1. Net Ionic equations §A solution of Barium Chloride reacts with a solution of sodium sulfate §BaCl 2 + Na 2 SO 4 → §BaCl 2 + Na 2 SO 4 → BaSO 4 + 2NaCl §Ba 2+ + 2Cl 1- + 2Na 1+ + SO 4 2- → BaSO 4 + 2Na 1+ 2Cl 1- §Ba 2+ + SO 4 2- → BaSO 4

2. Using Stoichiometry in the Laboratory Na 2 SO 3(aq) + 2AgNO 3(aq)  Ag 2 SO 3(s) + 2NaNO 3(aq) §20 ml 20 ml §0.12 M 0.35 M §1. How many grams of silver sulfite would be produced? §2. What would be the molarity of the sodium nitrate? §3.What volume of silver nitrate is needed to completely precipitate all of the sulfite ions.

3. Patterns of Reactivity Substance + Oxygen gas  Oxide of element Active metals and oxygen: Li(s) + O 2  Li 2 O(s) sulfides and oxygen: ZnS(s) + O 2  ZnO + SO 2 (g) CS 2 (s) + O 2  CO 2 (g) + SO 2 ammonia and oxygen: NH 3 + O 2  NO 2 + H 2 O(g) 

4. Substances reacting with water Metallic oxide reacting with water  Base Li 2 O + HOH  Li + + OH - MgO (s) + HOH  Mg(OH) 2 Metal hydrides reacting with water  Base + Hydrogen LiH(s) + HOH  Li + + OH - + H 2 CaH 2 (s) + HOH  Ca 2+ + OH - + H 2

5. Nonmetal halides reacting with water  weak acid + strong acid (Keep ox # the same) PCl 5 (s) + H 2 O  H 3 PO 4 (aq) + H + + Cl - PBr 3 (s) + H 2 O  H 3 PO 3 (aq) + H + + Br - Nonmetal oxide reacting with water  Acid (Keep ox # the same) SO 2 (g) + H 2 O  H 2 SO 3 (aq) N 2 O 5 + H 2 O  H + + NO 3 1- SO 3 + H 2 O  H + + HSO 4 - N 2 O 3 + H 2 O  HNO 2 (aq)

6. Hydrocarbon and oxygen: CH 3 OH(s) + O 2  CO 2 + H 2 O Substances reacting with nonmetal oxide  salt Metal oxide and nonmetal oxide yields a salt CaO(s) + CO 2  CaCO 3 (s) MgO(s) + SO 2  MgSO 3

7. Base reacting with nonmetal oxide: OH - +CO 2 (g)  CO 3 2- +H 2 O (base is aqueous) NaOH(aq) + CO 2 (g)  Na + + HCO 3 - + H 2 O Ca(OH) 2 (aq) + SO 2 (g)  Ca + + HSO 3 - + H 2 O Ca(OH) 2 (s) + SO 2 (g)  CaSO 3 (s) + H 2 O

8. Metal sulfides reacting with water  Base + gas Fe 2 S 3 (s) + HOH  Fe(OH) 3 (s) + H 2 S(g)

9. Metal carbides reacting with water  Base + hydrocarbon Na 2 C 2 (s) + H 2 O  Na + + OH - + C 2 H 2 Salt of an amphoteric metal reacting with water  complex Al(NO 3 ) 3 (s) + H 2 O  Al(H 2 O) 6 3+ + NO 3 -

10. Substances reacting with an acid metals reacting with acid  salt + hydrogen Zn(s) + H +  Zn 2+ + H 2 Cu(s) + H + + HSO 4 -  Cu 2+ + SO 2 + H 2 O (hot conc. sulfuric acid) Ca(s) + H +  Ca 2+ + H 2 (cold sulfuric acid) Ag(s) + H + + NO 3 -  Ag + + NO + H 2 O (dilute (6M) nitric acid) Ag(s) + H + + NO 3 -  Ag + + NO 2 + H 2 O (conc. nitric acid)

