3. Effects of chemical reactions: Chemical reactions rearrange atoms in the reactants to form new products. The identities and properties of the products are completely different from that of the reactants. Production of gases and color changes are signs of chemical reactions.

5. Energy and Reactions zEnergy must be ADDED to BREAK bonds. zEnergy is RELEASED when bonds are FORMED. zChemical energy is CONSERVED in chemical reactions.

6. Exo- vs. Endo- zEXOTHERMIC REACTIONS: release energy (More energy is released as the products form bonds than is absorbed to break the bonds in the reactants.) zENDOTHERMIC REACTIONS: absorb energy

7. A. Reaction Pathway zShows the change in energy during a chemical reaction

8. B. Exothermic Reaction zreaction that releases energy zproducts have lower PE than reactants 2H 2 (l) + O 2 (l)  2H 2 O(g) + energy energy released

9. C. Endothermic Reaction zreaction that absorbs energy zreactants have lower PE than products 2Al 2 O 3 + energy  4Al + 3O 2 energy absorbed

10. Reaction Rate (How fast?) Chemical Reactions

11. A. Collision Theory zReaction rate depends on the collisions between reacting particles. zSuccessful collisions occur if the particles... collide with each other have the correct orientation have enough kinetic energy to break bonds

12. A. Collision Theory zParticle Orientation Required Orientation Successful Collision Unsuccessful Collisions

13. A. Collision Theory zActivation Energy (E a ) yminimum energy required for a reaction to occur Activation Energy

14. A. Collision Theory zActivation Energy ydepends on reactants ylow E a = fast rxn rate EaEa

15. B. Factors Affecting Rxn Rate of a reaction 1. Surface Area yhigh surfaced area = fast rxn rate ymore opportunities for collisions yIncrease surface area by… xusing smaller particles (crushing particles) xdissolving in water

16. B. Factors Affecting Rxn Rate of a Reaction 2. Concentration high conc = fast rxn rate more opportunities for collisions

17. B. Factors Affecting Rxn Rate of a Reaction 3. Temperature yhigh temp = fast rxn rate yhigh KE xfast-moving particles xmore likely to reach activation energy

18. B. Factors Affecting Rxn Rate of a Reaction 3. Temperature Analogy: 2-car collision 5 mph “fender bender” 50 mph “high-speed crash”

19. B. Factors Affecting Rxn Rate 4. Catalyst substance that increases rxn rate without being consumed in the rxn lowers the activation energy

20. B. Factors Affecting Rxn Rate of a Reaction z5. Stirring or agitating xMakes particles come in contact more frequently by adding KE.

21. Chemical equations are used to represent or describe chemical reactions. For example when hydrogen H 2 burns, it reacts with oxygen, O 2, in the air to form water. We write the chemical equation for this reaction as follows: 2H 2 + O 2 —> 2H 2 O

22. Chemical Equations An equation shows…  Formulas of reactants  Formulas of products  Molar ratios of all compounds in the reaction.

23. Chemical Equations The (+) sign as “reacts with” The arrow (  ) means “yields” or “reacts to produce”. 2H 2 + O 2  2H 2 O Reactants Products

24. To show physical states of each substance: (s) solid (l) liquid (g) gas (aq) aqueous aqueous means dissolved in water

25. To show physical states of each substance: Consider the reaction of iron with oxygen to form iron (III) oxide, or rust. Fe (s) + O 2 (g)  Fe 2 O 3 (s)(unbalanced)

26. Coefficients & Subscripts COEFFICIENTS: numbers in front of compound that represents the number of molecules/moles of that compound SUBSCRIPTS: small numbers that help define the compound. 2H 2 SO 4 Coefficient Subscript

27. H 2 OOne molecule of water 2H 2 OTwo molecules of water H 2 O 2 One molecule of Hydrogen Peroxide

28. During a chem. rxn.; atoms are rearranged (NOT created or destroyed!) Chemical equations must be balanced to show the relative amounts of all substances. Balanced means: each side of the equations has the same # of atoms of each element. CH 4 + O 2 —> H 2 O + CO 2 Unbalanced CH 4 + 2O 2 —> 2H 2 O + CO 2 Balanced

29. A. Balancing Steps 1.Write the unbalanced equation. 2.Count atoms on each side. 3.Add coefficients to make #s equal. Coefficient  subscript = # of atoms 4.Reduce coefficients to lowest possible ratio, if necessary. 5.Double check atom balance!!!

