Download CHEM 20 FINAL EXAM: STUDY HEADINGS Jan 2012

CHEM 20 FINAL EXAM:
STUDY HEADINGS Jan/June 2012
A.
Introduction to Chemistry (Ch. 1,2,3)
physical and chemical change; physical and chemical properties;
accuracy and precision in measurements; properties of mixtures, elements, compounds
calculations using significant digits and unit analysis method
B.
Atomic Structure and Theory (Ch 4,5)
early theories of atomic structure; properties of subatomic particles: proton, neutron, electron
atomic number and average atomic mass; isotope symbols;
Dalton, Thomson and Rutherford atomic models
The Bohr atomic model; energy levels; emission and absorption spectra
Quantum mechanics model: Schrödinger equation, Pauli exclusion principle; Heisenberg uncertainty
principle quantum numbers; orbital types and shapes; electron configurations; stability
C.
Nuclear Chemistry (Ch. 28)
types of nuclear radiation and radioactive decay: alpha, beta and gamma
balance nuclear equations: transmutations, fissions, fusions, decays,
decay series, half life equations and problems
Einstein’s equation and mass changes in nuclear reactions
uses of nuclear energy
D.
The Periodic Table/ Periodic Trends (Ch. 6,10,11)
classification of elements; metals, semi-metals, non-metals; liquids, gases, solids;
main groups 1 ,2, 13 -18: properties of; valence electron configurations
transition metals: properties of; Lewis diagrams for atoms
periodic trends: ionization energy, metallic character, atomic radius, electron affinity, electronegativity;
characteristic properties of the chemical families
predict type of ion formed, number of valence electrons, ionization energy, by position in periodic table
E.
Nomenclature/Chemical Formulae (Ch. 7)
names of elements, symbols of elements; ionic compounds, molecular compounds
ionic charge and oxidation states; naming ionic compounds (formula units)
predicting formula units from ion charges; naming molecular compounds (molecules);
formulas of hydrated salts
F.
Chemical Reactions and Stoichiometry (Ch. 8,9)
finding the molar mass of elements and compounds
using Avogadro’s number: converting moles to number of particles, and vice versa
finding moles of a substance from a given mass and vice versa
determining empirical formula and molecular formula from experimental data
percentage composition and formulas of hydrated crystals
writing balanced equations; using subscripts to indicate phases
energy changes in reactions: endothermic and exothermic reactions
naming types of chemical reactions: synthesis, decomposition, combustion, dissolving( ionizations and
dissociations), single displacements, double displacements
predicting the products of single displacements using the activity series for metals and halogens
predicting the products of double displacements using solubility tables
predicting the products of combustion reactions
mass stoichiometry questions: moles to moles, mass to mass, mass to moles, moles to mass
energy questions using stoichiometry
limiting reagent problems – theoretical yield, excess reagents
G.
Chemical Bonding (Ch.12, 13, 14.1,17.2)
use difference in electronegativity to identify types of bonds: covalent, polar covalent, ionic
types of bonds: ionic, molecular, polar covalent; bond dipole moments
comparison of molecular radii , covalent radii, atomic radii, vanderwaals radii
drawing Lewis dot structures for elements and molecules; expanded valence
predicting bond angles, molecular geometry based on lewis dot diagrams
drawing correct shapes for simple molecules
hybridization orbitals of the central atom: sp, sp2, sp3, sp3d and sp3d2
molecular bonds: sigma and pi bonds; delocalized pi bonds in benzene, C6H6
determining and indicating direction of dipole within a covalent bond
predicting the polarity of molecules from dipole moments and molecular geometry
intermolecular forces: van der waals, hydrogen bonding, dispersion forces, ionic forces,
metallic bonding, covalent network solids
I.
