1. Unit 5

2. This unit will introduce the following tools that scientists use to “understand reactions”
Section 1 Chemical Reactions & Types
Section 2 Balancing Equations
Section 3 Predicting Products of Synthesis & Decomposition Reactions
Section 4 Predicting Products of Single, Double Replacement & Combustion Reactions

3. Section 1 What is a Chemical Reaction?

4. What is a Chemical Reaction?
Bonds are broken and atoms are rearranged to form new compounds.
2 H2 + O2  2 H2O O H H
Watch as 2 H2 and 1 O2 undergo a chemical reaction

5. Chemical Reactions
Bonds and atoms are rearranged to form new compounds.
2 H2 + O2  2 H2O H O O H O H
The compounds in the end are different from those in the beginning H
Bonds are broken and formed between different atoms

6. What is a Chemical Reaction?
Also known as a chemical change

7. Observations (Indicators) of a Chemical Reaction(change)
Gas production Odor change Color change
Precipitate formation: insoluble solid that is formed when 2 aqueous solutions are reacted
Energy change(change in temperature)
Exothermic Reactions: reactants release heat & temp. of
the reaction increases
Endothermic Reactions: reactants absorb heat & temp. of
the reaction decreases

8. What are Chemical Equations?
The “sentence” of chemistry
Uses chemical formulas & other symbols instead of words to show a chemical reaction
Reactants
Products
Starting Materials
New substances formed
NaCl(aq) + AgNO3(aq)  AgCl(s) + NaNO3(aq)
Arrow
: Read as: Yields, Produces, Forms
Physical State
s = solid l = liquid
g = gas aq = aqueous (dissolved in water)

9. Examples of Chemical Equations
Word equation:
Magnesium metal is reacted with aqueous hydrochloric acid to produce aqueous magnesium chloride and hydrogen gas 
Formula equation:
Mg (s) + 2 HCl (aq) MgCl2 (aq) + H2 (g)
Visualization

10. Writing Chemical Equations
You must write the correct chemical formula for each reactant & product
“and”, “is mixed with” or “reacts with” = +
“yield”, “produces” and “forms” = 
Remember “BrINClHOF”: to represent the 7 diatomic elements that have a subscript of 2 in their chemical formula
Br2 I2 N2 Cl2 H2 O2 F2

11. Write the word equation into symbol form
Let’s Practice #1
Potassium metal is reacted with calcium bromide to form potassium bromide and calcium
Example:
Write the word equation into symbol form

12. Write the word equation into symbol form
Let’s Practice #1
Potassium metal is reacted with calcium bromide to form potassium bromide and calcium
Example:
Write the word equation into symbol form
K + CaBr2  KBr + Ca

13. Write the word equation into symbol form
Let’s Practice #2
Copper (II) nitrate and sodium hydroxide form copper (II) hydroxide and sodium nitrate
Example:
Write the word equation into symbol form

14. Write the word equation into symbol form
Let’s Practice #2
Copper (II) nitrate and sodium hydroxide form copper (II) hydroxide and sodium nitrate
Example:
Write the word equation into symbol form
Cu(NO3)2 + NaOH  Cu(OH)2 + NaNO3

16. Synthesis(Combination): when 2 or more substances combine to form a SINGLE substance
Element + Element  Compound
Cu S  CuS
Compound + Compound  Compound
Li2O H2O  LiOH

18. Decomposition: when 1 substance breaks apart into 2 or more substances; requires Energy to initiate
Compound  Element Element
CaS  Ca S
Compound  Compound + Compound
H2O  H2O O2

20. Single Replacement: when a more reactive element replaces a less reactive element in a compound
Compound + Element  Element Compound
AgI K  Ag KI

22. Double Replacement: when the positive ions of each aqueous compound replace each other+ + 
Compound + Compound  Compound + Compound
KI AgNO3  AgI KNO3

25. Combustion: reaction of an element or compound with OXYGEN to form an oxide & produce heat hydrocarbons combust to always produce carbon dioxide,water & heat
Mg + O2  MgO
CH4 + O2  CO2 + H2O

26. Section 2—Balancing Equations
We need to finish writing those equations we started!

27. Law of Conservation of Mass
Matter cannot be created nor destroyed just changed in form. Thus mass must remain constant through changes.