30. B. Helpful Tips zBalance one element at a time. zUpdate ALL atom counts after adding a coefficient. zIf an element appears more than once per side, balance it last. zBalance polyatomic ions as single units. y“1 SO 4 ” instead of “1 S” and “4 O”

31. Al + CuCl 2  Cu + AlCl 3 Al Cu Cl 1 1 1 1 2 3 2  3  6   3 33 2 C. Balancing Example Aluminum and copper(II) chloride react to form copper and aluminum chloride. 2  2  6

32. Balancing Equations NOTE: When balancing equations, you may change coefficients as much as you need to, but you may never change subscripts because you can’t change what substances are involved.

33. Fe (s) + O 2 (g)  Fe 2 O 3 (s)(unbalanced) 4Fe (s) + 3O 2 (g)  2Fe 2 O 3 (s)(balanced)

34. Sample Problem 1 Water is decomposed (broken down) to form the gaseous products hydrogen, H 2, and oxygen, O 2. Write the balanced equation for this reaction. H 2 O  H 2 + O 2 2H + 1O  2H + 2O O is not balanced 2H 2 O  2H 2 + O 2 4H + 2O  4H + 2O The equation is balanced!

35. Sample Problem 2 Chlorine gas, Cl 2, reacts with potassium bromide, KBr, to form potassium chloride and bromine, Br 2. Write the balanced equation for this reaction, Cl 2 + KBr  KCl + Br 2 2Cl + 1K + 1Br  1Cl + 1K +2Br Cl and Br are not balanced Cl 2 + 2KBr  2KCl + Br 2 2Cl + 2K + 2Br  2Cl + 2K +2Br The equation is balanced!

36. Balancing equations involves a great deal of “trial and error” at first, but there are some tricks…

37. For example….. Sodium metal reacts with water to produce sodium hydroxide and hydrogen gas. Na + H 2 O —> NaOH + H 2 Note that on the product side (right side) there are an odd number of hydrogens. On the reactant side (left side) there is an even number. This implies there must be an even coefficient in front of the NaOH. Lets start with 2

38. _Na + _H 2 O —> 2NaOH + _H 2 Now lets balance sodium; we need a 2 in front of the Na… 2Na + _H 2 O —> 2NaOH + _H 2 Now consider hydrogen… 2Na + 2H 2 O —> 2NaOH + H 2

39. Check to see if it balances… 2 Na on the left2 Na on the right 4 hydrogen 2 + 2 = 4 hydrogen2 oxygen  the equation is balanced.

40. Examples CuCl 2 (aq) + Al(s)  Cu(s) +AlCl 3 (aq) 3CuCl 2 (aq) + 2Al(s)  3Cu(s) +2AlCl 3 (aq) (3:2:3:2)

41. Examples Propane, C 3 H 8, burns in oxygen, O 2, to form carbon dioxide and water. C 3 H 8 + O 2  CO 2 + H 2 O Balance C – then H – then O C 3 H 8 + 5O 2  3CO 2 + 4H 2 O (1:5:3:4)

42. Examples Pentane, C 5 H 12, burns in oxygen, O 2, to form carbon dioxide and water. C 5 H 12 + O 2  CO 2 + H 2 O Balance C – then H – then O C 5 H 12 + 8O 2  5CO 2 + 6H 2 O (1:8:5:6)

43. Examples Silver nitrate reacts with copper to produce silver and copper (II) nitrate. AgNO 3 + Cu  Ag + Cu(NO 3 ) 2 2AgNO 3 + Cu  2Ag + Cu(NO 3 ) 2 (2:1:2:1)

44. Examples Phosphorus reacts with oxygen gas to produce diphosphorus pentoxide. P + O 2  P 2 O 5 4P + 5O 2  2P 2 O 5 (4:5:2)

45. Examples C 7 H 14 + O 2  CO 2 + H 2 O Balance C – then H – then O 2C 7 H 14 + 21O 2  14CO 2 + 14H 2 O (2:21:14:14)

46. WORKSHEET (class work) zDO THE SIDE THAT SAYS: Reaction Types first zONLY write the abbreviation for the type of reaction. (some are already balanced) zWS Equation Practice: do the whole question: yWrite out the full equation from the words yBalance yDon’t Identify the type of reaction…….Yet

47. You will need your reference tables today! Get them out along with your note packet and homework.

48. What we have Learned

49. Types of Chemical Reactions Synthesis / Formation Decomposition Single Replacement Double Replacement Combustion

50. B. Synthesis zthe combination of 2 or more substances to form a compound zonly one product A + B  AB

52. Al(s)+ Cl 2 (g)  AlCl 3 (s) 2 3 2 B. Synthesis zProducts: yionic - cancel charges ycovalent - hard to tell because they use prefixes with their names

53. Synthesis / Combination Reactions A + B  AB Examples: 2K (s) + Cl 2 (g)  2KCl (s) SO 2 (g) + H 2 O (l)  H 2 SO 3 (aq) NH 3 + HCl  NH 4 Cl ** one to remember

