Download Review for Final Exam - Short Answer and Problems

Chem 1721/1821: Final Exam Review
Short Answer and Problems
1.
Provide the missing name, elemental symbol, or chemical formula for each of the following:
Tc
__________________________________
____________________________
____________________________
Y(NO3)3
______________________________
manganese
Ag2SO4
______________________________
ammonium chloride ____________________________
SnI2
__________________________________
CCl4 __________________________________
2.
copper (I) oxide
potassium bromate
____________________________
nitrous acid
____________________________
Balance the following chemical equation:
______ IBr (s) + ______ NH3 (g)  ______ NI3 (s) + ______ NH4Br (s)
3.
Write the balanced chemical equation that describes the combustion of acetylene gas, C2H2.
4.
Consider the following elements:
Cr Be N Co Ar Rb Br In
Identify which of these elements correspond to the following descriptions. NOTE: Each of the following descriptions is
written in the plural form for consistency only. Some blanks will have only one answer, others may have more than one answer.
5.
Those which are alkali metals:
______________________
Those which are non-metals:
______________________
Those which are in the 4th period of the periodic table:
______________________
Those which are noble gases:
______________________
Those which exist as diatomic molecules in their elemental state:
______________________
Those which are transition metals:
______________________
Provide the missing pieces of information in the following table:
Symbol
(including mass number and ion
number of protons number of neutrons
charge)
number of electrons
69
Ga3+
34
6.
Perform the following unit conversions:
72 mi/h = __________ m/s
7.
46
36
(1 mi = 1.6093 km, 1 h = 3600 s, 1 L = 10–3 m3)
3.68 mg/mm3 = ____________ µg/m3
10.449 dm = ___________ nm
Give the answer to the following calculation with the proper number of significant figures:
(6.167 + 83)/5.20 = _______________________
8.
Consider the following statement; then circle the physical properties and underline the chemical properties of
mercury (II) oxide.
Mercury (II) oxide is an orange-red solid with a density of 11.1 g/cm3. It decomposes when
heated to give mercury and oxygen. Mercury (II) oxide does not dissolve in water.
9.
Consider the following list of compounds, and then identify which of these compounds correspond to the following
descriptions. NOTE: All descriptions are written in the plural form. Some blanks will require only one answer,
others more than one. If none of the compounds match the description, write NONE.
P4O10 LiOH CH3OH NH3 H2SO4 MgCl2 HC2H3O2
10.
those which are strong electrolytes
_________________________
those which are weak electrolytes
_________________________
those which are non-electrolytes
_________________________
those which are strong acids
_________________________
those which are strong bases
_________________________
those which are salts
_________________________
2.00 mol of yttrium metal and 2.00 mol of molecular oxygen are allowed to react according to the following equation:
4 Y (s) + 3 O2 (g)  2 Y2O3 (s)
In this reaction the limiting reactant is _______________ . If 0.670 mol of Y2O3 is actually recovered in an
experiment, the percent yield is ________________.
11.
Predict possible products (include physical states) for the following reactions. If no reaction will occur, write NONE.
a.
LiOH (aq) + LiHSO4 (aq) 
___________________________________________________
b.
AgNO3 (aq) + CdCl2 (aq) 
___________________________________________________
c.
CO2 (g) + KOH (aq) 
___________________________________________________
12.
Write the balanced chemical, full ionic, and net ionic equations that describe the reaction that occurs when an aqueous
solution of calcium hydroxide reacts with chlorous acid (HClO2). Be sure to indicate the charges and phases throughout.
balanced chemical equation:
full ionic equation:
net ionic equation:
13.
Indicate the oxidation number for:
a. S in HSO3– __________
14.
b. Na in NaClO2 __________
c. Se in K2Se __________
Write the balanced redox equation for the following reaction that occurs in acidic solution. Use the
method of balancing the half reactions. Circle the reducing agent.