Physical Chemistry: Solutions and Gases (Ch 14, 20, 21, 15, 18, 19)
characteristics of the 3 phases of matter: solids, liquids, gases
characteristics of phase changes: fusion, vaporization, liquefaction
qualitative aspects of the kinetic theory: behavior of gases – pressure, volume
Avogadro’s hypothesis; molar volumes, STP and SATP, gas stoichiometry questions
qualitative properties of solutions: solvent, solute, solubility, effect of temperature and pressure on
solubilities of solids and gases in water;
Colligative properties of solutions: boiling point elevation, freezing point depression, vapor pressure
lowering and osmotic pressure of solutions
effects of molecular structure, intermolecular forces on solubility like dissolves like…principle,
hydrogen – bonding, polarity of molecules
determining the concentration of solutions using: Molarity (M, mol/L)
using c = n/v to solve for mass of solute, volume of solution, or Molarity of solution
preparation of solutions; diluting solutions; using the dilution formula C1V1 = C2V2
writing dissolving equations for ionic compounds, use mole ratios to determine molar concentrations of
ions in solution;
H
Organic Chemistry (Ch. 29)
naming organic compounds; draw isomers for given organic compounds;
identifying the basic hydrocarbons: alkanes, alkenes, alkynes, and cyclic hydrocarbons;
use prefixes to identify substituted groups; alkyl groups, halogens, and a few others
saturated and unsaturated hydrocarbons: general formulas for the main groups;
identify, name and draw structural formulas for aromatic compounds: benzene and its derivatives
identify selected functional groups by structure and name: halocarbons, alcohols,
carboxylic acids and esters
complete substitution, addition and esterification reactions
outline some of the use of hydrocarbons in modern society
FINAL EXAM REVIEW: 2012 January
The first part of this review are typical multiple-choice style questions. There is only one correct answer per
question:
1.
The theoretical value for the mass of one mole of a substance is 39.23 g. Irene obtained the
following experimental data: 32.24 g, 32.26 g, and 32.25 g. Her results could be described as:
a) precise b) accurate c) both precise and accurate d) neither precise nor accurate
2.
When Cesium metal is gently heated, it quickly melts. This would be an example of a(n):
a) physical change b) chemical change c) acid-base reaction d) precision
3.
Given the isotope symbol, Br – 81, which of the following statements is FALSE?
a) this isotope represents an atom of the element bromine
b) one atom of this isotope contains 81 protons
c) one atom of this isotope contains 46 neutrons
d) one atom of this isotope contains 35 electrons
4.
The electron configuration: 1s22s22p63s23p64s23d104p5 represents which element in its ground state?
a) Zinc
b) Bromine c) Chlorine d) Manganese
5.
Which of the following sets would represent isotopes of the same element?
a)
c)
28
A,
14
12
Q
7
29
B,
14
30
C
14
b)
16
R
8
20
S
11
d)
1
2
X,
3
Y,
Z
1
2
3
40
T
20
40
V
19
40
W
18
6.
Which of the following statements about subatomic particles is False?
a) A proton and a neutron have approximately the same mass.
b) A neutron has no electrical charge.
c) The mass of the nucleus is determined primarily by the protons and electrons present.
d) The proton and electron carry opposite electrical charges.
7.
When 2.00 moles of Uranium-234 undergoes alpha decay to form thorium-230, the mass converted to
energy is 0.001156 g. The energy released by this nuclear reaction is:
a) 5.20 x 1010 J b) 1.04 X 1011 J c) 5.20 X 1013 J d) 1.02 X 1014 J
8.
Tritium nuclei are a radioactive form of hydrogen. The isotope symbol for tritium is H – 3. How many
neutrons are present in one atom of tritium?
a) 0 b) 1
c) 2
d) 3
e) 4
9.
Which of the following elements would have the lowest electron affinity?
a) Magnesium
b) Sulfur
c) Silicon
d) Phosphorus
10.
Which of the following 4th row ions would have the smallest ionic radius?
a) V2+
b) K+
c) Sc3+
d) Se2e) Br-
11.
Considering the periodic table trends, which of the following statements is FALSE?
a) Sodium has a higher metallic character than Lithium.
b) Astatine is the halogen with the largest atomic radius.
c) Oxygen has a lower first ionization energy than sulfur.
d) Noble gases have very stable electron configurations.
12.
Which of the following group 13 ( 3A) elements would have the most metallic character?
a) Thallium b) Indium c) Gallium d) Aluminum e) Boron
13.
