2. Types of Reactions Chemistry Types of Reactions Chemistry

3. Learning Objective l TLW identify and quantify changes that occur during chemical reactions (TEKS 8)

4. Agenda l Review of Basics l Define Types of Chemical Reactions and Real-World Examples l Examples of the Types of Chemical Reactions l Group Practice l Individual Practice l Labs

5. Chemical Reactions – General Stuff l Reactants are added together to make products l Reactants are found on left side of chemical equation and products are on the right side l Compounds are formed – substances made up of two or more elements that chemically combined (not easily separated or impossible to return to original element)

6. Chemical Reactions – General Stuff l Valence electrons and periodic trends determine how various atoms of elements react (or don’t react) l Catalysts are “outside agents” that may be added to speed up reactions or cause reactions to perform differently under different conditions (such as varying temperatures). –They are not used up in the reaction

7. Chemical Reactions – General Stuff l Don’t forget about the laws of conservation of mass and energy… neither can be created nor destroyed –Reactants are transformed –Products are formed –Chemical energy converts to heat, light, electricity, sound

8. Chemical Reactions – General Stuff l Also – Law of Definite Proportions, which states in any sample of a chemical compound the elements are always combined in the same proportion by mass H 2 O H = 1 g x 2 = 2 g O = 16 g Total 18 g H = 2 g / 18 g = 11%O = 16 g / 18 g = 89%

9. Chemical Reactions – General Stuff l From Dalton’s Laws (remember him?) l Law of Multiple Proportions states whenever two elements form more than one compound different masses of one element that combine with the same mass of the other element are in the ratio of small whole number H 2 OH 2 O 2

10. Chemical Reactions – General Stuff l Law of Definite Proportions – describes composition of one compound l Law of Multiple Proportions – compares composition of two different compounds containing same elements

11. Types of Reactions l Combination (a.k.a. Synthesis, Addition) l Decomposition l Single Replacement (a.k.a. Single Displacement) l Double Replacement (a.k.a. Double Displacement) l Combustion l Neutralization l Precipitation l Reduction/Oxidation Reactions (Redox) l Energy Producing – Exothermic, Endothermic, Light

12. Combination (aka Synthesis, Addition) l Two l Two or more elements or substances combine to form a new compound. l A l A + B AB l Examples l Examples – –Formation of rust –Air pollutant sulfur dioxide –Polymerization (plastics) – Photosynthesis Photosynthesis (plants)

13. Combination l A l A + B AB l Where l Where A and B are elements and AB is a compound l Note l Note that only one compound exists on the RIGHT RIGHT SIDE…

14. Combination l4l4l4l4Fe(s) + 3O2(g)  2Fe2O3(s) lSlSlSlS(s) + O2(g)  SOx(g) l2l2l2l2Na + Cl2  2NaCl Examples

15. l Teacher Demo – need a better one…..

16. Decomposition l A single compound is broken down to produce two or more smaller compounds and/or elements. l AB  A + B l Example – – Water with electricity into hydrogen and oxygen (electrolysis) – Baking soda with heat

17. Decomposition l AB  A + B l Where AB is a compound and A & B are elements or other compounds l Note that only one compound exists on the LEFT SIDE…

18. Decomposition l Examples: l 2H 2 O (l)  2H 2(g) + O 2(g) l NaHCO 3(s)  2H 2(g) + NaCO 3 heat

19. l Teacher demo – baking soda and heat

20. Single Replacement (Single Displacement) l One l One element replaces (displaces) a similar element in a compound l Produces l Produces heat (is exothermic) l A l A + BC AC + B l Example l Example – –If –If you place an iron nail into a beaker of copper (II) chloride you will begin to see reddish copper forming on the iron. –Iron –Iron replaces replaces (displaces) (displaces) copper in the solution and the copper falls out of solution as a metal

21. Single Replacement l A l A + BC AC + B l Where l Where A and B are elements and BC and AC are compounds l Can l Can have more than 2 reactants and/or products l Example: 2HCl (l) 2HCl (l) + Zn (s) Zn (s)  ZnCl 2(l) ZnCl 2(l) + H 2(g) Fe (s) Fe (s) + CuCl 2(l) CuCl 2(l)  Cu (s) Cu (s) + FeCl 2(l)

22. Single Replacement l All single replacement reactions are exothermic l They give off heat and occur rapidly

23. l Teacher demo – Zinc plus hydrochloric acid or iron nail in copper(II)chloride

24. Double Replacement (Double Displacement) l Ions from two compounds in solution exchange to produce two new compounds l AB + CD  AD + CB l One compound usually forms a precipitate that settles out of the solution, a gas that bubbles out, or a molecular compound like water l The other compound formed often remains dissolved in the solution l Examples – – Baking soda and vinegar – Dried fruit

25. Double Replacement l AB + CD  AD + CB l Where AB, CD, AD, & CB are all compounds l Can have more than 2 reactants and/or products l Examples: l 2HCl + 2NaOH  2NaCl + 2H 2 O l Na 2 SO 3(aq) + 2HCl (aq)  2NaCl (aq) + H 2 O (l) + SO 2(g)

