1. Chemical Reactions Unit
Balance and Write Chemical Equations
Predict Products of Chemical Reactions

2. Overview
the core of chemistry – how chemicals react with one another

3. Overview Two parts: 1. Language of chemical reactions
a) formulas (words/vocabulary) – from last unit
b) chemical equations (sentences)

4. Language of chemical reactions: b) chemical equations
Sequence:
i) balance skeleton formula equations,
ii) word equations  balanced formula equations,
iii) Sentence descriptions  balanced formula equations
iv) write complete, balanced formula equations to describe chemical reactions observed in lab or as demos

5. Overview 2. Different types of reactions a) classify
b) predict products
c) acids and bases, pH

6. Writing Chemical Equations
Example: Iron and chlorine gas react to produce the salt iron(III) chloride. Word equation: describes the reaction in words, using equation symbols Iron(s) + chlorine(g)  iron (III) chloride(s) Skeleton formula equation: describes the reaction using equation symbols and correct chemical formulas; unbalanced Fe(s) + Cl2(g)  FeCl3(s)

7. Writing Chemical Equations
Balanced formula equation: coefficients are used to equalize numbers of each atom on each side of the reaction
2 Fe(s) Cl2(g)  2 FeCl3(s)
Iron atoms on left (2) iron atoms on right (2)
Chlorine atoms on left (3 x 2 = 6) chlorine atoms on right (2 x 3 = 6)

8. Writing Chemical Equations
Total and net ionic equations: describe the reaction using equation symbols and chemical formulas, shown as ions if the compound is aqueous (dissolved in water) (10.3)
later

9. Steps for Writing and Balancing Chemical Equations
Write a word equation with the names for all reactants on the left and all products on the right:
Reactants  Products
Convert the word equation to a skeleton formula equation by writing the correct formulas for all reactants and products. Be sure that your formulas correctly represent the particles in the reaction!
3. Use coefficients in front of formulas to balance the equation. Do NOT change the formulas!
Begin balancing with an element that occurs only once on each side of the arrow.

10. Steps for Writing and Balancing Chemical Equations
Multiply coefficient x subscript to determine the # of atoms of a specific element in one "term" of the equation:
e.g. 4 H2O molecules:
4 x 2 = 8 H atoms x 1 = 4 O atoms
Balance one type of atom at a time.
Balance H and O last, especially for combustion reactions.

11. Balancing Hints 8. Even/Odd rule e.g. Fe + O2  Fe2O3
Since the number of oxygen atoms on the left must always be even.,
start by making the number of oxygen atoms on the right even, then balance the iron atoms.
4 Fe O2  2 Fe2O3

12. Balancing Hints 9. Intact polyatomic ions
Fe + Pb(C2H3O2)2  Fe(C2H3O2)3 + Pb
The acetate ion (C2H3O2) stays together as a group, so balance the ion as a group.
2 Fe Pb(C2H3O2)2  2 Fe(C2H3O2) Pb
One reason it is useful to know your ions!

13. We’ll focus on these at a later time.
Balancing Hints
Combustion reactions
We’ll focus on these at a later time.
Honors – check these hints out – two questions on WS #2

14. Chemical Equation Symbols
“yields” indicates the products of the reaction
Indicates a reversible reaction
A reactant or product in aqueous solution (dissolved in water)
A reactant or product in the liquid state
A reactant or product in the solid state
A reactant or product in the gaseous state
Reactants are heated
Pressure exceeding normal atmospheric pressure
Temperature at which reaction is carried out
Formula of catalyst used to alter the rate of the reaction

15. Practice – identify errors
For each of the following, explain why the equation is not properly balanced, then write the correctly balanced equation.   
1. Mg(NO3)2(aq) K(s)  Mg(s) + K2NO6(aq)
2. AlCl3(aq) + AgNO3(aq)  AgCl(s) + Al(NO3)3(aq)

16. Practice – identify errors
For each of the following, explain why the equation is not properly balanced, then write the correctly balanced equation.   
1. Mg(NO3)2(aq) K(s)  Mg(s) + K2NO6(aq)
  Formula for KNO3 is not written correctly = 2 KNO3
2. AlCl3(aq) + AgNO3(aq)  AgCl(s) + Al(NO3)3(aq)
Needs balancing  1, 3, 3, 1

17. Practice – Writing and Balancing Chemical Equations (p. 4)
Write the word equation, skeleton formula equation, and balanced formula equation for each of the following reactions:
Solid magnesium metal and solid silver sulfide react to form solid magnesium sulfide and solid metallic silver.
 2. Aqueous nitric acid and calcium hydroxide solutions react to form water and aqueous calcium nitrate
For each write the
Word equation:
 Skeleton formula equation:
 Balanced Formula equation:

