CHEMICAL REACTIONS: MAKING CHEMICALS SAFELY WITHOUT DAMAGING THE ENVIRONMENT

1
CHAPTER 4 CHEMICAL REACTIONS MAKING CHEMICALS
SAFELY WITHOUT DAMAGING THE ENVIRONMENT
From Green Chemistry and the Ten Commandments of
Sustainability, Stanley E. Manahan, ChemChar
Research, Inc., 2006 manahans_at_missouri.edu
2
4.1. DESCRIBING WHAT HAPPENS WITH CHEMICAL
EQUATIONS
In our own bodies A chemical reaction occurs
Glucose sugar reacts with oxygen to give carbon
dioxide, water, and energy Represented by a
chemical equation C6H12O6 6O2 ? 6CO2
6H2O ( energy) (4.1.1) Reactants
Products Balanced Reactants 6 C,
12 H, Products 6 C, 12 H,
6 12 18 O 12
6 18 O
3
Chemical Reactions and Equations
States of matter of reaction participants CaCO3(s
) 2HCl(aq) ?
CO2(g) CaCl2(aq) H2O(l) (4.1.2) (s) for
solid, (aq) for a substance in solution, (g) for
gas, and (l) for liquid ??? denotes a reversible
reaction Example NH3(aq) H2O(l) ??
NH4(aq) OH-(aq) (4.1.3) ? for heat added

4
4.2. BALANCING CHEMICAL EQUATIONS
A balanced chemical equation shows the same
number of each kind of atom on both sides of the
equation. Balancing a chemical equation
(different example from one shown in book) MnO2
C2H6 ? Mn CO H2O There will have
to be at least 2 C atoms and 6 H atoms on the
right MnO2 C2H6 ? Mn 2CO
3H2O This gives 5 O atoms on the right, so the
number of O atoms on the left must be a multiple
of 5 5MnO2 C2H6 ? Mn 2CO
3H2O This means that there must be 10 O atoms on
the right, so multiply CO and H2O on the right by
2 5MnO2 C2H6 ? Mn 4CO 6H2O
5
Balancing Chemical Equations (Continued)
There must be 4 C atoms and 12 H atoms on the
left 5MnO2 2C2H6 ? Mn 4CO 6H2O
The products must have 5 Mn atoms 5MnO2
2C2H6 ? 5Mn 4CO 6H2O The equation
should be balanced, check the results
Reactants 5 Mn, 10 O, 4 C, 12 H Products 5
Mn, 4 6 10 O, 4 C, 12 H
6
4.3. JUST BECAUSE YOU CAN WRITE IT DOES NOT MEAN
THAT IT WILL HAPPEN
The following reaction occurs Fe(s)
H2SO4(aq) ? H2(g) FeSO4(aq) (4.3.1) The
following reaction does not occur Cu(s)
H2SO4(aq) ? H2(g) CuSO4(aq) (4.3.2) Alterna
tive way to make copper sulfate,CuSO4 2Cu(s)
O2(g) ? 2CuO(s)
(4.3.4) CuO(s) H2SO4(aq) ? CuSO4(aq)
H2O(aq) (4.3.5)
7
Alternate Reaction Pathways
Alternative reactions pathways for maximum
safety, minimum byproduct, and utilization of
readily available materials. Two ways to prepare
iron (II) sulfate, FeSO4. First method Fe(s)
H2SO4(aq) ? H2(g) FeSO4(aq) (4.3.1) Could
use scrap iron and waste sulfuric
acid Generates elemental H2, which is
explosive But H2 could be used in a fuel
cell Second pathway FeO(s) H2SO4(aq) ?
FeSO4(aq) H2O(aq) (4.3.6) No dangerous
H2 Could also use scrap iron and waste sulfuric
acid
8
4.4. YIELD AND ATOM ECONOMY IN CHEMICAL REACTIONS
Yield is the percentage of the degree to which a
chemical reaction or synthesis goes to
completion. Atom economy is defined as the
fraction of reactants that go into final
products. Consider yield and atom economy for the
preparation of HCl gas By reaction of sodium
chloride with sulfuric acid accompanied by
heating to drive off HCl gas 2NaCl(s)
H2SO4(l) ? 2HCl(g) Na2SO4(s) (4.4.1) When
all of the NaCl and H2SO4 react, there is 100
yield. Byproduct Na2SO4 gives less than 100 atom
economy.
9
Atom Economy (Continued)
Given the atomic masses H 1.0, Cl 35.5, Na 23.0,
and O 16.0 gives the following Mass of desired
product 2 ??(1.0 35.5) 73.0
(4.4.3) Total mass product 2 ??(1.0 35.5)
(2 ? 23.0 32.0 4 ??16.0) 215

