1. 1 Unit 4 “Chemical Reactions” Adapted from: Pre-AP Chemistry Charles Page High School Stephen L. Cotton

2. 2 Describing Chemical Reactions  OBJECTIVES: –Describe how to write a word equation.

3. 3 Describing Chemical Reactions  OBJECTIVES: –Describe how to write a skeleton equation.

4. 4 Describing Chemical Reactions  OBJECTIVES: –Describe the steps for writing a balanced chemical equation.

5. 5 All chemical reactions…  have two parts: 1.Reactants = the substances you start with 2.Products = the substances you end up with  The reactants will turn into the products.  Reactants  Products

6. 6 Reactants Products

7. 7 Symbols in equations?  the arrow (→) separates the reactants from the products (arrow points to products) –Read as: “reacts to form” or yields  The plus sign = “and”  (s) after the formula = solid: Fe (s)  (g) after the formula = gas: CO 2(g)  (l) after the formula = liquid: H 2 O (l)

8. 8 Symbols in equations (cont)  (aq) after the formula = dissolved in water, an aqueous solution: NaCl (aq) is a salt water solution   ↑ used after a product indicates a gas has been produced: H 2 ↑   used after a product indicates a solid has been produced: PbI 2 ↓

9. 9 Symbols in equations (cont)  double arrow indicates a reversible reaction (more later)  shows that heat is supplied to the reaction  is used to indicate a catalyst is supplied (in this case, platinum is the catalyst)

10. 10 Determining Phases of Substances  (g) – non-metal oxides & diatomics  (l) – water, Mercury, some hydrocarbons  (s) – Metal elements, insoluble substances (see rules)  (aq) – soluble substances (see rules)

11. 11 In a chemical reaction  Atoms aren’t created or destroyed (according to the Law of Conservation of Mass)  A reaction can be described several ways: #1. In a sentence every item is a word Copper reacts with chlorine to form copper (II) chloride. #2. In a word equation some symbols used Copper + chlorine  copper (II) chloride

12. 12 #3. The Skeleton Equation  Uses formulas and symbols to describe a reaction –but doesn’t indicate how many; this means they are NOT balanced  All chemical equations are a description of the reaction.

13. 13 Write a skeleton equation for: 1. Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas. 2. Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

14. 14 Now, read these equations: Fe (s) + O 2(g)  Fe 2 O 3(s) Cu (s) + AgNO 3(aq)  Ag (s) + Cu(NO 3 ) 2(aq) NO 2(g) N 2(g) + O 2(g)

15. 15 #4. Balanced Chemical Equations  Atoms can’t be created or destroyed in an ordinary reaction: –All the atoms we start with we must end up with (meaning: balanced!)  A balanced equation has the same number of each element on both sides of the equation.

16. 16 Rules for balancing: 1)Assemble the correct formulas for all the reactants and products, using “+” and “→” 2)Count the number of atoms of each type appearing on both sides 3)Balance the elements one at a time by adding coefficients (the numbers in front) where you need more - save balancing the H and O until LAST! Double-Check to make sure it is balanced.

17. 17  Never change a subscript to balance an equation (You can only change coefficients) –If you change the subscript (formula) you are describing a different chemical. –H 2 O is a different compound than H 2 O 2  Never put a coefficient in the middle of a formula; they must go only in the front 2 NaCl is okay, but Na 2 Cl is not.

18. 18 Balancing Equations  Count atoms. –Reactants: 2 atoms N and 2 atoms H –Products: 1 atom N and 3 atoms of NH 3 N 2 + H 2 NH 3 Nitrogen + hydrogen ammonia

19. 19 Balancing Equations  Nothing is balanced.  Balance the nitrogen first by placing a coefficient of 2 in front of the NH 3. N 2 + H 2 2NH 3

20. 20 Balancing Equations  Hydrogen is not balanced.  Place a 3 in front of H 2.  Reactant side: 2 atoms N, 6 atoms H  Product side: 2 atoms N, 6 atoms H N 2 + 3H 2 2NH 3

21. 21 Balancing Equations  Count atoms in the reactant:  Ca – 3 atoms, P – 2 atoms, O – 8 atoms; H – atoms, S – 1 atom, O – 4 atoms Ca 3 (PO 4 ) 2 + H 2 SO 4 CaSO 4 + H 3 PO 4

22. 22 Balancing Equations  Side note on Ca 3 (PO 4 ) 2  The subscript after the phosphate indicates two phosphate groups.  This means two PO 4 3- groups with two P and eight O atoms.