18. Practice - KEY
Solid magnesium metal and solid silver sulfide react to form solid magnesium sulfide and solid metallic silver.
 Word equation:
 magnesium(s) + silver sulfide(s)  magnesium sulfide(s) + silver(s)
Skeleton formula equation:
 Mg(s) + Ag2S(s)  MgS(s) + Ag(s)
Balanced Formula equation:
 Mg(s) + Ag2S(s)  MgS(s) + 2 Ag(s)

19. Practice - KEY
 2. Aqueous nitric acid and calcium hydroxide solutions react to form water and aqueous calcium nitrate
Word equation:
 nitric acid(aq) + calcium hydroxide(aq)  water(l) + calcium nitrate(aq)
Skeleton formula equation:
 HNO3(aq) + Ca(OH)2(aq)  H2O(l) + Ca(NO3)2(aq)
Balanced Formula equation:
  2 HNO3(aq) + Ca(OH)2(aq)  2 H2O(l) + Ca(NO3)2(aq)

20. Warm Up (p. 5 – Wr & Bal Notes)
3. Aluminum metal reacts with oxygen in the air to form aluminum oxide
When solid mercury (II) sulfide is heated with oxygen, liquid mercury metal and gaseous sulfur dioxide are produced
Oxygen gas can be made by heating potassium chlorate in the presence of the catalyst manganese dioxide. Potassium chloride is left as a solid residue.

21. Warm Up (p. 5 – Wr & Bal Notes)
3. Aluminum metal reacts with oxygen in the air to form aluminum oxide
Aluminum(s) + oxygen(g)  aluminum oxide(s)
Al(s) O2(g)  Al2O3(s)
4 Al(s) O2(g)  2 Al2O3(s)

22. Warm Up (p. 5 – Wr & Bal Notes)
4. When solid mercury (II) sulfide is heated with oxygen, liquid mercury metal and gaseous sulfur dioxide are produced
mercury (II) sulfide(s) + oxygen(g)  mercury(l) + sulfur dioxide(g)
HgS(s) + O2(g)  Hg(l) + SO2(g) D D D

23. Warm Up (p. 5 – Wr & Bal Notes)
5. Oxygen gas can be made by heating potassium chlorate in the presence of the catalyst manganese dioxide. Potassium chloride is left as a solid residue.
potassium chlorate(s)  potassium chloride(s) + oxygen(g)
KClO3(s)  KCl(s) O2(g)
2 KClO3(s)  2 KCl(s) O2(g)
D, MnO2
D, MnO2
D, MnO2

24. Intro to Chemical Reactions Lab
Prepare lab notebook for this lab:
Read the first page of the instructions carefully before you begin
Include:
Overall purpose
List of materials and flowchart for each of the 8 stations (recommendation is 1 station per page) – minimum of 4 for Th/F
Include safety information and notes to yourself in your flowchart – see Flinn Scientific catalog, first section, on lab benches
Check out the lab setups themselves
Note checklists for each station (in the instructions) – use these to help you prepare your lab notebook so that you can focus on the lab itself

25. Intro to Chemical Reactions Lab Add the following info to your lab notebook
Splint test:
flame goes out:
carbon dioxide
flame flares up:
oxygen
“pop” and flame goes out:
hydrogen

26. Intro to Chemical Reactions Lab
Full lab gear – apron and goggles
Complete the experiment – observations, reaction predictions, etc. at each station before you continue to another station
When finished, begin working on WSs

27. Solubility Chart - used to predict the state of matter of products in DR reactions S = soluble = aqueous (aq); I = insoluble = solid (s)
Aluminum acetate
vs.
aluminum hydroxide

28. Warm up
Identify the reactants and products in the following word equations:
Water decomposes to produce hydrogen and oxygen gases.
Sodium chloride is produced when sodium metal reacts with chlorine gas. 
Methane gas reacts with oxygen gas to form carbon dioxide and water.
Briefly describe the differences between a, b and c (how the sentence is written) and how you identified the reactants and products.

29. Warm up
Write and balance the following reactions: 1. Aqueous solutions of silver nitrate and potassium iodide are mixed. Silver iodide and potassium nitrate are produced. Use your solubility chart to figure out which one is the precipitate. 2. When nitrogen dioxide is bubbled through water it produces nitric acid and nitrogen monoxide. What are the states of matter?

