1
Unit 2 Quantities in Chemical Reactions
Day 1
-
atomic mass definition( pg 162) ; average atomic mass calculations based on isotopic
abundance. Do Practice problems page 167 ( 1, 3,4 ). Assign page 170 (2)
Mole definition, avogadro constant, distinguish between atoms, molecules and formula
units, Define molar mass. Calculate molar mass of compounds. Do practice problems
(24-26) on pg 184 .
Day 2
-
Relate molar mass to the mole to avogadros #. Convert from moles to mass.. Do page
186 (27b,d ,28,30). Convert from mass to moles. Do page 187(31a,32a,d,34
Day 3
-
Convert moles to # of particles . Page 177 (13,15,17,18).
Convert # of particles to moles . Do page 178 (19-22). Assign page 179
(4,5a,6,7,8,10)
Day 4/5-
Convert between moles, mass and # of particles.for atoms : mass ----a----> moles ----b-----> # of atoms
a=divide by molar mass b= divide by avogadros #
read chemistry bulletin page 188 . Try question # 2.
page 190 # 35 b,c. Page 191 #39 ( a.d ) 41,42. Page 192(2-5). Assign page
193 ( 3,9,11,13 a,c,e, 14(a,c,e),15,18.
Day 6
-
Quiz on chapter 5; Define law of definite proportions , percentage composition. Page
201 (2,4)page 204 (6-8). Assign page 205(3,4,6)
Day 7
-
Define empirical and molecular formula. determine the empirical formula given %
composition.
Do pg 211(13-16) pg 214(2-4,6-8)
Day 8
-
do lab empirical formula of MgO.Assign unit project Read page 216: Careers in chemistry
Day 9
-
Determining molecular formula given empirical formula. Do practice problems page 218
(17,19,20). Page 218 (2,4,5). Finding empirical and molecular formulas experimentally.
Page 221(21,22)
Day 10-
Determining formulas of hydrated compounds. Pg. 225 (23-25). Lab : Determining the
formula of a hydrate page 226. Assign page 228 (6,7). Page
229(5,7,8,10,12,14,17,19,21)
Day 11 -
Chapter 6 quiz. Define stoichiometry. Do page 237 (1-3). Mole-mole problems.. Do page
238(4,5,7). Page 240 (8,10)
Day 12 -
Mass - mass problems. page 244 (11,13,14). Page 246 (15,17,18). Mass - particles
problems. Page 248(19,20,21,22). Assign page 250 (4,5,6,7)
Day 13-
Limiting reagent problems. Page 254(23,24,26). Page 257(27,28,30). Assign page 259 (36)
Day 14 -
Percentage yield problems.. Page 262 (32,33) page 264 (35,36). Read page 265 Careers
in Chemistry.
Day 15 -
Lab percent yield page 266
Day 16-
Work period page 271 ( 5,6-15,17 ).Homework review for test page 276 (1-29),33-35,38.
Day 17 -
Test review
Day 18 -
Test
2
Day 1 -
atomic mass definition( pg 162) ; average atomic mass calculations based on isotopic
abundance. Do Practice problems page 167 ( 1, 3,4 ). Assign page 170 (2)
Mole definition, avogadro constant, distinguish between atoms, molecules and formula
units, Define molar mass. Calculate molar mass of compounds. Do practice problems
(24-26) on pg 184 .
ATOMIC MASS
 the mass of an atom is expressed in _______ ______________ __________ (amu or u)
 an atomic mass unit is a relative measure, defined by the mass of one atom of
__________one atom of _______________is assigned the mass of ______ u
 thus 1 u = 1/12 the mass of a C-12 atom
 all other atoms have a mass which is relative to this
ISOTOPIC ABUNDANCE
 the relative amount in which each isotope is present in an element
 expressed as a percent or as a decimal fraction
 e.g magnesium is 79% Mg-24, 10% Mg-25 and 11% Mg-26
AVERAGE ATOMIC MASS
 this is the mass shown for an element on the periodic table
 it is a weighted average of the atomic masses of all of the naturally occurring isotopes of
an element
 a mass spectrometer is the instrument used to determine mass and abundance of the
isotopes (see page 166)
A.A.M. = [ (abundance of isotope 1 X mass of isotope 1) + (abundance of isotope 2 X
mass of
isotope 2) + etc ]
For Lithium: 92.58% Li-7 with a mass of 7.015 u
7.42% Li-6 with a mass of 6.015 u
 average atomic mass of lithium= [(.9258 x 7.015 u) + (.0742 x 6.015 u)]= 6.94 u
Do Practice problems page 167 ( 1,3,4 ). Assign page 170 (2)
3
Mole(n)
-
represents amount of substance; 1 mole of anything consists of Avogadro’s #
(Na) of particles
-
1 mol = 6.02 x 1023 particles ( atoms, ion, molecules, formula units, people,
money )
-
For example, if we stack 1 mole of pennies on top of one another, how tall would the stack be?
The answer is it would be so tall that the stack of pennies could reach the sun and back almost
500 million times!

ions, atoms and molecules are too small to see so we would need very large quantities to “see”
anything--- instead of using large numbers we will work with atomic masses and the mole ( a unit
of measurement)
Number
Quantity
Example
2
Pair
12
Shoes
Dozen
Eggs
144
Gross
Eggs
500
Ream
Paper
6.02 x 1023
MOLE
Atomic mass
-
Molecules
mass of 1 atom ( amu ) e.g. 1 atom of C = 12µ ( amu )
Molecular mass -
Mass of 1 molecule ( amu ) e.g. 1 molecule of H2O = 18 µ
Molar mass
(M)
mass of 1 mole of particles ( g/mol) e.g. 1 mole of C atoms = 12 g
M = 12.0 g/mol
-
Gram formula mass
-
mass of 1 mole of formula units (g/mol)
e.g. M for NaCl = 23.0 + 35.5 = 58.5 g/mol
To Calculate Molar Mass (M)
1.
2.
3.
4.
Write the chemical formula.
Determine the number of atoms (or ions, etc) of each element.
Use the atomic mass from the periodic table and the amounts in moles to determine the molar mass.
Communicate the molar mass in the correct units.
Sample Problem:
Q: What is the molar mass of Ammonium Phosphate, (NH4)3PO4?
A: M(NH4)3 PO4 = 3(N +4H) +P+ 4O
= 3N + 12H P +4O
= (3 x 14.01) + (12 x 1.01) + (1 x 30.97) + (4 x 16.00)
= 149.12g/mol
6.02 X 1023 molecules of CO2 = _________ mol of CO2 = ____________ g of CO2
4
Pg 184 ( 24-26 )
Day 2 -
Relate molar mass to the mole to mass. Convert from moles to mass.. Do page
186 (27b,d ,28,30). Convert from mass to moles. Do page 187(31a,32a,d,34
Fomulas :
1.
# of moles =
Mass given
------------------Molar mass
ii.
Mass = # of moles x molar mass
Ex. 1
Given 36.0 g of H2O : Calculate # of moles
n = m/M
m=nxM
n=
Ex. 2
Given 2.5 moles of Ca3(PO4)2 . Calculate mass
m=
Do page 186 (27b,d ,28,30). Convert from mass to moles. Do page 187(31a,32a,d,34)
5
Problem set :
1. Determine the # of moles present in each of the following.
a.
17.4 g Na ( ans. 0.756 moles )
b.
60.0 g Na2SO4 ( 0.420 moles )
c.
93.5 g CO2 ( 2.12 moles )
d.
25.6 g NaNO3 ( 0.301 moles )
e.
33.2 g HCl ( 0.909 moles )
2. Calculate the mass of the following.
a.
0.750 mol Ca(OH)2 ( 55.5 g )
b.
0.200 mol KCl ( 14.9 g )
c
2.45 mol Cu(NO3)2 ( 459 g )
d.
0.500 mol H2O2 ( 17.0 g )
e.
1.00 mol H2O ( 18.0 g )
6
S IGNIFICANT FIGURES & ROU NDING
A. Indicate the number of significant figures then round each to the number of significant figures
indicated.
For example:
1.234
has ______4___ significant figures and, rounded to 2 significant figures, is ___1.2____
1.
0.6034
has ________ significant figures and, rounded to 2 significant figures, is _____
2.
12,700
has ________ significant figures and, rounded to 2 significant figures, is _____
3.
12,700.00
has __________ significant figures and, rounded to 1 significant figures, is ________
4.
0.000983
has __________ significant figures and, rounded to 2 significant figures, is ________
5.
123342.9
has __________ significant figures and, rounded to 5 significant figures, is ________
6.