The most common ion formed by the nitrogen atom would have a charge of:
a) +5 b) +3 c) -5 d) -3
14.
Use the following characteristics to identify the unknown element:
Properties:
I) it is a colorless gas at room temperatures
II) it has a very low density
III) it is very explosive
IV it has a bonding capacity of one
a) Helium
b) Oxygen
c) Hydrogen d) Nitrogen
15.
Which of the following is the correct molecular formula for the compound, tetraphosphorus
heptasulfide?
a) K4S7
b) P4S7
c) P4S2
d) P4(SO4)7
16.
The correct formula unit for the hydrated salt, potassium carbonate dihydrate, would be:
a) K2CO3.2H2O
b) P2CO3.2H2O
c) K2CO3.2H
d) K2 CO3. 2OH
17.
What is the correct formula unit for Lead IV oxalate?
a) Pb4C2O4
b) Pb(C2O4)2
c) PbC2O4
d) Pb(C2O4)4
18.
Which of the following compounds has the smallest molar mass?
a) NH3
b) H2O2
c) CH4
d) SO3
19.
The molar mass of acetone, CH3COCH3, is:
a) 58.0 g/mol
b) 31.0 g/mol
c) 74.0 g/mol
d) none of these are correct
20.
A certain compound has a molecular mass of 42.0 g/mol. The simplest formula for the compound is
most likely:
a) NH3
b) CH4
c) C2H2O
d) BF
21.
The empirical formula of a compound which contains 75.0 g of carbon, 12.5 g of hydrogen and
100.0 g of oxygen is:
a) CH3O
b) C2H4O2
c) C2H3O2
d) CH2O
22.
In an experiment an aqueous solution of Ca(NO3)2 reacts with an aqueous solution of Na3PO4 in a
replacement reaction. Indicate the formula for the resulting precipitate:
a) Ca2(PO4)3
b) Ca3(PO4)2
c) NaNO3
d) Na2NO3
23.
Substances which, added to a reaction vessel, to make reactions happen at a faster rate are known as:
a) solutions b) hydrocarbons c) catalysts d) reagents
24.
Which of the following is an example of combustion?
a) sodium metal + chlorine gas produces sodium chloride solid
b) propane + oxygen produces carbon dioxide and water
c) sulfuric acid decomposes into hydrogen gas, oxygen gas and sulfur
d) solid lead metal reacts with a silver nitrate solution to produce lead II nitrate and silver metal
25.
Which of the following chemical equations is an example of a decomposition reaction?
a) 4 P4 + 5 S8 → 4 P4S10
b) 2 C2H6 + 7 O2 → 2 CO2 + 3 H2O
c) KClO3 → KCl + 3 O2
d) P4 + 5 O2 → P4S10
26.
Dinitrogen oxide can be prepared by the thermal decomposition of ammonium nitrate according to
the following equation: NH4NO3(s)
→ N2O(g) + 2 H2O(g)
When 1.000 g of ammonium nitrate is decomposed in this way, 0.550 g of dinitrogen oxide is
produced. What mass of water would also be produced in this reaction?
a) 1.10 g
b) 0.450 g
c) 0.900 g
d) 0.550 g
27.
Which of the following electron dot diagrams is not likely to exist in nature?
a) H – C ≡ C – H
28.
c) H – F
d) H – O – O – H
Which is the correct Lewis dot structure for AlF3 ?
a) F – Al – F
29.
b) He
b) F – Al – F
c)
F – Al – F
d)
F – Al – F
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|
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F
F
F
F
Which of the following is a network covalent solid?
a) iron
b) sodium chloride
c) graphite
d) ice
30.
In which of the following molecules is the central atom NOT expanding its valence?
a) SF6
b) PF5
c) AsBr5
d) PCl3
31.
Which one of the following bonds would be expected to have the highest ionic character?
a) CsF
b) MgBr
c) AlO
d) HI
32.
Which of the following sets of molecules, are all the shapes linear?
a) CO2 , HCN and N2 b) HC≡CH, XeF2 and O3 c) H2S, CO and CO2 d) H2O, Cl2 and NO2
33.
