1. CHEMICAL REACTIONS

2. SO FAR...  So far we have looked at the properties of individual atoms and molecules  Atomic number  Atomic mass  Ions  Lewis dot structures  Molecular shapes  Naming  Moles

3. HOW DO THESE ATOMS AND MOLECULES INTERACT?  The way that these atoms and molecules interact with each other is through chemical reactions Chemical reaction: Change of one or more substances into a different substance.

4. BEFORE WE BEGIN  Important concept (by Antoine Lavosier):  Conservation of mass: Matter is neither created or destroyed in a chemical reaction  In other words, you must have the same components in the beginning and end or a chemical reaction

5. GENERAL TYPES OF REACTIONS 1. Gas evolution reactions: these are reactions that occur in liquids and produce a gaseous product  For example – Alka Seltzer in H 2 O 2. Oxidation-reduction reactions: electrons are transferred from one substance to another  For example – rust formation, batteries

6. GENERAL TYPES OF REACTIONS 3. Precipitation reactions: reactions that occur in liquid that produce a solid substance

7. IN MORE DEPTH...  We will discuss these reactions more in depth at a later time

8. HOW DO WE KNOW THAT A CHEMICAL REACTION HAS TAKEN PLACE?  Many chemical reactions produce easily detectable changes  Color change  Formation of a solid  Formation of a gas  Heat absorption or emission (uses up or produces heat)  Light absorption or emission

9. HOW DO WE KNOW THAT A CHEMICAL REACTION HAS TAKEN PLACE?  The previous are only INDICATIONS of a chemical reaction  Chemical analysis is needed to verify if a reaction is actually occurring.  For example: Boiling water looks like a gas is being produced, but this is a physical change, NOT A CHEMICAL CHANGE  H 2 O(l)  H 2 O(g) NO NEW SUBSTANCE

10. SHOWING CHEMICAL REACTIONS  Chemical equations are a way that chemists summarize a chemical reaction that takes place  Reactions take the following general format Reactants  Products  A set of chemicals react to form a set of products

11. REACTIONS AND ENERGY  Chemical reactions either absorb energy/heat from their surroundings or they release energy/heat:  EXOTHERMIC: a chemical reaction that produces heat (ie. Burning gas in your car)  ENDOTHERMIC: a chemical reaction that absorbs heat (ie. Icy hot)

12. SHOWING CHEMICAL REACTIONS  In chemical reactions, the state of matter is usually described: 1. Solid (s) 2. Liquid (l) 3. Gas (g) 4. Aqueous (aq)

13. PRACTICE  For each of the following, translate the sentence into a chemical reaction: 1. Solid zinc is mixed with aqueous hydrochloric acid to form zinc chloride and hydrogen gas 2. Aqueous sulfuric acid is combined with solid gold to produce aqueous gold(I) sulfate and hydrogen gas

14. ANSWER 1. Zn(s) + HCl(aq)  ZnCl 2 (aq) + H 2 (g) 2. H 2 SO 4 (aq) + Au  Au 2 SO 4 (aq) + H 2 (g)

15. HOW DO WE APPLY CONSERVATION OF MATTER TO CHEMICAL REACTIONS?  C 3 H 8 (l) + O 2 (g)  CO 2 (g) + H 2 O(l)  Let’s examine this reaction closer:  How many atoms of each element are on each side of the reaction? Reactants ProductsCHO

16. BALANCING THE REACTION  According the law of conservation of matter, you must have the same number of each atom on each side  NO MATTER is created or destroyed

17. BALANCING THE REACTION  Therefore, we must BALANCE the reaction to get the same number of atoms on the reactant side and the product side  To balance we can add more MOLECULES, but you cannot change the molecule

18. BALANCING THE REACTION  C 3 H 8 (l) + O 2 (g)  CO 2 (g) + H 2 O(l)  Therefore, we CAN add more C 3 H 8  For example: we can have 3 C 3 H 8  But we CANNOT change the molecule itself  Cannot make C 3 H 8 into CH 8

19. TRY TO BALANCE THE EQUATION  C 3 H 8 (l) + O 2 (g)  CO 2 (g) + H 2 O(l)

20. ANSWER  C 3 H 8 (l) + 5O 2 (g)  3CO 2 (g) + 4H 2 O(l)  Let’s look at each atom on the reactant and product side: ReactantProduct C3C3 H8H8 O10O10  Same number on each side = BALANCED

21. WHAT ARE THE NUMBERS IN FRONT OF EACH COMPOUND?  C 3 H 8 (l) + 5O 2 (g)  3CO 2 (g) + 4H 2 O(l)  The numbers in the front of each compound tells you how many MOLES you combine.  How many moles of each compound do you have in this reaction?