6.023 x 1023 has __________ significant figures and, rounded to 2 significant figures, is ________
7.
.005600
8.
10000.5006 has __________ significant figures and, rounded to 5 significant figures, is ________
9.
2.0 x 10-3
has __________ significant figures and, rounded to 1 significant figures, is ________
10.
3.456110
has __________ significant figures and, rounded to 3 significant figures, is _______
has __________ significant figures and, rounded to 1 significant figures, is ________
B. Given calculations with the calculator answer, write the answers with the appropriate number of
significant figures.
Example:
6.00 x 3.00
= 18
The answer should be ______18.0_____
1.
23 + 46
= 69
The answer should be
_______________
2.
23.0 + 46.0
= 69
The answer should be
_______________
3.
253 + 345.8
= 598.8
The answer should be
_______________
4.
56 – 35
= 21
The answer should be
_______________
5.
56.00 – 35.0
= 21
The answer should be
_______________
6.
46 x 12
= 552
The answer should be
_______________
7.
3.24 x 5.63
= 18.2412
The answer should be
_______________
8
(2.355 + 2.645) x 10.00 = 50
The answer should be
_______________
9
654  32
= 20.4375
The answer should be
_______________
10.
.024 x .063
= 1.512 x 10-03
The answer should be
_______________
7
Day 3 -
Convert moles to # of particles . Page 177 (13,15,17,18).
Convert # of particles to moles . Convert molecules to atoms or formula units
to ions.Do page 178 (19-22). Assign page 179 (4,5a,6,7,8,10).
# of moles = # of molecules
----------------Avogadro’s #
n = N/Na
# of molecules or formula units = # of moles x Avogadro’s #
# of atoms = # of atoms
------------ x
molecule
Ex. 1
N=
# of molecules
Given 2.00 x 1024 molecules of H2O : Calculate # of moles. Calculate # of H atoms. Calculate
total # of atoms
n=
# of H atoms =
# of atoms =
Ex. 2
Given 2.5 moles of Ca3(PO4)2 . Calculate the # of formula units (N). Calculate total # of ions
N=
# of ions =
Page 177 (13,15,17,18) page 178 (19-22). Assign page 179 (4,5a,6,7,8,10).
8
Day 4/5-
Convert between moles, mass and # of particles. for atoms : mass ----a---->
moles -----b-----> # of atoms
a=divide by molar mass b= divide by
avogadros #
read chemistry bulletin page 188 . Try question # 2.
page 190 # 35 b,c. Page 191 #39 ( a.d ) 41,42. Page 192(2-5). Assign page
193 ( 3,9,11,13 a,c,e, 14(a,c,e),15,18.
Conversions :
Mass to moles ----->
Moles to mass ------>
Moles to molecules ---->
Molecules to moles ----->
Mass to molecules ------>
Molecules to atoms ------>
divide mass by molar mass
multiply moles x molar mass
multiply # of moles by avogadro’s #
divide # of molecules by avogadro’s #
Divide mass by molar mass then multiply by Avogadro’s #
multiply # of molecules by the # atoms per molecule
Complete table below :
Substance
# of moles
# of
molecules
C
O2
# of atoms
1.0 moles
1.2 x 1024
0.50 moles
Ca3(PO4)2
NO
# of formula units
24.0 g
H2O
CH3OH
mass
1
1.2 x 1023
# of ions
9
Mole Worksheet
1.
Aspartame is an artificial sweetener that is 160 times sweeter than sucrose (table sugar) when
dissolved in water. It is marketed as Nutra-Sweet. The molecular formula for aspartame is
C14H18N2O5.
a) Calculate the molar mass of aspartame.
b) How many moles of molecules are present in 10.0 g of aspartame?
c) Calculate the mass in grams of 1.56 mol aspartame.
d) How many molecules are in 5.0 mg aspartame?
e) How many atoms of nitrogen are in 1.2 g aspartame?
f)
What is the mass in grams of 1.0 x 109 molecules of aspartame?
g) What is the mass in grams of one molecule of aspartame?
2.
Dimethylnitrosamine, (CH3)2N2O is a carcinogenic (cancer causing) substance that may be formed
in foods, beverages or gastric juices from the reaction of nitrite ion (used as a food preservative)
with other substances.
a) What is the molar mass of dimethylnitrosamine?
b) How many moles of (CH3)2N2O molecules are present in 250 mg dimethylnitrosamine?
c) What is the mass of 0.050 mole dimethylnitrosamine?
d) How many atoms of hydrogen are in 1.0 mole dimethylnitrosamine?
e) What is the mass of 1.0 x106 molecules of dimethylnitrosamine?
f)
What is the mass in grams of one molecule of dimethylnitrosamine?
Answers: 1a) 294.34 g/mol
b) 3.40 E-2 mol C14H18N2O5 c) 4.59 E2 g C14H18N2O5
d) 1.0 E19 molecules C14H18N2O5 e) 4.9 E21 atoms N f) 4.9 E-13 g C14H18N2O5
g) 4.9 E-22 g
C14H18N2O5
#2 a) 74.10 g/mol
b) 3.4 E-3 mole C2H6N2O c) 3.7 g C2H6N2O d) 3.6 E24 atoms H
e) 1.2 E-16 g C2H6N2O
f) 1.23 E-22 g C2H6N2O /1 molecule
10
Calculate the # of molecules in each of the following.
1.
15.0 g SO2 ( 1.41 x 1023 )
2.
2.5 mol CO ( 1.5 x 1024 )
3.
0.400 mol CH3COOH ( 2.40 x 1023 )
4.
0.700 g C6H12O6 (2.00x1021)
How many atoms in the following ?
1.
22.0 g NH3 ( 3.1 x 1024 )
2.
45.5 g C3H8 ( 6.85 x 1024 )
How many ions in :
3.
2.28 mol Ca3(PO4 )2 ( 1.78 x 1025 )
4.
0.200 mol Na2S2O3 ( 8.40 x 1023 )
11
Assignment: Moles
Complete the table
Substance
Molar
mass
(g/mole)
Ne
Name _______________________
Mass
(g)
# of
moles
# of
molecules
3.01 x 1023
Al2(SO4)3
0.240
CO2
1.25 x 1022
Zn
2.25
120.0
NH3
1.00
6.25
H2O2
Day 6 -
# of ions
2.20
H2O
X2Y3Z
# of atoms
28.60
Na
Ca3(PO4)2
# of
formula
units
100.0
1.20 x 1024
Quiz on chapter 5; Define law of definite proportions , percentage composition.
Page 201 (2,4)page 204 (6-8). Assign page 205(3,4,6)
12
Law of Definite Proportions (pg 198)
(Law of Constant Composition)
 The elements in a chemical compound are always present in the same proportions by _________

proposed by ____________in late 18th century after testing many samples of CuCO3 and finding
that all samples contained the same proportions of Cu, C and O by mass
Percentage Composition
 mass percent of an element is the mass of an element in a compound, expressed as a percent of
the total mass of the compound
 percentage composition is a statement of the mass percent of every element in a compound
e.g. % composition of water
 mass percent of oxygen is always 88.8%  mass percent of hydrogen is always 11.2%
O = 16/18 x 100% = _________
H = 2/18 x 100 % = _____________
Calculations can be done using different sources of data:
A
Using Mass Data
A 15.0 g sample of a compound is analyzed and found to contain 9.8 g of element X, with the rest being
element Y. What is the percent composition of this compound?
Mass percent of element X = mass of element_X x 100%
mass of the compound
=
Mass percent of element Y =
B
Using the Chemical Formula
 the formula of a compound states the number of atoms of each element
 assuming that there is one mole of a compound, its composition by mass can be determined by
using the molar mass of each element times the number of atoms present
Ex. 1
Calculate the % composition of Al2S3
i.
ii.
2 Al = 2 x 27g/mol =
3 S = 3 x 32 =
54 g/mol
96 g/mol
% Al = 54/150 x 100 % = _________
% S = 96/150 x 100% = __________
-----------------
150 g/mol
Ex. 2
Calculate % composition of Cu(ClO4)2
____ Cu = _______________=
____ Cl = ________________=
____ O = ________________=
% Cu =
% Cl =
%O=
Page 201 (2,4)page 204 (6-8). Assign page 205(3,4,6)
13
Day 7 -
Define empirical and molecular formula. determine the empirical formula given
% composition. Do pg 211(13-16) pg 214(2-4,6-8)
Empirical Formula ( Simplest formula ) (pg 207)
represents the smallest whole # ratio of atoms of an element in a compound
compound ethyne has a molecular formula of C2H2 ; the empirical formula is ____________
Ex. 1
Step i.