28. Conservation of Mass CaCl2 + Na2SO4  CaSO4 + 2NaCl
Mass before(reactants) = mass after(products)
# atoms before (reactants) = # atoms after(products)

29. How Does the Law Lead to Balancing?
Matter cannot be created nor destroyed
during a chemical or physical change.
The number of atoms of each element
must be the same on both sides
To ensure this we must balance the
equation .

30. How Do We Begin? Subscripts tell you how many atoms
Start by Counting the Atoms of each Substance
Subscripts tell you how many atoms
Coefficients are always multiplied by subscripts
With a parentheses, inner & outer subscripts are always multiplied together

31. Try These! Count the atoms in each compound
(a) 2 (NH4)3PO N= ___ H= ___ P= ___ O= ___
(b) 4 KC2H3O K= ___ C= ___ H= ___ O= ___

32. How Do We Balance Equations?
Practice Problems: How many atoms of each type are indicated in
the following compounds?
(a) 2 (NH4)3PO N= ___ H= ___ P= ___ O= ___
(b) 4 KC2H3O K= ___ C= ___ H= ___ O= ___
(c) 3 Ca(NO3) Ca= ___ N= ___ O= ___ 6 24 2 8 4 8 12 8 3 6 18

33. How to Balance Chemical Equations:1
Make a table of elements & count the atoms of each substance
_____ C H4 +
_____ O2 
_____ H2 O +
_____ C O2
Reactants
Products C H O

34. How to Balance By Inspection:2
Count the number of each element or ion on the reactants and products side.
Don’t forget to add all the atoms of the same element together—even if it appears in more than one compound!
_____ C H4 +
_____ O2 
_____ H2 O +
_____ C O2
Reactants
Products C 1 1 H 4 2 O 2 3

35. How to Balance By Inspection:2
Use coefficients to balance the numbers
Each time you add a coefficient, update your table with the new quantities of each atom.
_____ C H4 +
_____ 2 O2 
_____ 2 H2 O +
_____ C O2
Reactants
Products C 1 1 H 4 2 4 O 2 4 3 4

36. How to Balance By Inspection:
OPTIONAL Place a “1” in any empty coefficient location 3
Filling each coefficient location lets you and the grader know that you finished the problem rather than you left some blank because you weren’t done!
_____ 1 C H4 +
_____ 2 O2 
_____ 2 H2 O +
_____ 1 C O2
Reactants
Products C 1 1 H 4 2 4 O 2 4 3 4

37. Coefficients are added to change the number of atoms in a substance!
Total:
4 H
2 O
Total:
4 H
2 O
The equation is balanced.

38. Do You Remember Polyatomic Ions?
Polyatomic ion – Group of atoms that together has a net charge
e.g Nitrate NO31-
Carbonate CO32-
A polyatomic ion is a group of atoms that together have a charge.

39. Balancing with Polyatomic Ions:1
Make a table of elements
Hint: Chunk the polyatomic ions—IF they appear on both sides
OH is a polyatomic ion that is sometimes “hidden” in H2O.
Re-write H2O as HOH to “see” the OH polyatomic ion. H OH
_____ H3
PO4 +
_____ Ca
(OH)2 
_____
Ca3
(PO4)2 +
_____ H2 O
Reactants
Products H
PO4 Ca OH

40. Balancing with Polyatomic Ions:
Count the number of each element or ion on the reactants and products side. 2 H OH
_____ H3
PO4 +
_____ Ca
(OH)2 
_____
Ca3
(PO4)2 +
_____ H2 O
Reactants
Products H 3 1
PO4 1 2 Ca 1 3 OH 2 1