22. BALANCING TIPS  Just balancing the equation is fairly difficult  Therefore there are some tips to help you balance the equation

23. TIP #1  The first thing to do is to write the skeletal reaction with just the compounds involved:  For example:  Aqueous sodium hydroxide and aqueous sulfuric acid mix to form liquid water and aqueous sodium sulfate

24. TIP #2  H 2 SO 4 (aq) + NaOH(aq)  H 2 O(l) + Na 2 SO 4 (aq)  If an element occurs in only 1 compound on both sides of the reaction, balance it first. Otherwise, balance metals before nonmetals

25. TIP #3  H 2 SO 4 (aq) + 2 NaOH(aq)  H 2 O(l) + Na 2 SO 4 (aq)  Not balanced yet  Save diatomic molecules and/or water for last  H 2 SO 4 (aq) + 2 NaOH(aq)  2 H 2 O(l) + Na 2 SO 4 (aq)

26. SIDE NOTE  SIDE NOTE: Most times, when you have a polyatomic ion, do not break it up. Think of the polyatomic ion as one atom for balancing. It will save you some grief.  H 2 SO 4 (aq) + 2 NaOH(aq)  2 H 2 O(l) + Na 2 SO 4 (aq)  The SO 4 in the reactant is balanced by the SO 4 in the product. Don’t balance all the S and the O

27. TIP #4  If a balanced equation contains coefficient fractions, clear these by multiplying the entire equation by the appropriate number  C 8 H 18 + O 2  CO 2 + H 2 O  C 8 H 18 + O 2  8 CO 2 + 9 H 2 O  C 8 H 18 + (25/2) O 2  8CO 2 + 9H 2 O  2 C 8 H 18 + 25 O 2  16 CO 2 + 18 H 2 O

28. TIP #5  Check to make sure the equation is balanced by summing the atoms on each side of the reaction Reactant Product C16C16 H36H36 O50O50

29. TRY THESE 1. Solid aluminum and aqueous sulfuric acid combine to form aqueous aluminum sulfate and hydrogen gas 2. Solid iron combines with gaseous oxygen to form solid iron(III) oxide 3. Liquid ethanol (C 2 H 6 O) combines with gaseous oxygen to form gaseous carbon dioxide and liquid water

30. ANSWER 1. 2Al(s) + 3H 2 SO 4 (aq)  Al 2 (SO 4 ) 3 (aq) + 3H 2 (g) 2. 4Fe(s) + 3O 2 (g)  2Fe 2 O 3 (s) 3. C 2 H 6 O(l) + O 2 (g)  2CO 2 (g) + 3H 2 O(l)

31. SPECIAL NOTE  All reactions are reversible. They can go forwards or backwards.  For example:  4Fe(s) + 3O 2 (g)  2Fe 2 O 3 (s)  2Fe 2 O 3 (s)  4Fe(s) + 3O 2 (g)  To show this, we write reactions with double arrows  4Fe(s) + 3O 2 (g)  2Fe 2 O 3 (s)

32. SPECIFIC TYPES OF CHEMICAL REACTIONS  There are 4 specific types of chemical reactions. 1. Synthesis (combination) reactions 2. Decomposition reactions 3. Displacement reactions 4. Double displacement reactions 5. Combustion reactions (special)

33. SYNTHESIS REACTIONS  A synthesis reaction is when you take simpler molecules and combine them to form more complex molecules  They have the following general set up  A + B  AB  Example: 2Na(s) + Cl 2 (g)  2NaCl(aq)

34. DECOMPOSITION REACTIONS  These are the opposite of the synthesis reactions. A more complex compound breaks down to form simpler compounds.  They have the following general set up  AB  A + B  Example: 2H 2 O(l)  2H 2 (g) + O 2 (g)

