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Reaction Predictions

Reaction Predictions. Rev. 01/23/12. Objectives. SWBAT Review their knowledge of cations/anions, salts, diatomic molecules, strong acids and bases, solubility rules and the activity series. Review the proper format required to write molecular, and net ionic equations.

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Reaction Predictions

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  1. Reaction Predictions Rev. 01/23/12

  2. Objectives • SWBAT • Review their knowledge of cations/anions, salts, diatomic molecules, strong acids and bases, solubility rules and the activity series. • Review the proper format required to write molecular, and net ionic equations. • Practice using Solubility Rules to determine if a reaction occurs.

  3. New Format • You will have a reaction prediction quiz every Wed. until the AP exam. • The first quiz will be a group quiz (groups of 3) • The next two quizzes will be pairs. • The remaining quizzes will be individual.

  4. The AP Way • Always write net ionic equations. • Balance the equations. • There are 3 reaction prediction problems on the exam. • There is an associated question with each reaction prediction question.

  5. Most Commonly Used Cations and Anions • Hydrogen H+ • Sodium Na+ • Potassium K+ • Calcium Ca+² • Magnesium Mg+² • Iron (Ferrous) Fe+² • Iron (Ferric) Fe+³ • Hydroxide OHˉ • Chloride Clˉ • Sulfide Sˉ² • Bicarbonate HCOзˉ • Carbonate COзˉ² • Sulfate SO4ˉ² • Phosphate PO4ˉ ³

  6. Tip • Group 1A cations are spectator ions.

  7. Cations/ Anions, contd. • You can figure out the charge of an ion by using the periodic table. • Ex: • Alkali metals such as Lithium can easily lose an electron to become stable (just like a Noble gas) so taking away an electron give Lithium a +1 charge. • On the other hand Halogens can easily accept an electron to become stable. Accepting an electron gives halogens a -1 charge.

  8. Practice • What is the oxidation state of Oxide? • What is the oxidation state of Iodide? • What is the oxidation state of a Calcium ion? • What is the oxidation state of a Lithium ion?

  9. Answers • -2 • -1 • +2 • +1

  10. Net Ionic Equation • To create a net ionic equation, you break apart all ionic molecules in a balanced molecular equation into their ions if they are soluble. • Do not break up gases or solids !!! • Solids, pure liquids and gases have no charges. • If there are spectator ions, ions that appear on both sides of the equation, they cancel each other.

  11. Net Ionic Example • Silver nitrate is mixed with potassium chromate Molecular Equation2AgNO3 + K2CrO4→ Ag2CrO4 + 2KNO3Complete ionic equation 2Ag+ + 2NO3ˉ + 2K+ + CrO4-2→ Ag2CrO4 + 2K+ + 2NO3-2Net Ionic Equation 2Ag+ + CrO4-2→Ag2CrO4

  12. Solubility Rules • NO3-all nitrates are soluble • CH3COO- or C2H3O2- all acetates are soluble except AgCH3COO-1 • ClO3-all chlorates are soluble • Cl-all chlorides are soluble except AgCl, Hg2Cl2, PbCl2 • Br-all bromides are soluble except AgBr, PbBr2, Hg2Br2, and HgBr2 • I-all iodides are soluble except AgI, Hg2I2, HgI, and PbI2

  13. Solubility Rules, contd. • SO4-2all sulfates are soluble except BaSO4, PbSO4, Hg2SO4, CaSO4, AgSO4 and SrSO4 • Alkali metal cations, and NH4+1 are all soluble • H+all common inorganic acids and low molecular mass organic acids are soluble

  14. (In)Solubility Rules, contd. • CO3-² all carbonates are insoluble except those of alkali metals and NH4 • CrO4-² all chromates are insoluble except those of alkali metals, NH4, CaCrO4, and SrCO4 • OH-all hydroxides are insoluble except those of the alkali metals, NH4, Ba(OH)2, Sr(OH)2, and Ca(OH)2 • PO4-³ all phosphates are insoluble except those of alkali metals and NH4 • SO3-² all sulfites are insoluble except those of alkali metals and NH4 • S-² all sulfides are insoluble except those of alkali metals and NH4

  15. Strong Acids HCl HNO3 HBr H2SO4 HI HClO4

  16. Strong Bases • Generally, strong bases are made up of a metal ion and a hydroxide ion. LiOH Ca(OH)2 NaOH Sr(OH)2 KOH Ba(OH)2

  17. Strong Acids and Bases • Strong acids dissociate completely in water, so the reaction goes to completion and they never reach equilibrium with their conjugate bases. • Because there is no equilibrium, there is no equilibrium constant, so there is no dissociation constant for strong acids and bases.