30. Warm up Write and balance the following reactions:
Aqueous solutions of silver nitrate and potassium iodide are mixed. Silver iodide and potassium nitrate are produced. Use your solubility chart to figure out which one is the precipitate.
silver nitrate(aq) + potassium iodide(aq)  silver iodide + potassium nitrate
AgNO3(aq) + KI(aq)  AgI(s) + KNO3(aq)

31. Warm up
Write and balance the following reactions:
2. When nitrogen dioxide is bubbled through water it produces nitric acid and nitrogen monoxide. What are the states of matter of nitrogen dioxide, nitric acid and nitrogen monoxide?
Nitrogen dioxide + water(l)  nitric acid + nitrogen monoxide
Nitrogen dioxide(g) + water(l)  nitric acid(aq) + nitrogen monoxide(g)
NO2(g) + H2O(l)  HNO3(aq) + NO(g)
3 NO2(g) + H2O(l)  2 HNO3(aq) + NO(g)

32. Evidence of a Chemical Reaction (Chemical Reactions notes)
Observation
What it means
Bubbles, inflates bag, disappears, steam, smell, smoke
Gas Formation/Release
Clear to cloudy, solid appears on bottom
Precipitate (solid) Formation
Flame, light, temperature change
Change in energy
Color change
Color change, e.g. acid/base indicator
We can only know the products of a chemical reaction by carrying out the reaction in the laboratory, BUT, we can make general predictions about the products of a reaction based on “types” of chemical reactions.

33. Types of Chemical Reactions (not all reactions fit these categories)
Synthesis (or Combination)
Decomposition
Single Replacement
Double Replacement (or Metathesis)
Combustion
6. Oxidation-Reduction (Redox) (Honors)

34. Energy Changes in Chemical Reactions
1.Exothermic Reaction
Releases energy (energy = one of the products)
C(s) + O2(g)  CO2(g) kJ DH = kJ

35. Energy Changes in Chemical Reactions
1.Endothermic Reaction
absorbs energy (energy = one of the reactants)
CaCO3(s) kJ  CaO(s) + CO2(g) DH = kJ

36. Synthesis (combination)
A + B → AB (one product)
usually produces energy – exothermic
Examples:
metal + nonmetal (often oxygen)
metal oxide + water  base
Cu + O2 → Cu2O
MgO + H2O → Mg(OH)2
Cu + O2 → CuO
Na + Br2 → NaBr
nonmetal oxide + water  acid
nonmetal + nonmetal
CO2 + H2O → H2CO3
N2 + O2 → NO2
SO2 + O2 → SO3
(Honors)
(Honors)

37. Synthesis reactions are usually Exothermic
Releases energy (energy is a product)
C(s) + O2(g)  CO2(g) kJ ΔH = kJ

38. Decomposition AB A + B (one reactant, > 1 product)
D or 7 or light
AB A + B (one reactant, > 1 product)
usually requires energy - endothermic
Examples
Compound elements
Ag2O  Ag + O2
reactant elements (electrolysis)
H2O  H2 + O2
Metal carbonate metal oxide + carbon dioxide (Honors)
MgCO3  MgO + CO2
Acid nonmetal oxide + water (Honors)
H2CO3  CO2 + H2O 7

39. Decomposition reactions are usually Endothermic
Absorbs energy (energy is a reactant)
CaCO3(s) kJ  CaO(s) + CO2(g) ΔH = kJ

40. Demonstration Electrolysis of Water
Record your observations in your notes.

41. Warmup
Write balanced chemical equations for the following synthesis and decomposition chemical reactions.
Figure out what the states of matter should be at room temperature.
Predict products according to the type of reaction.

42. Warmup Synthesis: 1. carbon + oxygen  2. gallium + oxygen 
3. nitrogen + hydrogen 
(H) 4. phosphorus (V) oxide + water 
(H) 5. calcium oxide + water 

43. Warmup Synthesis: carbon + oxygen  C(s) + O2(g)  CO2(g)
or 2 C(s) + O2(g)  2 CO(g)
gallium + oxygen 
4 Ga(s) + 3 O2(g)  2 Ga2O3(s)
nitrogen + hydrogen 
N2(g) H2(g)  2 NH3(g)

44. Warmup Honors Synthesis: (H) 4. phosphorus (V) oxide + water 
P2O5(s) + 3 H2O(l)  2 H3PO4(aq)
(H) 5. calcium oxide + water 
CaO(s) + H2O(l)  Ca(OH)2(s)

45. Warmup Decomposition: 1. sodium oxide  aluminum chloride 
(H) 3. calcium carbonate 
(H) 4. potassium chlorate 

46. Warmup Decomposition: sodium oxide  2 Na2O(s)  4 Na(s) + O2(g)
aluminum chloride 
2 AlCl3(s)  2 Al(s) + 3 Cl2(g) D D

47. Warmup Decomposition: (H) 3. calcium carbonate 
CaCO3(s)  CaO(s) + CO2(g)
(H) 4. potassium chlorate 
2 KClO3(s)  2 KCl(s) O2(g) D D

