Download honors chemistry harvard-westlake first semester final exam review

HONORS CHEMISTRY
HARVARD-WESTLAKE
FIRST SEMESTER
FINAL EXAM REVIEW
FIRST SEMESTER REVIEW 1 FIRST SEMESTER REVIEW 2 The Dreaded Nine-bottle Problem
Almost all salts [a salt is an ionic compound] exist in water solutions as separated ions. When solutions of different salts are
mixed, the ionic species present may remain in solution as separate entities [i.e., no reaction occurs], may combine to form a
precipitate or may react chemically to produce gases or precipitates which slowly change into other compounds upon standing in
water.
In this experiment you will be provided with nine different salt solutions, each numbered but otherwise unidentified. All of
the solutions are colorless. They are:
NaBr
Na2SO4
NH4Cl
NaOH
KClO3
BaCl2
Na2S
KIO3
Pb(NO3)2
The object of the experiment is to identify each solution, using either the interactions occurring (if any) when small quantities of the
solutions are mixed with one another, or any physical evidence.
To carry out the experiment successfully, you must obviously know what interactions are expected for each mixture.
Intelligent preliminary study is therefore essential, and as part of your pre-laboratory preparation you will determine the expected
behavior when pairs of solutions are combined.
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Solubility of salts
A salt is classified as soluble (no precipitate) if more than 1 g will dissolve in 100 mL of water. It is insoluble (forms a precipitate)
if no more than 0.1 g will dissolve. Between these is the category called "sparingly soluble" which may or may not give a visible
precipitate under some conditions. For the problem at hand we note the following:
[This is NOT a complete list; these are only addenda to the rules you have been studying--you will need to use all of this information.]
All chlorates are soluble
All iodates are soluble except Ba(IO3)2 and Pb(IO3)2
Ba(OH)2 is more or less soluble but it often appears to precipitate due to the presence of dissolved CO2 in the hydroxide
solution (BaCO3 is very insoluble)
Other chemical reactions
1. Two gases, recognizable by their odors, may be produced on mixing some of the solutions:
a. ammonia, NH3, is produced when ammonium salts react with bases:
NH4Cl + NaOH  NH3 + NaCl + H2O
b. hydrogen sulfide, H2S, is generated by reaction of sulfide salts with ammonium salts:
2 NH4Cl + Na2S  H2S + 2 NH3 + 2 NaCl
[in this case, the odor of H2S is generally more noticeable than that of NH3]
2. While Pb(OH)2 is insoluble, addition of excess hydroxide ions will cause the precipitate to dissolve! This can give some confusing
results. Remember, precipitation is generally instantaneous while dissolution is much slower.
FIRST SEMESTER REVIEW 3 3. Compounds of the iodate ion may oxidize compounds of the sulfide ion to form elemental sulfur, a fine white precipitate, and
iodide ion:
IO3- + S2-  I- + S
(OH-)
[not balanced]
In terms of what you will be able to see, there is no difference between elemental sulfur and any other white precipitate.
Preparing to experiment
AS PART OF YOUR PRE-LAB: Prepare a grid such as the example shown in the technique section and fill it in with P if the
reaction represented by a given box is expected to form a precipitate (including sulfur), H2S if the reaction is expected to produce the
gas, NH3 if the reaction is expected to produce ammonia gas, or leave it blank if the mixture is not expected to react. You will have a
similar grid to work on in the lab, using a plastic sheet and one drop of each solution.
Technique
On the shelf above where you work you will find nine bottles. Write down the SET NUMBER which appears on each
bottle.
BE SURE TO OBSERVE AN ODOR EACH TIME YOU MIX TWO SOLUTIONS SINCE YOU WILL NOT BE
ABLE TO TELL WHERE THE ODOR IS COMING FROM IF YOU HAVE MORE THAN ONE MIXTURE ON THE
PLASTIC SHEET! Once you have identified one odor, it is a good idea to remove the drops so that they will not interfere with later
observations. An easy way to do this is to use a small amount of paper towel to blot up the mixture.