Cr2O72– (aq) + H2O2 (aq)  Cr3+ (aq) + O2 (g)
oxidation half reaction:
reduction half reaction:
overall equation:
15. Match each one of the following species with the appropriate descriptive chemical property:
Na+
__________
a. a proton donor
BaO
__________
b. an oxidizing agent
P2O5
__________
c. an acidic oxide
__________
d. a spectator in precipitation reactions
__________
e. a basic oxide
HC2H3O2
–
MnO4
16. Sodium chloride dissolves in water according to the following thermochemical equation:
NaCl (s)  Na+ (aq) + Cl– (aq);
17.
∆H = +3.90 kJ
a.
Is the dissolution of NaCl an endothermic or exothermic process?
_______________________
b.
Is energy released or absorbed during this process?
_______________________
2.50 mol of NO2 (g) is placed in a 15.0 L container at 35°C. The average speed of the NO2 molecules in this container
is __________________ . If the temperature is changed to 100°C, the average speed of the NO2 molecules
18.
will
_________________ (increase, decrease or remain constant), and the pressure exerted by the gas
will
____________________ (increase, decrease or remain constant).
a.
For a given compound, ΔH°vap is the enthalpy change that corresponds to the physical change from the
_________________ state to the _________________ state.
b.
The phase change described above in a is _____________________ (endothermic, exothermic).
c.
For mercury, ΔH°fus = 2.292 kJ/mol and ΔH°vap = 59.30 kJ/mol. What is the value of ΔH°sub for mercury? _____
d.
Which of the following will not have ΔHf° equal to zero? K (l), F2 (g), Ar (g), O2 (g) ____________________
19.
Write the chemical equation of formation for methanol, CH3OH (l).
20.
There are 2 compounds of the formula Pt(NH3)2Cl2.
Cl
NH3
Cl
Pt
Cl
Cl
Pt
NH3
NH3
NH3
The compound on the right, cisplatin, is used in cancer therapy. Both compounds have square planar geometry. Circle
the compound that has a non-zero dipole moment.
21.
Consider the acetonitrile molecule:
H
H
C
C
N
H
a.
What are the bond angles around each carbon?
__________
__________
b.
What is the hybridization on each C atom?
__________
__________
c.
Determine the total number of σ and π bonds.
σ __________
π __________
22.
23.
a.
Draw the Lewis structure of XeF2.
b.
What is the electron pair geometry?
_________________________
c.
What is the molecular shape?
_________________________
d.
What orbitals on Xe and F participate in the bonding?
_________________________
e.
Is XeF2 polar or nonpolar?
_________________________
a.
Will potassium nitrate be more soluble in water or hexane (C6H14)?
______________________
b.
Will propane (C3H8) be more soluble in water or carbon tetrachloride (CCl4)?
______________________
c.
Urea - a nonvolatile, molecular compound - is dissolved in pyridine. Will the vapor pressure of the solution
increase, decrease, or remain constant relative to pure pyridine?
d.
______________________
If NH4NO3 is dissolved in pyridine, would you expect the ∆Psoln to be larger or smaller than the ∆Psoln
observed in the solution prepared in part c above?
24.
25.
26.
______________________
A solution is prepared by dissolving 30.0 g of propanol (C3H7OH, molar mass = 60.10 g/mol) in 80.0 g of water
(molar mass = 18.02 g/mol) resulting in a solution with a total volume of 108 mL.
a.
Calculate the molarity of this solution.
______________________________
b.
Calculate the mole fraction of propanol in this solution.
______________________________
c.
Calculate the mass % propanol in this solution.
______________________________
a.
Will CCl4 be more soluble in benzene or water?
______________________
b.
Will C2H5OH be more soluble in hexane or ammonia?
______________________
a.
A solution is prepared by dissolving 1 mol of ammonium phosphate (NH4)3PO4 in water. What is the theoretical
value of i (# mol particles in solution) for this reaction?
b.
______________________
Will the vapor pressure of the solution described in a be higher or lower than the vapor pressure of pure water?
______________________
c.
Will the vapor pressure of the solution described in a be higher or lower than a solution prepared by dissolving
1 mol of potassium chloride (KCl) in water?
27.
______________________
a.
Write the balanced chemical equation for the dissociation of benzoic acid,HC7H5O2.
b.
What is the conjugate base of benzoic acid?
______________________
c.
Write the Ka expression for this dissociation.