48. Use activity series to predict – higher replaces lower
Single Replacement
A + BX → AX + B
Use activity series to predict – higher replaces lower
Cation Replacement
Anion replacement
Metal replaces Metal
Halogen replaces Halogen
Mg + SnCl2  MgCl2 + Sn
Br2 + SrI2  SrBr2 + I2
Metal replaces H in HOH (H2O)
K + HOH  KOH + H2
Metal replaces H in an acid
Fe + H2SO4  FeSO4 + H2
(s)
(aq)
(aq)
(s)
(l)
(aq)
(aq)
(s)
(s)
(l)
(aq)
(g)
(s)
(aq)
(aq)
(g)

49. See your reference sheet
Activity Series (tt34, 10.2) Used to predict Single Replacement Reactions
See your reference sheet

50. Demonstration – Single Replacement Reaction
Use the activity series on your reference sheet to predict the products of the following reactions:
Cu(s) + AgNO3(aq) 
Ag(s) + Cu(NO3)2(aq) 
During the demo, record your observations

51. The Activity Series
For each of the following pairs of elements, circle the one that would replace the other element in a compound.
a. calcium, tin e. iron, copper
b. bromine, fluorine f. iodine, chlorine
c. aluminum, potassium g. silver, lead
d. zinc, calcium

52. The Activity Series
For each of the following pairs of elements, circle the one that would replace the other element in a compound.
a. calcium, tin e. iron, copper
b. bromine, fluorine f. iodine, chlorine
c. aluminum, potassium g. silver, lead
d. zinc, calcium

53. The Activity Series a. Li(s) + Fe(NO3)3(aq)  _________
For each of the following reactants, use the activity series to determine whether the reaction would take place or not. If no reaction takes place, write NR in the blank. If a reaction does take place, write the formulas for the products of the reaction. (Hint: If an active metal replaces the hydrogen in water, the hydroxide of the active metal forms. H-OH)
a. Li(s) + Fe(NO3)3(aq)  _________
b. Au(s) + HCl(aq)  __________
c. Cl2(g) + KBr(aq)  ___________
d. Cu(s) + Al(NO3)3(aq)  ________
e. Ag(s) + HBr(aq)  _________
f. Ni(s) + SnCl2(aq)  ___________

54. The Activity Series a. 3 Li(s) + Fe(NO3)3(aq)  3 LiNO3(aq) + Fe(s)
For each of the following reactants, use the activity series to determine whether the reaction would take place or not. If no reaction takes place, write NR in the blank. If a reaction does take place, write the formulas for the products of the reaction. (Hint: If an active metal replaces the hydrogen in water, the hydroxide of the active metal forms.)
a. 3 Li(s) + Fe(NO3)3(aq)  3 LiNO3(aq) + Fe(s)
b. Au(s) + HCl(aq)  NR
c. Cl2(g) KBr(aq)  2 KCl(aq) + Br2(l)
d. Cu(s) + Al(NO3)3(aq)  NR
e. Ag(s) + HBr(aq)  NR
f. Ni(s) + SnCl2(aq)  Sn(s) + NiCl2(aq)

55. The Activity Series
3. Magnesium metal can be used to remove tarnish from silver items. Silver tarnish is the corrosion that occurs when silver metal reacts with substances in the environment, especially those containing sulfur. Why would magnesium remove tarnish from silver?

56. The Activity Series
Magnesium metal can be used to remove tarnish from silver items. Silver tarnish is the corrosion that occurs when silver metal reacts with substances in the environment, especially those containing sulfur. Why would magnesium remove tarnish from silver?
Mg is more active than Ag and will replace it in Ag compounds, restoring the Ag metal.

57. The Activity Series
Use the activity series for metals to explain why copper metal is used in plumbing where the water might contain compounds of many different metals.

58. The Activity Series
Use the activity series for metals to explain why copper metal is used in plumbing where the water might contain compounds of many different metals.
Cu is not an active metal, compared with the metals dissolved in water such as Ca2+ and Mg2+. It does not replace these metals in the compounds dissolved in water. The Cu pipes remain intact.

59. The Activity Series
5. The last four metals in the activity series of metals are commonly referred to as the “coinage metals”. Why would these metals be chosen over more active metals for use in coins? Why do you think some more active metals, such as zinc or nickel, are sometimes used in coins?

60. The Activity Series
The last four metals in the activity series of metals are commonly referred to as the “coinage metals”. Why would these metals be chosen over more active metals for use in coins? Why do you think some more active metals, such as zinc or nickel, are sometimes used in coins?
Cu, Ag, Pt and Au would be more durable because they will not react as readily with substances in the environment. However, the more active metals, such as Ni and Zn are somewhat durable and cost much less.