26. l Teacher demo – the ever popular baking soda and vinegar or cleaning pennies with vinegar using salt as a catalyst

27. Combustion Carbon substances combine with oxygen, releasing large amounts of energy, in the form of heat, light, etc. Carbon dioxide and water are also typical products C x H y + O 2  CO 2 + H 2 O Examples – – Natural gas to heat a house – Hydrogen powered cars

28. General Formula C x H y + O 2  CO 2 + H 2 O hydrocarbon oxygen carbon water dioxide dioxide Examples - CH 4 + 2O 2  CO 2 + 2H 2 O CH 4 + 2O 2  CO 2 + 2H 2 O 2H 2 (g) + O 2 (g)  2H 2 O(l) 2H 2 (g) + O 2 (g)  2H 2 O(l)

30. l Teacher demo – lighting a match

31. Neutralization l Double Replacement Reaction where wn acid and a base react to form water and a salt l General formula HA + BOH  H 2 O + BA acid base water salt Examples – HCl(aq) + NaOH(s)  H 2 O(l) + NaCl(s) H 2 SO 4 (aq) + Ca(OH) 2 (aq)  H 2 O(l) + CaSO 4 (s)

32. Precipitation l When solutions are saturated, adding additional ions will cause a precipitate to form l Solid – usually sinks to bottom, but can float l Solubility product (K sp ) can be used to predict formation of precipitates –If ion-product concentration > K sp then a precipitate will form –If ion-product concentration < K sp then a precipitate will not form

33. l Teacher demo – baking soda and vinegar… again (how boring)

34. l Practice Worksheet – Identify 5 basic reactions Practice Worksheet – Identify 5 basic reactions

35. Discovery Video - Electrochemistry Reduction/Oxidation Reactions A. Reaction in which electrons are transferred B. Commonly called the redox reaction C. One element is reduced – it gains electrons One element is oxidized – it loses electrons Reduction/Oxidation

36. Redox reactions have radicals fragments of molecules with at least one electron for bonding Ex. of radicals = Styrofoam

37. Example of redox reaction Rust = iron reacts with oxygen Fe 2 O 3 Fe loses 3 electrons O gains 2 electrons

38. Rules for Assigning Oxidation Numbers

39. Rule 1 The oxidation number of any uncombined element is 0

40. Example: The oxidation number of Na (s) is 0.

41. Rule 2 The oxidation number of a monatomic ion equals the charge on the ion.

42. Example: The oxidation number of Cl - is -1.

43. Rule 3 The more electronegative element in a binary compound is assigned the number equal to the charge it would have if it were an ion.

44. Example: The oxidation number of O in NO is -2.

45. Rule 4 The oxidation number of fluorine in a compound is always -1.

46. Example: The oxidation number of F in LiF is -1.

47. Rule 5 Oxygen has an oxidation number of -2 unless it is combined with F, when it is +2, or it is in a peroxide, such as H 2 O 2, when it is -1.

48. Example: The oxidation number of O in NO 2 is -2.

49. Rule 6 The oxidation state of hydrogen in most of its compounds is +1 unless it is combined with a metal, in which case it is -1.

50. Example: The oxidation number of H in LiH is -1.

51. Rule 7 In compounds, Group 1 and 2 elements and aluminum have oxidation numbers of +1, +2, and +3, respectively.

52. Example: The oxidation number of Ca in CaCO 3 is +2.

53. Rule 8 The sum of the oxidation numbers of all atoms in a neutral compound is 0.

54. Example: The oxidation number of C in CCO 3 is +4.

55. Rule 9 The sum of the oxidation numbers of all atoms in a polyatomic ion equals the charge of the ion.

56. Example: The oxidation number of P in H 2 PO- 4 is +5.

57. Exothermic Reactions l In many reactions, less energy is required to break the bonds in the reactants than is released when bonds form to make new products l In these reactions some type of heat or light is released and they are called exothermic l Exothermic reactions can be detected by a rise in temperature

58. Exothermic Reactions l What are examples of exothermic reactions you are familiar with? l Demonstration

59. Endothermic Reactions l Sometimes more energy is required to break bonds in the reactants than is released to form new products l The are called endothermic reactions l You can detect these reactions by a decrease in temperature

60. Endothermic Reactions l What are examples of endothermic reactions you are familiar with? l Demonstration

61. Group Practice ~ Name Those Reactions

62. A Group Activity l Types of Chemical Reactions – Sorting Matslinklink

63. Individual Practice l Identifying Types of Reactions Worksheetlinklink l Crossword Puzzle

64. Looking Ahead l Labs – –Conservation of Mass –Types of Chemical Reactions –Empirical Formula Determination –Predicting the Amount of Product in a Reaction –Identifying Relationships between Reactants and Products in a Reaction –Predicting the Products of a Reaction –Precipitation Reactions –Energetic Reactions – exothermic and endothermic experiments –Qualitative and Quantitative Analysis l More on calculating definite proportions and multiple proportions