Chapter 8: Chemical Equations and Reactions

1
Chapter 8Chemical Equations and Reactions

• pp. 260-297

The evolution of energy as light and heat is an
indication that a chemical reaction is taking
place.
2
(No Transcript)
3
Describing Chemical Equations

• chemical equations vs. chemical formulas ?
• H2 O2 H2O
• 2H2 O2 2H2O

4
Heat Concepts

• DE (DH)
• exothermic
• endothermic
• Example of a thermochemical equation
• If DH is (-) or written into the equation on the
  right side, the reaction is exothermic
• If DH is () or written into the equation on the
  left side, the reaction is endothermic.

change in energy (DE) or change in heat (DH)
heat is released (given off) during a reaction
(HOT)
heat is absorbed (taken in) during a reaction
(COLD)
2H2(g) O2(g) 2H2O(g) 483.6 kJ 2H2(g)
O2(g) 2H2O(g) DH -483.6 kJ
Exothermic
5
Exothermic or Endothermic?
2H2O(g) 483.6 kJ 2H2(g) O2(g)
6
Good Evidence of a Chemical Reaction
(These sometimes can be just a physical change)

• Color change Example

rusting

• Energy change Example

glow stick

• Gas released Example

vinegar and baking soda

• Precipitate formed Example

solid formed from two liquids
Which piece(s) of evidence is represented in each
picture?
7
Equation Terminology

• reactants on left
• products on right
• yield
• reversible reaction
• (s), (l), (g), (aq)
• gas - , precipitate - ONLY FOR PRODUCTS
• D heat written above the arrow
• elect. electricity written above the arrow
• catalyst written above the arrow (Ni, Pt, MnO2)

8
Phase Information

• Hg and Br2 are liquids at room temperature (l)
• metals are usually solid (s)
• nonmetals, use your prior knowledge
• ionic compounds are usually solid (s)
• covalent compounds, use your prior knowledge

9
Five Types of Chemical Reactions

• Synthesis (Direct Combination, Composition)
• Decomposition
• Combustion
• Single Replacement
• (Single Displacement)
• Double Replacement
• (Double Displacement)

10
Synthesis
One product formed during the reaction.
also known as combination

• element or element or compound
• compound compound
• A B AB

Example 2C O2 2CO


11
Decomposition
One reactant used during the reaction.
also known as analysis

• compound two or more elements
• or compounds
• AB A B

Example 2H2O 2H2 O2


12
Combustion
EXOTHERMIC
anything oxygen heat
hydrocarbon oxygen carbon dioxide water
CxHy O2 CO2 H2O
Example CH4 2O2 CO2 2H2O



13
Single Replacement
1 element and 1 compound in the reactants with 1
element and 1 compound in the products.

• A BC AC B
• Example Zn H3PO4 Zn3(PO4)2 H2
• 3Zn 2H3PO4 Zn3(PO4)2 3H2

14
Double Replacement
1 compound and 1 compound in the reactants with 1
compound and 1 compound in the products.

• AB CD AD CB
• Example Pb(NO3)2(aq) 2NaI(aq) PbI2(s)
  2NaNO3(aq)

precipitate
15
Checking for Understanding
Classify the following reactions as one of the 5
types

• 3Pb(NO3)2(aq) 2K3PO4(aq) ? Pb3(PO4)2(s)
  6KNO3(aq)

double replacement
2C8H18(l) 25O2(g) ? 16CO2(g) 18H2O(g)
combustion
Cl2(g) 2NaBr(aq) ? 2NaCl(aq) Br2(l)
single replacement
C12H22O11(s) ? 12C(s) 11H2O(l)
decomposition
2H2O(l) O2(g) ? 2H2O2(aq)
synthesis and combustion
16
You Try It.

• Do the Classification of Chemical Reactions
  Worksheet.

17
Characteristics of Chemical Equations

• The equation must represent known facts.
• The equation must contain the correct formulas
  for the reactants and products.
• The law of conservation of mass must be satisfied
  (balanced).

18
Balancing Equations
1770s

• Law of Conservation of Matter
• matter cannot be created nor destroyed

Antoine Lavoisier
19
Balancing Equations

• CH4 O2 ? H2O CO2
• 16g 32g ? 18g 44g
• 48g ¹ 62g

CH4 2O2 ? 2H2O CO2 16g 64g ? 36g
44g 80g 80g
coefficient a small, whole number that appears
in front of a formula in a chemical equation
20
Coefficients
coefficient a small, whole number that appears
in front of a formula in a chemical equation

• Coefficients express relative, not absolute
  amounts of reactants and products.
• Coefficients represent individual particles of
  matter.
• Because particles and moles are directly
  proportional, the coefficients also represent the
  number of moles of each substance.
• Because mass and moles are NOT directly
  proportional, the coefficients DO NOT represent
  the mass of each substance.