23. 23 Balancing Equations  Count atoms in the product.  Ca atoms – 1, S atom – 1, O atoms – 4; H atoms – 3, P atom – 1, O atoms – 4 Ca 3 (PO 4 ) 2 + H 2 SO 4 CaSO 4 + H 3 PO 4

24. 24 Balancing Equations  In this equation, the ion groups do not break up.  Instead of counting individual atoms, ion groups may be counted. Ca 3 (PO 4 ) 2 + H 2 SO 4 CaSO 4 + H 3 PO 4

25. 25 Balancing Equations  Reactants: Ca 2+ – 3, PO 4 3- - 2, H + – 2, SO 4 2+ - 1  Products: Ca 2+ - 1, SO 4 2- - 1, H + - 3, PO 4 3- - 1 Ca 3 (PO 4 ) 2 + H 2 SO 4 CaSO 4 + H 3 PO 4

26. 26 Balancing Equations  Balance the metal first by placing a coefficient of 3 in front of CaSO 4.  Products: Ca – 3 atoms, SO 4 2- - 3 groups Ca 3 (PO 4 ) 2 + H 2 SO 4 3CaSO 4 + H 3 PO 4

27. 27 Balancing Equations  Three sulfate groups are needed on the reactant side so place a coefficient of 3 in front of H 2 SO 4.  3H 2 SO 4 gives 6 H + and 3 SO 4 2-.  Neither phosphate nor calcium is balanced. Ca 3 (PO 4 ) 2 + 3H 2 SO 4 3CaSO 4 + H 3 PO 4

28. 28 Balancing Equations  A coefficient of 2 placed in front of H 3 PO 4 which balances both hydrogen and phosphate. Ca 3 (PO 4 ) 2 + 3H 2 SO 4 3CaSO 4 + 2H 3 PO 4

29. 29 Balancing Equations  Balancing hints: –Balance the metals first. –Balance the ion groups next. –Balance the other atoms. –Save the non ion group oxygen and hydrogen until the end.

30. 30 Practice Balancing 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 _CH 4 + _O 2  _CO 2 + _H 2 O 2 2 3 4 5 22 2 22

31. 31 Types of Chemical Reactions  OBJECTIVES: ‒ Describe the five general types of reactions.

32. 32 Types of Chemical Reactions  OBJECTIVES: –Predict the products of the five general types of reactions.

33. 33 Types of Reactions  Chemical Reactions fall into several categories.  We will learn:  a) the 5 major types.  b) to be able to predict the products.  How?  We recognize them by their reactants

34. 34 Steps to Writing Reactions  Some steps for doing reactions 1.Identify the type of reaction 2.Predict the product(s) using the type of reaction as a model 3.Balance it  Don’t forget about the diatomic elements! (BrINClHOF) Ex. Oxygen is O 2 as an element.  In a compound, it can’t be a diatomic element because it’s not an element anymore, it’s a compound!

35. 35 #1 - Synthesis Reactions  2 substances combine to make one compound (also called “combination”)  reactant + reactant  1 product  Basically: A + B  AB Ca + O 2  CaO SO 3 + H 2 O  H 2 SO 4  We can predict the products, especially if the reactants are two elements.