41. Balancing with Polyatomic Ions:3
Add coefficients to balance the numbers H OH
_____ 2 H3
PO4 +
_____ 3 Ca
(OH)2 
_____
Ca3
(PO4)2 +
_____ 6 H2 O
Reactants
Products H 3 6 1 6
PO4 1 2 2 Ca 1 3 3 OH 2 6 1 6

42. Balancing with Polyatomic Ions:5
Place a “1” in any empty coefficient location H OH
_____ 2 H3
PO4 +
_____ 3 Ca
(OH)2 
_____ 1
Ca3
(PO4)2 +
_____ 6 H2 O
Reactants
Products H 3 6 1 6
PO4 1 2 2 Ca 1 3 3 OH 2 6 1 6

43. Let’s Practice : without a chart
Example:
Balance the following equation
__ HCl + __ Ca(OH)2  __ CaCl2 + __ H2O

44. Balance the following equation
Example:
Balance the following equation
Did you see the “OH” polyatomic ion & change H2O to HOH? 2 1 1 2
__ HCl + __ Ca(OH)2  __ CaCl2 + __ H2O
HOH

45. Balance the following equation
SELF CHECK
Example:
Balance the following equation
_____ Li + _____ H2 O  ____ LiOH _____ H2

46. Balance the following equation
Example:
Balance the following equation 2 2
__ Li + __ H2O  __ LiOH + ___ H2 2 1

47. Balance the following equation
SELF CHECK
Example:
Balance the following equation
__ Fe + __ O2  ___ Fe2O3

48. Balance the following equation
Example:
Balance the following equation 4 3 2
__ Fe + __ O2  ___ Fe2O3

49. BALANCING HINTS
Never start balancing by placing a 1 as a coefficient. Do that last.
Leave H & O for last when balancing.
Chuck the polyatomic Ions
Change H2O to HOH if the hydroxide ion is present on both sides
If there is an odd number of 1 type of atom on one side and an even number of the same atom on the other side, make the odd number even by multiplying by 2.
Make sure the set of coefficients are in the simplest ratio
When finished balancing, always re-check that all of your atoms are equal on both sides.
In a combustion reaction, always start balancing with the C and H of the hydrocarbon

50. 1. 2 5 4 2 combustion 2. 1 1 1 1 SR 3. 1 2 2 1 SR 4. 2 2 1 decomp
Practice Problems
combustion SR SR
decomp
synthesis
synthesis DR SR

51. Continuation of AnswersDR
synthesis
synthesis
decomp SR DR DR

52. Predicting Products of a Chemical Reaction
Section 3
To be successful in this section, you will need to use the reaction models found on the Reference Sheet

53. Steps to Predicting Products of a Chemical Reaction
Identify the type of reaction.
Use the reference sheet to determine which model to use.
Create products for your reaction following the model as a guide.
As you form new ionic compounds, don’t forget to check charges.
Make sure to recognize the elements of BrINClHOF & format them correctly.
Balance the chemical equation.

55. Synthesis Reaction: Model 1A Formation of a Binary Ionic Compound
What do you do?
Bring symbols together and check charges
Example:
K Br2 

56. Synthesis Reaction: Model 1B Metal oxide & water forms a base
MOH
What do you do?
Combine metal symbol with hydroxide (OH) &
check charges
Example:
CaO H2O 

57. Synthesis Reaction: Model 1C Nonmetal oxide & water forms an acid
What do you do? Add up all atoms starting with H, then the nonmetal, ending with O. Example: CO2 + H2O 

58. Decomposition Reaction: Model 2A Break down of a Binary Compound
What do you do?
Break binary compound apart into the 2 elements that make it up Look for diatomic
elements!
Example: CaCl2 

59. Decomposition Reaction: Model 2B Metal Carbonates
What do you do?
Place metal next to oxygen & check charges
Then add carbon dioxide, CO2
Example:
K2CO3 