The molecular shape of the CCl3+ ion would best be described as:
a) linear
b) trigonal planar c) pyramidal d) square planar
34.
Which of the following bond pairs has the highest ionic character?
a) Li and O b) Se and F c) Sr and Br d) Ca and S
35.
Which of the following is not a property of metals?
a) electrical conductors in all phases
b) lustrous – reflects light
c) low melting points
d) delocalized electrons
36.
Which of the following processes would strengthen very soft metals?
a) add gases
b) temper the metal with high temperatures
c) add elements from the middle of the periodic table such as Carbon
d) allow the metals to oxidize
37.
Which of the following molecules would have an octahedral shape?
a) H2CO b) BF3 c) SeO2 d) SF6 e) S2Cl2 f) SF4
38.
Which of the following ions would have a tetrahedral shape?
a) NO3- b) SiF62- c) SO32- d) NH3+ e) PO43-
39.
Which of the following bonds is(are) polar covalent?
a) B – F
b) C – O
c) Si – F
d) N – H
e) O – H
40.
Which of the following molecules is most likely to contain a triple bond?
a) O3
b) HCN
c) CH4
d) SCl6
41..
Which of the following would have the highest melting point?
a) NaCl b) GeO2
c) H2
d) ZnCu
e) C2H5OH
42.
Which of the following compounds is ionic ?
a) CuSn(brass) b) NH3
c) H2O
d) KCl
43.
Which is true regarding intermolecular forces?
a) A hydrogen bond is a specialized type of dispersion force.
b) Hydrogen bonds can form with any atom in Groups 15, 16, or 17.
c) The only forces holding nonpolar compounds together are dispersion forces.
d) Substances with strong intermolecular forces tend to have very low boiling points.
44.
Which is NOT an assumption of matter made by the kinetic molecular theory of gases?
a) Gas particles are small and take up little volume relative to the volume of the space around them.
b) The kinetic energy of particles is determined by their mass and their velocity.
c) When particles collide, their total energy is decreased.
d) Gas particles are widely dispersed and therefore are not affected significantly by attractive or repulsive forces.
45.
A sample of four mixed gases is at 760 mmHg. The partial pressures of gases A and B are shown. If gases C and
D have equal partial pressures, what is the partial pressure of gas C?
a) 43.6 mmHg
b) 21.8 mmHg
46.
c) 165.7 mmHg
d) 331.4 mmHg
A helium-filled balloon has a volume of 50.0 L at 25°C and 1.08 atm. What volume will it have at .855 atm and
10.0°C?
a) 15.0 L
b) 30.0 L
c) 60.0 L
d) 120.0 L
47.
Hospital patients are administered oxygen from an pressurized hyperbaric oxygen chamber. 600.0 L of oxygen is
compressed in a cylinder at 160.0 atm. What volume of oxygen can a cylinder supply at the given pressure?
a) 11 L
48.
d) 32  103 L
What principle is illustrated in the figure?
a)
b)
49.
c) 11  103 L
b) 32 L
Boyle’s Law
Charles’s Law
c)
d)
Ideal Gas Law
Scientific Theory
What volume does .0685 mol of gas occupy at STP?
a) .3707 mol
b) .7515 mol
c) 1.53 L
d) 3.06 L
50.
Four liters of gas at atmospheric pressure is compressed into a 0.85L cylinder. What is the pressure of the
compressed gas if its temperature remains constant?
a) 0.15 atm
b) 0.21 atm
c) 3.4 atm
d) 4.7 atm
51.
Calculate the volume occupied by 35.2 g of methane gas (CH4) at 25°C and 101.3 kPa
a) 0.0186 L b) 4.5 L c) 11.2 L d) 49.2 L e) 53.5 L
52.
Which of the following properties of a gas is explained by the collision of its molecules with
the walls of its container?
a) volume b) pressure c) temperature d) mass
53.
Samples of helium and sulfur dioxide are released at the same time into opposite ends of a long tube. Which
position shows the location where these two gases will collide?
a) A
b) B
c) C
d) D
The next three questions refer to the diagram below:
54.
