2. Why we care about chemical reactions

3. Types of Chemical Reactions A. Combination or synthesis Reactions B. Decomposition Reactions C. Combustion Reactions D. Single-Replacement Reactions E. Double replacement Reactions A. Precipitation Reactions B. Acid-base Reactions (Neutralization)

4. Combination aka Synthesis Reactions A + B = AB A chemical change in which two or more substances react to form a single new substance Frequently these can be ionic Often multiple products are possible Examples 2Na(s) + Cl 2 (g)  2NaCl(s) S(s) + O 2 (g)  SO 2 (g) 2S(s) + 3O 2 (g)  2SO 3 (g) Fe(s) + S(s)  FeS(s) 2Fe(s) + 3S(s)  Fe 2 S 3 (s) Both involve S and O Both involve Fe and S

5. Combination/Synthesis Li + O 2  O Li 2 O Li O

6. Decomposition Reaction AX = A + X Opposite of synthesis reactions Chemical reactions in which a single compound breaks down into two or more simpler products Often seen when ionic substances are dissolved in water Examples: 2HgO(s)  2Hg(l) + O 2 (g) 2H 2 O(l)  2H 2 (g) + O 2 (g) NaCl (s)  Na + (aq) + Cl - (aq) Δ Δ Electricity

7. Decomposition KClO 3  KCl+ O 2 K Cl O O O K O O

8. Combustion Reactions C x H y + O 2 = CO 2 + H 2 O A carbon-based compound reacts with oxygen to release a large amount of energy in the form of heat or light Examples: 2C 8 H 16 (l) + 25O 2 (g)  16CO 2 (g) + 18H 2 O(l) CH 4 + 2O 2  CO 2 + 2H 2 O C 6 H 12 O 6 +6O 2  6CO 2 + 6H 2 O Cellular respiration! Burning natural gas!

9. Combustion CH 4 + 2O 2  CO 2 + 2H 2 O O O O O

10. Just because O 2 is involved and light/heat is produced, you do NOT necessarily have a combustion reaction! 2Mg(s) + O 2 (g)  2MgO(s) S(s) + O 2 (g)  SO 2 (g) NOT combustion reactions – synthesis reactions

11. Single Replacement Reactions A + BC = AB + C A chemical change in which one element replaces a second element in a compound Reactants AND products have an element and a compound One metal or halogen may NOT always displace another Examples: 2K(s) + 2HOH(L)  2KOH(aq) + H 2 (g) Br 2 (aq) + 2NaI(aq)  2NaBr + I 2 (aq) Zn(s) + Cu(NO 3 ) 2 (aq)  Cu(s) + Zn(NO 3 ) 2 (aq)

12. Example: Zn + CuCl 2  Zn Cl Cu + General: AB + C  AC + B Cl Zn Cu +

13. Single Replacement Reactions Displacement of one metal by another depends upon the relative reactivities of the two metals Example Zn(s) + Cu(NO 3 ) 2 (aq)  Cu(s) + Zn(NO 3 ) 2 (aq) Cu(s) + Zn(NO 3 ) 2 (aq)  No reaction

14. Activity Series of Metals Notice that zinc is higher on the activity series of metals than copper A reactive metal will replace any metal listed below it in the activity series Zn + Cu(NO 3 ) 2  Cu + Zn(NO 3 ) 2 Cu + Zn(NO 3 ) 2  No reaction Decreasing reactivity Single Replacement Reactions

15. Halogens higher in the periodic table will replace halogens lower in the periodic table Example: Br 2 + 2NaI  2NaBr + I 2 NaBr + I 2  No reaction Decreasing reactivity

16. Mixed Practice State the type & predict the products. Balance the reactions 1. KBr + F 2  2. C 6 H 12 + O 2  3. Al + CuSO 4  4. Cs + Br 2  5. FeCO 3  (hint: form an iron oxide) 6. Ni(PO 4 ) 2 + Pb  2KBr + F 2  2KF + Br 2, sgl rep C 6 H 12 + 9O 2  6CO 2 + 6H 2 O, combustion 2Al + 3CuSO 4  Al 2 (SO 4 ) 3 + 3Cu, sgl rep 2Cs + Br 2  2CsBr, synthesis FeCO 3  FeO + CO 2, decomposition Ni(PO 4 ) 2 + Pb  no reaction

18. Double Replacement Reactions AB + CD = AC + BD The ions of two compounds exchange places in an aqueous solution to form two new compounds Often produce a precipitate, a gas, or molecular compound such as water Examples: Precipitation (forms a solid or gas as a product) Na 2 S(aq) +Cd(NO 3 ) 2 (aq)  CdS(s) + 2NaNO 3 (aq) 2NaCN(aq) + H 2 SO 4 (aq)  2HCN(g) + Na 2 SO 4 (aq) Acid/Base (forms water as product) Ca(OH) 2 (aq) + 2HCl(aq)  CaCl 2 (aq) + 2H 2 O(l)