Step ii.
Step iii.
A compound consists of 36.5 % Na, 25.4 % S, and 38.1 % O. calculate it’s empirical formula
Assume 100g of compound, then calculate the # of moles of each element
Divide each by the smallest # of moles
If fractions exist multiply each by a whole # so that ratios are whole #’s (pg 210)
# mole Na = ___________=
ii.
Na = ---------- =
# mole S = ____________=
S = ----------- =
# mole O =____________=
O = ---------- =
iii.
Na : S : O = _____ : _____ : ______ -----> empirical formula is ___________________
Ex. 2
A compound consists of 53.73 % Fe and 46.27 % S . Calculate the empirical formula.
Find the chemical formulas based on the given percent compositions.
1.
75% carbon
25 % Hydrogen
1. ___________
2.
52.7% Potassium
47.3% Chlorine
2. _________
3.
22.1% Aluminum
25.4% Phosphorous
52.5% Oxygen
3. ___________
4.
13% Magnesium
87% Bromine
4. ___________
14
Find the chemical formulas based on the given percent compositions.
1. 69.6% Barium, 6.1% Carbon, 24.3% Oxygen
1. __________
2. 40.5% Zinc, 19.9% Sulfur, 39.6% Oxygen
2. __________
3. 88.8% Copper, 11.2% Oxygen
3. __________
4. 79.9% Copper, 20.1% Oxygen
4. __________
5. 36.7% Potassium, 33.3% Chlorine, 30.0% Oxygen
5. __________
6. 28.2% Potassium, 25.6% Chlorine. 46.2% Oxygen
6. __________
7. 40.2% Potassium, 26.8% Chromium, 33.0% Oxygen
7. __________
8. 26.6% Potassium, 35.3% Chromium, 38.1% Oxygen
8. __________
Do pg 211(13-16) pg 214(2-4,6-8)
15
Day 8 Empirical Formula Lab
Names ______________________________
Purpose: To obtain the empirical formula of a Magnesium, Oxygen compound
Material: Bunsen burner, Magnesium ribbon, crucible and lid, ring stand, triangle, scale, ruler
Procedure:
1.
clean and dry completely a crucible and lid
2.
measure five centimeters of magnesium ribbon and cut
3.
with sand paper, clean the magnesium ribbon to a shine
4.
mass the crucible and lid
5.
roll ribbon into a loose coil and place inside the crucible
6.
mass the crucible, lid and ribbon
7.
place the crucible on the ring stand inside the clay triangle
8.
tilt the lid so that air can get into the crucible
9.
heat the crucible with the bunsen burner with a strong blue flame
10.
the magnesium ribbon should turn to a gray color upon oxidation
11.
when all of the magnesium is gray, take away the heat and allow to cool
12.
remass the crucible, lid and contents
13.
clean all equipment and lab station
DATA
1. Mass of crucible and lid
2. Mass of crucible, lid and Mg ribbon
3. Mass of Mg ribbon
4. Moles of Mg used
5. Mass of crucible, lid and compound
6. Mass of compound
7. Mass of oxygen used
8. Moles of oxygen used (diatomic!!!)
9. Which is smaller: Moles of Magnesium or Oxygen?
10. Divide number of moles by the smaller to get ratio
11. Ratio of Magnesium to Oxygen
12. Empirical formula of compound
_____________
_____________
_____________
_____________
_____________
_____________
_____________
_____________
_____________
_______ : ___
_____________
16
Day 9 -
Determining molecular formula given empirical formula. Do practice problems
(17,19,20). Page 218 (2,4,5).
Molecular Formulas
-
represents the actual ratio of atoms in a molecule of a compound
whole # multiple of the empirical formula
molecular formula of glucose is C6H12O6 ----> empirical formula is (CH2O)x ---> simple multiplyer
is 6
Ex. 1
A compound consists of 92.25 % C and 7.75 5 H . If the molar mass is 78 g/mol; What is the
molecular formula ?

Determine empirical formula :
v.
# of moles C =
b.
Divide by the smallest :C =
# of moles H =
H=
w.
C : H = _________ -----> empirical formula is _____________________
x.
( CH )x = 78
( 12 + 1 )x = 78 x = __________(CH)x = ________ --> molecular formula is ___________
Ex. 2
A compound consists of 43.6 % P and 56.4 % O. Molar mass is 284 g/mol. Determine the
molecular formula.
Ex. 3
The molar mass of a compound is 92 g/mol. It consists of 0.608 g N and 1.388 g O. Determine
the molecular formula.
Do practice problems page 218 (17,19,20). Page 218 (2,4,5).
18
Worksheet on empirical formula calculation/determination. Making Connections
1.
l-Dopa, a drug used for the treatment of Parkinson’s disease, is 54.82% carbon, 5.62% hydrogen,
7.10% nitrogen, and 32.46% oxygen. How many moles of each element are present in a 100.0 g
sample of l-dopa?
2.
What is the empirical formula for l-dopa?
3.
Salicylic acid, used in the synthesis of aspirin, has a mass composition of 60.87% C, 4.38% H,
and 34.75% O. Determine the empirical formula for this compound.
4.
Lindane, used as an insecticide, has a mass composition of 24.78% C, 2.08% H, 73.14% Cl. What
is the empirical formula for Lindane?
5.
Lindane has a molar mass of 290.85 g mol-1. What is the molecular formula for lindane?
6.
7.
A chemical analysis of a complex carbohydrate is 40.0% C, 6.72% H, and
53.5% O. What is the empirical formula?
The molar mass for the carbohydrate in problem 6 is 870.8 g mol-1. What is the molecular
formula?
19
DETERMINATION OF SIMPLEST (EMPIRICAL) & MOLECULAR FORMULA (SET
ONE)
1. A molecule of table sugar (sucrose) contains 12 atoms of carbon, 22 atoms of hydrogen and 11 atoms
of oxygen. Write the formula of sucrose.
2. In an experiment, 109.5 g of hydrogen chloride were formed by the chemical union of 3.00 g of
hydrogen and 106.5 g of chlorine. Calculate the simplest formula of this compound.
3. An oxide of mercury with a mass of 78 g was decomposed by heating and 3.0 g of oxygen were
evolved. Find the simplest formula of this oxide.
4. When 2.61 g of aluminum were burned in oxygen, the resulting oxide had a mass of 4.921 g.
Determine the simplest formula of aluminum oxide.
5. Phosphorus forms two oxides containing 43.6% and 56.5% phosphorus respectively. Calculate the
simplest formula for each of these two oxides.
6. A gas known to be one of the five oxides of nitrogen was found to contain 36.8% nitrogen. Find the
simplest formula of this oxide. What further data would you require to determine the molecular (true)
formula of this gas?
7. Calculate the molecular formula of a compound containing 25.93% nitrogen, and 74.07% oxygen, if
the molar mass is 216 g.
8. Upon analysis a compound was found to have the following composition by mass: hydrogen 6.67%,
carbon 40.00% and oxygen 53.33% Find the molecular formula of this compound if the molar mass is
120 g/mol.
9. A compound contains 64.84% carbon, 13.51% hydrogen, and 21.62% oxygen. What is the simplest
formula of this compound? What is the molecular (true) formula of this compound given a molecular
mass of 148 u.
10. A compound contains sodium (43.4%) and carbon (11.3%), with the remainder being oxygen.
Determine the molecular formula of this compound if one mole of it has a mass of 106 g.
11. Calculate the simplest formula of each of the compounds whose composition is listed:
a) sodium 39.31% chlorine 60.69%
b) carbon 37.5% hydrogen 12.5% oxygen 50.0%
c) sodium 27.38% hydrogen 1.19% carbon 14.29% oxygen 57.14%
d) sodium 19.17% hydrogen 0.83% sulphur 26.66% (the remainder is oxygen)
1. C
H O
12 22 11
2. HCl
3. Hg O
2
4. Al O
2 3
5. P O P O
2 5 2 3
6. N O
2 3
7. N O
4 10
8. C H O
4 8 4
9. C H O
8 20 2
10. Na CO
2 3
11. NaCl
CH OH
3
NaHCO
3
NaHSO
4
20
DETERMINATION OF SIMPLEST (EMPIRICAL) & MOLECULAR
FORMULA (SET TWO)
1. A compound contains 21.6% sodium, 33.3% chlorine and 45.1% oxygen. Derive its empirical formula. Ans.
NaClO
3
2. When 1.010 g of zinc vapour is burned in air, 1.257 g of oxide is produced. What is the empirical formula of the
oxide?