36. 36 Practice  Predict the products. Write and balance the following synthesis reaction equations.  Sodium metal reacts with chlorine gas Na (s) + Cl 2(g)   Solid Magnesium reacts with fluorine gas Mg (s) + F 2(g)   Aluminum metal reacts with fluorine gas Al (s) + F 2(g) 

37. 37 #1 – Synthesis Reactions  Additional Important Notes: a) Some nonmetal oxides react with water to produce an acid: SO 2 (g) + H 2 O (l)  H 2 SO 3 (aq) (This is what happens to make “acid rain”) b) Some metallic oxides react with water to produce a base: CaO (s) + H 2 O (l)  Ca(OH) 2 (aq)

38. 38 #2 - Decomposition Reactions  One reactant breaks apart into two or more elements or compounds.  1 Reactant  Product + Product  In general: AB  A + B NaCl Na + Cl 2 CaCO 3 CaO + CO 2  Note that energy (heat, sunlight, electricity, etc.) is usually required

39. 39 #2 - Decomposition Reactions  We can predict the products if it is a binary compound (which means it is made up of only two elements)  It breaks apart into the elements: H2OH2O HgO

40. 40 Decomposition Exceptions  Carbonates and chlorates are special case decomposition reactions that do not go to the elements. See Packet for more rules

41. 41 Practice  Predict the products. Then, write and balance the following decomposition reaction equations:  Solid Lead (IV) oxide decomposes PbO 2(s)   Aluminum nitride decomposes AlN (s) 

42. 42 #3 - Single Replacement Reactions  One element replaces another  Reactants must be an element and a compound.  Products will be a different element and a different compound.  Na + KCl  K + NaCl  F 2 + LiCl  LiF + Cl 2 (Cations switched) (Anions switched)

43. 43 #3 Single Replacement Reactions  A metal can replace a metal (+) OR a nonmetal can replace a nonmetal (-).  element + compound  element + compound A + BC  AC + B (if A is a metal) OR A + BC  BA + C (if A is a nonmetal) (remember the cation always goes first!)

44. 44 #3 Single Replacement Reactions  Write and balance the following single replacement reaction equation:  Zinc metal reacts with aqueous hydrochloric acid Zn (s) + HCl (aq)  ZnCl 2 + H 2(g) Note: Zinc replaces the hydrogen ion in the reaction 2

45. 45 #3 Single Replacement Reactions Sodium chloride solid reacts with fluorine gas NaCl (s) + F 2(g)  NaF (s) + Cl 2(g) * Note that fluorine replaces chlorine in the compound Aluminum metal reacts with aqueous copper (II) nitrate Al (s) + Cu(NO 3 ) 2(aq)  2 2

46. 46 The “Activity Series” of Metals Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead Hydrogen Bismuth Copper Mercury Silver Platinum Gold 1)Metals can replace other metals, provided they are above the metal they are trying to replace (for example, zinc will replace lead) 2)Metals above hydrogen can replace hydrogen in acids. 3)Metals from sodium upward can replace hydrogen in water. Higher activity Lower activity

47. 47 The “Activity Series” of Halogens Fluorine Chlorine Bromine Iodine Halogens can replace other halogens in compounds, provided they are above the halogen they are trying to replace. 2NaCl (s) + F 2(g)  2NaF (s) + Cl 2(g) MgCl 2(s) + Br 2(g)  ??? No Reaction! ??? Higher Activity Lower Activity

48. 48 #3 Single Replacement Reactions Practice:  Fe (s) + CuSO 4 (aq)   Pb (s) + KCl (aq)   Al (s) + HCl (aq) 

49. 49 #4 - Double Replacement Reactions  Two things replace each other.  Reactants must be two ionic compounds, in aqueous solution Compound + compound  compound + compound AB + CD  AD + CB  NaOH + FeCl 3  Fe(OH) 3 + NaCl  The positive ions change place.

50. 50 #4 - Double Replacement Reactions  Think about it like “foil”ing in algebra, first and last ions go together + inside ions go together Example: AgNO 3(aq) + NaCl (s)  AgCl (s) + NaNO 3(aq) Another example: K 2 SO 4(aq) + Ba(NO 3 ) 2(aq)  KNO 3(aq) + BaSO 4(s) 2

51. 51 Practice  Predict the products. Balance the equation 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) 

52. 52 Complete and balance:  Assume all of the following reactions actually take place: CaCl 2 (aq) + NaOH (aq)  CuCl 2 (aq) + K 2 S (aq)  KOH (aq) + Fe(NO 3 ) 3 (aq)  (NH 4 ) 2 SO 4 (aq) + BaF 2 (aq) 