The entire plastic sheet should be rinsed thoroughly at the conclusion of the experiment and dried completely.
[the record for this experiment is 24 drops used in a total of 6 minutes with the correct solution given at completion of mixing!]
The grids on the attached page should be used to make your predictions and record your observations in the lab. They can be part of
your pre-lab.
The chemicals
Sodium bromide is a white, crystalline solid, freely soluble in water, with a saline, feebly bitter taste. It is used in photography. It has
medical and veterinary applications as a sedative, hypnotic, and anticonvulsant.
Sodium sulfate is used for standardizing dyes, in printing textiles, and in the manufacturing of glass and paper pulp. It is used
medically as a diuretic.
Potassium chlorate is a very reactive oxidizing agent. In solid form it explodes with sulfuric acid and inflames with explosion when
exposed to organic materials. It is used in fireworks, matches, dyeing and as a source of oxygen in chemical reactions.
It is irritating to the gastrointestinal tract. A toxic dose is approximately 5 grams.
Sodium sulfide is extremely hygroscopic, discoloring on exposure to air. It is unstable and may explode upon percussion or rapid
heating. It has the odor of dihydrogen sulfide gas and its solutions are strongly basic.
Analysis
Your goal is to determine the identity of each solution. Once you have done this, write a balanced net-ionic equation for each
reaction which you observed in the lab. You should have one equation for each reaction marked on the grid [duplicates, if any, need
not be written].
Summarize by giving the identity of each solution by number and formula.
FIRST SEMESTER REVIEW 4 NaOH
Na2S
Na2SO4
KClO3
KIO3
NH4Cl
BaCl2
Pb(NO3)2
NaBr
P
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d
i
c
t
e
d
NaOH
Na2S
Na2SO4
KClO3
R
e
s
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s
KIO3
NH4Cl
BaCl2
8
7
6
5
4
3
2
1
9
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8
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2
FIRST SEMESTER REVIEW 5 FIRST SEMESTER REVIEW 6 First Semester Review
The best preparation for the final exam is consistent study habits and effort throughout the term. We
know that a lot of material has gone down the lab sink, so to speak, and so we also have some suggestions for
areas you should concentrate on as you prepare. Pay close attention to the listing that follows. The items are
given as things you should be able to do if you intend to do well on the exam.
A brief description of the final exam format would be "business as usual, but more of it". The exam
consists of 25 multiple choice questions, 4 major required problems, your choice of 4 out of 5 minor
problems, your choice of 3 out of 5 reactions to complete, balance and identify, and your choice of 1 out of 6
essays. The essays are somewhat more involved than usual. You will be provided with a periodic table, a
complete set of solubility rules, the activity series, values for the gas constant (R), Planck's constant (h), the
speed of light (c), standard pressure values in different units, and an electronegativity table. You will be using
your own calculator for the test. You MAY NOT have additional items stored in your calculator. Doing so is a
serious violation of the school Honor Code.
The final exam is worth 17% of your first semester grade.
Following the listing of topics, you will find a "sample" exam. Most of the questions have been drawn
from earlier work. The "exam" is intended for practice only and should not be used to gauge the specific content
of the real thing. On the other hand, no attempt has been made to select problem areas which will not appear on
the exam.
The Topics
Unit 1
-you should be able to name binary and ternary compounds
-you should be able to write chemical formulas from the names of binary
and ternary compounds
-you should be able to calculate molar mass (g/mol) (Periodic table provided)
-you should be able to change g to mol and mol to g
-you should be able to do simple % composition calculations
-you should be able to determine the empirical formula of a substance from % composition data
-you should be able to distinguish between an empirical and molecular formula for a substance
Unit 2
-you should be able to complete and balance chemical reactions
-you should be able to identify the type of reaction (precipitation, redox, acid/base)
Unit 3
-you should be able to do mass-mass stoichiometry (including limiting reagents)
Unit 4
-you should be able to convert freely among energy, frequency and wavelength for electron transitions
(h and c given)
-you should be able to describe early atomic models AND any experiments or devices that were used to
formulate them
Specifically: Rutherford and Bohr
-you should be able to explain the generation of lines in the emission spectrum of an atom
-you should be able to write correct regular electron configurations for elements up to Z = 71
-you should be able to compare major periodic properties for elements (size, ionization energy, etc.)