61. Double Replacement AX(aq) + BY(aq) → AY(aq) + BX(s, g or l)
If both products are aqueous, NR
Use solubility table to predict
Precipitate-forming (s) Pb(NO3)2(aq) + KI(aq)  KNO3(aq) + PbI2(s) Gas-forming (g) (e.g. CO2 from a 2o decomposition rxn) NaHCO3(aq) + HC2H3O2(aq)  NaC2H3O2(aq) + H2O(l) + CO2(g) Water-forming (l) H2SO4(aq) + NaOH(aq)  Na2SO4(aq) + HOH(l) acid + base  salt + water

62. Solubility Chart - used to predict the state of matter of products in DR reactions S = soluble = aqueous (aq); I = insoluble = solid (s)
Aluminum acetate
vs.
aluminum hydroxide

63. Practice – SR and DR (p. 7 in Chem Rxns notes)
Use the activity series or solubility table to predict the product of the following reactions:
Na(s) + SrBr2(aq) 
CrI3(aq) + KCl(aq) 
Zn(s) + H2SO3(aq) 
K2CO3(aq) + HI(aq) 

64. Warm Up – SR and DR
Use the activity series or solubility table to predict the product of the following reactions:
2 Na(s) + SrBr2(aq)  NR
CrI3(aq) + 3 KCl(aq)  CrCl3(s) KI(aq) (DR – ppt)
Zn(s) + H2SO3(aq)  ZnSO3(aq) H2(g) (SR – metal + acid)
K2CO3(aq) + 2 HI(aq)  2 KI(aq) + H2CO3(aq) (DR – gas)
H2O(l) + CO2(g)

65. Practice – SR and DR (p. 7 in Chem Rxns notes)
Use the activity series or solubility table to predict the product of the following reactions:
Na(s) + H2O(l) 
HC2H3O2(aq) + (NH4)2S(aq) 
Fe(s) + CuCl2(aq) 
HBr(aq) + Ba(OH)2(aq) 

66. Warm Up – SR and DR
Use the activity series or solubility table to predict the product of the following reactions:
2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g) (SR – metal + H2O)
2 HC2H3O2(aq) + (NH4)2S(aq)  H2S(g) NH4C2H3O2(aq) (DR – gas)
Fe(s) + CuCl2(aq)  Cu(s) + FeCl2(aq) (SR – metal/metal)
2 HBr(aq) + Ba(OH)2(aq)  BaBr2(aq) + 2 H2O(l)
(DR – acid-base neutralization)

67. Total and Net Ionic Equations (Writing and Balancing Equations notes)
Dissolved potassium iodide and lead (II) nitrate react in aqueous solution to produce solid lead (II) iodide and dissolved potassium nitrate Word equation: Potassium iodide(aq) + lead (II) nitrate(aq)  lead (II) iodide(s) + potassium nitrate(aq) Formula (Skeleton) Equation KI(aq) + Pb(NO3)2(aq)  PbI2(s) + KNO3(aq) Balanced formula equation: 2 KI(aq) + Pb(NO3)2(aq)  PbI2(s) + 2 KNO3(aq)

68. Total and Net Ionic Equations
Balanced formula equation:
2 KI(aq) + Pb(NO3)2(aq)  PbI2(s) KNO3(aq)
Total ionic equation. Rewrite the equation so that all dissolved compounds (aq) (see solubility chart) are separated into their constituent ions:
2 K+(aq) + 2 I-(aq) + Pb2+(aq) + 2 NO3-(aq)  PbI2(s) + 2 K+(aq) + 2 NO3-(aq)
Spectator ions are those that appear on both sides of the equation and as such do not participate in the reaction. In the above example, the spectator ions are 2 K+(aq) and 2 NO3-(aq).
Net ionic equation:
Spectator ions cancel, and are not included in the net ionic equation:
Pb2+(aq) + 2 I-(aq)  PbI2(s)

69. Net Ionic Equations (p. 9 in Chem Rxns notes)
Complete Balanced Equation:
2 KOH(aq) + H2SO4(aq)  K2SO4(aq) + H2O(l)
Total Ionic Equation:
2 K+(aq)+ 2 OH-(aq)+ 2 H+(aq) + SO42-(aq) 2 K+ (aq) + SO42-(aq)+ 2 H2O(l)
Net Ionic Equation:
OH-(aq) + H+(aq)  H2O(l)

70. Net Ionic Equations What is a spectator ion?
What are the spectator ions in this reaction?
Compare and contrast each pair below.
a. Complete balanced equations, total ionic equations
b. Total ionic equations, net ionic equations

71. Net Ionic Equations For the reaction between aqueous silver
nitrate and aqueous sodium chloride, write
each of the following. The products of the
reaction are aqueous sodium nitrate and solid
silver chloride.
a. complete balanced equation
b. total ionic equation
c. net ionic equation

72. Net Ionic Equations
5. What is the net ionic equation for the reaction between aqueous calcium hydroxide and nitric acid? The products of this reaction are aqueous calcium nitrate and water. How does this net ionic equation compare to the net ionic equation shown on the earlier slide?

