Advanced Chemistry: 2012 – 2013
Summer Assignments
June 1, 2012
Dear ’12 -‘13 Advanced Chemistry Students,
I have put together the enclosed short syllabus and materials to hone your chemistry skills over
the summer. It is comprised of a reading assignment and questions through chapter 11. Below I
have provided a link to the Brown, Lemay, & Bursten’s A Central Science in case you would like
to use it as a reference to answer the questions I have assigned. We estimate that a thoughtful
effort would require about six hours.
These summer assignments are required. The deadline for mailing me your completed work is
August 20. You may send it by regular mail or e-mail scanned copies. Please keep the posted
materials – ONLY SEND YOUR WRITTEN WORK.
We thank you in advance for your efforts in this regard and look forward to seeing you in
Advanced Chemistry in the fall.
Best wishes,
Dr. Johnson
Syllabus
Chapter 1. Introduction
Assign.
Reading and videos
Questions
I.
Text available at the library website
Ch 1 through 11 (below)
http://standrews-de.libguides.com/content.php?pid=260281&sid=2261619
II.
Feynman 'Fun to Imagine': Jiggling Atoms
http://www.youtube.com/watch?v=v3pYRn5j7oI
Feynman 'Fun to Imagine': Fire
http://www.youtube.com/watch?v=ITpDrdtGAmo
III.
Feynman Lectures, Chapter 1
I. Advanced
Chemistry Summer Review
Instructions:
1.
Write your solutions on separate paper and start each chapter problem set with a new sheet.
2.
Number the problems clearly and skip a two lines between problems.
3.
Use complete sentences when verbal responses are called for. Sentences should contain the
question in the context of the answer
Example: Question: What color is the house? Answer: The color of the house is blue
4.
Show your calculations. Include the units of with every measured quantity. In a formula, bracket
every quantity and its unit with parentheses – Report answers with the correct number of
significant digits - Use scientific notation to unambiguously indicate the number of significant
digits.
Example: Calculate the density of a cube having a mass 32.234 g and side length 2.9 cm.
d =
5.
( 32.234 g)
( 4.9 cm) 3
= 0.273984... = 2.7 x 10−1 g /cm 3
When answering a question by dimensional analysis, start by writing the unit of the desired
answer to the left of the equal sign and use parentheses around every term.
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Example: If a car gets 25 miles to the gallon, and a gallon of gas costs $2.85, and it is moving
along a highway at 55 miles per hour, what is the cost in gas in cents per minute that it burns?
⎛ 55 miles ⎞⎛ 1 hours ⎞⎛ 1 gal ⎞⎛ 285 cents ⎞
cents
1 cents
= ⎜
⎟⎜
⎟⎜
⎟⎜
⎟ = 1.0 x 10
⎝ 1 hour ⎠⎝ 60 min ⎠⎝ 25 miles ⎠⎝ 1 gal ⎠
min
min
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Chapter 1.
1.1 The molar mass of oxygen is 32.0 g/mole. Its density at standard temperature and pressure is
22.4 moles per liter. Use dimensional analysis to convert this density into grams per cubic inch
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given that 1 Liter = 1000 ml, 1 ml = 1 cm , and 2.54 cm = 1 inch
1.2 The figure to the right shows a graduated cylinder filled
with a colored solution. The numbers refer to milliliters.
What is the volume of this sample? Report the
measurement with the proper number of significant digits.
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1.3 A 32.65-g sample of a solid metal is placed in a flask.
Liquid toluene, a substance in which the solid is insoluble, is
added to the flask until the total volume of solid and liquid
together is 50.00 ml. The solid and toluene together weigh
58.58 g. The density of toluene at the temperature of the
experiment is 0.864 g/ml. Calculate the density of the solid.
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Chapter 2
2.1. The mass ratio of nitrogen to hydrogen in ammonia is 4.67. Taking the atomic mass of hydrogen
to be 1 atomic mass unit (amu), determine the atomic mass of nitrogen if the molecular formula
of ammonia is assumed to be (a) NH, (b) NH2, and (c) NH3.
2.2. A common isotope of chromium would be identified verbally as “chromium-51”
(a) How many protons are in this isotope?