21
Balancing Chemical Equations

• When balancing, you do not rewrite the formulas.
  You only place coefficients in the equation.
• NEVER change the subscripts after the equation is
  written!

22
Balancing Equations

• Example
• C6H12O6 O2 ? H2O CO2

6
6
6
Balance the oxygen LAST. 6 C on the left1 C on
the right. So, place a 6 in front of CO2. 12 H
on the left2 H on the right. So, place a 6 in
front of H2O. 6 O 2 O 8 O on left6 O 12 O
18 O on right. So, place a 6 in front
of O2.
Double check to be sure you are done. Are the
coefficients in the lowest ratio? 1666 - YES
23
Balancing Equations

• Example
• CaCl2 NH4OH Ca(OH)2 NH4Cl

2
2
Since these are ionic compounds, balance the ions
as whole items. 1 Ca2 on each side. So, do
nothing. 2 Cl- on the left1 Cl- on the right.
So, place a 2 in front of the NH4Cl. 1 NH4 on
the left2 NH4 on the right. So, place a 2 in
front of NH4OH. 2 OH- on the left2 OH- on the
right. So, do nothing.
Double check to be sure you are done. Are the
coefficients in the lowest ratio? 1212 - YES
24
You Try It.

• Do the Balancing Chemical Equations Worksheet.

25
Word Equations

• word equation an equation in which the
  reactants and products in a chemical reaction are
  represented by words.
• Chemical equations are expressed in formulas.
• Chemists use words to communicate verbally.
• It is necessary to be able to take the words and
  express them as equations and vice versa.

26
Diatomic Molecules
H2, O2, N2, F2, Cl2, Br2, I2
27
The Hungry Dragon

• Potassium chlorate decomposes with the use of
  heat to form potassium chloride and oxygen.
• Upon placing glucose in the reaction mixture, the
  glucose reacts with the oxygen exothermically.

28
You Try It.

• Do the Word Equations Worksheet.

29
Synthesis
One product formed during the reaction.
also known as combination

• element or element or compound
• compound compound
• A B AB

Example 2C O2 2CO


30
Synthesis (Binary)
also known as combination

• If a metal and a nonmetal come together, form the
  ionic compound according to the charges of the
  ions formed.
• Example Na(s) Cl2(g)

NaCl(s) NOT NaCl2

• Balance 2Na(s) Cl2(g) 2NaCl(s)

31
Synthesis (Binary)
also known as combination

• If two nonmetals come together, form the simplest
  compound from the two.
• Example C(s) O2(g)

• CO(g)
• NOT CO2(g)

Balance 2C(s) O2(g) 2CO(g)
32
Synthesis Reactions with Oxides

• Oxides of active metals react with water to
  produce metal hydroxides.
• CaO(s) H2O(l) Ca(OH)2(s)
• Oxides of nonmetals react with water to form
  oxyacids.
• SO2(g) H2O(l) H2SO3(aq)
• 2H2SO3(aq) O2 2H2SO4(aq)

33
Synthesis Reactions with Oxides

• Certain metal oxides and nonmetal oxides react
  with each other to form salts (ionic compounds).
• CaO(s) SO2(g) CaSO3(s)

34
You Try It.

• Predict the products of the following reactions
• Na2O(s) H2O(l)
• N2O5(g) H2O(l)
• MgO(s) CO2(g)

2NaOH(aq) 2HNO3(aq) MgCO3(s)
35
Decomposition
One reactant used during the reaction.
also known as analysis

• compound two or more elements
• or compounds
• AB A B

Example 2H2O 2H2 O2


36
Decomposition
also known as analysis

• Decomposition often needs a push. It is not
  uncommon to see symbols over the yield sign,
  indicating that a catalyst, heat, or electricity
  was used.
• catalyst (write the specific catalyst)
• heat
• electricity

Pt
D
elect.
37
Decomposition
also known as analysis

• Table sugar will decompose to form carbon and
  water when sulfuric acid is used as a catalyst.

38
Decomposition
also known as analysis

• When a compound decomposes or breaks down, it
  does not immediately go to pure elements. First,
  it breaks down into smaller compounds if
  possible.
• Example H2O2(aq) H2O(l) O2(g)
• Balance 2H2O2(aq) 2H2O(l) O2(g)
• Example H2O(l) H2(g) O2(g)
• Balance 2H2O(l) 2H2(g) O2(g)

MnO2
MnO2
elect.
elect.
39
You Try It.

