1 / 78

1. Electrolytes – Completely Break up into ions in water (Arrhenius, 1884 (Nobel Prize, 1903)).

Types of Solutions. 1. Electrolytes – Completely Break up into ions in water (Arrhenius, 1884 (Nobel Prize, 1903)). a. Many Ionic Compounds and strong acids (HCl, HBr, HI, HNO 3 , H 2 SO 4 ,HClO 4 ) b. Different than decomposition because ions are produced. Types of Solutions.

vivien
Télécharger la présentation

1. Electrolytes – Completely Break up into ions in water (Arrhenius, 1884 (Nobel Prize, 1903)).

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Types of Solutions 1. Electrolytes – Completely Break up into ions in water (Arrhenius, 1884 (Nobel Prize, 1903)). a. Many Ionic Compounds and strong acids (HCl, HBr, HI, HNO3, H2SO4,HClO4) b. Different than decomposition because ions are produced.

  2. Types of Solutions c. Examples: NaCl(s)  Na+(aq) + Cl-(aq) CaCl2(s)  Ca2+(aq) + 2Cl-(aq) Al2(SO4)3(s)  2Al3+(aq) + 3SO42-(aq)

  3. Types of Solutions d. Examples Na2CO3(s)  (NH4)2Cr2O7(s)  HCl(l)  FeCl3(s)  e. Hydration Sphere for NaCl

  4. Types of Solutions 2. Weak Electrolytes a. Weak Acids b. Examples HC2H3O2, HF, HNO2

  5. Types of Solutions 3. Non-Electrolytes- Do not break up into ions in water a. Many Molecular Compounds C12H22O11(s)  C12H22O11(aq)

  6. Types of Reactions

  7. Types of Solutions

  8. Solubility Rules • Provide a rough idea of whether something will dissolve in water • DO NOT GIVE ACTUAL, NUMERICAL SOLUBILITIES (you must look in a book or do an experiment)

  9. Solubility Rules

  10. Solubility Rules Examples: Na2CO3(s) Zn(OH)2 (s)  Na2S (s)  CaCl2(s)  AgCl(s)  CuCO3 (s)

  11. Solubility Rules Examples: PbSO4(s) Ag2SO4 KCl(s)  Fe(OH)3(s) FeSO4(s)

  12. Net Ionic Rxns • Can be Double Replacement Rxns • Spectator Ions – Ions present in soln, but do not take part in the rxn Pb(NO3)2(aq) + KI(aq) 

  13. BaCl2(aq)+K2SO4(aq)

  14. Net Ionic Rxns Practice: AgNO3(aq) + NaCl(aq) 

  15. CaCl2(aq) + Na2CO3(aq)  NaNO3(aq) + NH4OH(aq)  Na2SO4(aq) + BaBr2(aq)  Fe2(SO4)3(aq) + LiOH(aq)  Na2S(aq) + CuCl2(aq)  Pb(C2H3O2)2(aq) + NH4OH(aq) 

  16. Net Ionic Rxns • The driving force for many reactions is the formation of a: a) Solid (precipitate) b) Liquid c) Gas

  17. Water forming Rxns • A special type of double replacement - neutralization • Acids a. produce H+ b. Often start with H (HCl) • Bases a. produce OH- b. Hydroxides (Drano, NaOH)

  18. NaOH (aq) + HCl (aq) HClO4 (aq) + LiOH (aq)  Ca(OH)2(aq) + HNO3 (aq)  Mg(OH)2(s) + HCl(aq) 

  19. Gas forming Rxns • Carbonates plus acids • Carbonic acid – unstable (in soda) H2CO3(aq)  H2O(l) + CO2(g) 3. Examples: CaCO3(s) + HCl(aq)  MgCO3(s) + HNO3(aq) 

  20. NaHCO3(aq) + HNO3(aq)  Na2S(aq) + HCl(aq) 

  21. Net Ionic Rxns Mixed Types Fe(NO3)3(aq) + Na2CO3(aq) SrCO3(s) + HCl(aq)  HBr(aq) + LiOH (aq) 

  22. Two Types of Chemical Rxns • Exchange of Ions – no change in charge/oxidation numbers • Acid/Base Rxns NaOH + HCl

  23. Precipitation Rxns Pb(NO3)2(aq) + KI(aq) • Dissolving Rxns CaCl2(s) 

  24. Exchange of Electrons – changes in oxidation numbers/charges Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s) Remove spectator ions Cu(s) + 2Ag+(aq)  Cu2+(aq) + 2Ag(s)

  25. Oxidation Numbers • Involves taking compounds apart • Oxidation numbers – Pretend charges for all compounds (as if they exist as a monoatomic ion) • Rules

  26. Fe H2 P4 Cl2 Elements = 0 Monoatomic Ions = Charge Na+ O2- Al3+ Use “bankables” to calculate the rest H2S Cl2O Na2SO4 Fe2O3 PO43- NO3- CaCr2O7 SnBr4 Gr I Gr II O-2 H+ F- “the higher the oxidation #, the more oxidized the element”

  27. Review of Oxidation Numbers Calculate the oxidation numbers for: HClO Cr3+ S8 Fe2(SO4)3 Mn2O3 SO32- KMnO4 NO3- HSO4-

  28. Oxidation • Classical Definition –addition of oxygen Fe + O2 Fe2O3 • Modern Definition – an increase in oxidation number Na + O2 Na2O 0 +1 Na was oxidized

  29. Reduction • Classical Definition –addition of hydrogen N2 + 3H2 2NH3 (Haber process) R-C=C-R + H2  | | H H (unsaturated fat) (saturated fat)

  30. Reduction • Modern Definition –decrease (reduction) in oxidation number N2 + 3H2 2NH3 0 -3 N was reduced

  31. Example In the following rxns, which element is oxidized, which is reduced? Al + HBr  AlBr3 + H2 Fe + Cu(NO3)2  Fe(NO3)2 + Cu H2 + O2  H2O

  32. Activity Series • Used to predict if a particular redox reaction will occur • Redox reactions - also called single replacement reactions • Not every element can replace every other • Higher elements get oxidized • Lower elements get reduced

  33. Will Copper metal replace silver in an aqueous solution of silver nitrate?

  34. Will aqueous iron(II)chloride oxidize magnesium metal?

  35. Can aluminum foil reduce Fe(NO3)2 to iron metal?

  36. Can aluminum foil react with HCl?

  37. Which of the following metals will be oxidized by Pb(NO3)2: Zn, Cu, and/or Fe?

  38. Will barium metal react with nickel(II)nitrate? Will iron(II)chloride react with calcium metal? Will aluminum chloride react with gold? Will calcium metal dissolve in HNO3?

  39. MgCO3(s) + HNO3(aq)  Zn(NO3)2 (aq) + Ag(s)  CuCl2(s)  (placed in water) K(s) + NiCl2(aq)  Sn(s) + CuCl2(aq)  PbSO4(s)  (placed in water) Fe(s) + HCl (aq)  Mg(OH)2(s) + HCl(aq) 

  40. Identifying Oxidizing/Reducing agents Oxidizing agent – gets reduced Reducing agent – get oxidized K + ZnCl2  AgNO3 + Ni Li + CaCl2  Cr(NO3)3 + Na 

More Related