1. 1 Chapter 10 Chemical Reactions Chemistry Tracy Bonza Sequoyah High School

2. 2 How do you tell if it is a chemical reaction??????? l Country Girls’ Pies Taste Luscious! l Color changechange l Gas Produced Produced l Precipitate (a solid that falls –like rain—out of a solution)(a solid that falls –like rain—out of a solution) l Temperature change (heat, cold)change (heat, cold) l Light given off given off

3. 3 Physical or Chemical Change l Growth of a tree. l Melting butter. l Fizzing soda l Use of food by body l Combustion of gas l Separation of crude oil l Freezing pond l Separation of water into Hydrogen and oxygen gas Chemical Physical Chemical Chemical Chemical Physical Physical Chemical

4. 4 All chemical reactions l have two parts l Reactants - the substances you start with l Products- the substances you end up with l The reactants turn into the products. Reactants  Products

5. 5 In a chemical reaction l The way atoms are joined is changed l Atoms aren’t created or destroyed. l Can be described several ways l In a sentence l Copper reacts with chlorine to form copper (II) chloride. l In a word equation Copper + chlorine  copper (II) chloride

6. 6 Symbols used in equations l the arrow separates the reactants from the products l Read “reacts to form” l The plus sign = “and” l (s) after the formula -solid l (g) after the formula -gas l (l) after the formula -liquid

7. 7 Symbols used in equations l (aq) after the formula - dissolved in water, an aqueous solution.  used after a product indicates a gas (same as (g))  used after a product indicates a solid (same as (s))

8. 8 Symbols used in equations l indicates a reversible reaction (More later) l shows that heat is supplied to the reaction l is used to indicate a catalyst used supplied, in this case, platinum.

9. 9 What is a catalyst? l A substance that speeds up a reaction without being changed by the reaction. l Enzymes are biological or protein catalysts.

10. 10 Skeleton Equation l Uses formulas and symbols to describe a reaction l doesn’t indicate how many. l All chemical equations are sentences that describe reactions.

11. 11 Convert these to equations l Solid iron (III) sulfide reacts with aqueous hydrogen chloride to form iron (II) chloride and hydrogen sulfide gas. Fe 2 S 3 (s) + HCl (aq)  FeCl 2 + H 2 S (g)

12. 12 l Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water. HNO 3(aq) + Na 2 CO 3  H 2 O (l) + CO 2(g) +NaNO 3(aq)

13. 13 The other way Fe(g) + O 2 (g)  Fe 2 O 3 (s) Cu(s) + AgNO 3 (aq)  Ag(s) + Cu(NO 3 ) 2 (aq) l NO 2 N 2 (g) + O 2 (g)

14. 14 Balancing Chemical Equations

15. 15 Balanced Equation l Atoms can’t be created or destroyed l All the atoms we start with we must end up with l A balanced equation has the same number of each element on both sides of the equation.

16. 16 C + O 2  CO 2 l This equation is already balanced l What if it isn’t already? C + O O  C O O

17. 17 C + O 2  CO l We need one more oxygen in the products. l Can’t change the formula, because it describes what is C + O  C O O

18. 18 l Must be used to make another CO l But where did the other C come from? C + O  C O O O C

19. 19 l Must have started with two C 2 C + O 2  2 CO C + O  C O O O C C

20. 20 Rules for balancing  Write the correct formulas for all the reactants and products  Count the number of atoms of each type appearing on both sides  Balance the elements one at a time by adding coefficients (the numbers in front). Start with the highest subscript. Save oxygen for last and hydrogen for next to last.  Check to make sure it is balanced.

21. 21 Never l Change a subscript to balance an equation. l If you change the formula you are describing a different reaction. l H 2 O is a different compound than H 2 O 2 l Never put a coefficient in the middle of a formula l 2 NaCl is okay, Na2Cl is not.

22. 22 Example H 2 +H2OH2OO2O2  Make a table to keep track of where you are.

23. 23 Example H 2 +H2OH2OO2O2  Need twice as much O in the product RP H O 2 2 2 1

24. 24 Example H 2 +H2OH2OO2O2  Changes the O RP H O 2 2 2 1 2

25. 25 Example H 2 +H2OH2OO2O2  Also changes the H RP H O 2 2 2 1 2 2

26. 26 Example H 2 +H2OH2OO2O2  Need twice as much H in the reactant RP H O 2 2 2 1 2 2 4

27. 27 Example H 2 +H2OH2OO2O2  Recount RP H O 2 2 2 1 2 2 4 2

28. 28 Example H 2 +H2OH2OO2O2  The equation is balanced, has the same number of each kind of atom on both sides RP H O 2 2 2 1 2 2 4 2 4

29. 29 Example H 2 +H2OH2OO2O2  This is the answer RP H O 2 2 2 1 2 2 4 2 4 Not this

30. 30 Examples AgNO 3 + Cu  Cu(NO 3 ) 2 + Ag Mg + N 2  Mg 3 N 2 P + O 2  P 4 O 10 Na + H 2 O  H 2 + NaOH

31. 31 Types of Reactions Predicting the Products

32. 32 Types of Reactions l There are millions of reactions. l Can’t remember them all l Fall into several categories. l We will learn 5 types. l Will be able to predict the products. l For some we will be able to predict whether they will happen at all. l Will recognize them by the reactants

33. 33 Types of Chemical Reactions l Let’s look in more detail.