FIRST SEMESTER REVIEW 7 Unit 5
-you should be able to distinguish between the major types of bonding: ionic and covalent
-you should be able to draw/identify correct Lewis structures for atoms, simple molecules and ions
-you should be able to relate bond energy information to ∆Horxn
-you should be able to use Hess' Law, either by adding chemical reactions or by using ∆Hof data, to
determine the enthalpy change for a reaction (∆Hof data provided)
-you should be able to identify exothermic and endothermic reactions based on data given
-you should be able to do stoichiometric calculations involving enthalpy changes
-given enthalpy and entropy data, you should be able to determine which reactions in a group are more
likely to be spontaneous
-you should be able to determine the free energy change for a process given enthalpy and entropy data
Unit 6
-you should be able to identify and draw hybrid orbital diagrams for simple molecules
-you should be able to apply VSEPR to draw the 3D shape of a molecule or ion
-you should be able to determine the existence of a net dipole on a molecule or ion (electronegativity
table provided)
-you should be able to distinguish among the types of weak forces and determine which is most
significant in a given atom, molecule or ion
-you should be able to make qualitative judgments about the vapor pressure of a liquid
Unit 7
-you should be able to interpret simple phase diagrams
-you should be able to solve simple partial pressure problems
-you should be able to apply Avogadro's Principle to predict the pressure/volume behavior of a gas
system
-you should be able to describe those factors which contribute to non-ideality of a gas
-you should be able to apply the ideal gas law to solve problems, including stoichiometry
That's all there is!
_____________________________________________
First Semester Sample Final Exam--Some typical multiple choice
1. Consider this reaction:
H2 + Cl2  2 HCl
This reaction is an example of
a. displacement
b. decomposition
c. precipitation
d. synthesis or combination
e. none of the above
2. The correct formula for dinitrogen tetraoxide is
c. N2O4
d. N2O5
a. 2N2O
b. N4O2
FIRST SEMESTER REVIEW e. Ni2O4
8 3. The correct name for Fe3(PO4)2 is
a. iron phosphate
b. iron(II) phosphate
c. iron(III) phosphate
d. iron(II) phosphorus tetraoxide
e. iron(III) phosphorus tetraoxide
4. The empirical formula for a compound is known to be CH3 and the molar mass for the same compound is
about 60 g/mol. The molecular formula is:
a. CH3
b. C2H6
c. C3H9
d. C4H12
e. C5H15
5. The percent, by mass, of phosphorus in Ca3(PO4)2 is
a. 10%
b. 20%
c. 30%
d. 40%
e. 50%
6. As one moves down the periodic table, which of the following properties increases?
a. electron affinity
b. ionization energy
c. electronegativity
d. atomic radius
7. An atom of element X has the electron configuration 1s22s22p63s2 and element Y has the electron
configuration 1s22s22p63s23p5. The formula for a compound of these two elements is:
a. X2Y
b. XY
c. XY3
d. XY2
e. X12Y17
8. Which bond will be ionic? (table provided on exam)
a. Al-Si
b. C-H
c. Li-S
d. Ba-O
e. S-F
For the following question, compare the assertion (top) and the reason (bottom) and select the appropriate
answer from the options given below:
a. both are true and reason fits the assertion
b. both are true but reason does not explain assertion
c. only the assertion is true
d. only the reason is a true statement
e. both are false
9. The first ionization energy of Li is less than that of Na
BECAUSE
Li is a smaller atom than Na.