______________________
d.
If pKa of benzoic acid = 4.19, what is Ka of benzoic acid?
______________________
28.
The reaction 2 A + B  C obeys the following rate equation: Rate = k[A][B]1/2. The order of the reaction with respect
to B is ___________________, and the total order of the reaction is ____________________.
The units of k are
_____________________. If the rate of formation of C is 2.00 mol•L—1•s—1, then the rate of
consumption of A = ____________________.
29.
The rate constant (k) for a reaction was studied as a function of temperature (T). The activation energy for this
reaction can be found from the slope in a linear plot of _________________ vs. ___________________.
30.
Write the expression for the equilibrium constant K for the following reaction: C (s) + 1/2 O2 (g)  CO (g)
31.
Consider the following equilibrium: Ca2+ (aq) + H2O (l) + CO2 (g)  CaCO3 (s) + 2 H+ (aq); ΔH = +16 kJ/mol
Each of the following changes are made separately. Circle the predicted effect of each change on the quantity listed
(i = increase, d = decrease, nc = no change).
Change
Add catalyst
Quantity
Amount of H+
i
Add more CaCO3
Amount of H+
i
d
nc
Increase Volume
Amount of H
+
i
d
nc
Decrease temperature
K
i
d
nc
Increase temperature
Amount of CaCO3
i
d
nc
Add Ar
Amount of CO2
i
d
nc
K
i
d
nc
Add more Ca
32.
2+
Effect
d nc
KP = 0.250 at 1100 K for the following equilibrium: 2 SO2 (g) + O2 (g)  2 SO3 (g)
For this reaction K = ________________________, and for the following reaction: SO2 (g) + 1/2 O2 (g)  SO3 (g)
KP = __________________________ .
33.
Write chemical formulas for three strong acids: ______________ , ______________ , _____________
34.
Identify each of the following species as acid, base, conjugate acid, or conjugate base:
H2PO4– (aq) + SO32– (aq)  HPO42– (aq) + HSO3– (aq)
___________
35.
___________
___________
___________
Suppose that QP < KP for the reaction 2A + B  C at some time. As the reaction proceeds, will the partial pressures of the
substances change? If so, how? _______________________________________________________________
36.
The pH of a solution at the stoichiometric point in a titration of a ______________ acid and a _____________ base
is less than 7.00.
37.
Write the Ksp expression for Cr2(CO3)3
Ksp = ____________________________
38.
The pH of a 0.20 M solution of propylamine, (C3H7NH2 - a weak base) is 12.0. If 0.10 M propylammonium chloride,
C3H7NH3Cl, is added to the solution, will the pH increase, decrease, or remain constant? _______________________
39.
40.
Write the balanced net ionic equation ( include physical states) that describes each of the following:
a.
the dissolution/precipitation equilibrium for copper (I) sulfate, Cu2SO4:
b.
the neutralization reaction that occurs when KOH is added to a buffer solution composed of acetic acid and
potassium acetate
c.
the equilibrium that determines the pH of a solution of sodium hypochlorite, NaOCl (aq)
You need to prepare a buffer solution with pH = 11.0. Which of the following components, together with its
conjugate is the best choice for this buffer? Circle your choice.
(CH3)2NH, Kb = 5.4 x 10-4
41.
HF, Ka = 3.5 x 10-4
NH3, Kb = 1.8 x 10-5
Which of the following has the greatest molar solubility? Circle your choice:
Al(OH)3, Ksp = 1.9 x 10-33
Fe(OH)3, Ksp = 2.6 x 10-39
Cr(OH)3, Ksp = 6.7 x 10-31
42.
A reaction will occur spontaneously at constant temperature and pressure when its ________________ is negative.
43.
Circle the Lewis acid in the following reaction: Ag+ (aq) + 2 CN– (aq) ↔ Ag(CN)2– (aq)
44.
Aniline (C6H5NH2) is a weak base, Kb = 3.8 x 10–10. What is pKa for C6H5NH3+?
45.