11. metal oxide reacting with hydrogen gas  salt + water FeO(s) + H + + NO 3 -  Fe 3+ + NO 2 + H 2 O (conc. nitric acid) metal oxide reacting with acid  salt + gas + water Fe 2 O 3 (s) + H 2  Fe + H 2 O

12. Salt of a weak acid and strong acid  salt of strong acid + weak acid C 2 H 3 O 2 - + H +  HC 2 H 3 O 2 C 2 H 3 O 2 - + H + + HSO 4 -  HC 2 H 3 O 2 + HSO 4 - (equimolar sulfuric acid) SO 2 2- + H +  H 2 SO 3  H 2 O + SO 2 Na 2 S(s) + H +  Na + + Cl - + H 2 S

13. Substances reacting with an Base Al and Zn hydroxides with a strong base Al(OH) 3 + OH -  [Al(OH) 6 ] 3- Zn(OH) 2 + OH -  [Zn(OH) 4 ] 2- salt of an amphoteric metal reacting with a strong acid salt of a weak acid reacting with a strong base NH 4 Cl + OH -  NH 3 (g) + H 2 O + Cl - HCO 3 - + OH -  CO 3 -2 (g) + H 2 O Al(OH) 3 + H +  Al 3+ + H 2 O

15. Complex ions Zn 2+ + NH 3 (excess)  [Zn(NH 3 ) 4 ] 2+ Zn(OH) 2(s) + NH 3 (ex)  [Zn(NH 3 ) 4 ] 2+ + OH - Ag +1 + NH 3 (excess)  [Ag(NH 3 ) 2 ] + Cu(OH) 2 + NH 3 (ex)  [Cu(NH 3 ) 4 ] 2+ + OH - Fe +3 + NCS 1- → FeNCS 2+ orange/brown blood red

16. Oxidation-Reduction Single Replacement (metals) Al(s) + Cu 2+ (aq)  Cu(s) + Al 3+ (aq) Zn(s) + Sn 2+ (aq)  Sn(s) + Zn 2+ (aq) H 2 (g) + CuO(s)  Cu(s) + H 2 O (aq) Single Replacement (halogens) F 2 (g) + Cl - (aq)  Cl 2 (g) + F - (aq)

17. Oxidation-Reduction Redox in an Acid Environment Cr 2 O 7 2- + Fe 2+ + H +  Cr 3+ + Fe 3+ + H 2 O Cr 2 O 7 2- + I - + H +  Cr 3+ + I 2 + H 2 O MnO 4 - + Cl - + H +  Mn 2+ + Cl 2 + H 2 O H 2 O 2 + I - + H +  H 2 O + I 2 + H 2 O

18. Oxidation-Reduction Redox in a Basic Environment H 2 O 2 (aq) + MnO 4 -  O 2 + MnO 2 (s) + H 2 O

19. Oxidation-Reduction Important Oxidizers MnO 4 - in acidic solutionMn 2+ Formed in the Reaction MnO 2 in acidic solutionMn 2+ MnO 4 - in basic solutionMnO 2 Cr 2 O 7 2- in acidic solutionCr 3+ HNO 3, concentrated NO 2 HNO 3, dilute NO H 2 SO 4, hot SO 2

20. Oxidation-Reduction Important OxidizersFormed in the Reaction Metal-ic ionsMetal-ous ions Free halogenshalide ions Na 2 O 2 NaOH HClO 4 Cl - H 2 O 2 (in acidic sol’n)H2OH2O H 2 O 2 (in basic sol’n)OH -

21. Important ReducersFormed in the Reaction Halide ionsFree halogens Oxidation-Reduction Free metalsMetal ions Sulfite ions (or SO 2 )Sulfate ions Nitrite ionsNitrate ions Free Halogens (dil., basic)Hypohalides (e.g. ClO - ) Free Halogens (conc., basic)Halate ions (e.g. ClO 3 - ) Metal-ousMetal-ic H2O2H2O2 O2O2