73. Upcoming Lab Quiz
With the help of your lab notebook, you will be expected to:
Write balanced chemical equations for reactions at all lab stations
Identify the type of reaction at each station
Provide safety information
Provide information regarding the various tests you performed at each station
Answer extension-type questions about each station, including writing equations for similar reactions.

74. Hydrocarbon + O2  CO2 + H2O + energy
5. Combustion
Hydrocarbon + O2  CO2 + H2O + energy
complete - forms CO2 + H2O + energy CH4 + O2  CO2 + H2O + energy incomplete - forms CO + H2O + energy CH4 + O2  CO + H2O + energy

75. Combustion reactions C2H2 + O2  CO2 + H2O 2 C2H2 + O2  CO2 + 2 H2O
e.g. C2H2 + O2  CO2 + H2O
Balance hydrogens, but at the same time make the number of oxygens in H2O even by doubling the coefficient if necessary – do NOT balance the oxygens at this point!
(i.e. make the number of Hs in the reactants divisible by 4)
C2H2 + O2  CO2 + H2O
2 C2H2 + O2  CO H2O
2. Balance carbons
2 C2H2 + O2  4 CO H2O
3. Lastly, balance the oxygen molecules on the left
2 C2H O2  4 CO H2O 

76. Practice - p. 8 of notes Balance Combustion Reactions
Is # of H in hydrocarbon divisible by 4?
If yes, hydrocarbon coefficient = 1
if no, hydrocarbon coefficient = 2
Balance hydrogen
Balance carbon
Balance oxygen LAST
Balance the following complete combustion reactions:
___C3H8(g) + ___O2(g)  ___CO2(g) + ___H2O(g)
___C5H12(g) + ___O2(g)  ___CO2(g) + ___H2O(g)

77. Balance Combustion Reactions
Now balance the chemical reactions for the complete combustion of hexane(C6H14) and decane (C10H22): ___C6H14(g) + ___O2(g)  ___CO2(g) + ___H2O(g) ___C10H22(g) + ___O2(g)  ___CO2(g) + ___H2O(g)

78. Combustion Practice (p. 8 of Chem Rxns notes)
Write balanced chemical equations for the complete combustion of
propane (C3H8)
C3H8(g) O2(g)  CO2(g) H2O(g) + heat
C3H8(g) O2(g)  CO2(g) H2O(g) + heat
C3H8(g) O2(g)  CO2(g) H2O(g) + heat
C3H8(g) O2(g)  CO2(g) H2O(g) + heat

79. Combustion Practice (p. 8 of Chem Rxns notes)
Write balanced chemical equations for the complete combustion of
decane (C10H22)
C10H22(g) O2(g)  CO2(g) H2O(g) + heat
2 C10H22(g) O2(g)  CO2(g) H2O(g) + heat
2 C10H22(g) O2(g)  20 CO2(g) H2O(g) + heat
2 C10H22(g) + 31 O2(g)  20 CO2(g) H2O(g) + heat

80. Combustion Practice (p. 8 of Chem Rxns notes)
Write balanced chemical equations for the complete combustion of
Ethyl alcohol (C2H5OH) – note the single O
C2H5OH(l) O2(g)  CO2(g) H2O(g) + heat
C2H5OH (l) O2(g)  CO2(g) + 3 H2O(g) + heat
C2H5OH (l) O2(g)  2 CO2(g) + 3 H2O(g) + heat
C2H5OH (l) O2(g)  2 CO2(g) + 3 H2O(g) + heat

81. Redox Reactions Electrons are transferred from one atom to another.
Oxidation = - electrons from a substance
Reduction = + electrons to a substance
Mg(s) HCl(aq)  MgCl2(aq) + H2(g)
Note: The transfer of electrons during oxidation-reduction reactions often produce energy (when spontaneous), which can be in the form of electricity.

82. Oxidation-Reduction Reactions
Mnemonics to help remember which is which:
LEO the lion says GER (Loss of Electrons = Oxidation/Gain of Electrons = Reduction)
OIL RIG (oxidation is loss and reduction is gain)

83. Redox Reactions Examples: a) Combustion: “rapid oxidation reaction in
which a large amount of heat and usually light
are released”
C + O2  CO2, CO
S + O2  SO2
What was oxidized?
What was reduced?

