UNIT 3 – QUANTITIES IN CHEMICAL REACTIONS
Unit Expectations:
Page References
After completing this unit, the student should be able to:
1) Define accuracy (pg. 672), precision (pg. 677) and answer questions
related to them.
18-19
2) a) Identify the significant digits in a measurement.
b) Express, rounded off to the correct number of significant digits,
the results of calculations involving measurements.
15-17
20-22
3) Calculate the % error for an experimental result.
Labs
4) a) Define average atomic mass (pg. 164).
b) Given the isotopic composition of an element, calculate the average
atomic mass of the element.
c) Given the mass numbers of the isotopes and the average atomic mass,
calculate the % abundance of the isotopes of an element.
Note
Note
5) Define the following: mole (pg. 172), Avogadro constant (pg. 172),
molar mass (pg. 180).
6) Solve problems involving # of moles and # of particles. (n =
N
)
NA
7) Solve problems involving # of moles, mass and molar mass. (n =
Note, 175-178
m
)
M
Note, 180-192
8) a) Define the law of definite proportions (pg.198).
b) Calculate the percentage composition by mass of a compound.
Note, 200-201
9) Calculate the simplest (empirical) formula of a compound.
Note, 207-211
10) Calculate the molecular formula of a compound.
Note, 215-225
11) a) Define stoichiometry (pg. 242).
b) Interpret chemical equations in terms of moles, molecules, ions
and atoms.
Note, 234-240
12) Solve problems involving mass-mass and mass-particle
stoichiometry.
13) a) Define the following: theoretical yield (pg. 260), actual yield (pg. 260)
percentage yield (pg. 261).
b) Solve problems involving the theoretical yield and the percentage yield
of a product of a chemical reaction.
14) a) Define limiting reactant (pg. 676).
b) Solve stoichiometric problems involving limiting reactants.
Note, 241-248
Note, 260-269
Note, 251-258
ACCURACY & PRECISION
Accuracy Precision -
Example
CALCULATIONS AND SIGNIFICANT DIGITS
Significant Digit - __________________________________________________________
Identifying the number of sig digs in a calculation or measurement is important, especially when
dealing with calculators, which do not consider sig digs.
Every measurement that we make will have some number of “certain” digits, and a final
“estimated” digit, which is the result of rounding.
HOW CAN WE TELL WHICH DIGITS ARE SIGNIFICANT?
RULES
1.
EXAMPLES
7.866 g
19.4 m
527.266 992 cm3
2.
408 kPa
25 074 L
3.
0.0907 ºC
0.000 000 000 06 km
4.
704.0 m
8200 kg
SCIENTIFIC NOTATION -
e.g.
10 492 000 m to 5 sig digs
0.00003456 kg to 4 sig digs
CALCULATING WITH SIGNIFICANT DIGITS
We have to be aware of the degree of certainty of our measurements whenever we are doing
calculations in chemistry.
Remember: your calculator usually reports results with much greater certainty (more significant
figures) than your data warrant. You must make the decisions about the certainty of your
measurements.
RULE
EXAMPLES
1. Multiplying and Dividing
2.34 cm x 1.5 cm =
2. Adding and Subtracting
4.35 kg + 0.346 kg =
3. Rounding
2.346 to 3 sig digs
5.73 to 2 sig digs
18.35 to 3 sig digs
18.25 to 3 sig digs
EXERCISE: ACCURACY, PRECISION & SIGNIFICANT DIGITS
Complete the following questions from the text book:
pg 22 #3, pg 24 #4 (answers pg 31), pg 29 – 31 # 4,7,8b, 9abc, 10abc, 14bc (answers pg 644)
pg 154 – 156 # 3 – 5, 10, 13, 14, 42 (answers pg 644)
Also, complete the following questions:
1) What is the difference between accuracy and precision?
