2. Chemical Equations & Reactions

3. Chemical Equations This equation means: 4 Al(s) + 3 O 2 (g) 2 Al 2 O 3 (s) 4 Al atoms + 3 O 2 molecules yield 2 molecules of Al 2 O 3 4 Al moles + 3 O 2 moles yield 2 moles of Al 2 O 3 or 4 g Al + 3 g O 2 yield 2 g Al 2 O 3 4 mol Al@27g/mol 3 mol O 2 @32g/mol 2 mol Al 2 O 3 @102g/mol 108 g + 96 g = 204 g

4. ? Visualizing a Chemical Reaction Na + Cl 2 NaCl ___ mole Cl 2 ___ mole NaCl___ mole Na 2 105 2 5

5. Types of Chemical Reactions Synthesis (combination) reaction Decomposition reaction A Single-replacement reaction B Double-replacement reaction Combustion reaction (of a hydrocarbon) A + B  AB AB  A + B A + BC  AC + B AB + CD  AD + CB C x H y + O 2  CO 2 + H 2 O A use activity series to predict products/reactivity B use solubility chart to predict products/reactivity elementcompoundelementcompound All compounds…

6. Practice: Balance and Classify 1. Ca(OH) 2 + HCl  CaCl 2 + H 2 O________________________ 2. C 2 H 4 + O 2  CO 2 + H 2 O________________________ 3. N 2 + O 2  N 2 O________________________ 4. SrCO 3  SrO + CO 2 ________________________ 5. NaI + Br 2  NaBr + I 2 ________________________ 6. C 2 H 4 O + O 2  CO 2 + H 2 O________________________ 7. MgBr 2 + (NH 4 ) 2 SO 3  MgSO 3 + NH 4 Br________________________ 8. AgClO 3 + (NH 4 ) 2 Cr 2 O 7  Ag 2 Cr 2 O 7 + NH 4 ClO 3 ______________________ 9. Cs + H 2 O  CsOH + H 2 ________________________ 10. Fe + O 2  Fe 3 O 4 ________________________

7. Symbols Used in Chemical Equations “Yields”; indicates result of reaction Used to indicate a reversible reaction (equilibrium) A reactant or product in the solid state; also used to indicate a precipitate Alternative to (s), but used only to indicate a precipitate A reactant or product in the liquid state A reactant or product in an aqueous solution (dissolved in water) A reactant or product in the gaseous state (s)(s) (l)(l) (aq) (g)(g)

8. Additional Symbols Used in Chemical Equations Alternative to (g), but used only to indicate a gaseous product Reactants are heated Pressure at which reaction is carried out, in this case 2 atm Pressure at which reaction is carried out exceeds normal atmospheric pressure Temperature at which reaction is carried out, in this case 0 o C Formula of catalyst, in this case manganese (IV) oxide, used to alter the rate of the reaction 2 atm pressure 0 o C MnO 2 

9. Signs of Chemical Reactions There are five main signs that indicate a chemical reaction has taken place: change in colorchange in odorproduction of new gases or vapor input or release of energy difficult to reverse release input

11. Combustion O2O2 General form: C x H x + O 2 CO 2 H2OH2O + + carbon-hydrogen compound carbon dioxide C 4 H 10 + oxygen water CO 2 H2OH2O 4 5 13/2 13 8 10 2

12. 1. Write a word equation for the reaction Write a balanced equation for the reaction between chlorine and solid sodium bromide to produce bromine and solid sodium chloride. 2. Write the correct formulas for all reactants and products, (with correct phases of matter) 3. Balance the resulting equation chlorine + sodium bromide  bromine + sodium chloride Cl 2 (g) + NaBr(s)  Br 2 (l) + NaCl(s) Writing Equations Practice 2 2

