1. Ch. 4 Patterns of Chemical Reactivity
Demo: Pour together two clear colorless liquids
Did a chemical reaction occur? How do you know?
Demo: AlkaSeltzer in water, or calcium in water
Did a chemical reaction occur?
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2. Observing and Predicting Reactions
How do we know whether a reaction occurs? What clues does nature offer? Make a list.
Review photos of reactions
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3. What clues does nature offer that a chemical reaction occurred?
ppt
crystal
color
gas
fumes
smoke
temperature
flames
magnetic
sound
light
solid decomp
explosion
solid dissol.
odor
elect. cond.
pH change
density
electrictyy
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4. Precipitate Formation
Ag+ + Cl-  AgCl Cr3+ + 3OH-  Cr(OH)3
Ba2+ + CrO42-  BaCrO4
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5. Solid Decomposition (NH4)2Cr2O7(s)  Cr2O3(s) + 4H2O(g) + N2(g)
CuSO4.5H2O 
CuSO4 + 5H2O(g)
2NI3  N2 + 3I2
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6. Gas Bubbles Ca + 2H2O  Ca(OH)2 + H2(g) Cr + 2H+  Cr2+ + H2(g)
Mg + 2HCl  MgCl2 + H2
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7. Fumes/Gas Formation Cu + 4HNO3  Cu(NO3)2 + 2NO2 + 2H2O
2H2O2  2H2O + O2
Zn + I2  ZnI2
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8. Smoke
2Al + 3Br2  2AlBr3 2Na + Cl2  2NaCl
P4 + 5O2  P4O10
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9. Flames 2Na + Cl2  2NaCl 2CrO3 + 3C2H5OH  Cr2O3 + 3CH3CHO + 3H2O
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10. Light 2Fe + 3O2  Fe2O3 2CH3OH + O2  2CO2 + 4H2O 2H2 + O2  2H2O
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11. Temperature Change
Ba(OH)2.8H2O + NH4Cl
Thermite: Al + Fe2O3
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12. Color Change Cl2 + 2I-  2Cl- + I2 Cu + 4HNO3  Cu(NO3)2 + 2NO2 + 2H2O
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13. Crystal Formation/Solid Deposition
Cu + 2Ag+  Cu2+ + 2Ag
Zn + Sn2+  Zn2+ + Sn
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14. Solid Dissolution Mg(OH)2 + 2HCl  MgCl2 + 2H2O
AgCl(s) + 2NH3(aq)  Ag(NH3)2+ + Cl-
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15. Sound
Oxidation of sugar
Fireworks
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16. Explosion
Dynamite
Building Demolition
Whale Removal
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17. Odor
Certain molecules, especially those containing sulfur or nitrogen, have distinctive odors.
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18. Electrical Conductivity
Ba(OH)2 + H2SO4
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19. Density/Volume
Sugar + H2SO4
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20. pH Change
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21. Magnetic Properties
Fe + S8

22. Generate Electricity
Chemical reaction in the battery
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23. 4.4 Observing and Predicting Reaction Patterns
Predictions:
do an experiment
use periodicity
use classifications of reactions
example: combustion reactions involve the reaction of an element or a compound with oxygen, usually with the evolution of heat

24. Reaction Classifications
In the following particulate representations, a circle represents an atom and different circles represent different elements. Using these representations, draw pictures of all the different types of atomic/molecular changes these substances could undergo. D F C B B E A

25. Combination (or Synthesis)+ B B B B A D D + C C A A F D D C C E + E F

26. Decomposition D D C C + F F E E +

27. Single Replacement D D + C + C A A F F A + E A + E

28. Double Replacement D F F D C + E C + E
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29. Reactions in Solution Precipitation Reactions
compound 1 + compound 2  compound 3 + compound 4
Also called double replacement or metathesis reactions.
exchange of ionic partners AB + CD  AD + CB
Pb(NO3)2(aq) + K2CrO4(aq)  PbCrO4(s) + 2KNO3(aq)
And other related reactions
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30. Double-Replacement Reactions (Metathesis)
Reactants Products

31. Precipitation Reactions
Precipitation reactions: (An example of a "double replacement" or "metathesis" reaction).
Precipitation Reactions form a solid when two solutions are combined.
An example is the combining aqueous potassium chromate with aqueous lead nitrate to form the precipitate lead chromate (still used in school bus paint!!)

