Chemistry 1000
Fall 2005
First assignment
High School Review and The Atomic Theory
Responsible sections of Kotz, 6th edition: Chapter 1 and 2 in their entirety (this is almost entirely review of Chem 30
material).
Read the relevant sections of the chapter, including the Example Problems, and do the Exercises that are keyed to most of the
example problems. The material covered in lecture is the material that you will be directly tested on. However, the much
greater amount of material in the text forms the background required to master the problem solving that lies at the heart of
Chemistry 1000. This distinction should form the basis for the way you study this course. Remember that Chemistry 1000 is
by its very nature a problem solving course.
Self-Study Assignment
Study Questions from the end of Chapters 1 and 2
a) At the end of each chapter in the text are important tools. First is the Chapter goals revisited. This summarizes key
concepts that you should completely understand and know outright. Second are Key equations that you should also
memorize and know exactly how to use.
b) In the Study Questions section, the questions numbered in blue have answers provided at the back of the book and in
the student solutions manual. You should attempt all the blue problems that correspond to the sections of the chapter
you are responsible for (these are listed at the head of the problem set). Be sure that you can do at least a couple of
questions from each sub-topic correctly.
c) Advanced problems are indicated with red triangles. They make excellent review problems, and can also be used if you
think you know the stuff in a given chapter but wish to see if you know the material in depth. If you do not do these
problems now, they make excellent material for Test and Exam preparation.
Problem Set #1
The Problem Set consists of questions that I have selected and used over several years. An answer guide will be provided.
These include some fairly advanced problems as well as a selection of routine exercises. It is always good to get inside the
head of your instructor, and I hope these problems and the kinds of answers I provide will help you to prepare for the kinds
of test and exam questions you may expect later.
1.
Solve the equation below, and report the answer to the correct number of significant figures:
(0.0546)⎡⎢ (23.56 − 1.43 )⎤⎥
⎣ 1.345 × 10 ⎦
(Note that this is typical of a calculation that mixes addition/subtraction with multiplication/division, and thus is a
realistic example of calculations we will ask you to carry out later in this course.)
2.
The anesthetic procaine hydrochloride is often used to deaden pain during dental surgery. The compound is packaged
as a 10.% solution (by mass) in water. If your dentist injects 0.50 mL of the solution, how many milligrams of procaine
hydrochloride are injected?
3.
A red blood cell has a diameter of 7.5 μm (micrometer). What is this dimension in (a) meters, (b) nanometers, and (c)
Angstrom units (where 1 Å = 1 × 10–10 m)?
4.
Boron has two stable isotopes, 10B and 11B, with masses of 10.012937 u and 11.009305 u, respectively. Calculate the
percent abundances of these isotopes of boron. Note: you can easily look up the answer to this question, for example in
the Chemical Rubber Handbook! Therefore be sure that your answer shows enough work to demonstrate that you have
actually calculated the percent abundances.
5.
Demonstrate (using atomic number and average atomic mass data), for which pairs of elements, the original
Mendelevian order of the periodic table, in sequence of increasing atomic mass, does not work. Consider the elements
up to 109Mt.
6.
Several compounds containing only sulfur and fluorine are known. Three of them have the following compositions:
(i)
1.188 g of F for every 1.000 g of S
(ii)
2.375 g of F for every 1.000 g of S
(iii)
3.563 g of F for every 1.000g of S
How do these data illustrate the law of multiple proportions?
7.
If “Epsom salt,” MgSO4⋅xH2O, 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. How many molecules of water occur per formula unit of Epsom salt, i.e. what
is the value of x in the formula?
8.
Solve the equation below for V, and report the answer to the correct number of significant figures:
⎛ 234 ⎞
. + 215
.)
⎜
⎟ V = (0.000214)(0.0821)(27315
⎝ 760.0⎠
(Note that this is actually the calculation required to solve for the volume of a gas using the ideal gas law
approximation, and thus is a realistic example of calculations we will ask you to carry out later in this course.)
9.
The density of a solution of commercial perchloric acid (HClO4) is 1.67 g mL–1, and it is 70.5% acid by mass. What
volume of the reagent acid (in mL) do you need to supply 54.3 g of pure perchloric acid?
10. A dry black powder is placed in a clean, dry glass tube. Pure hydrogen gas [H2(g)] is passed down the tube in contact
with the powder and drives out all the air. With the H2(g) still flowing, the powder is heated. The powder turns red, and
water vapor can be detected coming out the end of the tube. Was the original black powder an element or a compound?
Explain your reasoning.
11. Copper has two stable isotopes, 63Cu and 65Cu, with masses of 62.939598 u and 64.927793 u, respectively. Calculate
the percent abundances of these isotopes of copper.
12. Consider an atom of 64Zn.
(a) Calculate the density of the nucleus in grams per cubic centimetre, knowing that the nuclear radius is 4.8 × 10–6 nm
and the mass of the 64Zn atom is 1.06 × 10–22 g. Recall that the volume of a sphere is (4⁄3)πr3.