13. 1) Write a word equation for the reaction 2) Write the correct formulas for all reactants and products 3) Balance the resulting equation aluminum sulfate + calcium chloride  calcium sulfate Write the balanced equation for the aqueous reaction between aluminum sulfate and calcium chloride to form a white precipitate of calcium sulfate. + aluminum chloride ? ? Al 2 (SO 4 ) 3 (aq) + CaCl 2 (aq)  CaSO 4 (s) + AlCl 3 (aq) 3 3 2

14. Oxidation-Reduction Reactions “Redox” reactions involve the transfer of electrons (e-) Reduction: gain e- Oxidation: lose e- “LEO the lion says, ‘GER’” “OIL RIG” Use oxidation states to keep track of the e-

15. Leo says Ger “Lose electron  oxidation” Zn  2e - + Zn 2+ “Gain electron  reduction” 2e - + Cu 2+  Cu My name is Leo. Grr-rrrr…

16. Assigning Oxidation States Specific rules for assigning Ox #’s Usually the same charge assigned by the PT H is almost always +1 O is almost always -2 F is always -1 in compounds For elements (H 2, O 2, F 2, Ca, K, etc ) the oxidation state always = 0 Some exceptions do exist!

17. Assigning Oxidation Numbers Overall charge = sum of the oxidation states of all atoms in it Neutral Compounds (e.g. H 2 O, CO 2, CH 4 ) H 2 O : The overall charge is 2(1) + -2 = 0 CO 2 : What is the oxidation state of C? Since C + 2 (O) = 0… C + 2(-2) = 0, thus… CH 4 : Is C still +4? H is always +1 To remain neutral… 4(1) + C = 0 C must = - 4 H = +1 and O = -2 C = +4

18. Assigning Oxidation Numbers Charged compounds (e.g. NO 3 -, CO 3 2- ) NO 3 - or (NO 3 ) - : What is the oxidation # of N? O is -2, and the overall charge is -1 So N + 3(O) = -1 or N + 3(-2) = -1 N = + 5 (CO 3 ) 2- : What is the oxidation # of C? O is -2, and the overall charge is -2 So C + 3(O) = -2 or C + 3(-2) = -2 C = +4 The oxidation # of ions = charge of ions Mn 3+ has an oxidation # of +3 S 2- has an oxidation # of -2

19. Assigning Oxidation # Practice Assign oxidation numbers to each atom Cl 2 Fe 2+ ClO 3 - ClO 4 - IO 2 - CrO 4 2- Fe 3 (PO 4 ) 2 CoSO 4 Cl: 0 (element) Fe: 2+ (ion) O: 2-, 3(2-) + Cl = 1-…Cl: 5+ O: 2-, 4(2-) + Cl = 1-…Cl: 7+ O: 2-, 2(2-) + I = 1-…I: 3+ O: 2-, 4(2-) + Cr = 2-…Cr: 6+ Fe: 2+ (ion) PO 4 :3- (ion)….O:2-, 4(2-) + P = 3-, P: 5+ Co: 2+ (ion) SO 4 :2- (ion)….O:2-, 4(2-) + S = 2-, S: 6+

20. Assigning Oxidation Numbers Review Try these…MnO 4 -, Cr 2 O 7 2-, C 2 O 4 2- (MnO 4 ) - O = -2, so [4(-2) + Mn = -1] Mn = +7 (Cr 2 O 7 ) 2- O = -2, so [7(-2) + 2Cr = -2] 2Cr = 12, therefore… (C 2 O 4 ) 2- O = -2, so [2C + 4(-2) = -2] 2C = 6, therefore… Cr = +6 C = +3

21. Oxidation-Reduction Reactions Two separate reactions occurring simultaneously Oxidation: oxidation # of an atom increases e.g. Fe(s) → Fe 3+ (aq) Reduction: oxidation # of an atom is “reduced” e.g. O 2 (g) → O 2- (aq) When occurring together… Fe(s) + O 2 (g) → Fe 3+ (aq) + O 2- (aq) This is the redox reaction responsible for rust! But, how do we balance this? (ox # goes from 0 → +3) (oxidation # goes from 0 → -2)

