1. Types of Chemical Reactions (rxns.)

2. – Chemical reactions occur when bonds (between the electrons of atoms) are formed or broken – Chemical reactions involve changes in the chemical composition of matter (the making of new materials with new properties) energy changes –Bond breaking absorbs Energy (endothermic process) –Bond making releases Energy (exothermic process) – Symbols represent elements – Formulas describe compounds – Chemical equations describe a chemical reactionIntroduction

3. Types of Reactions Reactions are classified by their products. There are five main types of chemical reactions we will talk about: 1. Synthesis reactions 2. Decomposition reactions 3. Single replacement reactions 4. Double replacement reactions 5. Combustion reactions You need to be able to identify the type of reaction and predict the product(s)

4. 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) For example, 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!

5. Synthesis Reactions Synthesis (meaning to make) are – Also called Direct combination or combination reactions Addition reactions – typified by their single product. If you have a reaction in which at least 2 elements or compounds are reacted and produce a single product, the reaction is a synthesis reaction.

6. Synthesis Reactions reactant + reactant  1 product Basically: A + B  AB Example: 2H 2 + O 2  2H 2 O Example: C + O 2  CO 2 Note: Single Product! This is your clue that this is a synthesis or combination reaction.

7. Synthesis Reactions Here is another example of a synthesis reaction

8. Examples of Synthesis Reactions 2Na + SNa 2 S – This one is an example of two elements in atomic form (Na and S) combining to form a compound (sodium sulfide). 2H 2 + O 2 2H 2 O – In this example, A and B are two elements in molecular form (hydrogen and oxygen molecules), and the product is water, which is simply the chemical combination of hydrogen and oxygen.

9. Examples of Synthesis Reactions 4Fe + 3O 2 2Fe 2 O 3 – In this example, substance “A” is an element in atomic form (Fe), and substance “B” is an element in molecular form (O 2 ). The result is a direct chemical combination of the two elements (FeO, iron oxide, which is “rust”). CuO + H 2 OCu(OH) 2 – This is an example where both substances going into the reaction are molecules. The result is what you get when you “add” all of the atoms in the reaction together.

10. 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)  22NaCl (s) MgF 2(s) AlF 3(s) 232 Balanced

11. Decomposition Reactions Decomposition reactions are really just the opposite of a synthesis reaction. Remember, if you can make a substance, you should be able to break it back apart into its components. A good way to remember decomposition reactions to to remember what happens when something decomposes. It falls apart!

12. Decomposition Reactions Decomposition Reactions Decomposition reactions occur when a compound breaks up into the elements or in a few to simpler compounds 1 Reactant  Product + Product Basically: AB  A + B Example: 2 H 2 O  2H 2 + O 2 Example: 2 HgO  2Hg + O 2 Note: Single Reactant! The single reactant is your clue that this is a decomposition reaction.

13. Decomposition Reactions Another view of a decomposition reaction:

14. Decomposition Exceptions Carbonates and chlorates are special case decomposition reactions that do not go to the elements. Carbonates (CO 3 2- ) decompose to carbon dioxide and a metal oxide Example: CaCO 3  CO 2 + CaO Chlorates (ClO 3 - ) decompose to oxygen gas and a metal chloride Example: 2 Al(ClO 3 ) 3  2 AlCl 3 + 9 O 2 There are other special cases, but we will not explore those in this class

15. 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)  2 Pb (s) + O 2(g) Al (s) + N 2(g) 2

16. Practice Identify the type of reaction for each of the following synthesis or decomposition reactions, and write the balanced equation: N 2 (g) + O 2(g)  BaCO 3(s)  Co (s) + S (s)  NI 3(s)  2 N O 2 (g) N 2 (g) + 3I 2 (s) CoS (s) BaO (s) + CO 2 (g) 2 2

17. Single Replacement Reactions Single replacement reactions occur when one chemical takes the place of another in a reaction. In the typical single replacement reaction, an element trades places with one of the ions in a compound.

18. Single Replacement Reactions Single Replacement Reactions: A metal can replace a metal (+) OR a nonmetal can replace a nonmetal (-). element + compound  product + product 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!) When H 2 O splits into ions, it splits into H + and OH - (not H+ and O -2 !!)