60. Decomposition Reaction: Model 2C Metal Hydrogen Carbonates (bicarbonates)
What do you do?
Place metal next to oxygen & check charges
Then add water, H2O & then carbon dioxide, CO2
Example:
NaHCO3 

61. Decomposition Reaction: Model 2D Metal Hydroxides
What do you do?
Place metal next to oxygen & check charges
Then add water, H2O
Example:
LiOH 

62. Decomposition Reaction: Model 2E Metal Chlorates
What do you do?
Place metal next to chlorine & check charges
Then add oxygen gas, O2
Example:
AgClO3 

63. Decomposition Reaction: Model 2F OxyAcids
What do you do?
Remove water (H2O) from the oxyacid to form the nonmetal oxide
Example:
H2SO3 

64. Using the Activity Series: The Activity Series is a chart that shows the reactivity of metals
found on your reference sheet
Elements on top of the chart are more reactive than elements on the bottom
More reactive elements can only replace less reactive elements within a compound

65. Self Check
Using the activity series, determine if the following replacements can happen:
K + CaCl2 
Cu + Zn(OH)2 
Na + H2O 
Ni + H2O 
Mg + HCl 

66. Single Replacement Reaction: Model 3A More Reactive Metal replacing Less Reactive Metal
What do you do?
Check Activity Series: Is the single metal
above the metal in the compound?
If YES, replace it & check charges of new
compound. Be on alert for diatomic elements!
If NO, write NR for “no reaction”.

67. A + BC  AC + B Example: Zn + Cu(NO3)2 
metal ionic comp  ionic comp metal
Example: Zn + Cu(NO3)2 

68. Single Replacement Reaction: Model 3B More Reactive Metal replacing Hydrogen from Water
What do you do?
Check Activity Series: Is the single metal above the Hydrogen? Look at the specific details about temperature of water.
If YES, replace it & check charges of new compound. Then add hydrogen gas, H2.
If NO, write NR for “no reaction”.

69. Examples: Na + H2O  Al + H2O  Al + H2O (g) M + H2O  MOH + H2
metal water  metal hydroxide hydrogen gas
Examples:
Na H2O 
Al H2O 
Al H2O (g)

70. Single Replacement Reaction: Model 3C More Reactive Metal replacing Hydrogen from an Acid
What do you do?
Check Activity Series: Is the single metal above the
Hydrogen?
If YES, replace it & check charges of new compound. Then
add hydrogen gas, H2.
If NO, write NR for “no reaction”.

71. M + HX  MX + H2 Example: Mg + HCl 
metal acid  ionic comp hydrogen
Example: Mg HCl 

72. Single Replacement Reaction: Model 3D More Reactive Halogen replacing Less Reactive Halogen
What do you do?
Check Activity Series: Is the single halogen above
the halogen in the chart?
If YES, replace it & check charges of new compound.
Be on alert for diatomic elements!
If NO, write NR for “no reaction”.

73. D + BC  BD + C Examples: Cl2 + KBr  I2 + NaCl 
Element ionic compound  ionic compound element
(halogen) (halogen)
Examples:
Cl2 + KBr 
I2 + NaCl 

74. Solubility Rules Chart

75. Solubility Rules Chart: The Solubility Rules Chart classifies substances as soluble or insoluble
found on your reference sheet
Categorized by Anion
SOLUBLE- means that the substance will dissolve in water
INSOLUBLE- means that the substance will not dissolve in water and so it is considered a PRECIPITATE when
on the product side (PP)
Insoluble = Precipitate

76. Let’s Practice #1: decide whether each is soluble or insoluble
NaNO3 Fe(C2H3O2)2
CaBr2 Ba(OH)2
Cu(OH)2

77. NaNO3
Fe(C2H3O2)2
CaBr2
Ba(OH)2
Cu(OH)2
Soluble
Not Soluble

78. Self Check
Using the Solubility Rules, determine if the following substances are soluble or insoluble(precipitate):
1. CaCl Zn(OH)2
2. K3PO NiSO4
3. MgCO AgNO3