Consider the following reaction: AgI (s) + 2 CN– (aq) ↔ Ag(CN)2– (aq) + I– (aq). What is the equilibrium constant
_________________________
for this reaction given that for AgI, Ksp = 1.5 x 10–16 and Kf = 5.6 x 10+8 for the following reaction:
Ag+ (aq) + 2 CN– (aq) ↔ Ag(CN)2– (aq).
46.
Classify each of the following solutions as either acidic (A), basic (B), or neutral (N):
a. CH3NH3I (aq) ______________
47.
K = _________________________
b. KI (aq) ______________
Identify the reducing agent in the following chemical reaction:
14 H+ (aq) + Cr2O72–(aq) + 3 Ni (s)  3 Ni2+ (aq) + 2 Cr3+ (aq) + 7 H2O (l)
48.
c. NaH2PO4 (aq) _______________
_________________________
Consider a Galvanic cell based on the following cell reaction:
Mg (s) + 2 Ag+ (aq)  Mg2+ (aq) + Ag (s)
The individual half reactions are:
Mg2+ + 2 e–  Mg; E° = –2.37 V
Ag+ + e–  Ag; E° = 0.80 V
Does this reaction proceed spontaneously in the direction written?
___________________________
How many moles of electrons are transferred in this reaction?
___________________________
49.
Consider the following Galvanic cell: Co(s) | Co2+ (aq) || Fe2+ (aq) | Fe (s)
Write the half-reaction that takes place at the cathode:
50.
Consider a Galvanic cell based on the following chemical reaction: Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s).
This cell operates at a potential of 0.95 V when [Cu2+] = 1.0 x 10–5 M and [Zn2+] = 1.0 M. Calculate E° for this cell at 25°C.
51.
Consider the following half-reactions:
Cu+ (aq) + e-  Cu (s); E° = 0.52 V
NO3- (aq) + 4 H+ (aq) + 3 e-  NO (g) + 2 H2O (l); E° = 0.96 V
a.
Write the overall cell reaction for this Galvanic cell.
b.
Write the line notation that describes this cell.
c.
Calculate ΔG° for the chemical reaction that occurs in this Galvanic cell. (1 V = 1 J/C)
52.
An electrolytic cell is used to plate zinc metal from Zn2+ ions in solution. What current is required to produce
25.0 g of Zn in 30.0 min? (1 amp = 1 C/s)
53.
a.
What type of emission occurs when Mn-49 decays according to the following equation:
49
Mn  49Cr + ???
_________________________________
b.
54.
Consider a 25.0 g sample of glucose, C6H12O6 (molar mass = 180.18 g/mol).
a.
How many glucose molecules are there in this sample?
b.
55.
Write the balanced nuclear equation that describes α-emission of Pt-170.
How many hydrogen atoms are there in this sample?
Styrene is a compound containing only carbon and hydrogen. Combustion analysis performed on a 2.78 g sample of
styrene yielded 9.40 g of CO2 and 1.92 g of H2O. The molar mass of styrene is 104.2 g/mol. Determine the empirical
and molecular formula of styrene.
note: Not all Chem 1711 instructors discuss combustion analysis for formula determination.
56.
Limestone is composed of calcium carbonate (CaCO3) as well as other compounds. In an analysis, a chemist takes a
sample of limestone which has a mass of 413 mg and treats it with oxalic acid (H2C2O4). A chemical reaction occurs
between the calcium carbonate and the acid producing calcium oxalate and other products.
CaCO3 (s) + H2C2O4 (aq)  CaC2O4 (s) + H2O (l) + CO2 (g)
The mass of CaC2O4 obtained is 472 mg. What is the percent by mass of calcium carbonate in the original sample of
limestone?
57.
Calcualte the volume of a 2.50 M potassium sulfate solution that would be required to react completely with 50.0 mL
of a 1.48 M barium chloride solution.
58.
A 28.2 g sample of a metal is heated to 95.2˚C and dropped in a calorimeter which contains 100.0 g of water at 25˚C.
The final temperature in the calorimeter is 31.0˚C. Assuming no heat is lost to the surroundings or the calorimeter,
calculate the specific heat of the metal. The specific heat of H2O (l) is 4.184 J/g•°C.
59.