(4.4.4)
Alternatively, the following occurs with 100
atom economy H2(g) Cl2(g) ?2HCl(g)
(4.4.2)



10
4.5. CATALYSTS THAT MAKE REACTIONS GO
Carbon monoxide burns in air 2CO O2 ? 2CO2
(4.5.1) CO is generated
by automobile engines and is an undesirable air
pollutant. CO is eliminated by reaction with
oxygen over an automotive exhaust catalytic
converter. The metals on the surface of the
catalytic converter act as a catalyst to enable
the above reaction to occur efficiently. A
catalyst speeds up a chemical reaction without
itself being consumed.
11
Enzymes as Catalysts
Enzymes are specialized proteins that act as
biological catalysts. For example, aerobic
respiration, in which glucose reacts with oxygen
in living organisms using enzyme catalysts and
producing energy. C6H12O6 O2 ? 6CO2
6H2O energy (4.5.2) Enzymes
participate in many life processes
including Protein synthesis Repair damaged
DNA Detoxification Chemical kinetics deals
with rates of chemical reactions.
12
4.6. KINDS OF CHEMICAL REACTIONS
Combination reaction or addition reaction in
which two substances come together to form a new
substance C O2 ? CO2 CaO SO2 ?
CaSO3 Addition reactions are 100 atom
economical. Decomposition reaction in which a
compound decomposes to two or more
products. Example is electrolysis of water to
produce elemental hydrogen and oxygen by passing
an electrical current through water made
electrically conducting with a dissolved salt,
such as Na2SO4 2H2O(l) 2H2(g)
O2(g) (4.6.3) Can be 100 atom economical, but
may be less than 100 because of side reactions.
13
Decomposition Reactions (Continued)
Decomposition reaction to make sodium carbonate,
Na2CO3, from sodium bicarbonate,
NaHCO3 2NaHCO3(s) ? Na2CO3(s) CO2(g)
H2O(g) (4.6.4) to produce sodium carbonate,
Na2CO3, commonly used as an industrial chemical
to treat water, in cleaning solutions, and as an
ingredient of glass.

14
Kinds of Chemical Reactions (Continued)
Substitution or replacement reaction is one such
as the reaction of iron and sulfuric acid, Fe(s)
H2SO4(aq) ? H2(g) FeSO4(aq)
(4.3.1) This reaction is also evolution of a
gas. A double replacement or metathesis reaction,
in which two compounds trade ions or other
groups. H2SO4(aq) Ca(OH)2(aq) ? CaSO4(s)
2H2O(l) (4.6.5) This is also a
neutralization reaction in which an acid and a
base react to produce water and a
salt. Precipitation reactions produce
precipitates of insoluble substance that come out
of water solution CaCl2(aq) Na2CO3(aq) ?
CaCO3(s) 2NaCl(aq) (4.6.5) Calcium
removal from water is water softening. Calcium
can cause scale in pipes Calcium precipitates
soap in a useless solid form
15
4.7. OXIDATION-REDUCTION REACTIONS AND GREEN
CHEMISTRY
Oxidation-reduction reactions, frequently called
redox reactions Use of oxidation to describe the
reaction of a substance with oxygen
(4.7.1)
2Ca O2 ??2CaO Calcium metal is
oxidized. Elemental oxygen is reduced to produce
the oxide ion, O2- in CaO. When a chemical
species loses electrons in a chemical reaction it
is oxidized and when a species gains electrons it
is reduced. Whenever a chemical species combines
with elemental hydrogen, it is reduced.

16
Oxidation-Reduction Reactions (Continued)
FeO H2 ???Fe H2O
(4.7.2) In this case the Fe in FeO is reduced to
iron metal and the hydrogen in elemental H2 is
oxidized to H2O. When elemental oxygen reacts to
produce chemically combined oxygen, it is acting
as an oxidizing agent and is reduced. Oxidation-re
duction in photosynthesis, 6CO2 6H2O
h? ? C6H12O6 6O2 (4.7.3)
h? represents light energy Oxidation-reduction in
respiration C6H12O6 6O2 ? 6CO2 6H2O
energy (4.1.1) Oxidation of
fossil fuel CH4 2O2 ? CO2 2H2O
energy (4.7.4)
17
Electrolysis of Water in Which H2O is Oxidized at
One Electrode and Reduced at the Other
18
Oxidation-Reduction Reactions in Green Chemistry
Oxidation of fossil fuels and other materials in
producing energy Hydrogen and carbon in
hydrocarbons are in reduced form, such as in
ethane, C2H6. Many raw materials are partially
oxidized hydrocarbons, such as ethanol, C2H6O,
which can be made by 2C2H6 O2 ? 2C2H6O

