rools OK Applied not remembered

31. 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