A sample of neon in a rigid 5.00 L flask exerts a pressure of 0.100 atm at 250 K. Argon is then added until the mole
fraction of argon is 0.333. The temperature remains at 250 K.
a.
Calculate the number of mol of Ne in the flask.
b.
Calculate the mass (in g) of argon that was added
c.
Calculate the total pressure in the flask after the argon was added.
60.
A student in lab titrates an unknown cobalt (II) oxalate dihydrate sample with 23.4 mL of 0.0203 M KMnO4 (aq). The
balanced net ionic equation for the titration is:
2 MnO4– (aq) + 5 C2O42– (aq) + 16 H+ (aq)  2 Mn2+ (aq) + 10 CO2 (g) + 8 H2O (l)
Calculate the mass, in g, of oxalate ion present in the unknown sample.
61.
50.0 mL of HCl (g) at 25°C and 800 torr is dissolved in enough water to produce 125.0 mL of a hydrochloric acid solution.
a.
What is the molar concentration of the resulting HCl (aq)?
b.
62.
Suppose that coal, of density 1.5 g/cm3, is carbon. The combustion of carbon is described by the equation
C (s) + O2 (g)  CO2 (g) ∆H = –394 kJ
a.
Calculate the heat produced when a lump of coal of size 7.0 cm x 6.0 cm x 5.0 cm is burned.
b.
63.
64.
This HCl (aq) is used in a titration to determine the molar concentration of a solution of Ca(OH)2. 31.44 mL
of HCl (aq) are required to reach the stoichiometric point in the titration. Calculate [Ca(OH)2 (aq)].
Estimate the mass of water that can be heated from 15°C to 100°C with this piece of coal.
For the following elements write the ground state electron configuration and show how the atomic orbitals would be
populated for n > 3. Be sure that your atomic orbitals are clearly labeled. (An example of what I want: For krypton,
the ground state electron configuration would be [Ar] 4s2 3d10 4p6, and the atomic orbitals for n > 3 would be
populated as follows: 4s ↑↓
3d ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
4p ↑↓ ↑↓ ↑↓ )
a.
cobalt
b.
bromine
a.
Using the Bohr model of the atom determine the energy associated with the transition of an electron from n = 3
to n = 7.
b.
Calculate the wavelength of light associated with this transition.
c.
What type of eletromagnetic radiation does this wavelength correspond to?
___________________________
d.
Is light absorbed or emitted during this transition?
___________________________
λ = ___________________________
65.
Draw the orbitals of the l = 1 subshell. Be sure that their orientation with respect to an x, y, z axis system is clear.
66.
Consider the following equilibrium: H2 (g) + I2 (g)  2 HI (g); K = 54.1 at 700 K. 2.00 mol of H2 and 1.50 mol of
I2 were put in a previously evacuated 10.0 L vessel, and the temperature was maintained at 700 K. Determine the
equilibrium concentrations of all species in the vessel.
[H2]eq _________________
[I2]eq _____________________
—7
—1 —1
at 283°C and 5.12 x 10—4 M—1s—1 at 413°C.
67.
The rate constant for the decomposition of HI is 3.52 x 10
Determine the activation energy for this reaction.
68.
The pH of a 0.100 M solution of iodic acid, HIO3, is 1.91 at 25°C. Determine Ka for iodic acid.
69.
Calculate the pH of a 0.0100 M solution of cocaine, a weak base (pKb = 5.59).
70.
Calculate the molar solubility of AgI (Ksp = 1.5 x 10–16) in 0.200 M MgI2 (aq).
71.
Consider the following reaction at 40°C: H2O (l) ↔ H+ (aq) + OH– (aq). At 40°C, ΔG°f = –237.1 kJ/mol for H2O (l);
ΔG°f = –156.0 kJ/mol for OH– (aq); and ΔG°f = 0 kJ/mol for H+ (aq).
a.
Calculate ΔG° for this reaction.
72.
b.
Calculate the equilibrium constant (Kw) at 40°C.
c.
Calculate the pH of pure water at 40°C.
M s
[HI]eq __________________
Determine ΔH for the following reaction: C2H4 (g) + 6 F2 (g)  2 CF4 (g) + 4 HF (g).