(b) Calculate the density of the space occupied by the electrons in the zinc atom, given that the atomic radius is 0.125
nm and the electron mass is 9.11 × 10–28 g.
(c) Having calculated the densities above, what statement can you make about the relative densities of the parts of the
atom?
13. Uranium is used as a fuel, primarily in the form of uranium(IV) oxide, in nuclear power plants. The question considers
some uranium chemistry:
(a) A small sample of uranium metal (0.169 g) is heated to between 800 and 900 °C in air to give 0.199 g of a dark
green oxide, UxOy. How many moles of uranium metal were used? What is the empirical formula of the oxide,
UxOy? What is the name of the oxide? How many moles of UxOy must have been obtained?
(b) The naturally occurring isotopes of uranium are 234U, 235U and 238U. Which is the most abundant?
(c) If the hydrated compound UO2(NO3)2•zH2O is heated gently, the water of hydration is lost. If you have 0.865 g of
the hydrated compound and obtain 0.679 g of UO2(NO3)2 after heating, how many molecules of water of hydration
are there in each formula unit of the original compound? (The oxide UxOy is obtained if the hydrate is heated to
temperatures over 800 °C in the air.)
Study Hints
It is very important that you be able to work out these questions without sneaking looks at the solutions. This requires a
lot of self-discipline. I recommend that you work out written answers to the whole assignment before you look at the
answers! Use the information provided in the text to gather the information and the methods you think will give you a
solution.
Then take time to correct yourself against the solutions. It is not necessary to use the detailed SSM if you find you get
the right answers as given in the end of the text. If your answer disagrees, look back at your problem on your own to see if
you can spot the error. If all else fails, consult the SSM to see where you went wrong.
Make a mark against all the problems that you initially got wrong. These are problems you should definitely try doing
again when you study for the tests and final exam. Remember, studying University level courses is serious business, and you
should develop methods that help you to succeed!
I am very willing to help you with difficulties that you may encounter when attempting to solve problems. I ask that
you bring me your attempted solutions in written form so that I can help you with your problem-solving, rather than just
giving you my methods once again! Finally, I encourage you to make use of the chemistry Student Resource Centre. Upperyear and Graduate students staff the Centre. They are familiar with this Problem Set and are able to provide help. When you
use the Resource Centre, be sure to bring along your own partial or incorrect solutions to problems. That is the only way that
effective help can be provided to your own particular difficulties.
Additional Questions – High School review material
Chem1000 – Fall 2005
1. Calculate the molar mass of each of the following compounds:
(a) Fe2O3, iron(III) oxide
(b) BF3, boron trifluoride
(c) N2O, dinitrogen monoxide (laughing gas)
(d) MnCl2 .4H2O, manganese(II) chloride tetrahydrate
(e) C6H8O6, ascorbic acid or vitamin C
2. How many moles are represented by 1.00 g of each of the following compounds?
(a) CH3OH, methanol
(b) Cl2CO, phosgene, a poisonous gas
(c) NH4NO3, ammonium nitrate
(d) MgSO4.7H2O, magnesium sulfate heptahydrate (epsom salt)
3. Acrylonitrile, C2H3CN, is used to make acrylic plastics. If you have 2.50 kg of acrylonitrile, how many moles of the compound are
present?
4. An Alka-Seltzer tablet contains 324 mg of aspirin (C9H8O4), 1904 mg of NaHCO3, and 1000. mg of citric acid (C6H8O7). (The last two
compounds react with each other to provide the ‘‘fizz,’’ bubbles of CO2, when the tablet is put into water.)
(a) Calculate the number of moles of each substance in the tablet.
(b) If you take one tablet, how many molecules of aspirin are you consuming?
5. Sulfur trioxide, SO3, is made in enormous quantities by combining oxygen and sulfur dioxide, SO2. The trioxide is not usually isolated
but is converted to sulfuric acid. If you have 1.00 kg of sulfur trioxide, how many moles does this represent? How many molecules? How
many sulfur atoms? How many oxygen atoms?
6. Calculate the mass percent of each element in the following compounds:
(a) PbS, lead(II) sulfide, galena
(b) C3H8, propane, a hydrocarbon fuel
(c) CoCl2 ? 6 H2O, a beautiful red compound.
(d) NH4NO3, ammonium nitrate, a fertilizer and an explosive.
7. Vinyl chloride, CH2CHCl, is the basis of many important plastics (PVC) and fibers.
(a) Calculate the molar mass.
(b) Calculate the mass percent of each element in the compound.
(c) How many grams of carbon are in 454 g of vinyl chloride?
8. The empirical formula of succinic acid is C2H3O2. Its molar mass is 118.1 g/mol. What is its molecular formula?
9. Acetylene is a colorless gas that is used as a fuel in welding torches, among other things. It is 92.26% C and 7.74% H. Its molar mass is
26.02 g/mol. Calculate the empirical and molecular formulas.