79. Double Replacement Reaction: Model 4A Formation of a Precipitate
What do you do?
Swap the positive ions with each other & check the charges on the new compounds
Using the Solubility Rules chart, identify the product that is a precipitate (INSOLUBLE)

80. Precipitation Reactions
A precipitation reaction is when 2 soluble substances are mixed together and they form an insoluble substance. This is called a precipitate.
Reactants
2 soluble
chemicals:
NaOH and Cu(NO3)2
NaOH
Cu(NO3)2

81. Precipitation Reactions(DR Rxns)
NaOH + CuNO3  NaNO3 + Cu(OH)2
Precipitation Reactions(DR Rxns)
Na+1
OH-1
Cu+2
NO3 -1
Cu(OH)2(S)
Products: 1 soluble chemical: NaNO3
1 insoluble chemical (the precipitate): Cu(OH)2
Na+1
NO3 -1

82. AB + CD  AC + BD Example: NaCl + AgNO3
metal ionic comp  ionic comp metal
Example: NaCl + AgNO3

83. Double Replacement Reaction: Model 4B Acid Base Reaction
What do you do?
Identify the reactants as an acid & base. No need to use solubility rules chart!
Swap the positive ions with each other & check the charges on the new compounds which will always be a salt & water.

84. AB + CD  AC + BD Example: KOH + H2SO4 
Acid Base  ionic comp water
Example:
KOH + H2SO4 

85. CombustionReaction: Model 5
CxHy O  CO H2O
Hydrocarbon + oxygen  carbon dioxide water
What do you do?
Identify the hydrocarbon and O2 as reactants
Products will always be CO2 and H2O no matter what the formula of the hydrocarbon
Examples:
C3H O2 
CH3OH + O2 

86. Section 3: Part C Net Ionic Reactions – Precipitation
Shows the details of aqueous reactions that involve ions
in aqueous solution
Molecular Equation: the typical equation you are use to writing keeping all molecules together
Complete Ionic Equation: shows all the particles in a solution as they really exist, as IONS or MOLECULES.
Anything aqueous needs to be split apart into the cation and anion
Anything solid stays intact
Coefficients need to be multiplied by subscripts to determine the exact amount of each cation and anion.

87. Spectator ions: ions that do not participate in a reaction; they are identical on both sides of the equation & are crossed out!
Net Ionic Equation: the final equation showing the major players. All spectator ions have been removed.

88. NET IONIC REACTIONS for Precipitation Reactions
Molecular equation:
KI(aq) + AgNO3(aq)  AgI(s) + KNO3(aq)
Complete Ionic equation:
K+1 + I-1 + Ag+1+ NO3-1  AgI K NO3-1
Spectator ions: ions that do not participate in a reaction; they are identical on both sides of the equation & are crossed out!
Net Ionic equation:
I-1 + Ag+1  AgI

89. NET IONIC REACTIONS for Precipitation Reactions
Molecular equation:
2 NaOH(aq) + CuCl2(aq)  2 NaCl(aq) + Cu(OH)2(s)
Complete Ionic equation:
2 Na OH-1 + Cu Cl-1  2 Na Cl-1 + Cu(OH)2
Net Ionic equation:
2 OH-1 + Cu+2  Cu(OH)2

90. Take Home Practice: Predict products and balance Iron (III) chloride reacts with sodium hydroxide
Molecular equation:
1 FeCl3(aq) + 3 NaOH(aq)  1 Fe(OH)3(s) +3 NaCl(aq)
Complete Ionic equation:
3 Na OH-1 + Fe Cl-1  3 Na Cl-1 + Fe(OH)3
Net Ionic equation:
3 OH-1 + Fe+3  Fe(OH)3