2) How many significant digits are in each of the following?
a) 624 students
b) 22.40 mL
c) 0.00786 g
3) Round the following measured quantities to the number of significant digits specified:
a) 9.276 x 103 m (2 sig digs)
b) 87.45 g (3 sig digs)
c) 93.951 kg (3 sig digs)
4) Express each calculation to the correct number of significant digits:
a) 10.25 g + 2.5 g
b) 5845 cm x 23.9 cm
c) 1106.9 mL – 783.687 mL
d) 64.3 g ÷ 9.52 mL
Worksheet Numerical Answers:
2. (a) 3 sig digs
(b) 4 sig digs (c) 3 sig digs
3. (a) 9.3 x 103 m
(b) 87.4 g
4. (a) 12.8 g (b) 1.40 x 105 cm2
(c) 94.0 kg
(c) 323.2 mL (d) 6.75 g/mL
THE MOLE
Atomic Masses
- Look at atomic masses on the periodic table. What do these numbers represent?
e.g.
C has an atomic mass of 12.
Protons and neutrons have roughly the same mass. So, C weighs ______
(atomic mass units).
-
What is the actual mass of a C atom?
Answer:
Why Don’t We Use Actual Masses of Atoms?
1.
2.
With These Problems, Why Use Atomic Mass At All?
1. Masses give information about # of p+, n0, e–
2. It is useful to know relative mass
e.g.
What ratio is needed to make H2O?
3. It is useful to associate atomic mass with a mass in grams.
4. It has been found that 1 g H, 12 g C, or 23 g Na have 6.02 x 10 23 atoms
THE MOLE:
An easy way to say _______________________________________________
Also known as _____________________
n=
𝑛=
𝑁
𝑁𝐴
N=
NA=
Try This:
There are 268 donuts. How many dozens are present?
There are 2.46 x 1024 donuts. How many moles of donuts are present?
EXAMPLE: There are 0.50 mol of methane, CH4, in a flask. How many molecules of
methane are in the flask?
EXAMPLE: In the same sample of methane, determine how many atoms of H are present.
Note: Covalent compounds (e.g. H2O) are composed of ___________________________
Ionic compounds (e.g. NaCl) are composed of ______________________________
EXERCISE: MOLES AND PARTICLES
1) A small pin contains 0.0178 mol of iron, Fe. How many atoms of iron are in there in the pin?
2) How many formula units are contained in 0.21 moles of magnesium nitrate, Mg(NO3)2?
3) A liter of water contains 55.6 mol of water. How many molecules of water are in this
sample?
4) Ethyl acetate, C4H8O2, is frequently used in nail polish remover. A typical bottle of nail polish
remover contains about 2.5 mol of ethyl acetate.
a) How many molecules are in the bottle of nail polish remover?
b) How many atoms are in the bottle?
c) How many carbon atoms are in the bottle?
5) A sample of bauxite ore contains 7.71 x 1024 molecules of aluminum oxide, Al2O3. How
many moles of aluminum oxide are in the sample?
6) A vat of cleaning solution contains 8.03 x 1026 molecules of ammonia, NH3. How many
moles of ammonia are in the vat?
7) A sample of cyanic acid, HCN, contains 3.33 x 1022 atoms. How many moles of cyanic acid
are in the sample? HINT: Find the number of molecules of HCN first.
8) A sample of pure acetic acid CH3COOH, contains 1.40 x 1023 carbon atoms. How many
moles of acetic acid are in the sample?
9) Imagine that $6.02 x 1023 were evenly distributed among six billion people. How much
money would each person receive?
10) A typical adult human heart beats an average of 60 times per minute. If you were allotted a
mole of heartbeats, how long, in years, could you expect to live? You may assume that each
year has 3.65 days.
ANSWERS
1) 1.07 x 1022 atoms
2) 1.3 x 1023 formula units
3) 3.35 x 1025 molecules
4) a) 1.5 x 1024
b) 2.1 x 1025 c) 6.0 x 1024
5) 12.8 mol
6) 1.33 x 103 mol
7) 1.84 x 10-2 mol
8) 0.116 mol
9) $1.00 x 1014
10) 1.9 x 1016 years
MOLAR MASS
Molar Mass:
e.g.
What is the molar mass of:
(a) Carbon, C?
(b) Calcium, Ca?
(c) Carbon dioxide, CO2?
(d) Sodium sulfate, Na2SO4?
(e) Calcium phosphate, Ca3(PO4)2?
If we know the mass of the substance that we have, we can figure out how many moles of the
substance that are present.
𝑛=
𝑚
𝑀
n=
m=
M=
EXAMPLE:
Calculate the number of moles in 675 g of calcium phosphate.
EXAMPLE:
If there are 450 mg of magnesium phosphate, how many formula units are
present?