35. DISPLACEMENT REACTIONS (SINGLE DISPLACEMENT)  One element displaces (or takes the place of) another element in a compound.  They have the following general set up  A + BC  AC + B  Example: Zn(s) + CuCl 2 (aq)  ZnCl 2 (aq) + Cu(s)

36. DOUBLE DISPLACEMENT REACTION  Two elements, in separate compounds, displace (or switch places with) each other.  They have the following general set up  AB + CD  AD +BC  Example: AgNO 3 (aq) + NaCl(aq)  AgCl(s) + NaNO 3 (aq)

37. COMBUSTION REACTIONS  When a compound combines with oxygen to form carbon dioxide, water and heat  They have the following general set up  Molecule + O 2  CO 2 + H 2 O + heat  Example: 2C 2 H 6 + 7O 2  4CO 2 + 6H 2 O + heat

38. TRY THESE  Translate, balance and tell what type of reactions the following are: 1. Chlorine gas combines with sodium bromide to form sodium chloride and bromine gas. 2. Mercury(II) oxide breaks apart to form mercury and oxygen gas.

39. ANSWER 1. Cl 2 + 2NaBr  Br 2 + 2NaCl  The chlorine displaces the bromine in sodium bromide  This is a SINGLE DISPLACEMENT reaction 2. 2HgO  2Hg + O 2  The mercury(II) oxide breaks apart into mercury and oxygen  This is a DECOMPOSITION reaction

40. GENERAL TYPES OF REACTIONS (EXPANDED)  Before we can describe the general types of reactions, in detail, we have to visit the concept of solubility  Soluble: the substance in question will dissolve in water  Insoluble: the substance in question will not dissolve in water

41. WHY SOLUBILITY  If you know what compounds are soluble, you can figure out if a reaction is:  Gas evolution reaction  Precipitation reaction  If soluble, the reaction will not form a precipitate or evolve a gas

42. SOLUBLE (VERY IMPORTANT – MAKE SURE YOU COPY)  Soluble 1. All acetates (C 2 H 3 O 2 - ) except Fe +3 2. All ammonium (NH 4 + ) compounds 3. All bromides (Br - ) except Ag +, Hg 2 +2, and Pb +2 4. All chlorates (ClO 3 - ) 5. All chlorides (Cl - ) except Ag + and Pb +2

43. SOLUBLE (VERY IMPORTANT – MAKE SURE YOU COPY) 6. All iodides (I - ) except Ag +, Hg 2 +2, and Pb +2 7. All nitrates (NO 3 - ) 8. All perchlorates (ClO 4 - ) 9. All sulfates (SO 4 -2 ) except Ca +2, Ba +2, Pb +2, Sr +2, Hg 2 +2

44. INSOLUBLE  Insoluble 1. Carbonates (CO 3 -2 ), except group 1 and NH 4 + 2. Chromates (CrO 4 -2 ), except group 1 and NH 4 + 3. Hydroxides (OH - ), except group 1, Ca +2, Sr +2, and Ba +2 4. Oxalates (C 2 O 4 -2 ), except group 1 and NH 4 +

45. INSOLUBLE 5. Oxides (O -2 ), except group 1, Ca +2, Sr +2, and Ba +2 6. Phosphates (PO 4 -3 ), except group 1 and NH 4 + 7. Sulfides (S -2 ), except those of group 1, group 2, and NH 4 +

46. SOLUBLE OR INSOLUBLE? 1. H 2 SO 4 2. Ca(C 2 H 3 O 2 ) 2 3. AgNO 3 4. Na 2 O 5. Li 2 CO 3 6. BaC 2 O 4 7. Mg(OH) 2

47. ANSWER 1. Soluble 2. Soluble 3. Soluble 4. Soluble 5. Soluble 6. Insoluble 7. Insoluble

48. WHY DO WE CARE ABOUT SOLUBILITY?  By knowing the solubility rules, you will be able to do further analysis to determine if a chemical reaction will form a precipitate  If a chemical does not precipitate out of solution (have the product insoluble), then the reaction will not be a precipitation reaction

49. EXAMPLE  Examine the reactions below, determine which reaction produces a precipitate or not 1. Ba(OH) 2 + K 2 O  BaO + 2KOH 2. 2H 3 PO 4 + 3BaO  Ba 3 (PO 4 ) 2 + 3H 2 O

50. ANSWER 1. Ba(OH) 2 (aq) + K 2 O(aq)  BaO(aq) + 2KOH(aq) 1. BaO is soluble 2. KOH is soluble 3. Since both products are soluble, there is no precipitate 2. 2H 3 PO 4 (aq) + 3BaO(aq)  Ba 3 (PO 4 ) 2 (s) + 3H 2 O(l) 1. Ba 3 (PO 4 ) 2 is insoluble. 2. This means that as the reaction progresses, the Ba 3 (PO 4 ) 2 is a precipitate that is produced.