(4.7.5) Alternate biosynthesis of ethanol by
fermentation of carbohydrates C6H12O6 ?
2C2H6O 2CO2
(4.7.6)

19
4.8. QUANTITATIVE INFORMATION FROM CHEMICAL
REACTIONS
Formula mass The sum of the atomic masses of
all the atoms in a formula unit of a
compound. Molar mass Where X is the formula
mass, the molar mass is X grams of an element or
compound, that is, the mass in grams of 1 mole of
the element or compound. Consider 2C2H6
7O2 ? 4CO2 6H2O (4.8.1) In terms of
moles, 2 moles of C2H6 react with 7 moles of O2
to yield 4 moles of CO2 and 6 moles of H2O. Given
the atomic masses H 1.0, C 12.0, and O 16.0 the
molar mass of C2H6 is 30.0 g/mol, that of O2 32.0
g/mol, that of CO2 is 44.0 g/mol, and that of H2O
18.0 g/mol.
20
Quantitative Information from Chemical Reactions
(Cont.)
For the reaction 2C2H6 7O2 ? 4CO2
6H2O In terms of the minimum whole number of
moles reacting and produced 2 moles of C2H6
with a mass of 2 ? 30.0 g 60.0 g of C2H6 7
moles of O2 with a mass of 7 ? 32.0 g 224 g of
O2 4 moles of CO2 with a mass of 4 ? 44.0 g
176 g of CO2 6 moles of H2O with a mass of 6 ?
18.0 g 108 g of H2O The total mass of reactants
is 60.0 g of C2H6 224 g of O2 284.0 g of
reactants and the total mass of products is
176 g of CO2 108 g of H2O 284 g of products
21
4.9. Stoichiometry by the Mole Ratio Method
The calculation of quantities of materials
involved in chemical reactions is addressed by
stoichiometry. Based upon the law of conservation
of mass which states that the total mass of
reactants in a chemical reaction equals the total
mass of products. Holds true because matter is
neither created nor destroyed in chemical
reactions. The mole ratio method of
stoichiometric calculations is based upon the
fact that the relative numbers of moles of
reactants and products remain the same regardless
of the total quantity of reaction.
22
Example of the Mole Ratio Method
2C2H6 7O2 ? 4CO2 6H2O
(4.9.1) At the mole level, this chemical equation
states that 2 moles C2H6 react with 7 moles of O2
to produce 4 moles of CO2 and 6 moles of H2O. For
10 times as much material, 20 moles C2H6 react
with 70 moles of O2 to produce 40 moles of CO2
and 60 moles of H2O. Suppose that it is given
that 18.0 g of C2H6 react. What is the mass of O2
that will react with this amount of C2H6? What
mass of CO2 is produced? What mass of H2O is
produced? To solve this problem, the following
mole ratios are used
To solve for the mass of O2 reacting use the
following steps
A. Mass of C2H2 reacting
B. Convert to moles of C2H2
C. Convert to moles of O2
D. Convert to mass of O2
23
Mole Ratio Calculation (Continued)
Given the molar mass of C2H6 as 30.0 g/mol, the
molar mass of O2 (18.0 g/mol), and the mole ratio
relating moles of O2 to moles of C2H6,
The masses of CO2 and H2O produced are calculated
as follows
Total mass of reactants, 18.0 g C2H6 67.2 g O2
85.2 g Total mass of products, 52.8 g CO2
32.4 g H2O 85.2 g
24
4.10. LIMITING REACTANT AND PERCENT YIELD
One of the reactants is almost always a limiting
reactant. Example Reaction of 100 g of
elemental zinc (atomic mass 65.4) and 100 g of
elemental sulfur (atomic mass 32.0) are mixed and
heated undergoing the following reaction Zn
S ? ZnS (4.9.1) What mass of
ZnS, formula mass 97.4 g/mol, is produced? If 100
g of zinc react completely, the mass of S
reacting and the mass of ZnS produced would be
given by the following calculations
Only 48.9 g of the 100 g of S react, so zinc is
the limiting reactant. The mass of Zn required to
react with 100 g of sulfur would be 204 g of Zn,
but only 100 g of Zn is available.