EXERCISE: MOLES TO MASS
Do Practice Problem 12 (pg. 175) and Qu. 2, 3 (pg. 179) (Answers on pg 195). Then, complete
the following questions:
1. For each of the following, which group has the largest mass?
a) 5.00 mol of C, 1.50 mol of Cl2, 0.50 mol of C6H12O6
b) 7.31 mol of O2, 5.64 mol of CH3OH, 12.1 mol of H2O
2. How many moles of compound are in each sample?
a) 39.2 g of silicon dioxide, SiO2
b) 7.34 g of nitrous acid, HNO2
c) 1.55 x 105 kg of carbon tetrachloride, CCl4
3. Determine the mass of each sample:
a) 6.02 x 1024 formula units of ZnCl2
b) 7.38 x 1021 formula units of Pb3(PO4)2
4. Vitamin B, C17H20N4O6, is also called riboflavin. What is the mass, in grams, of a single
molecule of riboflavin?
5. Determine the number of molecules in the sample:
a) 10.0 g of water, H2O
b) 52.4 g of methanol, CH3OH
ANSWERS
1.
2.
3.
4.
5.
a) 1.5 mol of Cl2 b) 7.31 mol O2
a) 0.652 mol b) 0.156 mol c)1.01 x 106 mol
a) 1.36 x 103 g b) 9.95 g
6.25 x 10 –22 g
a) 3.34 x 1023 molecules b) 9.84 x 1023 molecules
CHALLENGE:
1. Imagine a computer that is capable of counting at the rate of 1 x 10 9 numbers per
second. How many years would it take the computer to count to 6.02 x 10 23?
(A: 1.91 x 107 years)
2. How long would it take to spend a mole of $1 coins if they were being spent at a rate of 1
billion per second? (A: 19 million years)
PERCENTAGE COMPOSITION
The calculations that you will be doing involve determining the percentages of elements
in compounds.
EXAMPLE 1 - Using Actual Masses
A compound with a mass of 48.72 g is found to contain 32.69 g of zinc and 16.03 g of sulfur.
What is the percentage composition of the compound?
EXAMPLE 2 – Using Molar Mass Ratios
Calculate the percentage composition of C3H7NO2.
EXAMPLE 3 - An Alternative Type Question
Calculate the mass of aluminum in 28g of aluminum oxide
EXERCISE – PERCENTAGE COMPOSITION
1) A sample of a compound is analyzed and found to contain 0.90 g of calcium and 1.60 g of
chlorine. The sample has a mass of 2.50 g. Determine the percentage composition of the
compound.
2) Find the percentage composition of a pure substance that contains 7.22 g of nickel, 2.53 g of
phosphorus and 5.25 g of oxygen.
3) Bromine azide is a compound that is used in explosives. A sample of bromine azide was
found to contain 2.35g of Br and 1.24 g of N. Determine the percentage composition.
4) Determine using the periodic table, the molar masses of the following compounds.
a) Na3PO4
c) Mg3(PO4)2
b) Ca(NO3)2
d) CH3COOH
5) Sulphuric acid, H2SO4, is an important acid in laboratories and industries. Determine the
percentage composition of sulphuric acid.
6) A mining company wishes to extract manganese metal from pyrolusite ore, MnO 2. What is
the percentage composition of pyrolusite ore?
7) Pyridine, C5H5N is a slightly yellow liquid with a nauseating odour. It is used in the synthesis
of vitamins and drugs and has many other uses in industrial chemistry. Determine the
percentage composition of pyridine.
8) Indigo, C16H10O2, is the common name of the dye that gives blue jeans their characteristic
colour. Calculate the mass of oxygen in 25.0 g of indigo.
9) Potassium perchlorate, KClO4, is used extensively in explosives. Calculate the mass of
oxygen in a 24.5 g sample of potassium chlorate.
10) A piece of magnesium reacts with 16.0 g of oxygen to form 40.3 g of a compound.
a) Calculate the percentage of magnesium in the compound.
b) Calculate the mass of magnesium that would be present in 10 000 g of the compound.
11) a) Calculate the percentage composition of iron (III) oxide.
b) Calculate the mass or iron that could be obtained from 20 000 g of this compound.