22. Balancing by Half-Reactions *in acidic solution 1.Assign oxidation states for each element. 2.Write separate half-reactions for the reduction/oxidation reactions. 3.Balance all the atoms EXCEPT O and H. 4.Balance the oxygen with water (H 2 O). 5.Balance the hydrogen with hydrogen ions (H + ). 6.Balance the charge with electrons. 7.Multiply each half-reaction by an appropriate number to make the electrons equal. 8.Combine both reactions into one and cancel the e -

23. Balancing by Half-Reactions * in acidic solution CH 3 OH (aq) + Cr 2 O 7 2- (aq) → CH 2 O(aq) + Cr 3+ (aq) 1.Assign oxidation states. C -2 H 4 + O 2- + (Cr 2 6+ O 7 2- ) 2- → C 0 H 2 + O 2- + Cr 3+ 2. Write separate half-reactions for the reduction and oxidation reactions. (only keep charges that are changing…) Ox: C -2 H 4 O → C 0 H 2 O (C is going from -2 to 0) Red: (Cr 2 6+ O 7 ) 2- → Cr 3+ (Cr is being reduced from +6 to +3)

24. 3. For each half reaction, balance all atoms EXCEPT O and H. 4.Balance the oxygen by adding water (H 2 O). 5.Balance the hydrogen by adding hydrogen ions (H + ) 6.Balance the charge by adding electrons. …use the oxidation state as a guide 7.Multiply each half-reaction by an appropriate number to make the electrons equal. 8.Add the reactions together and cancel e-/simplify. Balancing the half reactions… + 2H + + 2e- Ox: C 2- H 4 O → C 0 H 2 O Red: (Cr 2 6+ O 7 ) 2- → Cr 3+ 2 + 7H 2 O 14H + + 6e- + 3 ( ) 3CH 4 O → 3CH 2 O + 6H + + 6e- 3 CH 4 O + + Cr 2 O 7 2- → 3 CH 2 O + 2 Cr 3+ + 7 H 2 O 8 H + Red: (Cr 2 6+ O 7 ) 2- → Cr 3+

25. Ox: C 2- H 4 O → C 0 H 2 O 3. For each half reaction, balance all the atoms EXCEPT O and H. 4. Balance the oxygen by adding water (H 2 O). 5.Balance the hydrogen by adding hydrogen ions (H + ) 6.Balance the charge by adding electrons. …use the oxidation state as a guide Balancing the Oxidation… Carbon is already balanced! + 2H + + 2e- Oxygen is already balanced!

26. On to the reduction… (Cr 2 6+ O 7 ) 2- → Cr 3+ 3.Balance all elements except H and O 4.Balance O by adding H 2 O, if necessary 5.Balance H by adding H +,if necessary 6.Balance charge by adding e- Remember, you only care about the charges that are changing… 2 + 7H 2 O 14H + + 6e- +

27. Adding Half-Reactions *in acidic solution Now add the two reactions together… Ox: CH 4 O → CH 2 O + 2H + + 2e- Red: 6e- + 14H + + Cr 2 O 7 2- → 2Cr 3+ + 7H 2 O 7. Multiply each half-reaction by an appropriate number to make the electrons equal. CH 4 O → CH 2 O + 2H + + 2e- 6e- + 14H + + Cr 2 O 7 2- → 2Cr 3+ + 7H 2 O 3CH 4 O → 3CH 2 O + 6H + + 6e- 3 ( )

28. 6e- + 14H + + Cr 2 O 7 2- → 2Cr 3+ + 7H 2 O 3CH 4 O → 3CH 2 O + 6H + + 6e- 3CH 4 O + + Cr 2 O 7 2- → 3CH 2 O + 2Cr 3+ + 7H 2 O …and the reaction is now balanced! 8H + Adding Half-Reactions *in acidic solution 8. Add the reactions together and cancel e-/simplify.