At what temperature will this substance melt if the pressure is held constant at 1.2 atm?
a) 130°C
b) 68°C
c) 88°C
d) The substance does not melt at this pressure
55.
If the temperature of this substance is held constant at 130°C while the pressure is changed from 1.4 atm to 0.8
atm, which is the change in state that will occur?
a) vaporization
b) condensation
c) deposition
d) melting
56.
What state of matter will this substance be in at a temperature of 28oC and a pressure of 0.7 atm?
a) solid
b) liquid
c) in equilibrium between solid and gas
d) gas
57.
A student dissolves 6.84 g of aluminum sulfate in enough water to form 250 ml of solution.
What is the molar concentration ( M ) of the solution?
a) 0.020 M
b) 0.080 M
c) 50.0 M
d) 200 M
58.
What is the concentration of chloride ions, in mol/L, if 28.5 g of magnesium chloride is dissolved in
enough water to form 900 ml of solution?
a) 0.67M b) 0.60M
c) 0.33M
d) 0.033m
59.
How many moles of acetate ions, CH3COO-, are in 46.0 ml of a 0.250M solution of Ca(CH3COO)2?
a) 11.5 moles b) 5.43 moles c) 1.15 moles d) 0.0115 moles e) 0.023 moles
60.
52.0 g of K2CO3 are dissolved in 518 g of distilled water, what is the concentration of the resulting
solution, in mol/L, M ?
a) 0.100M b) 7.27 x 10-4M c) 0.727M
d) 9.96M
61.
Given the reaction: Li2SO4(aq) → 2 Li+(aq) + SO42-(aq)
If the Molarity of the lithium sulfate solution is known to be 0.20 M, then what is the [Li+] ?
a) 0.10 M b) 5.0 M c) 0.20 M d) 0.40 M
62.
What volume of 12.0 M HCl(aq) is needed to prepare 5.00 L of a 0.500 M HCl solution?
a) 0.208 L b) 12.0 ml c) 0.208 ml
d) 12.0 L
63.
In a high school chemistry lab, a student accidentally mixed a solution of CaCl2(aq) with
a solution of Na2CO3(aq). A precipitate formed as a result.
What is the complete ionic equation for the reaction?
a) CaCl2(aq) + Na2CO3(aq) → CaCO3(s) + 2NaCl(aq)
b) Ca2+(aq) + Cl-(aq) + Na+(aq) + CO32-(aq) → CaCO3(s) + NaCl(aq)
c) Ca2+(aq) + 2Cl-(aq) + 2Na+(aq) + CO32-(aq) → CaCO3(s) + 2Na+(aq) + 2Cl-(aq
d) Ca2+(aq) + CO32-(aq) → CaCO3(s)
64.
Which of the following solutions would have the highest boiling point?
a) 1 g of NaCl dissolved in 1 000 g of water
b) 1 g of NaCl dissolved in 100 g of water
c) 1 g of C12H22O11 dissolved in 1 000 g of water
d) 1 g of C12H22O11 dissolved in 100 g of water
65.
If 100 ml of gasoline, a nonpolar mixture of alkanes, were added to a flask containing
100 ml of distilled water, what would happen?
a) the gasoline would catch on fire
b) the gasoline would evaporate immediately
c) the gasoline would dissolve in the water d) 2 separate layers of liquid would form
66.
Which of the following is NOT a property of aqueous solutions?
a) they create as osmotic pressure when separated from water by a permeable membrane
b) they have lower vapor pressure than pure water
c) they are homogenous mixtures
d) they are opaque to light rays
e) they have lower freezing points than pure water
67.
If 8.00 g sodium hydroxide, NaOH, are dissolved in 500.0 ml of distilled water, determine the
concentration of the solution, in mol/L:
of the solution:
a) 0.0160M
b) 16.0M
c) 4.00 x 10-4M d) 0.400M
68.
The IUPAC name for CH3 – C ≡ C – CH3 is:
a) butyne b) butene c) 2,3 – butyne d) 2 – butyne
69.