10. Select the geometry that best describes NH4+ :
a. linear
b. bent (v-shaped)
c. trigonal planar
d. trigonal pyramid
e. tetrahedron
FIRST SEMESTER REVIEW 9 11. Bonds in which electron pairs are shared between atoms are best characterized as
a. ionic
b. hydrogen
c. hybrid
d. covalent
12. A liquid will boil when
a. its vapor pressure is less than atmospheric pressure
b. its critical pressure exceeds atmospheric pressure
c. its critical temperature equals room temperature
d. its vapor pressure equals atmospheric pressure
e. its temperature reaches the triple point
13. Given these two reaction heats:
∆H = -297 kJ
S(s) + O2(g)  SO2(g)
3
S(s) + 2 O2(g)  SO3(g) ∆H = -438 kJ
What is the heat of reaction for
2 SO2(g) + O2(g)  2 SO3(g)
under the same conditions?
a. -282 kJ
b. -297 kJ c. -735 kJ
d. -890 kJ
e. -1470 kJ
14. Consider the reaction below:
1.7 kJ + NaCl(s) + H2O(ℓ)  Na+(aq) + Cl-(aq)
Which of the following combinations for ∆H and ∆S is correct?
a. +,b. -,+
c. -,d. +,+
15. Metals generally have
a. high ionization energies
b. few electrons in the outermost energy level
c. high electron affinities
d. a tendency to gain electrons
16. In which of the following pairs of particles is the radius of the second particle larger than the first?
a. Mg, Ca b. Cl-, Cl c. K, Ca d. K, K+
17. The volume occupied by 0.50 mole of propane gas, at a temperature of 27oC and a pressure of 202.6 kPa is
best expressed by which of the following? (values of R provided)
a. (0.50 x 8.31 x 27)/202.6
b. (0.50 x 8.31 x 300)/202.6
c. (0.50 x 8.31 x 273)/300
d. (0.50 x 8.31 x 300)/(202.6 x 760)
e. (0.50 x 8.31 x 27)/(202.6 x 760)
FIRST SEMESTER REVIEW 10 18.
HCl(g) + NH3(g)  NH4Cl(s)
If 3.0 moles of HCl gas and 5.0 moles of NH3 gas, each measured at 20oC and 101.3 kPa pressure, are allowed
to react according to the equation above, the final mixture will contain
a. 3 moles of solid NH4Cl only
b. 5 moles of solid NH4Cl only
c. 3 moles of solid NH4Cl + 2 moles of NH3 gas
d. 3 moles of solid NH4Cl + 2 moles of HCl gas
e. 2 moles of HCl gas, 4 moles of NH3 gas and 1 mole of solid NH4Cl
19. On a phase diagram, the curved line which defines the vapor pressure of the liquid phase is the line along
which the following phase change occurs:
a. freezing
b. melting
c. boiling
d. sublimation
e. none of the above
20. The most reactive metal of the alkali metal family is probably
a. sodium b. potassium c. cesium d. francium
21. The bond between B and N is best described as
a. ionic b. polar covalent c. non-polar covalent
22. The hybrid orbitals of the carbon atom in methane, CH4, are called
a. sp b. sp2 c. sp3 d. sp4
23. Of the following molecules, all of which have polar bonds, the non-polar molecule is
a. HCl b. H2O c. NH3 d. CCl4
24. The bonds within polyatomic ions like SO42- are generally
a. ionic b. covalent c. non-polar d. metallic
25. Weak forces involving interactions between only non-polar species are called
a. hydrogen bonds
b. dipole-dipole forces
c. dispersion forces
d. non-polar weak forces
FIRST SEMESTER REVIEW 11 Some typical required problems
26. A compound A, upon analysis gave 38.67% K, 13.85% N, and 47.48% O by mass. When A is heated, a
new compound, B, is formed (and oxygen gas is given off). B has the composition 45.85% K, 16.47% N and
37.66% O by mass. Find the empirical formulas of these two compounds and then write a balanced equation
for the reaction that occurs when A is heated.