25
Percent Yield
The mass of product calculated from the mass of
limiting reactant in a chemical reaction is
called the stoichiometric yield of a chemical
reaction. By measuring the actual mass of a
product produced in a chemical reaction and
comparing it to the mass predicted from the
stoichiometric yield it is possible to calculate
the percent yield. Suppose that a water solution
containing 25.0 g of CaCl2 was mixed with a
solution of sodium sulfate, CaCl2(aq)
Na2SO4(aq) ? CaSO4(s) 2NaCl(aq)
(4.10.2) Removed by filtration and dried, the
precipitate was found to have a mass of 28.3 g,
the measured yield. What was the percent yield?
26
Percent Yield (Continued)
The stoichiometric yield of CaSO4 calculated by
the mole ratio method is 30.6 g CaSO4
The percent yield is calculated by the following
(4.10.4)
27
4.11. Titrations Measuring Moles by Volume of
Solution
If the molar concentration of a solution is
known, the number of moles may be measured by the
volume of the solution (see measuring glassware
below
28
Titration
Titration uses a buret to measure the volume of a
solution with a known concentration of a reagent
required to react exactly with another substance
in solution Reagent added from the buret until a
measured end point is reached indicating that the
reaction is complete Volume used with
stoichiometry to measure amount of substance The
pertinent equations relating to solution
concentration and stoichiometry are the following
(4.11.1)
(4.11.2)
(4.11.3)
29
Example of Analysis by Titration (Titrimetric
Analysis)
Consider a sample consisting of basic lime,
Ca(OH)2, molar mass 74.1 g/mol, and dirt with a
total sample mass of 1.26 g. Using titration with
a standard acid solution it is possible to
determine the mass of basic Ca(OH)2 in the
solution and from that calculate the percentage
of Ca(OH)2 in the sample. Assume that the solid
sample is placed in water and titrated with 0.112
mol/L standard HCl (concentration designated
MHCl), a volume of 42.2 mL (0.0422 L) of the acid
being required to reach the end point. The dirt
does not react with HCl, but the Ca(OH)2 reacts
as follows with the mole ratio given
below Ca(OH)2 2HCl ? CaCl2 2H2O
1 mol Ca(OH)2
2 mol HCl
At the end point the number of moles of HCl can
be calculated by MolHCl LitersHCl x MHCl
30
Titrimetric Analysis of Ca(OH)2 (Cont.)
The calculation of the percentage of Ca(OH)2 in
the sample is given by the following
31
4.12. INDUSTRIAL CHEMICAL REACTIONS THE SOLVAY
PROCESS
The Solvay process consists of saturating a
sodium chloride solution (brine) with ammonia gas
(NH3), then with carbon dioxide, then cooling it
to precipitate solid NaHCO3 NaCl NH3 CO2
H2O ? NaHCO3(s) NH4Cl The sodium
bicarbonate product is heated to produce sodium
carbonate, Na2CO3, a chemical with many
industrial uses 2NaHCO3 heat ? Na2CO3
H2O(g) CO2(g) The CO2 from this reaction is
recirculated back to the first reaction
above. Ammonia is made by the following reaction,
which requires heat, high pressures and a
catalyst 3H2 N2 ? 2NH3 Ammonia is
reclaimed from the reaction solution by adding
lime, Ca(OH)2, made from heating limestone and
reacting the CaO product with water
32
The Solvay Process (Continued)
CaCO3 heat ? CaO CO2 (calcination of
limestone) CaO H2O ? Ca(OH)2 When Ca(OH)2
is added to the spent solution from which NaHCO3
has precipitated, the ammonia is evolved and
reclaimed Ca(OH)2 (s) 2NH4Cl(aq) ? 2NH3(g)
CaCl2(aq) 2H2O(l) Although this reaction
reclaims ammonia, it generates large quantities
of calcium, chloride, CaCl2, which has few
commercial uses and tends to accumulate as
waste The overall reaction for the Solvay process
is CaCO3 2NaCl ? Na2CO3 CaCl2 from which
the stoichiometric atom economy 48.8 (mass of
Na2CO3 product divided by total mass of
reactants) In practice, the yield is less due to
incomplete precipitation of NaHCO3 and other
factors
33
Degree to Which the Solvay Process is Green It
is green in that

1. It uses inexpensive, abundantly available raw
   materials in the form of NaCl brine and limestone
   (CaCO3). A significant amount of NH3 is required
   to initiate the process with relatively small
   quantities to keep it going.
2. It maximizes recycle of two major reactants,
   ammonia and carbon dioxide. The calcination of
   limestone provides ample carbon dioxide to make
   up for inevitable losses from the process, but
   some additional ammonia has to be added to
   compensate for any leakage.

The Solvay process is not green because it
requires extraction of non-renewable NaCl and
CaCO3 (although they are abundant), generates
excess, potentially waste CaCl2, uses relatively
large amounts of energy, has a relatively low
atom economy In the U.S. and some other countries
Na2CO3.NaHCO3.2H2O (trona) is mined and the
Solvay process is not used.