84. Redox Reactions Examples: b) Metal + acid (SR)
Zn(s) HCl(aq)  ZnCl2(aq) + H2(g)
Total ionic equation:
What was oxidized?
What was reduced?

85. Redox Reactions Examples: b) Metal + acid (SR)
Zn(s) HCl(aq)  ZnCl2(aq) + H2(g)
Total ionic equation:
Net ionic equation:
What was oxidized?
What was reduced?
Half-reactions:

86. Redox Reactions Examples: c) Metal + salt (SR)
Mg(s) + CoSO4(aq)  MgSO4(aq) + Co(s)
Total ionic equation:
What was oxidized? What was reduced?

87. Oxidation Numbers
help us keep track of what happens to the electrons in reactions

88. Oxidation Numbers Rules: 1. Elemental form – oxidation number = 0
2. Monatomic ion – oxidation number = charge on the ion
3. nonmetals – usually have negative oxidation numbers
a) Oxygen – usually -2 (exception is peroxide ion, oxygen has oxidation number of -1)
b) Hydrogen - +1 when bonded to nonmetals,
-1 when bonded to metals
c) Fluorine in all compounds (why?)
Other halogens in most situations,
positive when combined with oxygen
(e.g. oxyanions)
d) sum of oxidation numbers of all atoms in a neutral compound = 0
sum of oxidation numbers in a polyatomic ion = charge on the ion

89. Warm Up – SR  Redox
Write total and net ionic equations for the following reactions. Identify the redox pairs. Label oxidation numbers.
2 Na(s) + SrBr2(aq)  Sr(s) + 2 NaBr(aq)
Zn(s) + H2SO3(aq)  ZnSO3(s) H2(g)
2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g)
Fe(s) + CuCl2(aq)  Cu(s) + FeCl2(aq)

90. Warm Up Label redox pairs and oxidation numbers:
2 Na(s) + SrBr2(aq)  NR
Zn(s) + H2SO3(aq)  ZnSO3(aq) H2(g)
0 2 x
Zn is oxidized: Zn  Zn2+, H+ is reduced  H2

91. Warm Up Label redox pairs and oxidation numbers:
2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g)
0 4 x 1+, 2 x x 1+, 2 x 2-, 2 x
Na is oxidized: Na  Na+, H+ is reduced  H2
Fe(s) + CuCl2(aq)  Cu(s) + FeCl2(aq)
x 2+, no change 2+, no change 0
Fe is oxidized: Fe  Fe2+, Cu2+ is reduced Cu

92. Predicting Products Class of Reaction Reactants Probable Products
Synthesis
two or more substances
Decomposition
one compound
Single Replacement
a metal and a compound
a nonmetal (halogen) and a compound
Double Replacement (metathesis)
two compounds dissolved in water
Combustion
(restricted definition)
hydrocarbon + oxygen
Oxidation-Reduction (Honors)
one or more substances

93. Predicting Products Class of Reaction Reactants Probable Products
Synthesis
two or more substances
one compound
Decomposition
two or more elements or compounds
Single Replacement
a metal and a compound
a nonmetal (halogen) and a compound
a new compound and the replaced metal or hydrogen
a new compound and the replaced nonmetal (halogen)
Double Replacement (metathesis)
two compounds dissolved in water
two different compounds, one of which is a solid, water, or a gas
Combustion
(restricted definition)
hydrocarbon + oxygen
carbon dioxide (or carbon monoxide) and water
Oxidation-Reduction (Honors)
one or more substances
two or more elements or compounds with different oxidation numbers

94. Summary of Reaction Types
For each set of reactants listed below, identify the type of reaction that the reactants might undergo. List as many reaction types as may apply. Assume that all the reactants for the reaction are listed.
a. a compound and an element
b. two compounds
c. one compound

95. Summary of Reaction Types
For each set of reactants listed below, identify the type of reaction that the reactants might undergo. List as many reaction types as may apply. Assume that all the reactants for the reaction are listed.
a. a compound and an element
synthesis, combustion, SR
b. two compounds
synthesis, DR
c. one compound
decomposition

96. Summary of Reaction Types
For each set of reactant products listed below, identify the type of reaction that might have formed the products. List as many reaction types as may apply. Assume that all the productions for the reaction are listed.
a. a compound and an element
b. two compounds
c. one compound

97. Summary of Reaction Types
For each set of reaction products listed below, identify the type of reaction that might have formed the products. List as many reaction types as may apply. Assume that all the productions for the reaction are listed.
a. a compound and an element
decomposition, SR
b. two compounds
decomposition, combustion, DR
c. one compound
synthesis

98. Summary of Reaction Types
Classify each of the following examples according to the type of reaction involved. List as many reaction types as may apply.
a. A match burns
b. The carbonic acid found in soft drinks breaks down into bubbles of carbon dioxide and water
c. Phosphorus and oxygen react rapidly, forming diphosphorus pentoxide
d. An iron nail is placed into a copper sulfate solution. Copper metal appears on the nail.