19. Double Replacement CuCl 2 + Na 2 S  2NaCl + CuS Cl Cu Na S

20. Example: MgO + CaS General:AB + CD  AD + CB S O  Mg Ca + O S Mg Ca +

21. Complete Ionic Equations Many important reactions take place in water—that is in aqueous solution When in solution, many ionic compounds will split into their constitute ions (strong electrolytes) Others are only slightly soluble (weak electrolytes) in water and precipitate (clump/solidify) out of solution Equations in aqueous solution can be written as complete ionic equations AgNO 3 (aq) + NaCl(aq)  AgCl(s) + NaNO 3 (aq) Can be written as a complete ionic equation: Ag + (aq) + NO 3 - (aq) + Na + (aq) + Cl - (aq)  AgCl(s) + Na + (aq) + NO 3 - (aq)

22. Net Ionic Equations Ag + (aq) + NO 3 - (aq) + Na + (aq) + Cl - (aq)  AgCl(s) + Na + (aq) + NO 3 - (aq) Ions that appear unchanged on both sides of the equation are called spectator ions When you rewrite the ionic equation without the spectator ions, you get a net ionic equation Ag + (aq) + Cl - (aq)  AgCl(s) A net ionic equation shows only those particles involved in the reaction and is balanced with respect to both mass and charge

23. Net Ionic Equations (Continued) Complete equation Pb(s) + AgNO 3 (aq)  Ag(s) + Pb(NO 3 ) 2 (aq) Complete ionic equation Pb(s) + Ag + (aq) + NO 3 - (aq)  Ag(s) + Pb 2+ (aq) + 2NO 3 - (aq) Net ionic equation Pb(s) + Ag + (aq)  Ag(s) + Pb 2+ (aq) Balanced net ionic equation Pb(s) + 2Ag + (aq)  2Ag(s) + Pb 2+ (aq) How do you know what is going to precipitate (become solid)?

24. General Solubility Guidelines Always soluble Nitrates (NO 3 - ) – trumps EVERY other rule Anything with an alkaline/alkali metal other than barium Halogens unless they combine Ag/Pb/Hg Sulfates (SO 4 2- ) except with Ba/Pb Usually Insoluble Carbonates (CO 3 2- ), chromates (CrO 4 2- ), and phosphates (PO 4 2- ) unless with K/Na/NH 4 + Silver salts Hydroxides (OH - ) made with anything other than column 1 metals These are basic GUIDELINES only. There are more guidelines and rules and always exceptions.

25. Acid-Base reactions (also known as neutralization) HOH HX + MOH  MX + H 2 O Acid - a substance that produces H+ ions when dissolved in water Strong acids completely disassociate in solution and can be considered a strong electrolyte Base - is a substance that produces OH- ions when dissolved in solution Strong bases completely disassociate in solution and can be considered a strong electrolyte Complete ionic equation

26. Reactions in Aqueous Solution Acid-Base reactions Acids and Bases react to form water and a ionic compound called a salt H 3 PO 4 + 3Fe(OH) 2  Fe 3 (PO 4 ) 2 + 6H 2 O H 2 SO 4 + Ca(OH) 2  CaSO 4 + 2H 2 O

28. A + BX  B + AX 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

29. AB  A + B 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

30. AB + XY  AY + XB 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

31. CH + O 2  CO 2 + H 2 O 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

32. 2H 2 O  2H 2 + O 2 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

33. 2C + O 2  2CO 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

34. Mg + CdCl 2  Cd + MgCl 2 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

35. Al 2 (SO 4 ) 3 + 3Ca(OH) 2  2Al(OH) 2 + 3CaSO 4 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction Acid/Base reaction

36. 2C 6 H 14 + 19O 2  12CO 2 + 14H 2 O + heat 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

37. BaO + H 2 O  Ba(OH) 2 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

38. 2NaNO 3  2NaNO 2 + O 2 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

39. AgNO 3 + NaCl  AgCl + NaNO 3 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

40. 2Al + 3ZnCl 3  3Zn + 2AlCl 3 1. Synthesis Reaction 2. Decomposition Reaction 3. Single Replacement Reaction 4. Double Replacement Reaction 5. Combustion Reaction

41. Practice Predict the products. Balance the reaction. Label each type of reaction. 1. HCl (aq) + AgNO 3(aq)  2. CaCl 2(aq) + Na 3 PO 4(aq)  3. Pb(NO 3 ) 2(aq) + BaCl 2(aq)  4. FeCl 3(aq) + NaOH (aq)  5. H 2 SO 4(aq) + NaOH (aq)  6. KOH (aq) + CuSO 4(aq) 

42. Predict the products. 1. HCl (aq) + AgNO 3(aq)  2. 3CaCl 2(aq) + 2Na 3 PO 4(aq)  3. Pb(NO 3 ) 2(aq) + BaCl 2(aq)  4. FeCl 3(aq) + 3NaOH (aq)  5. H 2 SO 4(aq) + 2NaOH (aq)  6. 2KOH (aq) + CuSO 4(aq)  AgCl(s) + HNO 3 6NaCl + Ca 3 (PO 4 ) 2 (s) PbCl 2 (s) + Ba(NO 3 ) 2 (aq) 3NaCl + Fe(OH) 3 (s) Na 2 SO 4 (aq) + 2H 2 O Cu(OH) 2 (s) + K 2 SO 4 Precip. A/B Precip.