Ans. ZnO
3. Determine the simplest formula of a compound that has the following composition:
Cr= 26.52% S= 24.52% and O= 48.96%.
Ans. Cr (SO ) (Cr S O )
2
43
2 3 12
4. A 3.245 g sample of titanium chloride was reduced with sodium to metallic titanium. After the resultant sodium
chloride was washed out, the residual titanium metal was dried and the mass determined to be 0.819 g. What
is the empirical formula of titanium chloride?
Ans. TiCl
4
5. A compound contains 63.1% carbon, 11.92% hydrogen, and 24.97% fluorine. Derive its molecular formula if
the molar mass is 152 g/mol.
Ans. C H F
8 18 2
6. An organic compound was found on analysis to consist of 47.37% carbon and 10.59% hydrogen. The balance
was presumed to be oxygen. What is the molecular formula of the compound if the molecular mass is 228 u.
Ans. C H O
9 24 6
7. A borane (a compound containing only boron and hydrogen) was analyzed and found to be 88.45% boron.
What is its empirical formula?
Ans. B H
5 7
8. An experimental catalyst used in the polymerization of butadiene has the following composition: 23.3% Co,
25.3% Mo, and 51.4% Cl. What is its empirical formula?
Ans. Co Mo Cl
3 2 11
21
Day 10-
Determining formulas of hydrated compounds. Pg. 225 (23-25). Lab : Determining the formula
of a hydrate page 226. Assign page 228 (6,7). Page 229(5,7,8,10,12,14,17,19,21)
Some compounds exist in a “hydrated” state; some specific # of water molecules are present for each
molecule of the compound.
Example: oxalic acid (COOH)2 can be obtained in the laboratory as (COOH)2•2H2O.
(The dot shows that the crystals of oxalic acid contain 2 water molecules per (COOH)2 molecule.)
The molar mass of (COOH)2 =
The molar mass of (COOH)2•2H2O =
% mass of anhydrous salt=
% mass of water=
Water can be driven out of a hydrated compound by heating it to leave an “anhydrous” (without water)
compound.
Example: A 7.0 g sample of calcium nitrate, Ca(NO3)2•4H2O, is heated to constant mass. How much
anhydrous salt remains?
Hydration Number:
Some molecules attach themselves to ___________ molecules. This is done in set numbers, depending on
the molecule. For example, Magnesium sulfate attaches to 7 water molecules. We say it’s hydration
number is ____.
MgSO4·7H2O name:______________________________________
Anhydrides:
A compound that is normally a hydrate and has lost its hydration water is said to be anhydrous and is
called an anhydride.
BaCl2·2H2O name:______________________________________
BaCl2name:___________________________or ___________________________________
Finding the Hydration Number:
The hydration number can be conveniently found by __________ the compound and measuring its
____________________. This mass loss is usually due to the hydration _______________ molecules
being driven off.
For example…
A 15.35 g sample of Strontium nitrate Sr(NO3)2•nH2O is heated to a _______________________ of
11.45 g. Calculate the hydration number.
Sample Data:
Mass Hydrate 15.35g
Mass Anhydride 11.45g___
Mass of Water (mass loss) ______
Calculations:
Moles Anhydride
Moles Water
Divide both by smallest…
Hydration Number is _____
Sr(NO3)2·____H2O name:_________________________________
22
Problem Set
1
Cupric chloride, CuCl2, when heated to 100C is dehydrated. If 0.235 g of CuCl2 · x H2O gives 0.185 g of
CuCl2 on heating, what is the value of x?
2.
The “alum” used in cooking is potassium aluminum sulfate hydrate, KAl(SO4)2 · x H2O . To find the value of
x, you can heat a sample of the compound to drive off all of the water and leave only KAl(SO 4)2. Assume
you heat 4.74 g of the hydrated compound and that the sample loses 2.16 g of water. What is the value of
x?
3.
If “Epsom salt,” MgSO4 · x H2O is heated to 250C, all the water of hydration is lost. On heating a 1.687-g
sample of the hydrate, 0.824 g of MgSO4 remains. What is the formula of Epsom salt?
4.
When CaSO4 · x H2O is heated, all of the water is driven off. If 34.0 g of CaSO4 (molar mass = 136) is
formed from 43.0 g of CaSO4 · x H2O, what is the value of x?
Pg. 225 (23-25). Page 229(5,7,8,10,12,14,17,19,21)
23
FORMULA FOR A HYDRATE
Names ________________________________________
Introduction:
Often times when you open up a new pair of shoes or a bottle of vitamins you’ll see a package like the one
illustrated. What is this and why is it there??
This little satchel contains a substance that attracts water from the air. By taking the water out of the surrounding
air vitamins, shoe leather or delicate electronics will not be damaged by moisture while in the package.
The silica gel in this satchel is in its anhydrous form and has the formula SiO2. When it collects the water from the
air, one water molecule attaches itself (still intact) and the SiO 2 molecule becomes hydrated. The water can be
removed again by simply heating the silica gel.
In hydrates, the water molecules are tucked in between “host” molecules of another substance. For example, the
Epsom salt that people often bathe in is a hydrate of magnesium sulfate, MgSO 4. H2O molecules can cluster around
and fit in between the MgSO4 molecules that make up a salt crystal in a ratio of seven water molecules to every
one magnesium sulfate molecule.
Hydrate formulas indicate this ratio and are written in the following way:
MgSO47H2O
In this case the molecular and empirical formulas are the same.
Purpose of this Investigation:
In this lab you are going remove the water from copper sulfate hydrate by heating so you can experimentally
determine its empirical formula:
CuSO4 X H2O
Then you will compare your results against given information to determine the molecular formula.
Equipment Required:






Glass test tube
Electronic balance
2 – 4 g of blue copper sulfate hydrate
Test tube tongs
Bunsen burner
Water
Procedure:
1.
Determine the mass of the glass test tube:
g
2.
Add 2 to 4 g of the blue hydrate to the test tube and record the new weight:
3.
Determine the mass of copper sulfate hydrate:
4.
Light the Bunsen burner and prepare a relatively cool flame.
g
g
24
5.
Using the tongs, tilt the test tube nearly horizontal so the hydrate covers about half the length
of the test tube.
6.
Move the test tube back and forth through the flame until all water disappears and there is
only a fine light powder. This is the anhydride (no water).
7.
Allow the powder to cool and then find the new weight of the test tube and contents:
8.
Reheat the anhydride for another minute, cool, weigh:
g. If this mass is more than
0.05 g different from step 6, repeat the heating, cooling and weighing process.
a.
g
Determine the mass of the anhydride in the test-tube by reweighing and subtracting the known
test-tube mass :
g
9.
Add approximately 1 mL of water to the test tube.
10.
Clean up your lab station and wash your hands.
Analysis:
1. When you heated the blue hydrate what were the observations you made? Why did this happen? What is the
final substance in the test tube?
______________________________________________________________________________________________
______________________________________________________________________________________________
_________________________________________________
2. What happened when you added water to your test tube? Why did this happen?
______________________________________________________________________________________________
______________________________________________________________________________________________
_________________________________________________
3. Explain what you think happened at the molecular level to the copper sulfate hydrate CuSO 4XH2O
a) When it was heated _______________________________________________________
b) When water was added to it ________________________________________________
4.
Calculate the value for x in the formula CuSO4 X H2O
5. Calculate % water
6. Calculate the % error considering the fact the accepted value for % water is 36%
25
DATE:
NAME:
:
RUBRIC FOR INVESTIGATION 6-B
Determining the Chemical Formula of a Hydrate
In this investigation, students carry out an experiment to determine the molecular formula of the hydrate of copper(II) sulfate. Measurements of the mass of the
hydrate before the heating and the mass of solid left after heating are used to determine the mass percent of water contained by the hydrate. Using this data,
students can determine the chemical formula for the hydrate. Students may be asked to present a written summary of their observations, conclusions, and
applications. The teacher should observe the students performing the investigation.