(b) How many neutrons?
(c) How many electrons would be in a +3 ion of this isotope.
(d) What is this isotopes atomic number?
(e) What is this isotopes mass number?
(f) Write the proper symbol for this isotope.
2.4 Only two isotopes of copper occur naturally. One is copper-63 and it has an atomic mass of
62.9296 amu and has a relative abundance of 69.17%. The other is copper 65 and it has an
atomic mass of 64.9278 amu and a relative abundance of 30.83%. Calculate the average atomic
mass of copper.
2.5
Write formulas for these compounds and acids. Note the names that need Roman numerals.
(a) tin (IV) bromide
sulfate
(d) copper (II) carbonate
(g) dinitrogen tetroxide
(b) ammonium dichromate
(c) sodium hydrogen
(e) carbonic acid
(f) phosphoric acid
(h) triphosphorous hexachloride
2.6 Write the names for the following compounds and acids. Note that acids are indicated by (aq)
following the symbol
(a) H 2C 2O4 (aq)
(d) CCl 4
(g) Fe( NO3) 3
(b) HC 2 H 3O2 (aq)
(e) Ni 3 ( PO4 ) 2
(c) Cl 2O8
(f) NaC 2 H 3O2
(h) Fe( NO3) 2
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Chapter 3
3.1 The characteristic odor of pineapple is due to ethyl butyrate, a compound containing carbon,
hydrogen, and oxygen. Combustion of 2.78 g of ethyl butyrate produces 6.32 g of CO2 and 2.58
g of H 2O . What is the empirical formula of the compound? (b) What is its molecular formula if
its molar mass is found by mass spectroscopy to be 174.2?
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3.2 Octane, C 8 H 18 , is a major component of gasoline.
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(a) Write the chemical reaction equation for the complete combustion of octane.
(b) How many grams of O2 are needed to burn 10.0 grams of octane?
(c) How many pounds of carbon dioxide are “dumped” into the atmosphere for every pound of
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octane burned?
€ when an Alka Seltzer® tablet is dissolved in water is due to the reaction
3.3 The fizz produced
between sodium hydrogen carbonate (or sodium bicarbonate) and citric acid. The unbalanced
reaction equation is:
NaHCO3 + H 3C 6 H 5O7 (aq) → CO2 (g) + H 2O(l) + Na 3C 6 H 5O7 (aq)
In a certain experiment 1.00 g of sodium bicarbonate and 1.00 g of citric acid are allowed to
react.
(a) Balance the above reaction (Hint: C 6 H 5O7 3− can be treated like a unit),
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(b) What is the limiting reagent?
(c) How many grams of CO2 (g) form?
(d) How many grams of excess reagent remain after the limiting reagent is completely
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consumed?
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(Molar masses:
NaHCO3: 84.01 g/mol, H 3C 6 H 5O7 :192.12 g/mol. CO2 : 44.01 g/mol.
Chapter 4
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4.1 How many mL of 1.5 M Ca( NO3 ) 2 contain 10.0 grams of Ca( NO3 ) 2 ?
4.2 Pure glacial acetic acid is a liquid with a density of 1.049 g/mL at 25˚C. Calculate the molarity
of a solution made by adding 20.0 mL of glacial acetic acid to enough water to make 100.0 mL of
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solution
4.3 Use the solubility rules to predict whether or not a precipitate will form when the following pairs
of salt solutions are mixed. If a precipitate forms, write the net ionic equations for the reaction.
If no precipitate forms, write the left side of molecular equation with an arrow pointing to the
words “no reaction.”
(a) Sodium carbonate and silver nitrate
(b) Sodium nitrate and nickel (II) sulfate
(c) Iron (II) sulfate and lead (II) nitrate
(d) Nickel (II) nitrate and sodium hydroxide
(e) Sodium hydroxide and potassium sulfate
4.4 What volume of 0.115 M hydrochloric acid is needed to completely neutralize 50.00 mL of 0.0875
M sodium hydroxide)?
4.5 By titration, 15.00 mL of 0.1008 M sodium hydroxide is needed to neutralize a 0.2053-g sample of
a dry organic acid. (a) What is the molar mass of the acid if it is monoprotic? (b) Combustion
analysis of the acid indicates that it is composed of 5.89% H, 70.6% C, and 23.5% O by mass.