19. Single Replacement Reactions Another view:

20. The Activity Series Not all single replacement reactions will occur. This depends upon the location of the elements present in the activity series Elements above MAY replace elements below; elements below MAY NOT replace elements above them on the series

21. You will be given a copy of this!!!!

22. 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 [If ZnCl 2 + H 2(g)  Zn (s) + HCl (aq) the reaction WOULD NOT OCCUR because Hydrogen is below zinc on the activity series] 2

23. 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)  Cu (s) + Al(NO 3 ) 3(aq) 22 2 3 32

24. Double Replacement Reactions Double Replacement Reactions occur when a metal replaces a metal in a compound and a nonmetal replaces a nonmetal in a compound two ions trade places and forming new compounds. Compound + compound  product + product AB + CD  AD + CB Notice that one ion from compound AB replaces one ion from compound CD.

25. 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 (aq)  AgCl (s) + NaNO 3(aq) Another example: K 2 SO 4(aq) + Ba(NO 3 ) 2(aq)  KNO 3(aq) + BaSO 4(s) 2

26. Solubility For a double replacement reaction to have occurred, a solid (precipitate) MUST be formed There are rules to determine which of the materials formed is the solid If no solid is formed, there is said to be no reaction.

27. Figure 8.4: The forming of solid AgCl.

28. Solubility Tables Solubility tables help determine which materials are soluble in water and which are not In general, Solubility Rules can be summarized as follows 1. All compounds containing alkali metal cations and the ammonium ion are soluble. 2. All compounds containing NO 3 -, ClO 4 -, ClO 3 -, and C 2 H 3 O 2 - anions are soluble. 3. All chlorides, bromides, and iodides are soluble except those containing Ag +, Pb 2+, or Hg 2 2+. 4. All sulfates are soluble except those containing Hg 2 2+, Pb 2+, Sr 2+, Ca 2+, or Ba 2+. 5. All hydroxides are insoluble except compounds of the alkali metals, Ca 2+, Sr 2+, and Ba 2+. 6. All compounds containing PO 4 3-, S 2-, CO 3 2-, and SO 3 2- ions are insoluble except those that also contain alkali metals or NH 4 +. You will be given a copy of this!!!!

29. Practice Predict the products. Balance the equation 1. HCl (aq) + AgNO 3(aq)  2. Pb(NO 3 ) 2(aq) + BaCl 2(aq)  3. FeCl 3(aq) + 3NaOH (aq)  4. H 2 SO 4(aq) + 2NaOH (aq)  HNO 3(aq) + AgCl (s) PbCl 2(s) + Ba(NO 3 ) 2(aq) Fe(OH) 3(s) + 3NaCl (aq) 2 H 2 O (l) + Na 2 SO 4(aq)

30. Combustion Reactions Combustion reactions are the ones that burn (or explode!). There are two types of combustion reactions—complete or incomplete reactions. These reactions are identified by their products. They either produce carbon monoxide and water or carbon dioxide and water.

31. Complete Combustion Reactions These reactions burn “efficiently” which means they produce carbon dioxide and water. These reactions typically burn cleanly and leave very little residue behind.

32. Combustion Reactions Combustion reactions occur when a hydrocarbon reacts with oxygen gas. This is also called burning!!! In order to burn something you need the 3 things in the “fire triangle”: 1) A Fuel (hydrocarbon) 2) Oxygen to burn it with 3) Something to ignite the reaction (spark)

33. Combustion Reactions In general: C x H y + O 2  CO 2 + H 2 O Products in combustion are ALWAYS carbon dioxide and water. (although incomplete burning does cause some by- products like carbon monoxide) Combustion is used to heat homes and run automobiles (octane, as in gasoline, is C 8 H 18 )

34. Complete Combustion Reactions CH 4 + O 2  CO 2 + H 2 O They may also be written: CH 4  CO 2 + H 2 O With O 2 written above the arrow Clues: CO 2 (carbon dioxide) in the product along with water O2O2

35. Combustion Example C 5 H 12 + O 2  CO 2 + H 2 O Write the products and balance the following combustion reaction: C 10 H 22 + O 2  C 10 H 22 + O 2  56 8 CO 2 + H 2 O10 11 2 3120 22 CO 2 + H 2 O 15.5

36. Incomplete Combustion Reactions Incomplete combustion reactions occur when something does not burn efficiently. This can cause a lot of harm if the gases produced cannot escape. Carbon monoxide,an odorless and colorless gas, is dangerous. People poisoned by this gas usually become sleepy and can die due to exposure.