Specific
Achievement Chart
Expectations
Categories
Knowledge/Understanding
-distinguish between
the empirical formula
and the molecular
formula of a
compound
Specific
Expectations
Inquiry
- determine the
percentage
composition of a
compound through
experimentation, as
well as through
analysis of the
formula and a table of
relative atomic masses
(e.g., composition of a
hydrate)
- understanding of
concepts, principles,
laws, and theories
(e.g., identifying
assumptions;
eliminating
misconceptions;
providing
explanations)
Achievement Chart
Categories
Task Criteria
Level 1
Level 2
Level 3
Level 4
- demonstrates
limited
understanding of
the concepts of
chemical
formulas and
percent
composition by
mass
Level 1
- demonstrates
some
understanding of
the concepts of
chemical
formulas and
percent
composition by
mass
Level 2
- demonstrates a
considerable
understanding of
the concepts of
chemical
formulas and
percent
composition by
mass
Level 3
- demonstrates a thorough
understanding of the
concepts of chemical
formulas and percent
composition by mass
The student:
- demonstrates
an understanding
of the concepts
of chemical
formulas and
percent
composition by
mass
Task Criteria
Level 4
The student:
- application of the
skills and strategies of
scientific inquiry
(e.g., predicting,
performing and
recording, analyzing
and interpreting,
problem solving,
communicating
results )
- applies the
procedure for
determining the
percentage
composition of
the compound
accurately
- applies the
procedure for
determining the
percentage
composition of a
compound with
limited accuracy
- applies the
procedure for
determining the
percentage
composition of a
compound with
some accuracy
- applies the
procedure for
determining the
percentage
composition of a
compound with
considerable
accuracy
- applies the procedure for
determining the
percentage composition of
a compound with a high
degree of accuracy
- application of
technical skills and
procedures (e.g.,
procedures in using
electronic balance)
- performs the
experiment and
records mass
data accurately
- performs the
experiment and
records mass
data with limited
accuracy
- performs the
experiment and
records mass
data with some
accuracy
- performs the
experiment and
records mass
data with
considerable
accuracy
- performs the experiment
and records mass with a
high degree of accuracy
- analyzes and
interprets mass
data accurately
- analyzes and
interprets mass
data with limited
accuracy
- analyzes and
interprets mass
data with some
accuracy
- analyzes and
interprets mass
data with
considerable
accuracy
- analyzes and interprets
mass data with a high
degree of accuracy
Task Criteria
Level 1
Level 2
Level 3
Level 4
- communicates
observations,
and conclusions
with limited
clarity
- communicates
observations,
and conclusions
with moderate
clarity
- communicates
observations,
and conclusions
with
considerable
clarity
- communicates
observations, and
conclusions with a high
degree of clarity
Specific
Expectations
Communication
Achievement Chart
Categories
- communicate the
procedures, data, and
conclusions of an
investigation
involving the
determination of the
molecular formula of
the hydrate of
copper(II) sulfate,
using appropriate
means (e.g., oral and
written descriptions,
numerical analysis,
tables)
- communication of
information and ideas
The student:
- communicates
observations,
and conclusions
with clarity
- displays data in
complete, wellorganized charts
- displays data in
incomplete
charts that show
limited
organization
- displays data in
partially
complete charts
that show some
organization
- displays data in
mostly complete
charts that show
considerable
organization
- displays data in very
complete charts that show
a high degree of
organization
26
Day 11 1.
Independent research :
Analyse processes in the home, the workplace and the environmental sector that involve the use of
chemical quantities and calculations ( e.g mixing household cleaning solutions, calculating chemotherapy
doses, monitoring pollen counts. Consider the following issue : Health care professionals are expected to
calculate dosages of prescription drugs accurately and safely. This requires precision in applying fractions,
decimals, ratios, percentage and metric conversions. Despite the care taken improper medication use by
patients accounts for about 30% of hospital emergency visits. Answer the questions that follow :
1. Why is baking powder used in cake batter and what happens when too much or too little of that
ingredient is used.
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
2.
Why might two people on the same drug regimen not necessarily take the same dosage to treat the
same illness.
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
3.
How are carbon dioxide emissions calculated and why are they monitored.
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
_________________________________________________________________________
27
2.
1.
Assess on the basis of research, the importance of quantitative accuracy in industrial chemical processes and
the potential impact on the environment if quantitative accuracy is not observed. Consider the issue : Errors
in quantitative accuracy have played a role in many chemical disasters worldwide. Failing to adjust the
quantities of chemicals needed to produce different batch sizes of a product have created runaway reactions,
resulting in huge explosions. Such industrial accidents can have devastating short and long term effects on
the environment. Answer the questions that follow :
Why is it important to use the correct salt-sand mix on the highways during winter storms.
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
2.
Why is it important to correctly measure the chemicals used in water treatment plants and how might
incorrect measurements affect the environment.
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
3.
How and why are environmental contaminants monitored in soil, water, and air around a chemical
manufacturing plant.
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
28
Chapter 6 quiz.
Define stoichiometry. Do page 237 (1-3). Mole-mole problems.. Do page
238(4,5,7). Page 240 (8,10)
Stoichiometry is the quantitative description of the proportions by moles of the substances in a
chemical reaction
The octane present in gasoline burns according to the following equation:
2 C8H18 + 25 O2 ---------> 16 CO2 + 18 H2O
(a) How many moles of O2 are needed to react fully with 4 moles of octane?
(b) How many moles of CO2 can form from 1 mole of octane?
(c) How many moles of water are produced by the combustion of 6 moles of octane?
(d) If this reaction is to be used to synthesize 8 mole of CO2, how many moles of oxygen are needed? How many
moles of octane?
page 237 (1-3)
page 238(4,5,7)
Page 240 (8,10)
29
Day 12 -
Mass - mass problems. page 244 (11,13,14). Page 246 (15,17,18). Mass - particles
problems. Page 248(19,20,21,22). Assign page 250 (4,5,6,7)
mole ratio
molar mass
given
4NH3(g) +
4
17.04 g
34.0 g
5O2(g)
5
32 g
m

4NO(g) +
4
30.01 g
6H2O(ℓ)
6
18.02 g
3. Convert the given mass into moles.
4. Use the mole-to-mole ratio to find the required number of moles of the second substance.
5. Convert the number of moles of the second substance into the desired quantity (i.e., mass, molecules)
6. Write a concluding statement.
Below is a flow diagram showing the relationship between quantities in a balanced chemical equation.
M = molar mass
NA = Avogadro's constant
Mole ratio – from the balanced chemical equation
Example
During its combustion, ethane C2H6, combines with oxygen O2 to give carbon dioxide and water. A sample of
ethane was burned completely and the water that formed has a mass of 1.61 grams. How much ethane, in moles
and in grams, was in the sample? How many molecules of ethane were there ?
1. Set up the equation based on the words in the problem. Then balance it correctly.
2 C2H6
+ 7 O2
--------> 4 CO2 + 6 H2O

30
page 244 (11,13,14)
Page 246 (15,17,18)
Page 248(19,20,21,22).
Assign page 250 (4,5,6,7)
31
1.) How many grams of Ag2SO4 can be produced from the reaction of 13.6 grams of
AgNO3 with excess Na2SO4?
(12.5 grams)
___ AgNO3 + ___ Na2SO4  ___ Ag2SO4 + ___ NaNO3
2.) How many grams of NaOH are needed to react completely with 9.3 grams of CaCO3 in
a double replacement reaction?
(7.4 grams)
___ NaOH + ___ CaCO3  ___ Na2CO3 + Ca(OH)2
3.) How many grams of oxygen are needed for the complete combustion of 25.0 grams of
propane (C3H8)?
(90.9 grams)
___ C3H8 + ___ O2  ___ CO2 + ___ H2O
4.) How many grams of silver are produced in the reaction of 0.139 moles of AgNO3 with
copper?
(15.0 grams)
___ AgNO3 + ___ Cu  ___ Ag + ___ Cu(NO3)2
5.) How many moles of KNO3 can be produced when 24.0 grams of Ni(NO3)2 react with
excess KOH?
(0.263 moles)
___ Ni(NO3)2 + ___ KOH  ___ Ni(OH)2 + ___ KNO3
32
Case Study: Combustion and Carbon Monoxide Poisoning
Combustion (burning) is a chemical reaction in which oxygen reacts rapidly with a fuel to produce energy in the
form of heat and light. Common fuels include carbohydrates (wood), waxes (candles), and hydrocarbons (natural
gas and gasoline).
Hydrocarbon fuels contain primarily hydrogen and carbon. The simplest hydrocarbon fuel is methane, CH 4(g),
also known as natural gas. The combustion of hydrocarbons such as methane may be complete or incomplete,
depending on the amount of oxygen that reacts with the fuel. Complete combustion produces only carbon dioxide,
CO2 (g), and water vapour, H2O(g). Incomplete combustion produces carbon particles (soot), C(s),or carbon
monoxide, CO(g), and water vapour. The following equations represent the complete and incomplete combustion of
methane.