What is its molecular formula?
Chapter 5
5.1 The specific heat of iron is 0.450 J/g˚C. How many J of thermal energy is necessary to raise the
temperature of a 1.05-kg block of iron form 25.0˚C to 88.5˚C?
5.2 A 1.800-g sample of liquid phenol (C 6 H 5OH (l)) was combusted in pure excess oxygen in a bomb
calorimeter whose total heat capacity is 11.66 kJ/˚C. The temperature of the calorimeter plus
contents increases from 21.36˚C to 26.37˚C.
(a) Write a balanced equation
for this reaction. Assume that the water produced is liquid.
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(b) What is ΔEint per mole of phenol combusted? The molar mass of phenol is 94.11 g/mol
5.3 Consider the following enthalpies of reaction
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H 2 (g) + F2 (g) → 2HF (g) ΔH = − 537 kJ
C (s) + 2F2 (g) → CF4 (g)
ΔH = − 680 kJ
2C (s) + 2H 2 (g) → C 2 H 4 (g) ΔH = + 52.3 kJ
Use these to calculate the enthalpy change of the reaction
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Chapter 6
C 2 H 4 (g) + 6F2 (g) → 2CF4 (g) + 4HF (g)
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6.1 Identify the element that corresponds to each of the following electron configurations:
(a) 1s 2 2s 2 2 p 6 3s 2 ,
(b) [ Ne] 3s 2 3p1 ,
(c)
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(d)
[ Ar]
[ Kr]
4s1 3d 5 , (note this is an exceptional electron configuration)
5s 2 4d 10 5p 4
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€6.2 An isoelectronic series is a collection of an atom and ions that have the same number of electrons
and, therefore, the same electronic configuration. For example, the isolectronic series
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P 3− , S 2− , Cl − , Ar, K + , Ca 2+ all have the electronic configuration identical to Ar:
1s 2 2s 2 2 p 6 3s 2 3p 6 . (a) Write the isoelectronic series and electron configuration that
corresponds to the atom and ions that have 10 electrons, (b) Write the electron configuration for
this series.
6.3 What are the names for the following rules associated with the filling of orbitals by electrons?
(a) Two electrons can occupy the same orbital (i.e. have the same values of m, l, and ml ) only if
they have opposite spin.
(a) One electron enters each orbital of a sub-shell before “backfilling” with a second electron of
opposite spin.
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(b) Electrons fill orbital subshells in order of lowest energy to higher energy.
6.4 Write the electron configuration for an (a) Fe3+ ion, (b) copper atom (note this has an exceptional
electron configuration.)
Chapter 7
7.1 What is the approximate effective nuclear charge experienced by: (a) a 3s electron in calcium
atom, (b) a 3p electron in a phosphorous
7.2 In general terms, why do atomic radii tend to decrease as atomic number increases across a
period (row)?
7.3 Explain why 1st ionization energies tend to increase as atomic number increases across a period
(column).
7.4 Explain why aluminum has a smaller 1st ionization energy than magnesium despite the fact that
Al comes after Mg in the 3rd period
7.5 Explain why oxygen has a lower first ionization energy than nitrogen despite the fact that oxygen
comes after nitrogen in the 2nd period.
Chapter 8
8.1. Explain why the lattice energy of aluminum oxide is so much greater than the lattice energy of
sodium chloride.
8.2. Which groups of elements are most likely to form covalent bonds with each other?
8.3. Draw the Lewis structures for the following molecular compounds (i.e. compounds that form
molecules held together by covalent bonds rather than ionic bonds) (a) water, (b) carbon
dioxide, (c) methane, (d) hydrogen cynanide (HCN),
8.4. Draw the Lewis structures for the following molecules and molecular ions (a) sulfur trioxide, (b)
sulfite ion, (c) sulfate ion (be sure to use parentheses appropriately).