The IUPAC name for the hydrocarbon to the right is:
a) 2,4,4 – trimethylhexane
b) 2,3,3 – trimethylhexane
c) 2,3,5 – trimethyl-4-isopropylhexane
d) 2,3,5 – triethylhexane
CH3
|
CH3 – CH – CH – CH2 – CH – CH3
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CH3
CH3
70.
Which of the following compounds is not classified as an aromatic compound?
a) benzene b) propene
c) toluene d) 3 – bromophenol
71.
Two organic compounds have the following structural formulas:
H H H H H
H H H H
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H–C–C–C–C–C–H
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H OH H H
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H – C – C – C – C – OH
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H
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H CH3 H
|
H
Which one of the following statements is incorrect?
a) they are isomers
b) they have different melting points
c) they are both alcohols
d) they both have the same IUPAC name
72.
Identify the correct IUPAC name for the compound shown to the right:
a) 1,2 – dimethylcyclohexane
b) 1,2 – dimethylbenzene
c) 1,2 – diethylbenzene
d) 1,3 – dimethylcyclohexane
CH3
CH3
The second part of this review are short answer questions:
1.
Indicate the number of significant figures for the following measurement: 7.020 litres
2.
Perform the following calculation to the correct number of significant figures:
(4.30 mole)(1.34 L) / (5.963 kPa)
3.
Complete the following table.
Symbol
Atomic
Number
Protons
Electrons
Neutrons
Mass
Number
____
43
_____
43
56
_____
Sn – 119
_____
50
_____
____
_____
Rb+
37
37
_____
_____
86
Se2-
_____
34
_____
45
_____
_____
_____
______
115
_____
193
Pt
78
4.
Element Y has two isotopes, Y – 101 with an abundance of 59.25% in nature, and Y – 103 with an
abundance of 41.75%. Calculate the average atomic mass of element Y.
5.
Explain the differences/similarities in the Rutherford, Bohr, and Quantum models of the atom.
6.
An electron has the following quantum numbers:
n=3
l=1
m= 0
and
s = +1/2
what atomic orbital is the electron located in?
7.
A hydrogen electron is excited from the 1s orbital to the 4p orbital. What will happen next?
8.
Complete the following nuclear reaction:
249
Cf
98
11
+
1
B
5

5 n
0
+
__________________
9.
Write the nuclear reaction to represent the alpha decay of Gold – 199
10.
Iodine – 132 is used in the treatment of thyroid conditions. It has a half-life of 2.33
hours. How much of an initial 69.0 gram sample would remain after 5.8245 hours?
11.
Name the 4th row element that has the valence electron configuration of: s2p2
12.
The ion S2- would have the same electron configuration as which noble gas?
13.
An element has a ground state electron configuration of 1s22s22p3.
a) what type of ion would this element form to become stable? _______
b) what group (family) does this element belong to? ___________
14.
Explain why a chloride ion, Cl-, (181 pm) is so much larger than a chlorine atom (99 pm).
15.
Write valence electron configurations for the following atoms in their ground state:
a) Cu
b) Mo
c) Br
16.
Predict the most likely oxidation states for the elements in question 15 based on electron configurations
17.
Explain the following periodic trends: atomic radius, metallic character, and electron affinity.
18.
Determine the number of moles present in the following samples:
a) 4.44 g of carbon monoxide
b) 3.65 x 1027 molecules of carbon monoxide
19.
Determine the mass of the following:
a) 4.60 moles of F2O
b) 4.58 x 1024 molecules of CO2
20.
A sample contains, by mass, 40.0 g of sulfur and 60.0 g oxygen. Determine its empirical formula.
21.
A compound contains 25.8 g of potassium 42.4 g of sulfur, and 31.7 g of oxygen. Its molar mass is
known to be 453.9 g/mol. Determine its empirical and molecular formulas.
22.
Balance the following equations:
a) MnO2 + KOH + O2 + Cl2 → KMnO4 + KCl + H2O
b) C3H7OH(l) + O2(g)
→ CO2(g) + H2O(g)
c) Al2Cl3(s) → Al(s) + Cl2(g)
d) SO3(g) + HNO3(aq) → H2SO4(aq) + N2O5(g)
________________
23.