12) A sample of a compound with K, Cl and O in it, was placed in a test tube and decomposed
by heating. After heating, a compound containing K and Cl was left (called the residue) and
oxygen gas escaped into the room. From the given data, calculate the percent of oxygen in
the compound.
Mass of empty test tube
= 18.00 g
Mass of test tube and compound
= 24.13 g
Mass of test tube and residue
= 21.73 g
ANSWERS: 1) Ca = 36%, Cl = 64%; 2) Ni = 48.1%, P = 16.9%, O = 35.0%; 3) Br = 65.5%, N = 34.5%;
4) a) 163.94 g/mol b) 164.10 g/mol c) 262.87 g/mol d) 60.06 g/mol ; 5) H = 2.06%, S = 32.70%,
O = 65.25%; 6) Mn = 63.19%, O = 36.81%; 7) C = 75.91%, H = 6.38%, N = 17.71%; 8) 3.42 g; 9) 11.3 g ;
10) a) 60.3% b) 6.03 x 103 g; 11) a) Fe2O3 – 69.94%Fe, 30.06%O
b) 1.399 x 104g; 12) 39.2%
SIMPLEST (EMPIRICAL) AND MOLECULAR FORMULA
MOLECULAR FORMULA:
SIMPLEST FORMULA:
(EMPIRICAL)
EXAMPLE 1:
Given that a compound contains 12.7% C, 2.1% H, and 85.2% Br,
calculate its simplest (empirical) formula.
1. Assume that you have 100 g of the substance.
2. Mole Calculation: Using , 𝒏 =
3. Ratio Calculation:
Calculate moles
𝒎
𝒏= ,
𝑴
Ratio
𝑴
, calculate the moles for each element.
Divide the smallest number of moles by each other value to
determine the ratio between the atoms in the formula.
Carbon
Assume 100g
𝒎
Hydrogen
Bromine
EXAMPLE 2:
Calculate the molecular formula of the compound in the previous example
if the actual molar mass of the compound is 190 g/mol.
EXAMPLE 3:
What is the empirical formula of a compound that contains 69.9 g Fe
and 30.1 g O?
EXERCISE: SIMPLEST AND MOLECULAR FORMULA
1) What is the empirical formula of a compound that contains 46.3 % lithium and 53.7%
oxygen?
2) What is the empirical formula of a compound that contains 15.9 % boron and 84.1 %
fluorine?
3) Phosphorus reacts with oxygen to give a compound that is 43.7% phosphorus and 56.4%
oxygen. What is the empirical formula of the compound?
4) An inorganic salt is composed of 17.6% sodium, 39.7% chromium and 42.8% oxygen. What
is the empirical formula of this salt?
5) Compound X contains 69.9% carbon, 6.86% hydrogen and 23.3% oxygen. Determine the
empirical formula of X.
6) Oxalic acid has the empirical formula CHO2. Its molar mass is 90 g/mol. What is the
molecular formula of oxalic acid?
7) The empirical formula of codeine is C18H21NO3. If the molar mass of codeine is 299 g/mol,
what is its molecular formula?
8) A compounds molar mass is 240.28 g/mol. Its percentage composition is 75.0% carbon,
5.05% hydrogen and 20.0% oxygen. What is the compound’s molecular formula?
9) The wintergreen plant produces methyl salicylate, or oil of wintergreen. It can also be
prepared easily in the laboratory. Methyl salicylate is 63.1% carbon, 5.31% hydrogen and
31.6% oxygen. Calculate the empirical formula of this compound.
10) An inorganic salt is made up of 38.8% calcium, 20.0% phosphorus, and 41.2% oxygen.
a) What is the empirical formula of the compound?
b) On further analysis, each formula unit of this salt is found to contain two phosphate ions.
Predict the molecular formula of this salt.
Answers
1) Li2O
2) BF3
3) P2O5
4) Na2Cr2O7
5) C12H14O3
6) C2H2O4
7) C18H21NO3
8) C15H12O3
9) C8H8O3
10) a) Ca3P2O8 b) Ca3(PO4)2
HYDRATES
HYDRATE:
ANHYDROUS:
EXAMPLE 1
(a) Calculate the percentage of water in Na2S2O3 ∙ 5H2O
(b) Calculate the mass of water in 215 g of Na2S2O3∙ 5H2O.