51. GENERAL TYPES OF REACTIONS (REVISITED)  Precipitation reactions  Only nonsoluble compounds form precipitates  If any product is insoluble, it forms a precipitate

52. PRECIPITATION REACTION  How to determine: 1. Write the two compounds in question as reactants 1. Na 2 CO 3 + CuCl 2  2. Write the formulas of the potential products of the reaction 1. Na 2 CO 3 + CuCl 2  NaCl + CuCO 3

53. PRECIPITATION REACTION 3. Look at the solubility rules to determine the solubility of the products 1. NaCl  soluble 2. CuCO 3  insoluble 4. Write the equation with states of matter 1. Na 2 CO 3 (aq) + CuCl 2 (aq)  NaCl(aq) + CuCO 3 (s)

54. PRECIPITATION REACTION 5. Balance the equation 1. Na 2 CO 3 (aq) + CuCl 2 (aq)  2NaCl(aq) + CuCO 3 (s)

55. GAS EVOLUTION REACTIONS  Whenever a reaction produces a gas, that gas will leave the reaction (bubble forms)  If a product is a gas, then you have a gas evolution reaction  HCl(aq) + NaHCO 3 (aq)  H 2 O(l) + NaCl(aq) + CO 2 (g)  This is the reaction of Alka Seltzer with your stomach acid

56. OXIDATION-REDUCTION REACTIONS  These reactions (redox) involve the transfer of electrons from one element to another  In order to understand these reactions better, we must look at the charges of the elements in a chemical reaction

57. DEFINITIONS  Oxidation: when an element loses electrons (the atom becomes +)  Reduction: when an element gains electrons (the charge is reduced, becomes -)  Oxidation and reduction must occur together so the electrons can travel from one element to another

58. OXIDATION-REDUCTION REACTIONS  Example  4Na(s) + O 2 (g)  2Na 2 O  4 Na (s)  4 Na + + 4e-  4e- + 2O  2O -2  In this example, Na goes from a neutral charge to a +1, it needed to lose an electron  O goes from a neutral charge to a -2, it had to gain electrons

59. TRY THESE  In each of the following reactions, identify the element that is oxidized and the element that is reduced: 1. 2Na + Cl 2  2NaCl 2. 4Fe + 3O 2  Fe 2 O 3

60. ANSWER 1. 2Na + Cl 2  2NaCl 1. 2Na  2Na + + 2e- (oxidized) 2. 2e- + 2Cl  2Cl- (reduced) 2. 4Fe + 3O 2  2Fe 2 O 3 1. 4Fe  4Fe +3 + 12e- 2. 12e- + 6 O  6 O -2

61. SUMMARY  There are 3 general types of reactions  Precipitation  One of the products must be insoluble  Use the solubility rules  Gas evolution  If a gas is produced, it will come out of solution  Oxidation-reduction  Involves the movement of electrons  Something must be oxidized (lose electrons) and something must be reduced (gain electrons)

62. SUMMARY  There are 5 specific types of reactions  Synthesis  A + B  AB  Decomposition  AB  A + B  Displacement  A + BC  AB + C  Double displacement  AB + CD  AD + BC  Combustion

63. SUMMARY  With any sort of chemical reaction, we must follow the law of conservation of mass  Therefore, every chemical reaction must have the same number of atoms in the reactants and products  To balance, use the following tips

64. SUMMARY 1. Write the skeletal reaction with just the compounds involved. 2. If an element occurs in only 1 compound on both sides of the reaction, balance it first. Otherwise, balance metals before nonmetals 3. Save diatomic molecules and/or water for last 4. If a balanced equation contains coefficient fractions, clear these by multiplying the entire equation by the appropriate number 5. Check to make sure the equation is balanced by summing the atoms on each side of the reaction