29. Practice Balancing Redox Reactions Unbalanced reaction (in acid): MnO 4  + Fe 2+  Mn 2+ + Fe 3+ Balanced Reduction half-reaction: 8H + + MnO 4  + 5e   Mn 2+ + 4H 2 O Balanced Oxidation half-reaction: Fe 2+  Fe 3+ + e  Balanced overall reaction: 8H + + MnO 4  + 5Fe 2+  Mn 2+ + 5Fe 3+ + 4H 2 O 5( )

30. Balancing by Half-Reactions *in basic solution 1.Assign oxidation states. 2.Write separate half-reactions for the reduction/oxidation reactions. 3.Balance all the atoms EXCEPT O and H. 4.Balance the oxygen by adding water (H 2 O). 5.Balance the hydrogen by adding H +. 6.Balance the charge by adding electrons. 7.Multiply each half-reaction by an appropriate number to make the electrons equal. 8.Combine both reactions into one and cancel. 9.Add OH - to both sides to cancel out H + and create H 2 O. Simplify further, if necessary.

31. Balancing by Half-Reactions (in basic solution) Let’s balance a previous example in basic solution Remember, it is all the same steps up to this point 3CH 4 O + 8H + + Cr 2 O 7 2- → 3CH 2 O + 2Cr 3+ + 7H 2 O 3CH 4 O + + Cr 2 O 7 2- → 3CH 2 O + 2Cr 3+ + 7H 2 O + 8OH - 3CH 4 O + H 2 O + Cr 2 O 7 2- → 3CH 2 O + 2Cr 3+ + 8OH - + 8OH - 8H 2 O

32. Practice Balancing Basic Redox Rxns Unbalanced reaction: ClO  + Zn  Cl - + Zn 2+ Balanced Reduction half-reaction: 2e - + 2H + + ClO -  Cl - + H 2 O Balanced Oxidation half-reaction: Zn  Zn 2+ + 2e - Balanced overall reaction (acidic): 2H + + ClO  + Zn  Zn 2+ + Cl - + H 2 O Balanced overall reaction (basic): H 2 O + ClO  + Zn  Zn 2+ + Cl - + 2OH -

33. Ca Activity Series Foiled again: Aluminum is knocked out by Calcium Element Reactivity Li Rb K Ba Ca Na Mg Al Mn Zn Cr Fe Ni Sn Pb H 2 Cu Hg Ag Pt Au Halogen Reactivity F 2 Cl 2 Br 2 I 2 Printable Version of Activity Series Printable Version of Activity Series

34. Mg + AlCl 3 Al + MgCl 2 Predict if these reactions will occur… Al + MgCl 2 Can magnesium replace aluminum? Activity Series YES, magnesium is more reactive than aluminum. 2233 Can aluminum replace magnesium? NO, aluminum is less reactive than magnesium. Therefore, no reaction will occur. NR (No Reaction) MgCl 2 + Al No reaction We must determine if the lone element is more reactive than the bonded one… metals replace metals or non-metals replace nonmetals Order of reactants DOES NOT determine how they react.

35. More SR Reactions… FeCl 2 + Cu MgBr 2 + Cl 2 “Magic blue-earth” Zinc in nitric acid 2 A + BC AC + B General Form Zn(NO 3 ) 2 + H 2 Can Fe replace Cu? Yes Li Rb K Ba Ca Na Mg Al Mn Zn Cr Fe Ni Sn Pb H 2 Cu Hg Ag Pt Au F 2 Cl 2 Br 2 I 2 Can Zn replace H? Yes Can Br replace Cl? No NO REACTION Fe + CuCl 2 Zn + HNO 3 MgCl 2 + Br 2 Activity Series

36. Double Replacement Reactions K 2 CO 3 (aq) Potassium carbonate BaCl 2 (aq) Barium chloride 2 KCl (aq) Potassium chloride BaCO 3 (s) Barium carbonate ++