  18. Salts • Salts are generally formed from a cation (on the far left side of the periodic table) and an anion (on the far right side of the table) • Salts that are soluble in water include all of the salts of: CationsAnionslithium nitratesodium acetatepotassiumammonium Check the solubility rules for other soluble/insoluble salts

  19. Diatomic Molecules • You can use a mnemonic device to help you remember the diatomic molecules: • BrINClHOF (say “Brinkelhof”) • Help Our Needy Class Find Brains Immediately

  20. Synthesis • Synthesis occurs when two or more reactants combine to form a single product. • There are several common types of synthesis reaction.

  21. Synthesis Examples metal combines with a nonmetal to form a binary salt. A piece of lithium metal is dropped into a container of nitrogen gas. 6Li+ N2 2Li3N Metal oxide and water forms a base (metallic hydroxide)   Solid sodium oxide is added to water. Na2O + H2O 2NaOH

  22. Synthesis, contd. • Nonmetal oxide and water forms acids. Nonmetal retains its oxidation number.   -Carbon dioxide is burned in water. CO2 + H2O   H2CO3 • Metallic oxide and nonmetallic oxide form salt.   -Solid sodium oxide is added to carbon dioxide. CaO + SO2 CaSO3

  23. Decomposition • Occurs when a single reactant is broken down into two or more products. • The reactants react to form basic compounds or elements. • When a compound is heated or electrolyzed, it means that it is broken up into its ions. • AB A+B

  24. Decomposition Example Base  metal oxide + water Ca(OH)2 CaO + HOH Acid containing oxygen  non metal oxide + water H2CO3  CO2 + HOH Salt containing oxygen  metal oxide + nonmetal oxide CaCO3  CaO + CO2

  25. Examples of Decomposition • A sample of magnesium carbonate is heated. MgCO3 MgO + CO2 • Molten sodium chloride is electrolyzed. 2NaCl  2Na + Cl2 • A sample of ammonium carbonate is heated. (NH4)2CO3  2NH3 + H2O + CO2

  26. Single Replacement • Reactions that involve an element replacing one part of a compound. The products include the displace element and a new compound. An element can only replace another element that is less active than itself. (Look a activity series/ AP packet) • Generally, synthesis reactions occur when the two reactants come from the following types of compounds:acid, base, salt, water • A +BX B+AX

  27. Single Replacement Active metals replace less active metals from the less active metals’ compounds in aqueous solutions3Mg+ 2FeCl3 —> 2Fe + 3MgCl2 Active metals replace hydrogen in water2Na + 2H2O —> H2 + 2NaOH Active metals replace hydrogen in acids 2Li + 2HCl —> H2 + 2LiCl

  28. Single Replacement Rules, contd. Active nonmetals replace less active nonmetals from their compounds in aqueous solutionsCl2 + 2KI —> I2 + 2KCl If a less reactive element is combined with a more reactive element in compound form, there will be no reactionCl2 + KF —> no reaction* * On the AP test reactions will ALWAYS have products; it will never be “no reaction.”

  29. Activity Series (Single Replacement) • Metals Li, Ca, Na, Mg, Al, Zn, Fe, Pb, [H2], Cu, Ag, Pt  • Nonmetals F2, Cl2, Br2, I2, More active  Less Active

  30. Double Replacement • Two compounds react to form two new compounds. No changes in oxidation numbers occur. • Each cation pairs up with the anion in the other compound. • The “driving force” in these reactions is the removal of at least one pair of ions from solution. • This removal of ions happens with the formation of a precipitate, gas, or molecular species. • When a double replacement reaction doesn’t go to completion, it is a reversible reaction (no ions have been removed). • AX+ BY  AY+ BX

  31. How do you know a double replacement reaction occurs? • One of the products will be a(n): gas insoluble precipitate molecular species *Remember– on the AP test the reaction will always occur

  32. Examples of Dbl. Replacement • Solutions of potassium bromide and silver nitrate are mixed. KBr + AgNO3 AgBr + KNO3 • A solution of sodium sulfate is added to a solution of hydrochloric acid. Na2SO3 + 2HCl  2NaCl + H2SO3 Remember that one of your products must be a solid or a gas.

  33. Common Gases Released (Dbl. Repl.) • H2S Any sulfide plus any acid forms H2S and a salt. • CO2Any carbonate plus any acid form CO3, water, and a salt. • SO2Any sulfite plus any acid form SO2, water, and a salt. • NH3Any ammonium plus a soluble hydroxide form NH3, water, and a salt.