EXAMPLE 2
A hydrate of barium chloride (BaCl2∙ xH2O) has a mass of 1.500g. When heated to drive off the
water, the residue (BaCl2) has a mass of 1.279g.
(a) Calculate the percentage of water in the hydrate.
(b) Calculate the formula of the hydrate (i.e. solve for x)
Homework: pg 221, 225 # 21, 23, 24; pg 228 #6, 7; pg 229-230 #5, 7a, 21, 22b
STOICHIOMETRY
STOICHIOMETRY: Calculations involving chemical reactions;
Also refers to any calculation that involves the use of mole ratios.
EXAMINE:
NH3(aq)
O2(g) 
+
H2O (l)
+
NO(g)
-
We balance chemical equations to make sure that we have the same number of each
type of atom on each side of the equation.
(_______________________________________________)
-
In other words, the equation tells us that for every __________________ used, we need
_________________, and __________________ are produced, etc.
-
We CANNOT, however, convert directly from grams to grams. _________________
does not give us _____________________ because the particles have different
_____________.
EXAMPLE 1: 4NH3(aq)
+
5O2(g)

6H2O (l)
+
4NO(g)
How many moles of H2O (l) are produced if 0.176 mol of O2(g) are used?
EXAMPLE 2: How many moles of NO(g) are produced if 17 moles of H2O (l) are also
produced?
Notice that in order to correctly determine this, we need a ____________________________.
The fraction that we use to determine number of moles is called a ______________________.
Notice that we cannot directly convert from grams of one compound to grams of another.
Instead we have to go through moles.
Many stoichiometry problems follow a pattern:
Molar Mass of x
Molar Mass of y
grams(x)  moles(x)  moles(y)  grams(y)
Mole Ratio
We can start anywhere along this path depending on the question we want to answer.
Mass to Mass Stoichiometry
EXAMPLE 3:
Fe2O3 (s) + 3CO (g)  2Fe (s) + 3CO2 (g)
How many moles of carbon monoxide are required to react with 163.0 g of
iron(III) oxide?
HW: pg 237 – 246 # 1bc, 2 – 6, 10a, 13 – 16; pg 250 #4
(check answers on pg 273)
STOICHIOMETRY & PERCENTAGE YIELD
RECALL - STOICHIOMETRY:
RECALL:
Calculations involving chemical reactions.
Many stoichiometry problems follow a pattern:
Molar Mass of x
Molar Mass of y
grams(x)  moles(x)  moles(y)  particles*(y)
Mole Ratio
*can be mass
or other things
as well
We can start anywhere along this path depending on the question we want to answer.
We can use ___________________ in combination with any equation that uses ____________.
Therefore, we can also determine ______ through stoichiometry if we are interested in doing so.
EXAMPLE 1: Calculate the number of molecules of O2(g) that would be needed to react
with 50.0 g of C3H8 (g).
C3H8 (g) + 5O2(g)  3CO2 (g) + 4H2O (g)
EXAMPLE 2: How many grams of oxygen would be required to produce 5.45 x 1024
molecules of water vapour?
PERCENTAGE YIELD
-
In many reactions, the reaction will ________________________________________.
-
As a result of this, _______________________________ will be converted to products.
-
In order to determine how “far along” a reaction has proceeded, we use a concept called
“____________________________”
-
In % Yield, we compare the amount of product that was _________________________
_______________ (the ___________________) to the amount of product that ________
_______________________________, according to our stoichiometric calculations
(the __________________________).
% 𝑌𝑖𝑒𝑙𝑑 =
EXAMPLE 3:
𝐴𝑐𝑡𝑢𝑎𝑙 𝑌𝑖𝑒𝑙𝑑
𝑥 100
𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑌𝑖𝑒𝑙𝑑
In the fermentation process used to make wine, sugar (C6H12O6(s)) is
decomposed into ethanol (C2H5OH(aq)) and carbon dioxide gas.
C6H12O6(s)  2 C2H5OH(aq) + 2CO2(g)
(a) What is the theoretical yield of ethanol available from 10.0g of sugar?
(b) In a particular experiment, 10.0 g of sugar produces 0.664 g of ethanol. What is the
percentage yield of the experiment?
(c) If a student gets 95.0% yield, what mass of ethanol was produced?