32. Describing Reactions in Solution
To identify the precipitate or predict the formation of a precipitate the solubilities of compounds can be used. These rules should already be memorized!
Table 4.1, pg 144

33. Solubility Principles
Most nitrate and acetate salts are soluble.
Most salts containing the alkali metal ions (Li+, Na+, K+, Cs+, Rb+) and the ammonium ion (NH4+) are soluble.
Most chloride, bromide and iodide salts are soluble. Notable exceptions are salts containing the ions Ag+, Pb2+ and Hg22+.

34. Solubility Principles
Most sulfate salts are soluble. Notable exceptions are BaSO4, PbSO4, HgSO4 and CaSO4.
Most hydroxide salts are only slightly soluble. Important soluble hydroxides are NaOH and KOH. Ca(OH)2, Sr(OH)2, and Ba(OH)2 are somewhat soluble*.
Most sulfide, carbonate, chromate, and phosphate salts are only slightly soluble**.
* Note Group 2 trends : As you go down the group sulfate solubility decreases and hydroxide solubility increases.
** Slightly soluble compounds will form precipitates using "normal" concentrations.

35. Describing Reactions in Solution
For reactions involving ionic compounds, we can write the reaction as a molecular equation (or formula equation). This shows the normal (complete) formulas of all compounds:
Example:
K2CrO4(aq) + Pb(NO3)2(aq)  PbCrO4(s) KNO3(aq)

36. Describing Reactions in Solution
We can rewrite the same reaction as a complete ionic equation - Shows a picture of all that actually occurs in solution
strong electrolytes represented as ions in solution
weak and non- electrolytes still written in molecular (non-ionized) aqeuous state.
Example:
2K+(aq) + CrO42-(aq) + Pb2+(aq) NO3-(aq) 
PbCrO4(s) K+(aq) + 2 NO3-(aq)

37. Describing Reactions in Solution
A net ionic equation includes only the solution components involved in the reaction (spectator ions, which do not undergo change, are omitted)
Pb2+(aq) + CrO42-(aq)  PbCrO4(s)

38. Stoichiometry
Stoichiometry of Precipitation Reactions – based on Chapter 3 stoichiometry concepts, but using molarity (concentration) relationships.
Practice with Chapter 3 & Molarity!
Sample: How many grams of lead(II) hydroxide can be formed when 22.5 mL of M Pb(NO3)2 solution reacts with excess sodium hydroxide? (Hint: Use a BCA table).

39. Acid-Base Reactions Definitions: Arrhenius: Brønsted-Lowry :
Acid - forms H+ ions in solution (e.g HCl)
Base - forms OH- ions in solution (e.g. NaOH)
Brønsted-Lowry :
Acid - proton (H+) donor (e.g. HCl)
Base - proton acceptor
e.g. NH3: NH3 + H+  NH4+

40. Acid-Base Reactions General reaction :
Acid + base(metallic hydroxide)  salt + water
(neutralization reaction)
e.g. HCl and NaOH
molecular equation :
HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)
complete ionic equation. :
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) 
H2O(l) + Na+(aq) + Cl-(aq)
net ionic equation: H+ (aq) + OH- (aq)  H2O (l)

41. Acid-Base Titrations
Acid-base titrations (volumetric analysis) – determine an unknown quantity through titration.
Titration involves adding a precisely measured volume of a solution of known concentration (the titrant) into a solution containing the substance being analyzed (the analyte).
The titrant reacts with the analyte in a known manner, such as an acid-base reaction.

42. Acid-Base Titrations
An indicator marks the equivalence point (or stoichiometric point) where just the right amount of titrant has been added to completely react with the analyte.
The endpoint is where the indicator actually changes color, which hopefully occurs near the equivalence point.