  34. Solubility Computer Tasks • Task #1  • Go to the site listed below which is a tutorial on precipitation reactions. Follow all directions. • Record the complete molecular reaction for each combination of solutions you mix. Identify the precipitate (name and formula) produced. • http: www.wisc-online.com/objects/index_tj.asp?objID=GCH2904

  35. Solubility Computer Tasks • Task # 2 • Go to the site below and mix two solutions at a time. • Write the complete molecular equations for each combination of solutions. • Identify the precipitate (name and formula) that is produced. (You should have three reactions when you are finished.) • http:// www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/crm3s2_3.swf

  36. Acid/ Base Reactions (Dbl. Repl.) • An acid and a base will react and form water and a salt. • Hydrochloric acid is added to sodium hydroxide. HCl + NaOH  NaCl + H2O

  37. Acid/Base Reactions • This is a neutralization reaction. • Remember: • Weak acids & bases do not dissociate in water • Strong acids & bases do dissociate in water

  38. Hydrolysis (Dbl. Repl.) • It is the reverse of neutralization and results when a salt plus a water molecule yields an acid plus a base. • Salt + water   acid + base  • Key things to know about hydrolysis reactions: • Salts of a strong acid plus a weak base will hydrolyze into an acidic solution. • NH4+   +Cl-   +H2O   →    H+   +Cl-   +  (NH)4OH  • Salts of a weak acid and a strong base will always hydrolyze to give a basic solution. • K+   +F-  +H2O →  K+  +OH-   +HF 

  39. Hydrolysis • Salts of a strong acid and a strong base will never undergo hydrolysis and therefore make a neutral solution. Na+   + Cl-1  + H2O   →   Na+1  + OH-1   + H+1 +  Cl-1 • Salts of a weak acid plus salts of a weak base may hydrolyze as an acid, base, or a neutral solution; the final result depends on the Ka’s and Kb’s of the acids and bases formed during the hydrolysis process. • Disclaimer!! The spectator ions were not removed 

  40. Hydrolysis Sample Problems • Try these: • An aqueous solution of manganese (II) sulfate undergoes hydrolysis. • Ammonium fluoride and water are mixed together. • Answers are on the next slide.

  41. Hydrolysis answers • MnSO4 + 2H2O →  H2SO4 + Mn(OH)2 • NH4F + H2O → HF + NH4OH

  42. Gas Producing Reactions Acid + Metal carbonate → Salt + Water + Carbon dioxide H2SO4(aq) + MgCO3(s) → MgSO4(aq) + H2O(l) + CO2(g)

  43. Combustion (Organic Rxns.) • An organic compound reacts with O2 and an ignition source to form water and carbon dioxide. • If something is burned there is a combustion reaction. • Methanol is burned in oxygen gas. 2CH3OH + 3O2 4H2O + 2CO2

  44. Combustion • On the AP exam: • If the reaction is a combustion reaction and you don’t know the chemical formula for the hydrocarbon, make up a chemical formula and complete the reaction. You will earn partial points.

  45. Addition (Organic Rxns) • A halogen, or hydrogen, is added to an alkene or alkyne, breaking apart the double or triple bonds and forming one compound with single bonds. • This is an addition to multiple bonds. • Fluorine is added to ethene F2 + CH2=CH2 CH2F-CH2F

  46. Substitution (Organic Rxns.) • An atom attached to a carbon is removed and something else takes its place. Classic reaction for organic cpds with no multiple bonds. • This is a substitution of single bonds. • Bromine is added to methane Br2 + CH4 CH3Br + HBr

  47. Esterification (ether formation) • An organic acid reacts with an alcohol to form an ester and water. organic acid + alcohol  ester + water CH3CH2COOH + CH3CH2CH2CH2OH  CH3CH2COOCH2CH2CH2CH3 + H20

  48. Ether Formation alcohol + alcohol  ether + water C2H5OH + C2H5OH  C2H5OC2H5 + H2O

  49. Oxidizing Agents (Redox Rxns.) Common Oxidizing Agents MnO4¯ in acidic solution • MnO2 in acidic solution • MnO4¯ in neutral or basic solution • Cr2O7ˉ² in acidic solution • HNO3, concentrated • HNO3, dilute • H2SO4, hot, concentrated • Metallic ions (higher oxidation #) • Free halogens • Na2O2 • HClO4 • C2O4ˉ² • H 2O2 Products Formed Mn +² Mn +² MnO2(s) Cr +³ NO2 NO SO2 Metallous ions (lower oxidation #) Halide ions NaOH Clˉ CO2 O2

  50. Handy Hint • If the reaction prediction shows a chemical formula, such as, manganese (IV) oxide as a reactant, the reaction is probably an oxidation/reduction reaction.

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