54. Decomposition Reactions

55. C. Decomposition za compound breaks down into 2 or more simpler substances zonly one reactant AB  A + B

56. Decomposition Reactions AB  A + B Examples: 2H 2 O (l)  2H 2 (g) + O 2 (g) CaCO 3  CaO + CO 2

57. KBr(l)  K(s) + Br 2 (l) 2 2 C. Decomposition zProducts: ybinary - break into elements yothers - hard to tell must use reference tables. LiOH  Li 2 O + H 2 O 2

58. Li K Ca Na Mg Al Zn Fe Pb (H)* Cu Hg Ag Increasing Activity Any element will replace any element below it. Activity Series: (in ref tables) *Metals from Li to Na will replace H from acids and water; from Mg to Pb they will replace H from acids only

59. Single-Replacement Reactions Definition: Reaction where atoms of one element replace atoms of a second element in a compound. XA + B  BA + X Note: A reactive metal will replace any metal listed below it in the activity series. Generally, nonmetal replacement is limited to the halogens. The activity of the halogens decreases as you go down Group 17 of the periodic table. See reference tables. Examples: AgNO 3 + Mg  Mg+LiNO 3  no reaction Mg(NO 3 ) 2 + Ag 22

60. D. Single Replacement zProducts: ymetal  metal (+) ynonmetal  nonmetal (-) yfree element must be more active (check activity series)

61. For Example… Ca + MgO  The Ca will replace the Mg because Ca is more active than Mg. That is to say…Ca is above Mg on the activity list. CaO + Mg

62. Double-Replacement Reactions Definition: 2 positive ions “switch places” between two compounds to form 2 new ionic compounds. XA + BY  BA + XY Note: These reactions generally take place between two ionic compounds in aqueous solution, and are often characterized by one of the products coming out of solution in some way. Examples: 2NaCN (aq) +H 2 SO 4 (aq)  2HCN (g) +Na 2 SO 4 (aq) Na 2 S (aq) +Cd(NO 3 ) 2 (aq)  CdS (s) +2NaNO 3 (aq)

63. AB + CD  AD + CB E. Double Replacement nions in two compounds “change partners”. Also called Ionic rxn’s ncation of one compound combines with anion of the other

64. Pb(NO 3 ) 2 (aq)+ KI(aq)  PbI 2 (s)+ KNO 3 (aq) E. Double Replacement zProducts: yswitch negative ions yone product must be insoluble (check solubility rules on ref tables) NaNO 3 (aq)+ KI(aq)  N.R. 2 2

65. Combustion Reactions Definition: a hydrocarbon (containing C and H) or other substance burns in the presence of oxygen gas (O 2 ) to produce CO 2 and H 2 O Examples: CH 4 +2O 2  2Mg (s) +O 2 (g)  2MgO (s) ***exception*** CO 2 +2H 2 O

66. Now you try it…… Determine the products and type of reaction for: H 2 O 2 (l)  KI

67. Do you know?? zWhich type of gas was produced? Hydrogen gas (H 2 ) or Oxygen gas (O 2 )?

68. Therefore… z2 H 2 O 2 (l)  2 H 2 O (g) + O 2 (g) zWhat type of reaction is this? KI

69. Answer is…… zDECOMPOSITION

70. Copper Silver Nitrate Determine the products and type of reaction for: Cu (s) + AgNO 3 (aq) 

71. Will it react? zwhat happened… zCu (s) + AgNO 3 (aq)  produced what?

72. Cu (s) + 2 AgNO 3 (aq)  2 Ag (s) + Cu(NO 3 ) 2 (aq) What type of reaction is this?

73. zSINGLE REPLACEMENT

74. Ionic Equations When a soluble substance is dissolved in water, the substance often breaks into ions. This solution is said to be an aqueous solution. Pb(NO 3 ) 2 (aq)  Pb 2+ + 2NO 3 - NaI (aq)  Na + + I - An aqueous solution is ions dissolved in water

75. Ionic Equations Consider the reaction… Pb(NO 3 ) 2 (aq) + NaI (aq)  PbI 2 (s) + NaNO 3 (aq) What is really going on is… Pb 2+ + NO 3 - + Na + + I -  PbI 2 (s) + Na + + NO 3 - Note that the Na + ion and the NO 3 - ion are not reacting. They are said to be spectator ions.

76. Net Ionic Equations It is often useful to write an equation showing only the species that are actually reacting. This is called a net ionic equation. It does not show the spectator ions. Pb 2+ + NO 3 - + Na + + 2I -  PbI 2 (s) + Na + + NO 3 - becomes…. Pb 2+ + 2I -  PbI 2 (s)