Use the following information and apply Hess's Law of Heat Summation:
H2 (g) + F2 (g)  2 HF (g); ΔH = –537 kJ
C (s) + 2 F2 (g)  CF4 (g); ΔH = –680 kJ
2 C (s) + 2 H2 (g)  C2H4 (g);
ΔH = +52.3 kJ
73.
Determine the standard enthalpy change for the following reaction using ΔH°f data:
4 NH3 (g) + 5 O2 (g)  4 NO (g) + 6 H2O (g)
for NH3 (g); ΔH°f = –46 kJ/mol
for NO (g); ΔH°f = + 90 kJ/mol
for H2O (g); ΔH°f = –242 kJ/mol
74. How much energy will be released when 10.0 g of H2O (l) initially at 80°C is cooled to H2O (s) at –10°C?
For H2O (l); s = 4.18 J/g•°C. For H2O (s); s = 2.03 J/g•°C. For H2O at 0°C; ΔH°fus = 6.01 kJ/mol
75.
A biochemical engineer isolates a bacterial gene fragment and dissolves a 10.0 mg sample in enough water to make
30.0 mL of sol'n. The osmotic pressure of the sol'n is 0.340 Torr at 25°C.
What is the molar mass of the bacterial gene fragment?
76.
Consider an ideal solution of 200.0 mL of pentane (C5H12, d = 0.63 g/mL) and 200.0 mL of hexane (C6H14, d = 0.66
g/mL). The vapor pressures of the pure liquids are 511 Torr for pentane and 150 Torr for hexane.
77.
a.
b.
Which compound is more volatile: pentane or hexane?
Determine the total vapor pressure of this solution.
______________________
c.
Calculate the mole fraction of pentane in the vapor.
a.
Which of the following has the highest boiling point? PCl3, NH3, or He
______________________
b.
Which of the following has the lowest vapor pressure at 25°C? Cl2, Br2, or I2
______________________
c.
Which of the following is the most polarizeable? F, Cl, Br, or I
______________________
d.
Rank the following in order of increasing strength of intermolecular forces: C2H4, LiCl, H2CO, He, CH3OH
____________ < ____________ < ____________ < ____________ < ____________
e.
Put the following bonds in order of increasing bond polarity: O – Cl, Cl – Cl, Na – Cl
__________ < __________ < __________
78.
The normal boiling point of benzene is 80.1°C.
a.
What is the vapor pressure of acetone at 80.1°C?
b.
Calculate the vapor pressure of acetone at 25°C. For benzene, ΔHvap = 30.8 kJ/mol.
______________________
79.
80.
b.
Draw the molecular orbital diagram for N2+. Show the atomic orbitals on the left and right sides of the diagram
and the molecular orbitals in the middle.
Calculate the bond order for this ion.
__________________
c.
Is this species paramagnetic or diamagnetic?
d.
Write the molecular electron configuration for N2+.
e.
Which species would MO theory predict to be more stable, N2+– or C2+?
__________________
f.
Would C2 or N2+ have a longer bond?
__________________
g.
Would C2 or N2+ have a greater bond dissociation energy?
__________________
a.
__________________
______________________________________
Indicate whether each of the following statements is true or false:
a.
The C–O bonds in the carbonate ion are best described as delocalized π bonds.
b.
When one 2s orbital and three 2p orbitals mix, three equivalent sp3 hybrid orbitals are formed.
c.
If the central atom in a molecule is sp d hybridized, the bond angles in the molecule are ≈109°. ____________
d.
sp hybrid orbitals have two lobes with one lobe significantly larger than the other.
e.
The extent of parallel overlap in π bond formation is typically less than the orbital overlap in a σ bond.
Match the description on the left with the one best term on the right.
i.
hybridization that results in one unhybridized p orbital ________
a. π bond
b. sp2
ii.
hybridization resulting in bond angles of 120 and 90°
________
c. double bond
d. sp3
iii.
geometry around each C in C2H4
________
e. triple bond
f. sp3d
iv.
one σ bond + two π bonds
________
g. σ bond
h.
v.
e– density concentrated symmetrically about
i. trigonal planar
j. linear
tetrahedral
________
The electron configuration for an element is [Ne]3s23p2.
a.