37. Formation of a solid precipitate: AgNO 3 (aq) + KCl(aq)  KNO 3 (aq) + AgCl(s)

38. TABLE OF SOLUBILITIES IN WATER aluminum sssnsnissisd ammonium sssssssssss barium ssisisssiid calcium ssissssssi d copper (II) ssisiinsisi iron (II) ssisnissisi iron (III) ssnsiinsissd lead sssi ii siii magnesium ssississisd mercury (I) ssiii nisi i mercury (II) sssis iisidi potassium sssssssssss silver ssiii nisi i sodium sssssssssss zinc ssississisi acetate bromide carbonate chloride chromate hydroxide iodide nitrate phosphate sulfate sulfide i = insoluble ss = slightly soluble s = soluble d = decomposes n = not isolated SOLID AQUEOUS Legend

39. Solubility Rules 1. Most nitrates are soluble. 2.Most salts containing Group I ion and ammonium ion, NH 4 +, are soluble. 3.Most chloride, bromide, and iodide salts are soluble, except Ag +, Pb 2+ and Hg 2 2+. Ohn-Sabatello, Morlan, Knoespel, Fast Track to a 5 Preparing for the AP Chemistry Examination 2006, page 91 4. Most sulfate salts are soluble, except BaSO 4, PbSO 4, Hg 2 SO 4, and CaSO 4. 5.Most hydroxides except Group 1 and Ba(OH) 2, Sr(OH) 2, and Ca(OH) 2 are only slightly soluble. 6.Most sulfides, carbonates, chromates, and phosphates are only slightly soluble.

40. FeCO 3 Na + Fe 2+ iron (II) chloride + sodium carbonate Cl 2 Using the SOLUBILITY TABLE: sodium chloride is soluble iron (II) carbonate is insoluble CO 3 Fe 2+ Fe Na + Na 2 Cl - CO 3 2- Cl- CO 3 2- NaCl sodium chloride iron (II) carbonate + (aq) (s) 2 FeCl 2 Na 2 CO 3 NaCl FeCO 3 (aq) (s) ++ Predict if a reaction will occur when you combine aqueous solutions of iron (II) chloride and sodium carbonate… If the reaction does occur, write a balanced chemical equation showing it (be sure to include phase notation). (aq) Balanced chemical equation

41. KNO 3 Na + K+K+ potassium chloride + sodium nitrate KCl (aq) Using the SOLUBILITY TABLE: sodium chloride is soluble potassium nitrate is soluble NaNO 3 (aq) K+K+ Na + Cl - NO 3 - Cl- NO 3 - NaCl sodium chloride potassium nitrate + (aq) Predict if a reaction will occur when you combine aqueous solutions of potassium chloride and sodium nitrate… If the reaction does occur, write a balanced chemical equation showing it (be sure to include phase notation). Notice that nothing has really changed because the ions are still dissolved in water! NR

42. Pb 2+ NO 3 – Na + CI – Ions in Aqueous Solution Expt. Pb(NO 3 ) 2 (s) Pb(NO 3 ) 2 (aq) Pb 2+ (aq) + 2 NO 3 – (aq) add water NaCI(s) + H 2 O(l) Dissociation reactions: solids mixed with water dissociate into ions + H 2 O(l) Na + (aq) + CI – (aq) Mix them and get… Balance to get complete ionic equation… Cancel spectator ions to get net ionic equation… NaCI(aq) NO 3 – Pb 2+ NO 3 – in solution, aqueous, soluble, dissolved Na + CI – Chem Think

43. Mix them and get… Pb 2+ (aq) + 2 NO 3 1– (aq) + 2 Na 1+ (aq) + 2 CI – (aq) PbCI 2 (s) + 2 NO 3 1– (aq) + 2 Na 1+ (aq) Pb 2+ (aq) + 2 CI – (aq) PbCI 2 (s) Pb(NO 3 ) 2 (aq) + NaCI(aq) Balance to get complete ionic equation…separate anything (aq) Cancel spectator ions to get net ionic equation… Solubility Chart Solid (precipitate) in solution (aqueous) Pb 2+ NO 3 – Na + CI – NO 3 – Na + CI – Pb 2+ NO 3 – Na + CI – NO 3 – Na + CI – PbCI 2 + NaNO 3 (s) (aq) 22

