Chapter 3 Stoichiometry: Calculations with Chemical Formulas and Equations

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CHEMISTRY The Central Science 10th Edition
Chapter 3Stoichiometry Calculations with
Chemical Formulas and Equations
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Chemical reactions involve largely the outer
(valence) electrons of an atom
Nuclear reactions involve nuclear processes
(Chapter 21)
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Radioactive decay
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In nuclear reactions, the number of nucleons is
conserved
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Variety of mechanisms possible for radioactive
decay.
The three most common are called a, ß and ? decay
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a particle emission
An a particle is the nucleus of a 4He atom
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ß particle emission
A ß particle is an electron that comes from the
conversion of a neutron into a proton and an
electron
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Gamma ray emission
?-Radiation is the loss of a high-energy photon
from the nucleus. This has no effect on the
identity of the nucleus.
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Other radioactive decay mechanisms are now known
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Positron emission
A positron is a particle with the same mass as an
electron but a positive charge. It comes from
the decay of a proton into a neutron and a
positron
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Electron capture
An electron from one of the inner electron shells
of the atom can be captured by a proton in the
nucleus
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Initiated nuclear reactions
Nuclear reactions may be initiated by a nucleus
capturing a particle.
Examples
Complete
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This can lead to a chain reaction.
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Chemical Equations
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Chemical Equations
Chemical equations are descriptions of chemical
reactions.
There are two parts to an equation reactants and
products H2 O2 ? H2O

• To balance
• change only the stoichiometric coefficients
• cannot change the chemical species

Stoichiometric coefficients numbers in front of
the chemical formulas give ratio of reactants
and products
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Chemical Equations
H2 O2 ? H2O
Thus the equation H2 O2 ? H2O2 is a balanced
equation but for a different reaction.
We could have H2 1/2 O2 ?
H2O, or 2H2 O2 ? 2H2O, or 4H2 2O2 ?
4H2O, etc.
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These are all OK, but

• preferably, no fractions are used
• smallest possible coefficients used

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Example
Combustion (burning, oxidation) of hydrocarbons
produces CO2 and H2O.
Write a balanced equation for the combustion of
hexane, C6H14
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In combustion reactions
C, H, O ? CO2, H2O N ? N2 S ? SO2
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Example
Write a balanced equation for the combustion of
purine, C5H4N4
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Patterns of chemical reactivity
Elements in the same group of the periodic table
react in a similar manner
Alkali metal water ? metal hydroxide hydrogen
2Na 2H2O ? 2NaOH H2 2K 2H2O ? 2KOH H2
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Alkali earth metal water ? metal hydroxide
hydrogen
Mg 2H2O ? Mg(OH)2 H2 Ca 2H2O ? Ca(OH)2 H2
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Combination reactions have fewer products than
reactants
2Mg(s) O2(g) ? 2MgO(s)
Mg has combined with O2 to form MgO.
metal oxygen ? metal oxide
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non metal hydrogen ? hydride
C(s) 2H2(g) ? CH4(g)
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metal halogen ? metal halide
Ca(s) Cl2(g) ? CaCl2(s)
2Li(s) F2(g) ? 2LiF(s)
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Decomposition reactions have fewer reactants than
products
2NaN3(s) ? 2Na(s) 3N2(g)
(the reaction that occurs in an air bag in a
motor car)
The NaN3 has decomposed into Na and N2 gas
metal azide ? metal nitrogen
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metal carbonate ? metal oxide carbon dioxide
CaCO3(s) ? CaO(s) CO2(g)
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p. 85
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Redox reactions involve the transfer of
electrons. An example is the production of a
metal from its oxide using C as reducing agent.
2CuO C ? 2Cu CO2
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Molecular and Formula Weights

• Formula weights (Fr) sum of Ar for atoms in
  formula unit

Fr (KMnO4) 1Ar(K) 1Ar(Mn) 4Ar(O)
1(39.10) 1(54.94) 4(16.00) 158.04
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• Molecular weight (Mr) is the weight of the
  molecular formula.

MW(C6H12O6) 6(12.01) 12(1.008)
6(16.00) 180.16
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By definition
Atomic mass of 12C 12.0000 amu
Now express this in grams
12.0000 g of 12C contains
6.022 x 1023 atoms
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6.022 x 1023
Avogadros number
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DEFINE
Mole (convenient measure chemical quantities)
1 mole of something 6.022 ? 1023 of that thing
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Molar Mass
Molar mass mass in grams of 1 mole of substance
(units g.mol-1).
How?
Express the formula wt or molecular wt in grams
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1 mol of N2
1 mol of H2O
1 mol of NaCl
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Study this carefully!
p. 93
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Example
Calculate the number of H atoms in 20.00 g of
C6H12O6.
Strategy
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DEFINITIONS
AMOUNT number of mols
(QUANTITY applies to mass, volume, etc.)

• Percentage Composition
• Percent composition is the atomic weight for each
  element divided by the formula weight of the
  compound expressed as a percentage

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Example
What is the by mass ( composition) of O in
KMnO4?
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Empirical Formulas from Analyses

• Start with mass of elements (i.e.
  experimental data) and calculate a formula.

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Strategy
Mass of elements
Empirical formula
Assume have 100 g sample
Mols of each element
Grams of each element in sample
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Example (p. 98)
Ethylene glycol 38.7 C 9.7 H 51.6
O by mass Molar mass (from mass
spectrometry) 62.1 g mol-1
Calculate (a) empirical formula (b) molecular
formula
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Combustion Analysis

• The experimental data are often obtained from
  combustion
• analysis

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O in CxHyOz is determined by mass difference
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Example
Combustion of 2.78 mg of ethyl butyrate produced
6.32 mg CO2 and 2.58 mg of H2O. What is the
empirical formula of this compound?
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Quantitative Information from Balanced Equations
Example
What mass of CO2 is produced when 1.00 g of
propane, C3H8, is burned?
Balanced equation
C3H8 O2 ? CO2 H2O
? C3H8 5O2 ? 3CO2 4H2O
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STRATEGY
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Limiting Reactants

• If the reactants are not present in
  stoichiometric amounts, at end of reaction some
  reactants are still present.
• These reactants are said to be in excess.
• Limiting Reactant one reactant that is consumed
  completely in the reaction.

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2 H2(g) 1 O2(g) ? 2 H2O(l)
2 mol H2 reacts with 1 mol O2 4 mol H2 reacts
with 2 mol O2 13 mol H2 reacts with 6.5 mol
O2 etc
But suppose we have 10 mol H2 and 7 mol O2
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Unreacted O2 left over H2 was the limiting
reagent
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Strategy
Mols of product based on limiting reagent
mols reactants
Mass reactants
Identify limiting reagent
Mass of product
Example Al reduces Fe2O3 to Fe. How much Fe is
produced from the reaction of 30.0 g of Al and
100 g Fe2O3?
2Al Fe2O3 ? 2Fe Al2O3
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A footnote

• Some compounds have water of crystallisation.
  These are water molecules associated with the
  solid as it crystallises from solution

CuSO45H2O

• Can often be driven off

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End of Chapter 3Stoichiometry Calculations
with Chemical Formulas and Equations