37. Incomplete Combustion Reactions CH 4 + O 2  CO + H 2 O + CO 2 These reactions may also be written by: CH 4  CO + H 2 O + CO 2 Again, the O 2 is usually written over the arrow. Clue: CO (Carbon monoxide as a product.) O2O2

38. Combustion Reactions Edgar Allen Poe’s drooping eyes and mouth are potential signs of CO poisoning.

39. Mixed Practice State the type, predict the products, and balance the following reactions: 1. BaCl 2 + H 2 SO 4  2. C 6 H 12 + O 2  3. Zn + CuSO 4  4. Cs + Br 2  5. FeCO 3  BaSO 4 + HCl CO 2 + H 2 O ZnSO 4 + Cu CsBr FeO + CO 2

40. Predicting Products of Reactions Completing reactions requires knowledge of the different reaction types (sometimes called mechanisms). You must first identify the reaction type by the reactants. The only type of reaction that cannot be predicted this way is the combustion reaction since the products are very similar.

41. First Step: Identify reaction type Example: Al + O 2  Clue: 2 elements – Synthesis or combination reaction

42. Second Step: Write the net ionic equation for the reactants Al + O 2  becomes Al 3+ + O 2- 

43. Step 3 Using clues, complete reaction taking care to write each formula correctly by checking charges and “criss-crossing” if necessary. Al + O 2  Al 3+ O 2- Al + O 2  Al 2 O 3

44. Predicting Products of Reactions (cont.) For Single Replacement reactions, check activity series to make sure the reaction goes. Once you write the molecular equation, you should check for reactants and products that are soluble or insoluble. (Double Replacement only)

45. Reactions in Aqueous Solutions a.k.a. Net Ionic Equations Molecular Equations : shows complete formulas for reactants and products – Does not show what happens on the molecular level Total (or Complete) Ionic Equations : All substances that are strong electrolytes (are soluble and dissociate) are written as their ions. – Some ions participate in the reaction – Some ions do NOT participate in the reaction-called spectator ions. Net Ionic Equations : show only the ions that participate in the reaction

46. Writing Total Ionic Equations Once you write the molecular equation (synthesis, decomposition, etc.), you should check for reactants and products that are soluble or insoluble. We usually assume the reaction is in water We can use a solubility table to tell us what compounds dissolve in water. If the compound is soluble (does dissolve in water), then splits the compound into its component ions If the compound is insoluble (does NOT dissolve in water), then it remains as a compound

47. Writing Total Ionic Equations Molecular Equation: K 2 CrO 4 + Pb(NO 3 ) 2  PbCrO 4 + 2 KNO 3 SolubleSoluble Insoluble Soluble Total Ionic Equation: 2 K + + CrO 4 -2 + Pb +2 + 2 NO 3 -  PbCrO 4 (s) + 2 K + + 2 NO 3 -

48. Net Ionic Equations These are the same as total ionic equations, but you should cancel out ions that appear on BOTH sides of the equation Total Ionic Equation: 2 K + + CrO 4 -2 + Pb +2 + 2 NO 3 -  PbCrO 4 (s) + 2 K + + 2 NO 3 - (Spectator ions) Net Ionic Equation: CrO 4 -2 + Pb +2  PbCrO 4 (s)

49. Net Ionic Equations Try this one! Write the molecular, total ionic, and net ionic equations for this reaction: Silver nitrate reacts with Lead (II) Chloride in hot water. AgNO 3 + PbCl 2  Molecular: 2 AgNO 3 + PbCl 2  2 AgCl + Pb(NO 3 ) 2 Total Ionic: 2 Ag + + 2 NO 3 - + Pb +2 + 2 Cl -  2 AgCl (s) + Pb +2 + 2 NO 3 - Net Ionic: 2Ag + + 2Cl -  2AgCl (s)