44. Pre-lab: 1. What ions are present in the following solutions? NaCl(aq)  ____________________ AgNO 3 (aq)  ____________________ 2.When these solutions are mixed together, a precipitate is seen. What are the new combinations of ions that could have formed the precipitate? ____________________ and ____________________ 3.Using the solubility table, which new combination will form a precipitate? ____________________  4.Which new combination will remain in solution? ____________________ 5. Write the complete ionic equation for this reaction. Be sure to indicate the correct phase (reaction condition) for each reactant and each product. 6. Write the net ionic equation for this reaction by canceling out spectators. Again, include the phases (reaction conditions). 7. Explain why you would expect no reaction between solutions of KOH(aq) and NaOH(aq). Na + (aq) Cl - (aq)Ag + (aq) NO 3 - (aq) Na + (aq) Cl - (aq) Ag + (aq) NO 3 - (aq) Cl - (aq) Ag + (aq) Na + (aq)NO 3 - (aq) AgCl(s) Na + (aq) + Cl - (aq)+ Ag + (aq) + NO 3 - (aq) + Na + (aq)+ NO 3 - (aq) AgCl(s) + Cl - (aq)Ag + (aq) AgCl(s) When the cations switch places they end with a hydroxide (no new combination is formed)

45. Ba 2+ OH – NO 3 – Mix together Zn(NO 3 ) 2 (aq) and Ba(OH) 2 (aq): Zn 2+ (aq) + 2 NO 3 – (aq)Ba 2+ (aq) + 2 OH – (aq) Ba(OH) 2 (aq)Zn(NO 3 ) 2 (aq) Balance to get complete ionic equation… Zn 2+ Zn(NO 3 ) 2 (aq) + Ba(OH) 2 (aq) Zn(OH) 2 (s) + Ba(NO 3 ) 2 (aq)Zn 2+ (aq) + 2 NO 3 – (aq) + Ba 2+ (aq) + 2OH – (aq) Zn(OH) 2 (s) + 2 NO 3 – (aq) + Ba 2+ (aq) Mix them and get… Zn 2+ (aq) + 2 OH – (aq) Zn(OH) 2 (s) Cancel spectator ions to get net ionic equation… Solubility Chart

46. Separation of Cations You have a solution containing Fe 2+, Cu 2+, Ba 2+, Ag + and K + ions. By what means could you separate these ions from each other? In Chem I, we discussed various ways to separate things… Distillation Filtration Centrifugation Reactivity Will any of these work to separate aqueous ions?

47. Separation of Cations Fe 2+, Cu 2+, Ba 2+, Ag +, K + (aq) + Cl - (aq) AgCl(s) Fe 2+, Cu 2+, Ba 2+, K + (aq) BaSO 4 (s) Fe 2+, Cu 2+, K + (aq) CuCrO 4 (s) Fe 2+, K + (aq) + SO 4 2- (aq) + CrO 4 2- (aq) FeS, Fe 3 (PO 4 ) 2, Fe(OH) 2, or FeCO 3 (s) K + (aq) + S 2-, PO 4 3- OH - or CO 3 2- (aq)

48. Separation of Cations Pb 2+, Ca 2+, Zn 2+, NH 4 + (aq) + CrO 4 2- (aq) PbCrO 4 (s) Ca 2+, Zn 2+, NH 4 + (aq) CaSO 4 (s) Zn 2+, NH 4 + (aq) NH 4 + (aq) + SO 4 2- (aq) ZnCO 3, Zn 3 (PO 4 ) 2, Zn(OH) 2, or ZnS(s) + CO 3 2-, PO 4 3- OH - or S 2- (aq) Try this example on your own…

49. Summary of Classes of Reactions Chemical reactions Double Replacement reactions Acid-Base Reactions Oxidation-Reduction Reactions Combustion Reactions Single Replacement reactions Synthesis and Decomposition reactions