For each single displacement below, predict whether it occurs or not, and if it does, complete and
balance it:
a) Mg(s) + Pb(NO3)2(aq) →
b) Cl2(g) + NaBr(aq) →
24.
Write a balanced equation for the reaction and identify the precipitate formed:
Ca(NO3)2(aq) + Li2SO4(aq) →
25.
Complete and balance the equation for the following double displacement reactions. In each case
indicate whether a precipitate will be formed, a gas will be produced, or neutralization will occur.
a) Mg(OH)2 + H2SO4(aq) →
b) Na2CO3(aq) + H3PO4(aq) →
26.
Complete and balance the following combustion reactions:
a) C2H5OH(l) + O2(g)
→
b) CH3COCH3(l) + O2(g) →
27.
Given the reaction below:
CaCl2(aq) + AgNO3(aq) → AgCl(s) + Ca(NO3)2(aq)
i) what does the phase symbol, (aq) mean?
ii) balance the equation
iii) write an ionic equation for the reaction
iv) write a net ionic equation for the reaction; which ions are spectator ions?
28.
In a reaction between solid sulfur, S8, and oxygen gas, 160.0 g of sulfur dioxide is formed:
a) Write a balanced equation for this reaction
b) how many grams of sulfur were burned?
29.
Consider the reactants propene and oxygen: 2 C3H6(g) + 9 O2(g) → 3 CO2(g) + 3 H2O(g)
How many grams of oxygen are needed to react with 12.0 moles of propene?
30.
Consider the following reaction: C16H32(g) + 2 H2(g) → 2 C8H18(g)
How many grams of C8H18 can be made using 11.4 g of C16H32 ?
31.
If 35.0 g of Mg3(PO4)2 is present in a sample, determine the number of atoms of oxygen present
in the sample.
32.
Given the balanced equation which represents the reaction between butane and oxygen:
2 C4H10(g) + 13 O2(g) → 8 CO2(g) + 10 H2O(g) + 2 888 kJ
a) is this reaction endothermic or exothermic?
b) a student has 16.0 g of oxygen available to react with 40.0 g of butane. Determine the limiting
reagent and the theoretical yield of carbon dioxide gas. If the student obtains 11.0 g of carbon
dioxide, determine the percent yield of the experiment.
33.
Propane burns in the presence of oxygen to form carbon dioxide and water.
a) write a balanced equation for the reaction
b) how many grams of water can be produced from 10.0 grams of propane?
c) how many grams of carbon dioxide can be produced from 20.0 grams of propane?
d) how many moles of water can be formed from 16.2 moles of oxygen gas?
e) how many grams of carbon dioxide can be produced from 2.00 moles of propane?
34.
If a student has 15.0 g of nitrogen gas and combines it with 4.00 g of hydrogen gas, how much
ammonia can be produced? N2 + 3 H2 → 2 NH3
35.
What volume of oxygen gas at 310 K and 800 mmHg will react completely with 4.00 L of NO gas at the
same temperature and pressure?
2NO(g) + O2(g) → 2NO2(g)
36.
Draw the Lewis dot structures for, and predict the shape of the following:
a) phosphorous pentachloride , PCl5
b) sulfate ion, SO42c) Ozone, O3
37.
What is the difference between a bond dipole and a molecular dipole? Give an example of each.
Give an example of a molecule which is non-polar, but has bonds with a measurable dipole moment
38.
Construct Lewis diagrams and structural formulas for the following covalently bonded molecules:
identify the correct shapes.
a) H2S b) CO c) NO3d) PF3
e) TeCl4
f) XeF2
g) SO2
39.
Draw lewis structures, structural formulae for the following molecules:
indicate the number of pi bonds present in each molecule:
CO2 , SO3, and C2HF
40.
Convert the following Celsius temperatures into Kelvins:
a) 1500C
b) -300C
41.
Convert the following Kelvin temperatures into Celsius:
a) 457 K
b) 65 K
42.
Relate the strength of London dispersion forces to the size of the particles involved.
43.
Compare and contrast dispersion forces and dipole-dipole forces.
44.
Explain the difference between vaporization and evaporation.
45.