Complete combustion:
Incomplete combustion:
Incomplete combustion:
CH4(g) + 2 02(g)
2 CH4(g) + 3 02(g)
CH4(g) + 02(g)
C02(g) + 2 H20 (g)
2 CO(g) + 4 H20(g)
C (S) + 2 H20 (g)
The amount of heat produced is greatest in the complete combustion reaction and least in the reaction that
produces soot. The gaseous products of combustion are sometimes called flue gases, or exhaust gases. Flue
gases contain carbon dioxide, carbon monoxide, and water vapour. They may also contain small amounts of sulfur
dioxide, S02(g), sulfur trioxide, SO3(g), nitrogen, N2(g), and nitrogen oxides such as nitrogen monoxide, NO(g), and
nitrogen dioxide, N02(,,) . The sulfur oxides are produced by the reaction of sulfur impurities in the fuel. The
nitrogen oxides are produced by the combustion of nitrogen, which is the main constituent of air. (Air contains
approximately 21 % oxygen gas and 79% nitrogen gas.) Sulfur oxides and nitrogen oxides are the main causes of
acid precipitation in the atmosphere.
In the combustion reaction equations for methane, notice that the methane : oxygen mole ratio (CH 4(g) : 02(g) )
is higher for complete combustion than it is for the two forms of incomplete combustion.
Complete combustion:
CH4(g) : 02(g) = 1 : 2
Incomplete combustion (CO(g)): CH4(g) : 02(g) = 2: 3 or 1 : 1.5
Incomplete combustion (C(s)):
CH4(g) : 02(g) =1 : 1
The higher ratio means that more
oxygen (air) is required for complete combustion than for
incomplete combustion. Thus, to
ensure complete combustion, excess oxygen (excess air) must
be supplied beyond the amount theoretically required to combust the fuel.
In a stoichiometric calculation of the combustion of a hydrocarbon fuel, the fuel reacts with the exact amount of
oxygen that is required to combust all the hydrocarbon molecules to carbon dioxide and water vapour.
Theoretically, the exhaust gas contains no incompletely combusted fuel products (carbon monoxide and carbon)
and no unreacted oxygen. The percentage of carbon dioxide in the products of stoichiometric combustion is the
maximum attainable and is referred to as the stoichiometric carbon dioxide, ultimate carbon dioxide, or maximum
theoretical percentage of carbon dioxide.
Stoichiometric combustion is almost never possible in real life because of imperfect fuel-air mixing. For
economy and safety, most combustion equipment (such as heaters, furnaces, and engines) should operate with
excess air to ensure that fuel is not wasted and that combustion is
complete.
Conditions that promote incomplete combustion include
 insufficient air and fuel mixing;
 insufficient air supply to the flame;
 insufficient time for reactants to react;
 low flame temperature (slow combustion reactions).
Incomplete combustion is dangerous because carbon monoxide is a poison. It also uses fuel inefficiently
because complete combustion produces more heat. Finally, it contributes to air pollution because it produces
carbon particles and other undesirable combustion products.
Carbon Monoxide
Carbon monoxide is a colourless, odourless, tasteless gas. When it enters the lungs, it attaches itself to
hemoglobin in red blood cells. Carbon monoxide is more strongly attracted to hemoglobin than oxygen is, and it
forms stronger bonds with hemoglobin molecules. Inhaling high enough concentrations of carbon monoxide can be
fatal. Because of the strong bonds carbon monoxide forms with hemoglobin, the body takes many hours to rid
itself of carbon monoxide, even once the person is out of the poisoned atmosphere.
33
Furnaces, gas water heaters, fireplaces, wood stoves, charcoal grills, and gas-powered engines (including
automobile engines, lawmnower engines, and motorcycle engines) may produce carbon monoxide.
Symptoms of Carbon Monoxide Poisoning
Carbon monoxide concentrations are usually measured in parts per million (ppm). Table 1 lists the symptoms that
are associated with various amounts of carbon monoxide exposure.
Table 1 Symptoms of Carbon Monoxide Poisoning
CO(g) concentration
(ppm)
Exposure Time
200
2-3 h
400
1-2h
800
4 min
1600
20 min
6400
1-2 min
Symptoms
mild headache, fatigue, nausea, dizziness
severe headache, fatigue, nausea, dizziness, life threatening
after 3 h
dizziness, nausea, convulsions, death in 2-3 h
headache, dizziness, convulsions, death within 1 h
headache, dizziness, convulsions, death within 125 – 30 min
Since carbon monoxide attaches strongly to hemoglobin, many hours of exposure to low levels of carbon monoxide
can be just as deadly as shorter exposures to high levels.
Questions
1. (a) What is combustion? _________________________________________________________________
(b) Distinguish between complete and incomplete combustion of hydrocarbons.
___________________________________________________________________________________
___________________________________________________________________________________
2. (a) Calculate the mass of oxygen that is required for the complete combustion of 1.0 g of methane.
(b) Calculate the mass of oxygen that is required for the incomplete combustion of 1.0 g of methane to produce
carbon and water.
(c) Compare the masses of oxygen that are required for the complete and incomplete combustion of methane.
34
3. How does the fuel : oxygen mole ratio change between complete and incomplete combustion reactions?
4. (a) Why do flue gases include sulfur dioxide, sulfur trioxide, nitrogen monoxide, and nitrogen dioxide?
(b) What effect do these components of flue gas have on the environment?
5. (a) What is meant by the term “stoichiometric combustion”
(b) Why is there no carbon monoxide or soot in flue gas when stoichiometric combustion occurs?
(c) Why is stoichiometric combustion rare?
6. Describe three conditions that support incomplete combustion.
Day 13-
Limiting reagent problems. Page 254(23,24,26). Page 257(27,28,30). Assign page 259
35
(3-6)
Limiting Reactants
Suppose you are a chef preparing a breakfast for a group of people, and are planning to cook French toast. You make French
toast the way you have always made it: one egg for every three slices of toast. You never waiver from this recipe, because the
French toast will turn out to be either too soggy or too dry. There are 8 eggs and 30 slices of bread in the pantry. Thus, you
conclude that you will be able to make 24 slices of French toast and not one slice more.
This is a similar situation with chemical reactions in which one of the reactants is used up before the others - the reaction stops
as soon as one of the reactants is consumed. For example, in the production of water from hydrogen and oxygen gas suppose
we have 10 moles of H2 and 7 moles of O2.
Because the stoichiometry of the reaction is such that 1 mol of O2 : 2 moles of H2, the number of moles of O2 needed to react
with all of the H2 is ______________________
Thus, after all the hydrogen reactant has been consumed, there will be _____ moles of O 2 reactant left.
The reactant that is completely consumed in a chemical reaction is called the limiting reactant (or limiting reagent) because
it determines (or limits) the amount of product formed. In the example above, the _______ is the limiting reactant, and because
the stoichiometry is 2H2 : 2H2O (i.e. H2 :H2O), it limits the amount of product formed (H2O) to 10 moles. We actually have
enough oxygen (O2) to form ________ moles of H2O (1O2 : 2H2O).
One approach to solving the question of which reactant is the limiting reactant (given an initial amount for each reactant) is to
calculate the amount of product that could be formed from each amount of reactant, assuming all other reactants are
available in unlimited quantities. In this case, the limiting reactant will be the one that produces the least amount of potential
product.
Consider the following reaction: ___ Na3PO4 + ___ Ba(NO3)2 ---->___ Ba3(PO4)2 + __ NaNO3
Suppose that a solution containing 3.50 grams of Na3PO4 is mixed with a solution containing 6.40 grams of Ba(NO3)2. How
many grams of Ba3(PO4)2 can be formed?
___ Na3PO4 + ___ Ba(NO3)2 ---->___ Ba3(PO4)2 + __ NaNO3
Mole Ratio
_______
________
Molar Mass
_______
________
________
Given/Required
_______
________
________
1.
First we need to convert the grams of reactants into moles:
2.
Now we need to define the stoichiometric ratios between the reactants and the product of interest (Ba 3(PO4)2):
2 Na3PO4 : Ba3(PO4)2
3 Ba (NO3)2: Ba3(PO4)2
3.
We can now determine the moles of product that would be formed if reactant were to be consumed in its entirety
during the course of the reaction:
4.
The limiting reactant is the Ba (NO3)2 and we could thus make at most 0.0082 moles of the Ba3(PO4)2 product.
5.
The mass of Ba3(PO4)2 product would be equal to:
Page 254(23,24,26)
36
Page 257(27,28,30).