10. Nitrogen and oxygen form an extensive series of oxides with the general formula NxOy. One of them is a blue solid that comes apart,
reversibly, in the gas phase. It contains 36.84% N. What is the empirical formula of this oxide?
11. Mandelic acid is an organic acid composed of carbon (63.15%), hydrogen (5.30%), and oxygen (31.55%). Its molar mass is 152.14
g/mol. Determine the empirical and molecular formulas of the acid.
12. Cacodyl, a compound containing arsenic, was reported in 1842 by the German chemist Robert Wilhelm Bunsen. It has an almost
intolerable garlic-like odor. Its molar mass is 210 g/mol, and it is 22.88% C, 5.76% H, and 71.36% As. Determine its empirical and
molecular formulas.
13. Elemental sulfur (1.256 g) is combined with fluorine, F2, to give a compound with the formula SFx , a very stable, colorless colorless
gas. If you have isolated 5.722 g of SFx, what is the value of x ?
14. What mass of lead (II) sulfide (the mineral galena) is required to produce 2.00 kg of lead?
15. Elemental phosphorus is made by heating calcium phosphate with carbon and sand in an electric furnace. How many kilograms of
calcium phosphate must be used to produce 15.0 kg of phosphorus?
16. Metallic sodium and chlorine gas are obtained by decomposing sodium chloride with an electric current. What masses of sodium and
chlorine can be obtained from 2.00 metric tons of salt? (A metric ton is exactly 1000 kg.)
17. Write the molecular formula and calculate the molar mass for each of the molecules below. Which has the larger percentage of
nitrogen? Of carbon?
18. What is the mass, in grams, of one molecule of the naturally occurring acid, oxalic acid, C2H2O4?
19. Write the molecular formula and calculate the molar mass for the molecule shown here. What are the weight percentages of carbon,
hydrogen, and sulfur in the compound? If you have 10.0 g of the compound, how many moles does this represent? How many grams of
sulfur are in 10.0 g of the compound? Color code: carbon = gray; hydrogen = white; sulfur = yellow; and oxygen = red. (The molecule is
called dimethyl sulfoxide, and it is commonly used as a solvent.)
20. The compound shown here contains carbon, hydrogen and iron. Color code: carbon = gray; hydrogen = white; iron = red. (The
molecule is called ferrocene. It is representative of a large class of so-called organometallic compounds.)
(a) Write the molecular formula for the compound.
(b) How many grams of iron are in 0.150 g of the compound? How many iron
atoms? (c) Which element accounts for most of the mass of the molecule?
21. Azulene, a beautiful blue hydrocarbon, has 93.71% C and a molar mass of 128.16 g/mol. What are the empirical and molecular
formulas of azulene?
22. Direct reaction of iodine (I2) and chlorine (Cl2) produces an iodine chloride, IxCly , a bright yellow solid. If you completely used up
0.678 g of iodine, and produced 1.246 g of IxCly , what is the empirical formula of the compound? A later experiment showed the molar
mass of IxCly was 467 g/mol. What is the molecular formula of the compound?
23. Iron pyrite, often called ‘‘fool’s gold,’’ has the formula FeS2 (see photo). If you could convert 15.8 kg of iron pyrite to iron metal, how
many kilograms of the metal do you obtain?
24. Stibnite, Sb2S3, is a dark gray mineral from which antimony metal is obtained. If you have 1.00 kg of an ore that contains 10.6%
antimony, what mass of Sb2S3 (in grams) is in the ore?
25. Transition metals can combine with carbon monoxide (CO) to form compounds such as Fe(CO)5 and Co2(CO)8. Assume that you
combine 0.125 g of nickel with CO and isolate 0.364 g of Ni(CO)x . What is the value of x?
26. The mass of 2.50 mol of a compound with the formula ECl4, in which E is a nonmetallic element, is 385 g. What is the molar mass of
ECl4? What is the identity of E?
27. A mixture is made of two pure compounds: A (empirical formula = CH4O) and B (empirical formula = C2H6O). If an analysis of the
mixture indicates it is 49.9% carbon, how many grams of B are mixed with 1.00 g of A?
28. The elements A and Z combine to produce two different compounds A2Z3 and AZ2. If 0.15 mole of A2Z3 has a mass of 15.9 g, and 0.15
mole of AZ2 has a mass of 9.3 g, what are the atomic masses of A and Z?
29. The weight percent of oxygen in an oxide that has the formula MO2 is 15.2%. What is the molar mass of this compound? What element
or elements are possible for M?
30. A piece of nickel foil, 0.550 mm thick and 1.25 cm square, is allowed to react with fluorine, F2, to give a nickel fluoride.
(a) How many moles of nickel foil were used? (The density of nickel is 8.908 g/cm3.) (b) If you isolated 1.261 g of the nickel fluoride,
what is its formula? (c) What is its name?