53. 53 How to recognize which type?  Look at the reactants: E + E =Combination C =Decomposition E + C =Single replacement C + C =Double replacement

54. 54 Practice Examples:  H 2 (g) + O 2 (g)   H 2 O (l)   Zn (s) + H 2 SO 4 (aq)   HgO (s)   KBr (aq) + Cl 2 (g)   AgNO 3 (aq) + NaCl (aq)   Mg(OH) 2 (aq) + H 2 SO 3 (aq) 

55. 55 #5 – Combustion Reactions  Combustion means “add oxygen”  Normally, a compound composed of a hydrocarbon (only C, H, and maybe O) is reacted with oxygen – usually called “burning”  If the combustion is complete, the products will always be CO 2 and H 2 O.

56. 56 Combustion Reaction Examples:  C 4 H 10 (g) + O 2 (g)   C 6 H 12 O 6 (s) + O 2 (g)   C 8 H 8 (l) + O 2 (g) 

57. 57 SUMMARY: An equation...  Describes a reaction  Must be balanced in order to follow the Law of Conservation of Mass  Can only be balanced by changing the coefficients.  Has special symbols to indicate the physical state, if a catalyst or energy is required, etc.

58. 58 Summary: Reactions  Come in 5 major types.  We can tell what type they are by looking at the reactants.  Single Replacement happens based on the Activity Series  Double Replacement happens if one product is: 1) a precipitate (an insoluble solid), 2) water (a molecular compound), or 3) a gas.

59. 59 Reactions in Aqueous Solution  OBJECTIVES: –Describe the information found in a net ionic equation.

60. 60 Reactions in Aqueous Solution  OBJECTIVES: –Predict the formation of a precipitate in a double replacement reaction.

61. 61 Net Ionic Equations  Many reactions occur in water- that is, in aqueous solution  When dissolved in water, many ionic compounds “dissociate”, or separate, into cations and anions

62. 62 Net Ionic Equations  Write the molecular equation (synthesis, decomposition, etc.), then check for reactants and products that are soluble or insoluble.  We usually assume the reaction is in water  We can use a solubility table (see oxidation sheet) to tell us what compounds dissolve in water  If the compound is soluble (does dissolve in water), then the compound splits into its component ions  If the compound is insoluble (does NOT dissolve in water), then it remains as a compound

63. 63

64. 64 Net Ionic Equations  Example (needs to be a double replacement reaction) AgNO 3 (aq) + NaCl (aq)  AgCl (s) + NaNO 3 (aq) 1. This is the Molecular Equation 2. Next, write it as an ionic equation by splitting the compounds into their ions: Ag 1+ (aq) + NO 3 1- (aq) + Na 1+ (aq) + Cl 1- (aq)  AgCl (s) + Na 1+ (aq) + NO 3 1- (aq) Note that the AgCl did not ionize, because it is a “precipitate”

65. 65 Net Ionic Equations 3. Simplify by crossing out ions not directly involved (called spectator ions) Ag 1+ (aq) + Cl 1- (aq)  AgCl (s) This is called the net ionic equation

66. 66 Net Ionic Equations Molecular Equation: K 2 CrO 4 + Pb(NO 3 ) 2  PbCrO 4 + 2 KNO 3 SolubleSolubleInsoluble Soluble Net Ionic Equation: 2 K + (aq) + CrO 4 -2 (aq) + Pb +2 (aq) + 2 NO 3 - (aq)  PbCrO 4 (s) + 2 K + (aq) + 2 NO 3 - (aq) *Note it remains balanced ion pairs

67. 67 Net Ionic Equations  Net Ionic Equation:  You should cancel out ions that appear on BOTH sides of the equation 2 K + (aq) + CrO 4 -2 (aq) + Pb +2 (aq) + 2 NO 3 - (aq)  PbCrO 4 (s) + 2 K + (aq) + 2 NO 3 - (aq) Net Ionic Equation: CrO 4 -2 (aq) + Pb +2 (aq)  PbCrO 4 (s)

68. 68 Let’s do some examples together of net ionic equations, starting with these reactants: BaCl 2 + AgNO 3 → NaCl + Ba(NO 3 ) 2 →