• Do the Synthesis and Decomposition Worksheets.

40
Combustion
EXOTHERMIC
anything oxygen heat
hydrocarbon oxygen carbon dioxide water
CxHy O2 CO2 H2O
Example CH4 2O2 CO2 2H2O



41
You Try It.

• Do the Combustion Worksheet.

42
Single Replacement
1 element and 1 compound in the reactants with 1
element and 1 compound in the products.

• A BC AC B
• Example Zn H3PO4 Zn3(PO4)2 H2
• 3Zn 2H3PO4 Zn3(PO4)2 3H2

43
Single Replacement
I want to cut in, but am I more reactive than she
is?
cation/anion element
44
Single Replacement
I want to cut in, but am I more reactive than he
is?
cation/anion element
45
Single Replacement

• The metal element will only react with the
  compound if it is more reactive on the activity
  series.
• The nonmetal element will only react with the
  compound if it is more reactive.
• F2 gt Cl2 gt Br2 gt I2

46
SingleReplacement
lithium potassium calcium sodium magnesium aluminu
m zinc chromium iron nickel tin lead HYDROGEN copp
er mercury silver platinum gold
decreasing activity
47
Single Replacement Video
http//cwx.prenhall.com/petrucci/medialib/media_po
rtfolio/text_images/024_SILVERCRYSTA.MOV
48
You Try It.

• Do the Single Replacement Worksheet.

49
Double Replacement
1 compound and 1 compound in the reactants with 1
compound and 1 compound in the products.

• AB CD AD CB
• Example Pb(NO3)2(aq) 2NaI(aq) PbI2
  2NaNO3(aq)
• cation/anion cation/anion cation/anion
  cation/anion

precipitate
Most double replacement reactions are run in
water.
50
Double Replacement
cation/anion cation/anion
51
Double Replacement
precipitate a solid that is produced as a
result of a chemical reaction in solution and
that separates from the solution therefore it is
NOT soluble in water. (insoluble)
52
Double Replacement

• precipitate a solid that is produced as a
  result of a chemical reaction in solution and
  that separates from the solutiontherefore it is
  NOT soluble in water. (insoluble)
• How do you know which product is a precipitate,
  if any?
• Check if the products are soluble.

53
Double Replacement

• Memorize for dissolving in water
• Alkali metal compounds are soluble
• NH4 compounds are soluble
• NO3- compounds are soluble
• C2H3O2- compounds are soluble
• ClO3- compounds are soluble
• CO3-2, PO4-3, O-2, SiO3-2, S-2, SO4-2 are
  generally not soluble (insoluble)

54
You Try It.

• Do the Double Replacement Worksheet.

55
You Try It.

• Do the Predicting Products of Chemical Reactions
  Worksheet.

56
Oxidation-Reduction Reactions
Page 631
57
Oxidation-Reduction Reactions

• Redox Reactions
• oxidation processes reactions in which the
  atoms or ions of an element experience an
  increase in oxidation state (losing electrons)
• reduction processes reactions in which the
  oxidation state of an element decreases (gaining
  electrons)

58
L osing E lectrons O xidation
G aining E lectrons R eduction
the lion goes
59
Oxidation-Reduction Reactions

• 2Cu(s) O2(g) 2CuO(s)
• What is the oxidation number of Cu(s)?
• 0
• What is the oxidation number of Cu in the CuO(s)?
• 2
• If it goes from 0 to 2, it is losing electrons
• LEOtherefore it is being oxidized.

60
Oxidation-Reduction Reactions

• 2Cu(s) O2(g) 2CuO(s)
• What is the oxidation number of O2(g)?
• 0
• What is the oxidation number of O in the CuO(s)?
• -2
• If it goes from 0 to -2, it is gaining electrons
• GERtherefore it is being reduced.

61
Oxidation-Reduction Reactions

• 2Na(s) Cl2(g) 2NaCl(s)
• Oxidation numbers
• Na(s) 0
• Cl2(g) 0
• Na in NaCl(s) 1
• Cl in NaCl(s) -1

62
Oxidation-Reduction Reactions

• Acetylene gas is burned.
• 2C2H2(g) 5O2(g) 4CO2(g) 2H2O(g)
• What is happening to the O2?
• 0 -2 reduction
• What is happening to the C?
• -1 4 oxidation
• What is happening to the H?
• 1 1 nothing

63
You Try It.

• Do the Notes Oxidation-Reduction Reactions

64
Chemical Reactions