43. Acid-Base Reactions Note the similarities to precipitation reactions.
Acid-Base reactions are another variation of a double replacement reaction. The key is the production of water.
Other common double replacement reactions produce gases.

44. Acid-Base Stoichiometry
There are numerous variations on the acid-base reaction. Be sure to read through the many examples in Section We will consider these examples now from a “modeling” perspective.

45. Acid-Base Reactions
You first want to examine the acid-base reaction (similar to predicting a precipitation reaction). Here are some general steps (they can and should vary depending on the problem):
1. List the major species present in solution before the reaction occurs. Decide what reaction will occur (look for formation of water or gases)
2. Write a balanced equation. (leave space for a BCA table)
3. Calculate the moles of reactants. For solutions, use the volumes of the original solutions and their molarities (before mixing). Input into a BCA table.

46. Acid-Base Reactions 4. Determine the limiting reactant if appropriate.
5. Analyze the problem and find the moles of reactant or product asked for.
6. Convert to grams or volume of solution if asked for
*All problems are different. Don’t “force” a problem into a particular solution method.

47. Reaction Classes Combination Reactions
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element + element  compound
metal + nonmetal  ionic compound
2Na(s) + Cl2(g)  2NaCl(s)
nonmetal + nonmetal  covalent cmpd
2H2(g) + O2(g)  2H2O(l)
Draw a molecular diagram of this type of reaction
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48. Combination Reactions
Reactants Product

49. Combination: K + Cl2

50. Reaction Classes Addition Reactions
element + compound  compound
Cl2 + 2TiCl3  2TiCl4
Cl2 + C2H4  C2H4Cl2
Draw a molecular diagram of this type of reaction

51. Addition Reactions
Reactants Product

52. Reaction Classes Decomposition Reactions
Compound  2 elements or element + compound or 2 compounds
Oxides, peroxides  O2
Nitrates  NO2 or NO2-
Carbonates  CO2
Hydrates  H2O
Ammonium salts  NH3
Draw a molecular diagram of this type of reaction
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53. Decomposition of HgO

54. Decomposition Reactions
Reactant Products

55. Decomposition Reactions
Reactant Products

56. Reaction Classes Single-Displacement Reactions
element + cmpd  cmpd + element (The more metallic element in the compound is displaced.)
carbon + metal oxides
3C + Fe2O3  3CO + 2Fe
metals + water
Ca(s) + 2H2O(aq)  Ca(OH)2(aq) + H2(g)

57. Single Displacement: Li + H2O

58. Single-Displacement Reactions
Reactants Products

59. Reaction Classes Single-Displacement Reactions
metals + acids
Fe(s) + 2HCl(aq)  FeCl2(aq) + H2(g)
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60. Reaction Classes Single-Displacement Reactions
metals + metal salts
Zn(s) + SnCl2(aq)  ZnCl2(aq) + Sn(s)
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61. Single Displacement: Cu + AgNO3

62. Reaction Classes Single-Displacement Reactions
nonmetals + salts
Cl2(aq) + 2KI(aq)  2KCl(aq) + I2(aq)
What do all these types of reactions have in common???

63. Oxidation-Reduction Reactions
reactions in which electrons are transferred
causes a change in the charge of an ion or of oxidation state of an element in a molecule
Oxidation states - numbers assigned to elements
used to keep track of electrons (not the same as charge, but related)

64. Rules for Assigning Oxidation States (Table 4.2)
The oxidation state of an uncombined element is zero (includes diatomic elements H2,N2, O2, F2, Cl2, Br2 and I2).
The oxidation state of a monatomic ion is the same as its charge (e.g. the sulfide ion, S2-, has an oxidation state of -2).

65. Rules for Assigning Oxidation States (Table 4.2)
Oxygen has an oxidation state of -2 in covalent compounds (except in peroxides (O22-) where each oxygen is assigned an oxidation state of -1).
In covalent compounds hydrogen is assigned an oxidation state of +1. (Hydrogen has a -1 charge in hydrides such as lithium hydride (LiH) or sodium hydride (NaH).