Assign page 259 (3-6)
37
Day 14 -
Percentage yield problems.. Page 262 (32,33) page 264 (35,36). Read page 265
Careers in Chemistry.
Theoretical yield
The quantity of product that is calculated to form when all of the limiting reactant is consumed in a reaction is
called the theoretical yield.
The amount of product actually obtained is called the actual yield.
Actual yield < Theoretical yield
for the following reasons:
•
for some reason not all the reactants may react
•
there maybe some significant side reactions
•
physical recovery of 100% of the sample may be impossible (like getting all the peanut butter out of the
jar)
The percent yield of a reaction relates the actual yield to the theoretical yield:
Percent yield =________________ x 100
For example, in the previous exercise we calculated that 4.94 grams of Ba 3(PO4)2 product should be formed. This
is the theoretical yield. If the actual yield were 4.02 grams the percent yield would be:
Percent yield =________________
x 100 = _____________________
1. Given the following equation:
a) Balance the equation.
_____ K2PtCl4 + 2 NH3 --------> _____ Pt(NH3)2Cl2 + 2 KCl
b) Determine the theoretical yield of KCl if you start with 34.5 grams of NH3 ( 151 g KCl )
a) Starting with 34.5 g of NH3, and you isolate 76.4 g of Pt(NH3)2Cl2, what is the percent yield?
Theoretical yield = 303.9 g Pt(NH3)2Cl2; 25.14 % yield
2. Given the following equation:
H3PO4 +
3 KOH
------> K3PO4
+ 3 H2O
If 49.0 g of H3PO4 is reacted with excess KOH, determine the percent yield of K3PO4 if you isolate 49.0 g of
Theoretical yield = 106.1 g K3PO4 ; 46.17 % yield
38
3. Given the following equation:
Al2(SO3)3 + 6 NaOH ------> 3 Na2SO3
+ 2 Al(OH)3
If you start with 389.4 g of Al2(SO3)3 and you isolate 212.4 g of Na2SO3, what is your percent yield for this
reaction? Theoretical yield = 500.6 g Na2SO3; 42.43 % yield
4. Given the following equation:
Al(OH)3 (s) + 3 HCl (aq) -------> AlCl3 (aq) + 3 H2O (l)
If you start with 50.3 g of Al(OH)3 and you isolate 39.5 g of AlCl3, what is the percent yield?
86.0 g AlCl3; 45.93 % yield
Page 262 (32,33)
page 264 (35,36).
Theoretical yield =
39
Read page 265 Careers in Chemistry.
WS – Limiting Reactant and % Yield
Name__________________
Answer the following questions using dimensional analysis. Give you answers in scientific notation (except for
percentages). Make sure all answers reflect the correct number of significant figures.
1. If 50.0 grams of zinc metal reacts with 250. grams of aqueous copper I nitrate then a) how many grams of
copper metal will be produced and b) how much of the excess reactant is unused according to the following
equation.
Zn(s) + 2CuNO3 (aq)  2Cu(s) + Zn(NO3)2 (aq)
L.R. = _____________
Answers a) 9.72 x 101 g Cu produced
b) 5.8 x 101 g CuNO3 unused
2. Determine the actual yield of carbon tetrachloride if 150. g of chlorine gas reacts with excess methane. The
percent yield for the reaction is 75.4%.
Answer 6.14 x 101 g CCl4 actually
produced
CH4(g) +
4Cl2 (g)  CCl4 (l) + 4HCl (g)
40
Day 15
A Stoichiometric Analysis of Alka-Seltzer
Names ___________________________________
Alka-Seltzer is one of the world's best known antacids. Its main function is to absorb excess stomach acid. The
two ingredients that accomplish this are sodium bicarbonate (NaHCO3) and citric acid (C6H8O7). Aspirin is also
present in Alka-Seltzer tablets to reduce fever and relieve headaches, but in this lab, we are going to study the
reaction that takes place between the NaHCO3 and C6H807:
NaHCO3(aq) + C6H8O7(aq)
------>
H2O(l) + CO2(g) + Na3C6H5O7(aq)
Whenever an Alka-Seltzer tablet is placed in water, the above UNBALANCED equation is what is taking place.
Note that without water, that tablet stays unreacted. It is stored in a foil pouch to prevent water vapor in the air
from causing the tablet to start reacting. Na3C6H5O7 is called Sodium Citrate.
Purpose
The purpose of this lab is to determine what mass of carbon dioxide is produced from the reaction of one tablet
of Alka-Seltzer and water.
Hypothesis
Predict the mass of CO2 that will be produced by looking carefully at the ingredients of an Alka-Seltzer tablet.
You must balance the above equation and find out which reactant is the limiting reactant. Using the
amount of limiting reactant present, calculate the mass of CO 2 that should be produced. This is your
hypothesis.
To predict the mass of CO2 produced when one tablet reacts in water complete the following:
1.
Calculate moles of NaHCO3 in one unreacted tablet? (You must look at the back of the package to get
ingredients. The ingredient masses are for one tablet. Use the mass listed on the package to calculate
moles)
MOLES of NaHCO3 ___________________
2.
Calculate moles of C6H8O7 in one unreacted tablet? (You must look at the back of the package to get
ingredients. The ingredient masses are for one tablet. Use the mass listed on the package to calculate
moles)
MOLES of C6H8O7 _____________________
3.
There are several ways to determine which reactant limits a reaction. The following will show you the
steps of one method. If your balanced equation above shows a 3 to 1 ratio between NaHCO 3 and C6H8O7
respectively, you balanced the equation properly. This is the ratio we will use to determine which reactant
limits the reaction. According to this ratio, if we do not have 3 times as much NaHCO3 as C6H8O7 then
NaHCO3 limits the reaction. If we have more than 3 times as much NaHCO 3 as C6H8O7 then C6H8O7 limits the
reaction. So lets do some calculations:
Fill in the moles of NaHCO3 calculated from in step 1:
a. ______moles of NaHCO3
b.
*
1 C6H8O7
=
3 NaHCO3
_______moles C6H8O7
From here you simply compare the moles of C6H8O7 calculated in step 3a (theoretical
calculation) to the moles of C6H8O7 calculated in step two (actual moles). If the moles of
41
C6H8O7 in step 3a (theoretical calculation) are less than actual moles calculated in step 2
than NaHCO3 limits the reaction. This means there is enough C6H8O7 to react all the
NaHCO3 and we will have some left over. If the moles of C6H8O7 in step 3a (theoretical
calculation) are more than actual moles calculated in step 2 than C6H8O7 limits the reaction
because we do not have enough. This means that C6H8O7 will run out before NaHCO3.
c.
What reactant limits the reaction ________________
4. Now you use the moles of the limiting reactant calculated in step 1 or 2 (depending on which one limits
the reaction) and its ratio to CO2 to determine the amount of CO2 produced. This remember is your
hypothetical.
a.
___moles of limiting reactant * ratio of limiting reactant to CO 2 = ___moles of CO2 produced
Hypothetical grams of CO2 produced is ____________
Materials/Procedure
Fill a Styrofoam cup about half full of distilled water and weigh it. Weigh one Alka-Seltzer tablet. Add the tablet
to the water and observe the reaction. After the tablet has dissolved, weigh the cup and its contents. Wait
about 10 minutes and weigh the cup again. Record both measurements in your data table.
Mass of cup and water __________
Mass of one Alka-Seltzer_________
Mass of cup and contents after reaction________
Ten minutes later_______________
How much CO2 was released in the reaction (CO2 is the only gas) ____________
Conclusion
1.
How does you experimental value of CO2 compare to you theoretical value? Calculate the percent error.
2. What is the limiting reactant in an Alka-Seltzer tablet? Test this by adding either citric acid or sodium
bicarbonate to your plastic cup to see if you can make the reaction start again.
3. How many grams of the excess reactant is left over after an Alka-Seltzer tablet dissolves in water?
4. How many grams of which reactant would you have to add to your cup after the reaction is complete to
completely use up the excess reactant?
5. Using your data, what was the percent yield of CO2 in this reaction?
6. Suggest a reason why your percent yield is not 100% What could be happening? (Hint - do you think all of
the CO2 left the cup?)
42
Unit 2 : Mole and Mass Calculations
Mole/Mole, Mass/Mass, Limiting Reagents, % Yield
1.
For the reaction:
VO(s) +
a)
Balance the equation. ( 2,3,6,1 )
b)
Calculate the number of moles of Fe2O3 that are necessary to completely react with 1.00 mole of VO (
1.5 mol. )
c)
Calculate the number of moles of Fe2O3 necessary to produce 3.00 moles of V2O5.( 9 mol .)
