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Chapter 2 Part 1

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Title: Chapter 2 Part 1


1
Chapter 2 Part 1
  • Stoichiometry
  • Calculations with Chemical Formulae and Equations

2
Law of Conservation of Mass
Figure 2.1
We may lay it down as an incontestable axiom
that, in all the operations of art and nature,
nothing is created an equal amount of matter
exists both before and after the experiment.
Upon this principle, the whole art of performing
chemical experiments depends. --Antoine
Lavoisier, 1789
3
Chemical Equations
  • Concise representations of chemical
  • reactions.
  • They mainly give proportions phases
  • They do not imply mechanism

4
Anatomy of a Chemical Equation
Figure 2.4
CH4 (g) 2 O2 (g) CO2 (g) 2
H2O (g)
5
Anatomy of a Chemical Equation
CH4 (g) 2 O2 (g) CO2 (g) 2
H2O (g)
  • Reactants appear on the
  • left side of the equation.

Note that the equation contains molecules or
elements and the phases they exist in during the
reaction. e.g. Methane (gas)
6
Anatomy of a Chemical Equation
CH4 (g) 2 O2 (g) CO2 (g) 2
H2O (g)
Products appear on the right side of the
equation.
The Arrow indicates which direction the reaction
tends to go. If it points both ways it signifiies
equilibrium.
7
Anatomy of a Chemical Equation
CH4 (g) 2 O2 (g) CO2 (g) 2
H2O (g)
  • The states of the reactants and products
  • are written in parentheses to the right of
  • each compound. In this example all are gases.

8
Anatomy of a Chemical Equation
CH4 (g) 2 O2 (g) CO2 (g) 2
H2O (g)
Coefficients are inserted to balance the equation.
Balanced Equations have the same elements and
numbers of atoms of each element on both sides of
the equation.
Numbers of Molecules are proportionate to a
complete reaction, not absolute.
9
Subscripts and Coefficients Give Different
Information
Figure 2.2
  • Subscripts indicate the number of atoms of each
    element in a molecule
  • Coefficients indicate the number of molecules

10
  • Reaction Types
  • Combination
  • A B ? C
  • Decomposition
  • A ? B C , often with a gas
  • Combustion
  • Fuel Oxidant ? Ash, Gases

11
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12
Combination Reactions
Figure 2.5
  • Two or more substances react to form one product
  • AB?C
  • Examples
  • N2 (g) 3 H2 (g) 2 NH3 (g)
  • C3H6 (g) Br2 (l) C3H6Br2 (l)
  • 2 Mg (s) O2 (g) 2 MgO (s)

13
Combination Reactions
Figure 2.5
  • Nitrogen gas reacts with hydrogen gas to form
    ammonia gas (Nitrogen fixation).
  • Cyclopropane gas reacts with liquid bromine to
    form 1,5 dibromopropane
  • Magnesium metal reacts with oxygen gas to form
    solid magnesium oxide.
  • Examples
  • N2 (g) 3 H2 (g) 2 NH3 (g)
  • C3H6 (g) Br2 (l) C3H6Br2 (l)
  • 2 Mg (s) O2 (g) 2 MgO (s)

14
Combination Reactions
Figure 2.5
  • 2 Mg (s) O2 (g) 2 MgO (s)

48.61 g Mg 32.00 g O2 80.61 g MgO What
happens if you only have 16 g of O2 ?
15
Decomposition Reactions
One substance breaks down into two or more
substances
2 NaN3 (s) 2 Na (s) 3 N2 (g)
Figure 2.6
  • Examples
  • CaCO3 (s) CaO (s) CO2 (g)
  • 2 KClO3 (s) 2 KCl (s) 3 O2 (g)
  • 2 NaN3 (s) 2 Na (s) 3 N2 (g)

16
Decomposition Reactions
Figure 2.6
One substance breaks down into two or more
substances
2 NaN3 (s) 2 Na (s) 3 N2 (g)
  • Sodium azide solid reacts to form sodium metal
    and nitrogen gas.
  • (So what does it mean if somebody texts you
    U NaN3 ?)

17
You Airbag!
18
Decomposition Reactions
One substance breaks down into two or more
substances
CaCO3 (s) CaO (s) CO2 (g)
  • Calcium carbonate (limestone) decomposes to form
    solid unslaked lime and carbon dioxide gas.
  • This is the 1st step to make portland cement.

19
Decomposition Reactions
One substance breaks down into two or more
substances
POP QUIZ! 2 moles of solid _______decomposes to
form 2 moles of solid _ _ and 3
moles of gas.
  • 2 KClO3 (s) 2 KCl (s) 3 O2 (g)

20
Decomposition Reactions
One substance breaks down into two or more
substances
POP QUIZ! 2 moles of solid Potassium Chlorate
decomposes to form 2 moles of solid
Potassium Chloride and 3 moles of Oxygen gas.
  • 2 KClO3 (s) 2 KCl (s) 3 O2 (g)

21
Combustion Reactions
Figure 2.7
  • Rapid reactions that produce a flame
  • Most often involve hydrocarbons reacting with
    oxygen in the air
  • Examples
  • CH4 (g) 2 O2 (g) CO2 (g) 2
    H2O (g)
  • C3H8 (g) 5 O2 (g) 3 CO2 (g) 4
    H2O (g)

22
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23
Combustion Reactions
Figure 2.7
  • Most often involve hydrocarbons reacting with
    oxygen in the air
  • Which hydrocarbon fuel generates the least
    amount of greenhouse gas?
  • Examples
  • CH4 (g) 2 O2 (g) CO2 (g) 2
    H2O (g)
  • C3H8 (g) 5 O2 (g) 3 CO2 (g) 4
    H2O (g)

24
  • Formula Mass

25
Formula Mass ( or Weight)
  • Sum of the atomic weights of the atoms in a
    chemical formula
  • The formula mass of calcium chloride, CaCl2,
    would be
  • Ca 1 x 40.1 g 40.1 g
  • Cl 2 x 35.5 g 71.0 g
  • 111.1 g
  • According to Daltons Law what would the
    decomposition of 1 mole of Calcium Chloride
    generate?

26
Molecular Mass
  • Sum of the atomic weights of the atoms in a
    molecule
  • For the molecule ethane, C2H6, the molecular
    weight would be
  • C 2 x 12.0 g 24.1 g
  • H 6 x 1.0 g 6.0 g
  • 30.1 g
  • For the complete combustion of 1 mole of ethane,
    what weight of water would be produced?

27
Mass of Water from Ethane
  • For the complete combustion of 1 mole of ethane,
    what weight of water would be produced?
  • C2 H 6 (g) ? 2 CO2 (g) 3 H2O (g) .
  • 6 g H makes 3 moles of H2O,
  • each mole of water is 1 mole of Oxygen for 16 g
    and 2 moles of Hydrogen for 2 g, so 18 g.
  • But there are 3 moles of water so
  • 3 moles water X 18 g/mole 54g water

28
Percent Composition
  • One can find the percentage of the mass of a
    compound that comes from each of the elements in
    the compound by using this equation

29
Percent Composition
  • For example, the percentage of carbon in ethane,
    C2H6, is

If there are 3X as many Hydrogens, why is the
percentage of carbon so high?
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