27. When hydrochloric acid and zinc metal react, the gas produced is collected in an evacuated flask. After the
system has returned to room temperature (27.2oC), the pressure inside the flask is 350 mm Hg. The volume of
the flask is 255 mL.
a. How many moles of gas have been collected?
b. Write the chemical reaction.
c. Assuming that the zinc has been completely consumed, how many grams of zinc were used in this reaction?
28. For the burning of 1 mole of C2H4 to form carbon dioxide and liquid water
a. write the balanced equation
b. calculate ∆Horxn at 298 K. [you will need values from the reference sheet in unit 5 to do this--tables provided
on exam]
c. how many kJ of heat energy will be absorbed/released if 5.7 g of C2H4 is burned?
d. using values from the unit 5 reference page, calculate ∆Sorxn (J/mol K) for this reaction
e. use your values of ∆Horxn and ∆Sorxn to calculate ∆Gorxn; is this reaction spontaneous at room temperature?
29. Draw the Lewis structures, expected geometries, give the VSEPR classification, and indicate which neutral
species would be polar:
a. CHCl3 (carbon is cental)
b. IF5 c. NF3
d. NO3-
Some typical problems to choose from
30. If the quantum of energy emitted in an electron transition between two levels is 2.96 x 10-20 J, what is the
frequency of light emitted as a result of the transition? What is the wavelength?
31. Consider an atom of Chlorine. How many neutrons are in an atom of this element with A = 37? How many
electrons are in this atom? Give the electron configuration and orbital diagram.
32. What are the strongest attractive forces that must be overcome to:
a. melt Benzene, C6H6 (symmetrical ring of 6 carbons)
b. dissolve CH3OH in H2O
c. melt PCl3
d. boil SiH4
FIRST SEMESTER REVIEW 12 33. Draw the complete orbital diagrams for normal and hybridized aluminum in the compound AlCl3. What
type of hybrid is this? What geometry would you expect for the compound AlCl3?
34. The following information concerning four consecutive--by atomic number-- (not necessarily in the same
row) elements in the periodic table was collected as a result of laboratory experiments. The letters have been
assigned arbitrarily:
Element A reacts vigorously with water producing hydrogen gas and AOH
Element B is very unreactive and is monoatomic and gaseous at room temperature
Element C exists as a diatomic gas at room temperature; reacts with A to form AC
Element D reacts with C to form a compound DC2. It reacts slowly with water to form hydrogen gas and
D(OH)2.
a. Assuming that these elements are not in the same row, but are consecutive, put them
in order from lowest to highest atomic number.
b. Which of these elements is a halogen?
c. Which of these elements is an alkaline earth metal?
35. Using bond energy tables from unit 5 (tables provided on exam), determine ∆Hrxn for:
2 H2(g) + O2(g)  2 H2O(g)
Some typical reactions to choose from
[in each case you would be asked to complete the word equation, write a balanced chemical reaction (net-ionic
if appropriate) and identify the type of reaction as precipitation, redox, etc.]
[all reactions occur in aqueous solution]
36. silver nitrate + dihydrogen sulfide gas 
37. hydrochloric acid + sodium carbonate 
38. potassium + water 
39. nitrate ions react with zinc metal in an acidic solution; zinc ions and ammonium ions are among the
products
40. nitric acid + sodium hydroxide 
Some typical essay TOPICS to choose from
41. early atomic models
42. emission spectra and atomic structure
43. periodic properties and shielding
44. intermolecular forces
45. ideality of gases
46. enthalpy, entropy and spontaneity
[remember, prime candidates for essay topics include those things which we have spent time describing in some
detail and which do not lend themselves to problem-solving; diagrams may be very useful in some cases]
FIRST SEMESTER REVIEW 13 FIRST SEMESTER REVIEW 14