66. Rules for Assigning Oxidation States (Table 4.2)
In compounds, fluorine always has an oxidation state of -1.
The sum of the oxidation states of the elements in a neutral compound must equal zero.
The sum of the oxidation states of the elements in a polyatomic ion must equal the charge on the polyatomic ion.

67. Rules for Assigning Oxidation States (Table 4.2)
Oxidation states may be non-integers. For example in iron (III) oxide (Fe3O4), the iron has an oxidation state of 8/3 (eight-thirds).

68. Rules for assigning oxidation states
Practice: Identify the oxidation state of each atom in the following compounds:
Magnesium nitrate
Lithium nitride
Sodium nitrite

69. Characteristics of Redox Reactions
Oxidation
a loss of electrons
an increase in oxidation state
the substance oxidized is the reducing agent (gives electrons to another substance)
Reduction
a gaining of electrons
a decrease in oxidation state
the substance reduced is the oxidizing agent (takes electrons away from another substance)

70. Balancing Redox Reactions
By the half-reaction method :
In acidic solution
1. Write separate oxidation and reduction reactions for the reaction.
2. For each half reaction :
balance all the elements except hydrogen and oxygen
balance oxygen atoms using H2O
balance hydrogen atoms using H+
balance the charge using electrons

71. Balancing Redox Reactions
3. If necessary, balance electrons lost and gained in each half reaction by multiplying one or both half reactions by an integer.
4. Add the half-reactions and cancel out like species.
5. Check to make sure charges and elements are balanced.

72. Example 4.19
Potassium dichromate is a bright orange compound that can be reduced to a blue-violet solution of chromium(III) ions. In acidic conditions, potassium dichromate reacts with ethyl alcohol as follows:
Cr2O72-(aq) + C2H5OH(l)  Cr3+(aq) + CO2(g) + H2O(l)
Balance this equation using the half reaction method.

73. Balancing Redox Reactions
In Basic solution (see example 4.20):
1. Balance as in an acidic solution (see above).
2. Add a number of OH- ions equal to the H+ ions present to both sides of each half reaction to for H2O.
3. Eliminate the number of H2O molecules that appear on both sides of the equation.
4. Check to make sure charges and elements are balanced.

74. Group (Partner) Quiz
1. Give the oxidation state of each element in sodium chlorate (NaClO3)
2. In the following reaction, identify the oxidizing agent, the reducing agent, the substance being oxidized, and the substance being reduced
Br – (aq) + MnO4-(aq)  Br2(l) + Mn2+(aq)
3. Balance the above RedOx reaction that occurs in acidic solution.

75. Classify the following reactions, based on the changes happening at an atomic/molecular level.
1. AlF3(aq) + 3H2O(l)  Al(OH)3(s) + 3HF(aq)
2. BaCl2(aq) + Na2SO4(aq)  BaSO4(s) + 2NaCl(aq)
3. Ca(OH)2(s)  CaO(s) + H2O(g)
4. Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
5. CaO(s) + CO2(g)  CaCO3(s)
6. Cl2(aq) + 2NaI(aq)  2NaCl(aq) + I2(aq)
7. Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)
8. Fe(s) + 2HCl(aq)  FeCl2(aq) + H2(g)
9. H2SO3(aq)  H2O(l) + SO2(g)
10. 2HgO(s)  2Hg(l) + O2(g)
11. KOH(aq) + HNO3(aq)  KNO3(aq) + H2O(l)
12. 4Li(s) + O2(g)  2Li2O(s)
13. Na2S(aq) + 2HCl(aq)  2NaCl(aq) + H2S(g)
14. NH3(g) + HCl(g)  NH4Cl(s)
15. NiCO3(s)  NiO(s) + CO2(g)
16. P4(s) + 10F2(g)  4PF5(g)
double displacement
decomposition
single displacement
combination