HW: pg 248 – 249 #19 – 22; pg 262-264 #31 – 34, 35ab, 36
(check answers pg 273)
LIMITING REAGENTS
Balloon and Flask Demonstration
g of NaHCO3 (s)
mL of 3M HCl(aq)
Relative amount of
CO2(g) produced
(size of balloon)
1
2
3
Conclusions:
- The amount of NaHCO3 (s) _______________________, so it ___________________
the amount of gas produced
-
The amount of HCl (aq) was ______________, so it ______________ the amount of gas
produced.
What does this mean?
- the HCl (aq) ___________________, therefore it must ___________________________.
-
The NaHCO3 (s) must have been in ______________, meaning there must be some of it
______________ after the reaction has stopped.
CONSIDER:
4NH3 (aq) + 5O2 (g)  6H2O (l) + 4NO (g)
How many moles of NO (g) are produced when ___ moles of NH3 (aq)
and ___ moles of O2 (g) react?
(a) 4 moles NH3 (aq) & 5 moles O2 (g)
(b) 4 moles NH3 (aq) & 20 moles O2 (g)
(c) 8 moles NH3 (aq) & 20 moles O2 (g)
(d) 4 moles NH3 (aq) & 2.5 moles O2 (g)
We always use the quantity that _____________________________ in order to figure out
_______________________________.
If we used the __________________________________ to determine the amount of
product that we could make, it would be _____________.
EXAMPLE 1: Determine the number of grams of NO (g) produced when 1.75 mol
of NH3(aq) reacts with 0.975 mol of O2(g)
4NH3 (aq) + 5O2 (g)  6H2O (l) + 4NO (g)
EXAMPLE 2: Calculate the mass of aluminum chloride that can be produced from
20.0 g of aluminum and 30.0g of chlorine gas.
HW: pg 254 -258 #24, 25, 28ab, 30a; pg 259 #3 (answers on pg 273)
UNIT 3 – REVIEW QUESTIONS
(Please note that the purpose of this sheet is to give you some practice questions to help you
prepare for the test. It does not have every type of question on it. The complete list of what
you are responsible for knowing for the Unit 3 test is on the Unit 3 Expectations sheet.)
1) Calculate the number of moles of NO2 that are present in a sample that is made up of
4.74 x 1022 molecules.
2) In 12.5g of Mg3(PO4)2, a) calculate how many formula units there are.
b) calculate how many atoms there are.
c) calculate how many P atoms there are.
3) Calculate the percentage composition of:
a) Na2CO3
4) Calculate the empirical formula of a compound that is:
b) aluminum oxide.
a) 30.4% N and 69.6% 0
b) 43.6% P and 56.4% 0
5) Calculate the molecular formula of a compound that is 39.97%. C, 6.73% H, and 53.30% 0
It has a molar mass of 60.0 g/mol.
6) A certain compound has the following percent composition: 57.1% C, 4.8% H, and 38.l% O.
If the molar mass of this compound is 126 g/mol, what is the molecular formula?
7) (a) How many grams of lead(II) chloride should be produced when 6.7g of lead(II) nitrate
solution react with hydrochloric acid to form nitric acid and lead(II) chloride solution?
(b) If a reaction was conducted and a student produced 3.5 g of lead (II) nitrate from the
above reaction, calculate the percentage yield of the reaction.
8) A reaction involved in the production of iron from iron ore is:
Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g)
Calculate how many grams of CO must react to produce 3.5 kg of Fe.
9) Calculate the mass of sulphuric acid that can be prepared from 50.0g sulphur dioxide,
15.0g oxygen gas and an unlimited amount of water.
10) A hydrate MgSO4.xH2O was heated to drive off the water. The following experimental data
was collected:
Mass of crucible = 12.59 g
Mass of crucible and hydrate = 20.62 g
Mass of crucible and residue = 16.52 g
a) Calculate the percent of water in the hydrate
b) Calculate the formula of the hydrate.
Answers:
1) 0.0787 mol
2) a) 2.86 x 1022 formula units b) 3.72 x 1023 atoms c) 5.73 x 1022 P atoms
3) a) 43.38% Na, 11.33% C, 45.29% O b) 52.92% Al, 47.08% O
4) a) NO2 b) P2O5
5) C2H4O2
6) C6H6O3
7) (a) 5.6g (b) 63%
8) 2.6 x 103g
9) 76.5g
10) a) 51.1% b) MgSO4.7H2O