Identify this element.
____________________________
b.
What does [Ne] represent?
____________________________
c.
How many unpaired electrons are there in this atom?
d.
83.
____________
____________
the internuclear axis
82.
____________
3 2
Therefore a π bond is weaker than a σ bond.
81.
____________
–
Give the 4 quantum numbers for one of the e ‘s in the 3s orbital.
____________________________
n = ____; l = ____; ml = ____; ms = _____
–
e.
For this element a 3p e is best described as a(n) {core, valence, noble gas, or paramagnetic} e–._____________
f.
Identify the electron-occupied orbital of highest energy in this atom.
a.
The energy required to remove an e– from a gas phase atom or ion is the {ionization energy, electron affinity,
____________________________
enthalpy of formation, or lattice energy}.
b.
Write the balanced chemical equation that corresponds to the 2nd ionization energy for aluminum.
c.
For Al, the 2nd ionization energy will be {less than, greater than, or equal to} the 1st ionization energy.
84.
85.
86.
87.
a.
How many valence electrons does P3– have?
____________________________
3–
b.
Give one example of an atom or ion that is isoelectric with P .
a.
Which electromagnetic radiation has the longer wavelength? ν = 4.5 x 1012 Hz or ν = 6.4 x 1015 Hz
b.
Which one of the following has the longest wavelength? visible light, UV light, microwaves, IR
c.
Which light has the greater speed? i. λ = 20 nm, ii. λ = 800 nm, or iii. neither – both are the same
d.
Which electromagnetic radiation has greater energy per photon? i. λ = 2 x 10–2 nm, or ii. λ = 450 nm
a.
Give all possible values for ml for an electron in a 2p orbital.
b.
Is the following combination of quantum numbers allowed? n = 3, l = 3, ml = 1, ms = +1/2
c.
How many electrons can have the following quantum numbers? n = 3 and ml = 0.
Consider the phase diagram for substance Z shown to the right.
a.
Estimate the normal melting temperature of Z.
b.
What phase change (if any) occurs when Z is held at
200 K and undergoes a ΔP from 1 to 0.001 atm?
c.
What phase change (if any) occurs when Z is held at
200 atm and undergoes a ΔT from 300 to 100 K?
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Chem 1721/1821: Final Exam Review
1.
Short Answer and Problems ANSWERS
technetium
Cu2O
yttrium nitrate
Mn
silver sulfate
NH4Cl
tin (II) iodide
KBrO3
carbon tetrachloride
HNO2 (aq)
2.
3 IBr (s) + 4 NH3 (g)  1 NI3 (s) + 3 NH4Br (s)
3.
2 C2H2 (g) + 5 O2 (g)  4 CO2 (g) + 2 H2O (g)
4.
Rb
N, Ar, Br
Cr, Co, Br
Ar
N, Br
Cr, Co
5.
31, 38, 28
80
Se2–
6.
32 m/s, 3.68 x 1012 µg/m3, 1.0449 x 109 nm
7.
15
8.
physical properties: orange-red solid, d = 11.1 g/cm3
chemical properties: decomposes when heated, does not dissolve in water
9.
LiOH, H2SO4, MgCl2
NH3, HC2H3O2
P4O10, CH3OH
H2SO4
LiOH
LiOH, MgCl2
10.
Y, 67.9
11.
a.
Li2SO4 (aq) + H2O (l)
b.
AgCl (s) + Cd(NO3)2 (aq)
c.
K2CO3 (aq) + H2O (l)
12.
Ca(OH)2 (aq) + 2 HClO2 (aq)  Ca(ClO2)2 (aq) + 2 H2O (l)
Ca2+ (aq) + 2 OH– (aq) + 2 HClO2 (aq)  Ca2+ (aq) + 2 ClO2– (aq) + 2 H2O (l) (*because HClO2 is a weak acid)
OH– (aq) + HClO2 (aq)  ClO2– (aq) + H2O (l)
13.
a. +4, b. +1, c. –2
14.
ox: H2O2 (aq)  O2 (g) + 2 H+ (aq) + 2 e–
red. 6 e– + 14 H+ (aq) + Cr2O72– (aq)  2 Cr3+ (aq) + 7 H2O (l)
net: 3 H2O2 (aq) + 8 H+ (aq) + Cr2O72– (aq)  3 O2 (g) + 2 Cr3+ (aq) + 7 H2O (l);
15.
top to bottom: d, e, c, a, b
16.
a. endothermic, b. absorbed
17.