50. Acid-Base Reactions Acid: – produces hydrogen ions (H + ) in solution (Arrhenius) – proton donor (Lewis) Base: – produces hydroxide ions (OH - ) in solution (Arrhenius) – proton acceptor (Lewis) The reaction ALWAYS forms water and an ionic compound (mostly aqueous, known as a salt). – The actual definition of a salt is the ionic product of an acid-base neutralization reaction)

51. Acid-Base Neutralization Chemically the reaction looks like this: Acid + BaseSalt + Water A classic example: HCl + NaOHNaCl + H 2 O Hydrochloric Acid Sodium Hydroxide (Lye) Sodium Chloride (Table Salt) Water

52. Example HNO 3 (aq) + KOH (aq)  Molecular: HNO 3 (aq) + KOH (aq)  Total Ionic: H + (aq) + NO 3 - (aq) + K + (aq) + OH - (aq)  H 2 O ( l ) + K + (aq) + NO 3 - (aq) Net Ionic: H + (aq) + OH - (aq)  H 2 O ( l ) Acid-Base Reactions H 2 O ( l ) + KNO 3 (aq)

53. Acid-Base Neutralization Here’s the equation again: – HCl + NaOHNaCl + H 2 O – Chemically, this is a double replacement reaction: The H traded its Cl for an OH The Na traded its OH for a Cl.

54. Oxidation-Reduction Reactions a.k.a. Redox Equations Between a metal and a nonmetal forming an ionic compound – Electron transfer occurs – Oxidation numbers: assigning an excess or deficiency in electrons for each element (the charge based on the compound).

55. Rules for Oxidation Number (ox. #) Determination 1. The sum of the oxidation numbers add up to the charge a. all elements have an ox. # of 0 b. ions of elements, ox. # is the charge ( Cl - ) c. the sum of the ox. # of a complex ion equals the charge ( CO 3 -2 ) 2. H is 1+ when combined with a nonmetal and 1- with a metal H 3 PO 4 CaH 2 H= 1+ PO 4 = 3- Ca = 2+ H= 1-

56. 1. F is always 1-; Cl, Br, I are 1- except when combined with each other or O 2. O is 2- except when combined with F ( F 2 O ) 3. Group I is 1+ and Group II is 2+ in their compounds Rules for Oxidation Numbers (cont.)

57. Recognizing Redox Rxns. 2 HCl (aq) + Mg (s)  MgCl 2 (aq) + H 2 (g) Net: 2 H + (aq) + Mg 0 (s)  Mg 2+ (aq) + H 2 0 (g) Loss of electron = oxidation Gain of electron = reduction – “LEO the lion goes GER" reduction oxidation 1+ 0 2+ 0 1- 2(-1)

58. Half Reactions Separate the individual oxidation and reduction reactions. Look at electron movement Half rxn.: Mg 0 (aq)  Mg 2+ + 2 e- 2e- + 2 H +  H 2 0 Net: 2 H + (aq) + Mg 0 (aq)  Mg 2+ (aq) + H 2 0 (q) Oxidizing agent: the one reduced (H+) Reducing agent: the one oxidized (Mg 0 )

59. Recognition of Redox rxns.  Oxidation # changes  Reactions with oxygen  Reaction of any element (forms a new compound) Balancing  Balance by mass  Balance by charge  Balance net ionic equation

60. Fe (s) + Cl 2 (aq)  FeCl 3 (aq) 1. Balance by mass 2 Fe (s) + 3Cl 2 (aq)  2 FeCl 3 (aq) 2. Write the ionic equation 2 Fe (s) + 3 Cl 2 (aq)  2 Fe 3+ (aq) + 6 Cl - (aq) 3. Write half reaction 2 Fe 0 (s)  2 Fe 3+ + 6 e- 6e- + Cl 2  2 Cl - Example Problem :

61. 3. Balance by charge (want # of e- to cancel) 2 Fe 0 (s)  2 Fe 3+ + 6e- 6e- + 3Cl 2 0  6 Cl - 2 Fe 0 (s) + 6e- + 3Cl 2  2 Fe 3+ + 6e- + 6 Cl - Final eqn.: 2 Fe (s) + 3Cl 2  2 Fe 3+ + 6 Cl - +