2.
In the neutralization of sulfuric acid with sodium hydroxide:
H2SO4
+
Fe2O3(s) -------->
NaOH
-------->
FeO(s)
H2O
+
+
V2O5(s)
Na2SO4
a)
Balance the equation. ( 1,2,2,1 )
b)
If there are 8.61 moles of sulfuric acid, how many moles of sodium hydroxide are required? ( 17.2
mole )
c)
How many moles of sulfuric acid can be neutralized by 3.5 moles of sodium hydroxide? ( 1.75 moles
)
d)
How many moles of sodium sulfate are produced by the neutralization of 2.10 moles of sulfuric acid?
( 2.10 moles )
3.
If 6.4 g of methane (CH4) are burned according to the balanced equation:
CH4(g) + 2 O2(g)
-------->
CO2(g)
+
2 H2O(g)
a)
How many grams of oxygen are required? ( 25.6g )
b)
How many grams of carbon dioxide will be produced? ( 17.6g )
4.
Given the equation: FeS
a)
Balance the equation. ( 4,7,2,4 )
b)
How many grams of FeS will react with 7.81 g of oxygen? ( 12.06 g )
c)
If 180.0 g of FeS react with excess oxygen, what mass of Fe2O3 will be produced? ( 163.2 g )
+
O2
---------->
Fe2O3
+
SO2
43
d)
(L.R) How many grams of SO2 are produced when 20.9 g of FeS react with 14.1 g of O2? ( 15.23 g )
5.
Tungsten (W) metal is prepared by reacting tungsten (VI) oxide with hydrogen gas as shown in the
equation:
WO3 + H2 ---------->
W +
H2O
a)
Balance the equation. ( 1,3,1,3 )
b)
What mass of WO3, is required to prepare 1.00 kg of tungsten metal?
6.
Given the equation: Al + H2SO4 ------> Al2(SO4)3
a)
Balance the equation. ( 2,3,1,3 )
b)
(L.R) How many grams of aluminum sulfate form when 6.71 g of aluminum react with 12.95 g of sulfuric
acid? ( 15.06 g )
c)
If 28.0 g of aluminum react completely with excess sulfuric acid, how many grams of hydrogen gas are
produced. ( 3.11 g )
7.
Hydrogen gas is commonly prepared in the lab by the reaction of zinc and some acid such as sulfuric acid or
hydrochloric acid.
a)
How many grams of zinc are required to react with excess sulfuric acid in order to prepare 20.0 g of
hydrogen gas:
Zn + H2SO4 ---------> H2
+
+
H2
ZnSO4
b)
If hydrochloric acid is to be used, how many grams of pure acid will be required to produce 10.0g of
hydrogen gas, assuming excess zinc is used? Given the equation:
Zn + HCl -------->
H2
+ ZnCl2
8.
Ammonia gas and oxygen gas react according to the balanced equation:
4 NH3(g) + 5 O2(g) -------> 4 NO(g) + 6 H2O(g)
a)
(L.R)
If 25.0 g of ammonia react with 30.0 g of oxygen, how many grams of nitrogen monoxide result?
b)
How many grams of ammonia are required to react with excess oxygen, in order to produce 80.0 grams of
NO?
9.
Try these ones: Calculate percentage yield for the following.
44
a)
16.0 g of sulfur are burned in excess oxygen; 20.0 g of sulfur dioxide result. ( 63% )
b)
3.00 g of carbon are burned in excess oxygen; 8.00 g of carbon dioxide result. ( 73% )
c)
26.16 g of zinc are combined with excess sulfuric acid; 50.0 g of zinc sulfate are isolated. ( 77% )
d)
2.50 g of mercury (II) oxide are decomposed by heating. If 2.0 g of mercury are produced, determine the
theoretical, actual, and percentage yield. ( 2.3g, 2g, 87% )
10.
Given: 2 KMnO4 + 5 H2SO3 --------> K2SO4 + 2 MnSO4 + 2 H2SO4 + 3 H2O
a)
What mass of H2SO3 is needed to react with 3.16 g of KMnO4? (4.10 g )
b)
What mass of H2O can be produced by reacting 32.8 g of H2SO3 with excess KMnO4? (4.32 g )
c)L.R
What mass of MnSO4 can be produced by treating 7.90 g of KMnO4 with 13.0 g of H2SO3? ( 7.55 g
)
11. L.R & P.Y If 8.30 g of sodium and 14.0 g of chlorine are heated together, a total of 19.5 g of sodium chloride is
isolated. Determine the percentage yield of the process. ( 93% )
45
Day 16-
Work period page 271 ( 5,6,8,9, 11, 14, 15).Homework review for test page 276 (129),33-35,38,51.
UNIT 2 REVIEW
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
What is the mass of 2.5 X 1025 molecules of methanol (CH4O)?
What is the mass of 4.39 moles of Mg(NO3)2?
The average mass of a molecule of morphine is 4.74 X 10-22 g. Determine the molar mass of
morphine.
How many atoms of O are in 2.33 g of nitrogen dioxide?
Determine the mass of one CCl4 molecule.
What is the percent composition of C, H, and O in ASA (C9H8O4)?
A 2.68 g sample of a hydrate of sodium sulfate, (Na2SO4xH2O), is heated. The mass of water
released is determined to be 1.26 g. What is the empirical formula?
Orlon is 67.9% C, 5.70% H, and 26.4% N. Determine the empirical formula.
Lactic acid consists of C: 40%, H: 6.7%, O: 53.3%. The molar mass of lactic acid is 90 g/mol.
a) What is the empirical formula?
b) What is the molecular formula?
Consider this balanced equation:
2 Na(s) + 2 H2O(l) --- 2 NaOH(aq) + H2(g)
a) How many moles of sodium are needed to produce 0.85 mole of hydrogen gas?
b) How many moles of water are needed to produce 0.4 mole of sodium hydroxide?
11.
Consider the reaction N2 + 3H2 --- 2NH3. How many grams of NH3 are produced when 6.0
g of H2 completely react?
12.
A solution with 14.0 g of AgNO3 is mixed with a solution with 4.83 g of CaCl2. A precipitate of
AgCl is formed with Ca(NO3)2 remaining in solution.
a) Write the balanced equation for this reaction.
b) How many grams of silver chloride are formed?
c) How many grams of excess reagent are remaining?
13.
Water and nitrogen dioxide react to form nitrogen monoxide and nitric acid.
a) Write the balanced equation for the reaction.
b) If 100.0 g of nitrogen dioxide combine with 20.00 g of water, determine the limiting
reagent.
c) What is the theoretical yield of nitric acid in grams?
14.
In this reaction: 2 NaClO3 -- 2 NaCl + 3 O2, how many grams of oxygen are produced from
12.00 moles of NaClO3?
15.
A student burns a 5.00 g piece of magnesium ribbon in air to produce 8.10 g of magnesium
oxide. What is her percentage yield?
16.
Virtually all the nitric acid manufactured commercially is obtained by the ammonia oxidization
process. This involves this first step: Ammonia gas combines with oxygen to form nitrogen
monoxide and water vapour.
a) Write the balanced equation for the reaction.
b) The reaction of 36.61 g of NH3 produces a 6.11% yield of water. How many grams are
produced?
46
18.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
The processing of zinc begins with the roasting of "zinc blende" (ZnS) in oxygen to form zinc
oxide and sulfur dioxide.
a) Write the balanced equation for the reaction.
b) If the reaction of 33.31 g of ZnS and 11.11 g of O2 produces 12.54 g of SO2, what is the
percent yield?
1330 g
651 g
285 g/mol
6.10 Á 1022 atoms
2.55 Á 10-22 g
60.0% C, 4.4% H, 35.6% O
Na2SO4 ú 7H2O
C3H3N
a) CH2O
b) C3H6O3
C7H6O2
a) 1.7 moles
b) 0.4 mole
34 g NH3
a) 2 AgNO3(aq) + CaCl2(aq) Ģ AgCl(s) + Ca(NO3)2(aq)
b) 11.8 g
c) 0.260 g CaCl2 remain
a) 3 NO2(g) + H2O(l) Ģ 2 HNO3(aq) + NO(g)
b) NO2 is limiting
c) 91.29 g
576.0 g O2
97.6%
a) 4 NH3(g) + 5 O2(g) Ģ 4 NO(g) + 6 H2O(g)
b) 3.55 g H2O
a) ZnS(s) + 3 O2(g) Ģ 2 ZnO(s) + 2 SO2(g)
b) 84.56%