99. Summary of Reaction Types
Classify each of the following examples according to the type of reaction involved. List as many reaction types as may apply.
a. A match burns
combustion
b. The carbonic acid found in soft drinks breaks down into bubbles of carbon dioxide and water
decomposition
c. Phosphorus and oxygen react rapidly, forming diphosphorus pentoxide.
synthesis
d. An iron nail is placed into a copper sulfate solution. Copper metal appears on the nail. SR

100. Summary of Reaction Types
Classify each of the following examples according to the type of reaction involved. List as many reaction types as may apply.
e. The acid in baking powder reacts with baking soda (NaHCO3), forming carbon dioxide gas and other products.
f. Water and sulfur trioxide react to form sulfuric acid.
g. Copper wire is placed in a silver nitrate solution. The solution turns blue, which is the color of the copper ion, and solid silver forms on the wire.

101. Summary of Reaction Types
Classify each of the following examples according to the type of reaction involved. List as many reaction types as may apply.
e. The acid in baking powder reacts with baking soda (NaHCO3), forming carbon dioxide gas and other products.
DR – followed by decomposition, gas-forming,
acid-base neutralization
f. Water and sulfur trioxide react to form sulfuric acid.
synthesis
g. Copper wire is placed in a silver nitrate solution. The solution turns blue, which is the color of the copper ion, and solid silver forms on the wire. SR

102. Summary Write, Balance and Predict Chemical Reactions:
Write and balance chemical equations
Identify evidence of chemical reactions
Identify 6 types of chemical reactions
Predict products of reactions
Write total and net-ionic equations for single and double replacement reactions

103. Warm Up - Extra
Write complete balanced equations and name the type of reaction:
Aqueous solutions of aluminum chloride and potassium phosphate are mixed.
Methane (CH4) is burned in air.
Sodium and chlorine gas are mixed.
A piece of chromium is put into a solution of nickel sulfate.
Copper (I) sulfide is heated.
Aqueous solutions of chloric acid (HClO3) and potassium hydroxide are mixed.
A piece of lead is put into a solution of nitric acid (HNO3).

104. Warmup (extra practice) not in notes
Write the complete balanced equation, total ionic equation and net ionic equation for the reaction between solutions of aluminum chlorate and potassium oxide.
2. Write the complete balanced equation, total ionic equation and net ionic equation for the acid-base neutralization reaction that occurs when aqueous nitric acid is mixed with aqueous potassium hydroxide.

105. Warmup (not in notes)
Write the complete balanced equation, total ionic equation and net ionic equation for the reaction between solutions of aluminum chlorate and potassium oxide.
2 Al(ClO3)3(aq) K2O(aq)  Al2O3(s) + 6 KClO3(aq)
2 Al3+(aq)+ 6 ClO3-(aq)+ 6 K+(aq)+ 3 O2-(aq)  Al2O3(s) + 6 K+(aq)+ 6 ClO3-(aq)
2 Al3+(aq)+ 3 O2-(aq)  Al2O3(s)

106. HNO3(aq) + KOH(aq)  H2O(l) + KNO3(aq)
Write the complete balanced equation, total ionic equation and net ionic equation for the acid-base neutralization reaction that occurs when aqueous nitric acid is mixed with aqueous potassium hydroxide.
HNO3(aq) + KOH(aq)  H2O(l) + KNO3(aq)
H+(aq) + NO3-(aq) + K+(aq) + OH-(aq)  H2O(l) + K+(aq) + NO3-(aq)
H+(aq) + OH-(aq)  H2O(l)

107. Total and Net Ionic Equations Practice (extra)
When solutions of potassium hydroxide and aluminum chloride are mixed, they produce solid aluminum hydroxide and aqueous potassium chloride.
In your notes, write word, skeleton, balanced formula, total ionic and net ionic equations for this reaction.

108. Extra Warm up for early on
Name the following compounds:
Na2Cr2O7
Fe(C2H3O2)3
AlI3
KHSO4
SnO2
Co(NO2)2
Write the following as balanced chemical equations:
iron (III) chloride + calcium hydroxide  iron (III) hydroxide + calcium chloride
carbon + oxygen  carbon monoxide
potassium nitrate  potassium nitrite + oxygen