Explain why water has a higher boiling point than methane even though they have comparable molar masses.
46.
Explain why the molecules of cooking oil are not held together as tightly as the molecules of table salt.
47.
A gas at 110 kPa and 30.0°C fills a flexible container with an initial volume of 2.00 L. If the temperature is raised
to 80.0°C and the pressure increases to 440 kPa, what is the new volume?
48.
The scuba diver in the figure blows an air bubble 10 m underwater. As it rises to the surface, the pressure
decreases from 2.25 atm to 1.03 atm. What will the volume of air in the bubble at the surface?
49.
Calculate the number of moles of ammonia gas (NH3) contained in a 3.0 L vessel at 3.00  102 K with a pressure
of 1.50 atm.
50.
The complete combustion of ethane, C2H6, produces CO2 and water vapor as the only products.
Assume all substances are gases measured at the same temperature and pressure:
2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(g)
a) How many liters of ethane would produce 7.07 L of CO2 ?
b) How many liters of oxygen would be required for the reaction?
51.
Which liquid would you predict to be the best solvent for dissolving iodine crystals, I2 ? explain
a) water, H2O b) ethanol, CH3CH2OH
c) carbon tetrachloride, CCl4
52.
Write equations to represent the dissolving of the following ionic solids in water:
a) PbCl4(s) →
b) Mg(NO3)2(s) →
53.
A student prepares a solution by dissolving 2.40 moles of sodium hydroxide, NaOH, in enough
distilled water to produce 250 ml of solution. Calculate the concentration of this solution in
moles/L (M)
54.
A student is required to prepare 3.00 liters of a 4.00 x 10-2M solution of hydrochloric acid, HCl.
a) how many moles of HCl are required to prepare this solution?
b) what mass of HCl is required to prepare this solution?
55.
Describe the proper procedure for:
a) preparing 200.0 ml of a 0.064 M Na2CO3(aq) solution
b) preparing 460.0 ml of a 0.030 M NaI(aq) solution from a 1.0M stock solution of NaI(aq)
56.
A student dissolves 15.00 g of sodium sulfate in 500.0 ml of distilled water.
a) write an equation to represent the dissolving process
b) determine the concentration of the solution in mol/L
c) determine [Na+] in the solution
57.
Draw and name two isomers for:
a) C5H12
b) C4H9OH
58.
Name the following organic molecule:
59.
Show the structural or condensed formula for the following:
a) 2,3 dimethyl-2-pentene b) 4-ethyl-2-octyne
60.
Draw the structural and condensed formula for the following organic compounds:
a) 2 –methylpentane
b) 2,4-dimethyl-3-ethyl-1-hexene
c) 3-ethyloctane
d) bromocyclopropane
e) 4-methyl-1-pentanol
f) 2,3-dichloro-3-bromo-1-butene
g) 5-fluoro-2-hexyne
h) 3,4,4-trimethyl-2-decene-5-yne
i) 2-chlorophenol
j) 1,2,3-trichlorobenzene
k) 1,3,5-tribromobenzene
l) methylpentanoate
m) 3-methyl – 1 – butanol
n) ethylchloride
o) ethylbutanoate
p) 2-phenyl-4-chlorooctane
CH3CH2CH = C(CH3)CH2C(F)(CH3)CH3
61.
Match the name in the left colomn with the correct structure to the right
_________ cyclohexane
A. CH3 – CH – CH2 – CH2 –CH3
|
OH
_________ 3 – pentanol
B.
_________ 1, 4 – dimethylbenzene
C. CH3 – CH2 – CH – CH2 – CH3
|
OH
CH3
_________ 2 – pentanol
D.
CH3
________ 1,2 – dimethylchyclohexane
E.
CH3
CH3
F.
CH3
CH3
62.
Give the correct IUPAC name for the following hydrocarbons:
CH3
a) CH ≡ C – CH3
b)
|
CH2
|
CH3 – CH2 – CH – CH – CH2 – CH2 – CH2 – OH
|
CH3
____________________________
c)
________________________________________
d)
Cl
CH2 – CH3
Cl
Cl
CH2 – CH3
_______________________________
Good Luck!!
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