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Warm-up

Warm-up. Determine the type of reaction and predict the products: NaOH  Li + Br 2  C 2 H 4 + O 2 . Activity Series . How to use the activity series. Find the element in the compound on the table. Find the Solo element

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Warm-up

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  1. Warm-up • Determine the type of reaction and predict the products: NaOH Li + Br2  C2H4 + O2 

  2. Activity Series • How to use the activity series. • Find the element in the compound on the table. • Find the Solo element • If the solo element is above the element in the compound then the reaction will take place.

  3. 2Al(s) + 3ZnCl2(aq) To replace the Zinc, Aluminum must be higher on the series Cu(s) + 2NaCl(aq) Can copper replace sodium in the compound? 3Zn(s) + 2AlCl3(aq) NO REACTION

  4. Activity Series for Halogens • Above the activity series for metals, there is an activity series for Halogens. • If your solo element is a halogen, it will replace the bonded halogen as long as it is above it on the activity series. • Remember, every halogen on the series is a diatomic molecule, so when it’s by itself, there will be two of them (F2, Br2, …)

  5. Practice! Cr(NO3)2(aq) + Pb(s) • Cr(s) + Pb(NO3)2(aq) • Pt(s) + CaCl2(aq) • Ca(s) + FeO(aq) NO REACTION CaO(aq) + Fe(s)

  6. Warm-up • Determine the type of reaction and predict the products: NaOH Li + Br2  C2H4 + O2 

  7. Predicting Products:Double Displacement Unit 6, Day 5 Kimrey 1 November

  8. Remember Double Displacement • Anions switch places and are each bonded to a different cation • AB + CD  AD + CB

  9. Predicting the Products of Double Displacement • Involves determining charges, criss-crossing, and the solubility rules

  10. Why do solubility rules matter? • All double displacement reactions (in this unit) will produce a precipitate • A precipitate is a solid that’s produced during a chemical reaction in a solution • So, if a precipitate is not formed, then the reaction will not take place!! • We can determine if a precipitate is formed by looking at our solubility rules

  11. Solubility • If something is soluble, then it can be dissolved by what it’s bonded to • If something is insoluble, then it cannot be dissolved

  12. What does it mean for us? • If one of your products is insoluble, then its state of matter is solid and a precipitate has formed. • If one of your products is soluble, then its state of matter is aqueous and no precipitate has formed. • You must have at least one solid product for a reaction to occur.

  13. Solubility Rules: Soluble • Soluble • All Nitrates, Acetates, Ammoniums, and Group 1 salts. • All Chlorides, Bromides, and Iodides, except Silver, Lead, and Mercury (I) • All Fluorides except Group 2, Lead (II), and Iron (III) • All Sulfates except Calcium, Strontium, Barium, Mercury, Lead(II), and Silver

  14. Solubility: Insoluble • Insoluble • All Carbonates and Phosphates except Group 1 and Ammonium • All Hydroxides except Group 1, Strontium, Barium , and Ammonium • All Sulfides except Group 1, Group 2, and Ammonium • All Oxides except Group 1

  15. Steps • First break the reactants into their ions (find the charges!). • Next, swap partners for both (OI with a twist) • Check solubility rules to see if a solid (precipitate) has formed. • Write complete balanced equation with states of matter.

  16. Example • Sodium Hydroxide + Copper (II) Sulfate • What are the Ions? • What are the reactants? • What are the potential products? • Are any potential products insoluble? • What is the complete equation

  17. PracticePredict the products and determine if a precipitate forms. • Sodium phosphate + Nickel (II) chloride • NaCl and Ni3(PO4)2. • Lead (II) Nitrate + Potassium Iodide • PbI2 and KNO3 • Sodium Hydroxide + Potassium Chloride • NaCl and KOH • Sodium phosphate + Lead (IV) nitrate • Pb3(PO4)4and NaNO3

  18. Writing molecular equations • You already know how to do this!  • This is the chemical equation with the states of matter in it. • Make sure it’s balanced!

  19. Writing the Net Ionic equation • You almost know how to do this! • Start with the completely balanced equation. • Look at the solid product and make it the product of your Net Ionic equation. • For the reactants, put the ions that lead to the product

  20. Example • Na2SO4 + CaCl2 2NaCl+ CaSO4 • SO42-(aq)+ Ca2+(aq)CaSO4 (s) • 3NaOH + FeCl3 3NaCl + Fe(OH)3 • 3OH-(aq) + Fe3+(aq) Fe(OH)3 (s)

  21. Net Ionic equation • Take the complete ionic equation and remove the spectator ions. • Spectator ions are the ions not involved in the reaction. • Ex. Na2SO4 + CaCl2 2NaCl+ CaSO4 • CaSO4 (s)+ 2NaCl(aq) • SO42-(aq)+ Ca2+(aq)  CaSO4 (s)

  22. Practice • NaCl + AgNO3 AgCl +NaNO3 • Na+(aq) + Cl-(aq) + Ag+(aq) + NO3-(aq)AgCl(s) + Na+(aq) + NO3-(aq) • Ag+(aq) + Cl-(aq)AgCl(s) • 2NaOH + CuSO4 Cu(OH)2 + Na2SO4 • 2Na+(aq) + 2OH-(aq)+ Cu2+(aq)+ SO42-(aq) Cu(OH)2(s) + 2Na+(aq)+ SO42-(aq) • Cu2+(aq)+ 2OH-(aq) Cu(OH)2 (s)

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