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Net Ionic equations

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+ + 2Cl 1- + 2Na 1+ + SO 4 2- → BaSO 4 + 2Na 1+ 2Cl 1-

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Net Ionic equations

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  1. Net Ionic equations • A solution of Barium Chloride reacts with a solution of sodium sulfate • BaCl2 + Na2SO4→ • BaCl2 + Na2SO4→ BaSO4 + 2NaCl • Ba2+ + 2Cl1- + 2Na1+ + SO42-→ BaSO4 + 2Na1+ 2Cl1- • Ba2+ + 2Cl1- + 2Na1+ + SO42-→ BaSO4 + 2Na1+ 2Cl1- • Ba2+ + SO42-→ BaSO4

  2. Using Stoichiometry in the Laboratory Na2SO3(aq) + 2AgNO3(aq) Ag2SO3(s) + 2NaNO3(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) + O2 Li2O(s) sulfides and oxygen: ZnS(s) + O2 ZnO + SO2(g) CS2(s) + O2 CO2(g) + SO2 ammonia and oxygen: NH3 + O2 NO2 + H2O(g) 

  4. Substances reacting with water Metallic oxide reacting with water  Base Li2O + HOH  Li+ + OH- MgO (s) + HOH  Mg(OH)2 Metal hydrides reacting with water  Base + Hydrogen LiH(s) + HOH  Li+ + OH- + H2 CaH2 (s) + HOH  Ca2+ + OH- + H2

  5. Nonmetal oxide reacting with water  Acid (Keep ox # the same) SO2(g) + H2O  H2SO3(aq) SO3 + H2O  H+ + HSO4- N2O3 + H2O  HNO2(aq) N2O5 + H2O  H+ + NO31- Nonmetal halides reacting with water  weak acid + strong acid (Keep ox # the same) PCl5 (s) + H2O  H3PO4(aq) + H+ + Cl- PBr3 (s) + H2O  H3PO3(aq) + H+ + Br-

  6. Hydrocarbon and oxygen: CH3OH(s) + O2 CO2 + H2O Substances reacting with nonmetal oxide  salt Metal oxide and nonmetal oxide yields a salt CaO(s) + CO2 CaCO3(s) MgO(s) + SO2 MgSO3

  7. Base reacting with nonmetal oxide: OH- +CO2(g) CO32- +H2O (base is aqueous) Ca(OH)2 (s) + SO2(g)  CaSO3(s) + H2O NaOH(aq) + CO2(g)  Na+ + HCO3- + H2O Ca(OH)2 (aq) + SO2(g)  Ca+ + HSO3- + H2O

  8. Metal sulfides reacting with water  Base + gas Fe2S3(s) + HOH  Fe(OH)3(s) + H2S(g)

  9. Metal carbides reacting with water  Base + hydrocarbon Na2C2(s) + H2O  Na+ + OH- + C2H2 Salt of an amphoteric metal reacting with water  complex Al(NO3)3(s) + H2O  Al(H2O)63+ + NO3-

  10. Substances reacting with an acid metals reacting with acid  salt + hydrogen Zn(s) + H+ Zn2+ + H2 Ca(s) + H+  Ca2+ + H2 (cold sulfuric acid) Cu(s) + H+ + HSO4- Cu2+ + SO2 + H2O (hot conc. sulfuric acid) Ag(s) + H+ + NO3- Ag+ + NO + H2O (dilute (6M) nitric acid) Ag(s) + H+ + NO3- Ag+ + NO2 + H2O (conc. nitric acid)

  11. metal oxide reacting with hydrogen gas  salt + water Fe2O3(s) + H2 Fe + H2O metal oxide reacting with acid  salt + gas + water FeO(s) + H+ + NO3- Fe3+ + NO2 + H2O (conc. nitric acid)

  12. Salt of a weak acid and strong acid  salt of strong acid + weak acid Na2S(s) + H+ Na+ + Cl- + H2S SO22- + H+  H2SO3  H2O + SO2 C2H3O2- + H+  HC2H3O2 C2H3O2- + H+ + HSO4- HC2H3O2 + HSO4-(equimolar sulfuric acid)

  13. Substances reacting with an Base salt of an amphoteric metal reacting with a strong acid Al(OH)3 + H+  Al3+ + H2O 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 a weak acid reacting with a strong base NH4Cl + OH-  NH3(g) + H2O + Cl- HCO3- + OH-  CO3-2 (g) + H2O

  14. Complex ions Zn2+ + NH3(excess) [Zn(NH3)4]2+ Zn(OH)2(s) + NH3(ex) [Zn(NH3)4]2+ + OH- Ag+1 + NH3(excess) [Ag(NH3)2] + Cu(OH)2 + NH3(ex) [Cu(NH3)4]2+ + OH- Fe+3 + NCS1-→ FeNCS2+ orange/brown blood red

  15. Oxidation-Reduction Single Replacement (metals) Al(s) + Cu2+ (aq)  Cu(s) + Al3+ (aq) Zn(s) + Sn2+ (aq)  Sn(s) + Zn2+ (aq) H2(g) + CuO(s) Cu(s) + H2O(aq) Single Replacement (halogens) F2(g) + Cl- (aq) Cl2(g) + F- (aq)

  16. Oxidation-Reduction Redox in an Acid Environment Cr2O72- + Fe2+ + H+ Cr3+ + Fe3+ + H2O Cr2O72- + I- + H+ Cr3+ + I2+ H2O MnO4- + Cl- + H+ Mn2+ + Cl2+ H2O H2O2 + I- + H+ H2O + I2+ H2O

  17. Oxidation-Reduction Redox in a Basic Environment H2O2 (aq) + MnO4- O2 + MnO2(s) + H2O

  18. Oxidation-Reduction Important Oxidizers Formed in the Reaction MnO4- in acidic solution Mn2+ MnO2 in acidic solution Mn2+ MnO4- in basic solution MnO2 Cr2O72- in acidic solution Cr3+ NO2 HNO3, concentrated NO HNO3, dilute SO2 H2SO4, hot

  19. Oxidation-Reduction Important Oxidizers Formed in the Reaction Metal-ic ions Metal-ous ions Free halogens halide ions Na2O2 NaOH HClO4 Cl- H2O2 (in acidic sol’n) H2O H2O2 (in basic sol’n) OH-

  20. Oxidation-Reduction Important Reducers Formed in the Reaction Halide ions Free halogens Free metals Metal ions Sulfite ions (or SO2) Sulfate ions Nitrite ions Nitrate ions Free Halogens (dil., basic) Hypohalides (e.g. ClO-) Free Halogens (conc., basic) Halate ions (e.g. ClO3-) Metal-ous Metal-ic H2O2 O2

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