8.5. Using the bond enthalpies listed in Table 8.4 to estimate ∆H for the following gas phase reaction
Chapter 9
9.1 Draw the Lewis structures and state the electron-domain and molecular geometries of the following
molecules and ions
(a) HCN
2−
(b) SO3
(c) CF4
(d) SF4
(e) PF6−
(a) N 3−
9.2 Consider the Lewis structure for glycine, the simplest amino acid
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(a) What are the approximate angles between adjacent bonds of the carbon atom on the left?
(i.e. the one bonded to the N)
(b) What are the approximate angles between adjacent bonds of the carbon atom on the right?
(i.e. the one bonded to the two O.
(c) Is bond angle 1 (the H-N-H bond) greater or less that 109.5˚?
(d) What is the approximate bond angle of bond angle 2 (the C-O-H bond on the right side)
Chapter 10
10.1. What are the four major macroscopic physical behaviors or characteristics of gases?
10.2. A student inflates a balloon from an initial volume of 1.00 L of air to a final volume of 2.0L of
air using her breath. In the process the pressure increases from 1.1 atm to 1.3 atm and the
temperature increases from 23.0˚C to 29.3˚C. What is the percent increase in moles of air
added to the balloon?
10.3. (a) Calculate the number of molecules in a deep breath of air whose volume is 2.50 L at body
temperature, 37˚C. (b) How does this compare to the number of liters of air in the earth’s
entire atmosphere? Estimate this by taking the average height of the atmosphere to be 80 km,
the surface area of the earth to be 5.0x10 8 km 2 . Note that 1 km = 1000 m and 1 m 3 = 1000 L .
10.4. Hydrogen gas is produced when zinc reacts with sulfuric acid. If 159 mL of wet hydrogen gas
is collected over water at 24˚C
and 738 torr of total pressure, how many grams of Zn have been
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consumed? The partial pressure of water at 24˚C is 22.38 torr.
10.5. Calculate the pressure that carbon tetrachloride (CCl4) vapor will exert at 40˚C if 1.00 moles
occupies 28.0 L, assuming (a) CCl4 obeys the ideal gas law, (b) CCl4 obeys the van der Waals
equation (Values for the van der Waals constants are given in Table 10.3)
Chapter 11
11.1 Which type of intermolecular force . . .
(a)
(b)
(c)
(d)
(e)
(f)
Acts between all molecules and increases with molar mass?
Is responsible for the dissociation and hydration of sodium chloride (or, any soluble salt).
Is responsible for the high solubility of ammonia?
Is responsible for the solubility of sugar in water?
Is responsible for iodine being a solid at room temperature but still sublimating readily?
Are responsible for the attractive forces between hydrogen disulfide.
Chapter 11 (continued)
11.2. Use concepts of intermolecular forces to explain the following facts . . .
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(a)
(b)
(c)
(d)
CH 3OH boils at 65˚C and CH 3 SH boils at 6˚C
Xe is a liquid at 1 atm and 120K while Ar is a gas?
Cl 2 is a gas at 1 atm and 273K while Br2 is a liquid?
Acetone boils at 56˚C while 2-methyl propane boils at -12˚C (see Lewis structures below)
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11.3 Which member of the following pairs has the stronger intermolecular London dispersion
forces? Briefly explain your answer.
(a) CH 3CH 2 SH or CH 3CH 2CH 2 SH
(b) Br2 or O2
(c) CH 3CH 2CH 2Cl or (CH 3 ) 2 CHCl
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11.4 Refer to the phase diagram of carbon dioxide
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shown below to
(a) What are is the pressure and tenmperature of
the triple point? What does the triple point
signify?
(b) Does carbon dioxide have a normal boiling
point? Explain why or why not.
(c) At what approximate temperature does carbon
dioxide freeze when the pressure is 70 atm?
(d) Describe the phase changes, and the
temperatures at which they occur, when carbon
dioxide is heated from -80˚C to -20˚C at a
constant pressure of 3 atm.
11.5 Use the vapor pressure curve in the figure to the
right to…
(a) Estimate the boiling point of diethyl ether at
400 torr.
(b) Estimate the external pressure under which
ethyl alcohol will begin to boil at 70˚C.
(c) Estimate the temperature at which water
boils at the top of Mt. Everest where
atmospheric pressure is approximately 1/3
atm.
(d) Estimate the temperature at which water
boils inside a pressure cooker that reaches 1.5
atm.