34. 34 #1 Combination Reactions or Synthesis l Combine - put together l 2 elements, or compounds combine to make one compound. Ca +O 2  CaO SO 3 + H 2 O  H 2 SO 4 l We can predict the products if they are two elements. Mg + N 2 

35. 35 Write and balance Ca + Cl 2  Fe + O 2  iron (II) oxide Al + O 2  l Remember that the first step is to write the formula l Then balance

36. 36 #2 Decomposition Reactions l decompose = fall apart l one reactant falls apart into two or more elements or compounds. l NaCl Na + Cl 2 l CaCO 3 CaO + CO 2

37. 37 #2 Decomposition Reactions l Can predict the products if it is a binary compound l Made up of only two elements l Falls apart into its elements lH2OlH2O l HgO

38. 38 #2 Decomposition Reactions l When the reactant has a polyatomic ion they break apart in a special way l You have to know how 3 special polyatomic ions decompose l Carbonates, chlorates and hydroxides

39. 39 Special Decomposition Rxns l Carbonates –break apart into metal oxide + CO 2 Na 2 CO 3  Na 2 O + CO 2 MgCO 3  MgO + CO 2

40. 40 Special Decomposition Rxns l Chlorates –Break apart into metal chloride + O 2 KClO 3  KCl + O 2 Ca(ClO 3 ) 2  CaCl 2 + O 2 2 3 3

41. 41 Special Decomposition Rxns l Hydroxides –Break apart into the metal oxide and H 2 O Ca(OH) 2  CaO + H 2 O LiOH  Li 2 O + H 2 O 2

42. 42 #3 Single Replacement l One element replaces another l Reactants must be an element and a compound. l Products will be a different element and a different compound. Na + KCl  K + NaCl F 2 + LiCl  LiF + Cl 2

43. 43 #3 Single Replacement l Metals replace metals (and hydrogen) K + AlN  Zn + HCl  l Think of water as HOH l Metals replace one of the H, combine with hydroxide. l Na + HOH

44. 44 #3 Single Replacement l We can tell whether a reaction will happen l Some are more active than other l More active replaces less active l There is a list on your cruncher l Higher on the list replaces lower. l If the element by itself is higher, it happens, in lower it doesn’t

45. 45 #3 Single Replacement l Lithium l Potassium l Calcium l Sodium l Magnesium l Aluminum l Zinc l Iron l Nickel l Tin l Lead l Hydrogen l Copper l Silver l Gold Only the first 5 (Li - Na) react with water.

46. 46 #3 Single Replacement Iron (II) + Copper (II) sulfate  Lead (II) + potassium chloride  Aluminum + Hydrochloric acid 

47. 47 #3 Single Replacement l What does it mean that Au And Ag are on the bottom of the list? l Nonmetals can replace other nonmetals l Limited to F 2, Cl 2, Br 2, I 2 l The order of activity is that on the table. l Higher replaces lower. F 2 + HCl  Br 2 + KCl 

48. 48 #4 Double Replacement l Two things replace each other. l Reactants must be two ionic compounds or acids. l Usually in aqueous solution NaOH + FeCl 3  l The positive ions change place. NaOH + FeCl 3  Fe +3 OH - + Na +1 Cl -1 NaOH + FeCl 3  Fe(OH) 3 + NaCl

49. 49 #4 Double Replacement l Will only happen if one of the products –doesn’t dissolve in water and forms a solid –or is a gas that bubbles out. –or is a covalent compound, usually water.

50. 50 Complete and balance l assume all of the reactions take place. Calcium chloride + sodium hydroxide  Copper (II) nitrate + potassium sulfide 

51. 51 Ammonium sulfate + Barium fluoride 

52. 52 How to recognize which type l Look at the reactants l E + E Combination l CDecomposition l E + CSingle replacement l C + CDouble replacement

53. 53 Examples H 2 + O 2  H 2 O  Zn + H 2 SO 4  HgO  KBr +Cl 2  AgNO 3 + NaCl  Mg(OH) 2 + H 2 SO 3 

54. 54 Last Type: Combustion l A compound composed of only C H and maybe O is reacted with oxygen l If the combustion is complete, the products will be CO 2 and H 2 O. l If the combustion is incomplete, the products will be CO and H 2 O.

55. 55 Examples C 4 H 10 + O 2  (complete) C 4 H 10 + O 2  (incomplete) C 6 H 12 O 6 + O 2  (complete) C 8 H 8 +O 2  (incomplete)

56. 56 Chapter 10 Summary

57. 57 An equation l Describes a reaction l Must be balanced because to follow Law of Conservation of Energy l Can only be balanced by changing the coefficients. l Has special symbols to indicate state, and if catalyst or energy is required.

58. 58 Reactions l Come in 5 types. l Can tell what type they are by the reactants. l Single Replacement happens based on the activity series using activity series. l Double Replacement happens if the product is a solid, water, or a gas.

59. 59 The Process l Determine the type by looking at the reactants. l Put the pieces next to each other l Use charges to write the formulas l Use coefficients to balance the equation.