409 m/s, increase, increase
* H2O2 is the reducing agent
18.
a. l  g (or g  l), b. endothermic for l  g (exothermic for g  l), c. 61.59 kJ/mol, d. K(l)
19.
C (s) + 2 H2 (g) + 1/2 O2 (g)  CH3OH (l)
20.
compound on right – cis-Pt(NH3)2Cl2 is polar (has non-zero dipole moment)
21.
a.
109.5°, 180°
b.
sp3, sp
c.
σ 5, π 2
a.
should be 22 electrons or 11 e– pairs
b.
trigonal bipyramidal
c.
linear
d.
sp3d on Xe with sp3 on F
e.
nonpolar
22.
23
a. water, b. CCl4, c. decrease, d. larger, e. 4, f. lower, g. lower
24.
a. 4.62 M, b. 0.101, c. 27.3%, d 6.24 m
25.
a.
benzene
b.
ammonia
a.
4
b.
lower
c.
lower
a.
HC7H5O2 (aq) + H2O (l)  C7H5O2– (aq) + H3O+ (aq)
b.
C7H5O2–
c.
Ka = [C7H5O2–][H3O+]/[HC7H5O2]
d.
6.5 x 10–5
26.
27.
28.
0.5, 1.5, M–.5•s–1, 4.00 M•s–1
29.
ln k vs. 1/T
30.
K = [CO]/[O2]1/2
31.
nc, nc, d, d, i, nc, nc
32.
22.5, 0.500
33.
HCl, HBr, HI, HNO3, HClO4, H2SO4
34.
H2PO4– acid, SO32– base, HPO42– conjugate base, HSO3– conjugate acid
35.
reaction proceeds forward toward equilibrium therefore PC increases and PA and PB decrease
36.
strong, weak
37.
Ksp = [Cr3+]2[CO32–]3
38.
decrease (the equilibrium in question is C3H7NH2 (aq) + H2O (l)  C3H7NH3+ (aq) + OH– (aq))
39.
a.
Cu2SO4 (s)  2 Cu+ (aq) + SO42– (aq)
b.
HC2H3O2 (aq) + OH– (aq)  C2H3O2– (aq) + H2O (l)
c.
OCl– (aq) + H2O (l)  HOCl (aq) + OH– (aq)
40.
(CH3)2NH
41.
Cr(OH)3
42.
ΔG
43.
Ag+ (aq)
44.
4.58
45.
K = Ksp • Kf = 8.4 x 10–8
46.
a. acidic, b. neutral, c. acidic
47.
Ni
48.
yes, 2
49.
Fe2+ + 2 e–  Fe
50.
1.1 V
51.
a.
3 Cu (s) + NO3– (aq) + 4 H+ (aq)  NO (g) + 2 H2O (l) + 3 Cu+ (aq), E° = 0.44 V
b.
Cu (s) | Cu+ (aq) || NO (g) | H+ (aq), NO3– (aq) | inert conductor like Pt (s) or C (gr)
c.
–127 kJ
52.
41.0 A
53.
a. positron emission (β+), b.
54.
a. 8.36 x 1022 molecules, b. 1.00 x 1024 H atoms
55.
CH, C8H8
56.
89.3%
57.
29.6 mL
58.
1.39 J/g°C
59.
a. 0.0244 mol Ne, b. 0.487 g Ar, c. 0.150 atm
60.
0.105 g C2O42–
61.
a. 0.0172 M, b. 0.0180 M
62.
a. 1.03 x 104 kJ released, b. 2.5 x 104 g
63.
a. [Ar]4s23d7, b. [Ar]4s23d104p5
64.
a. 1.98 x 10–19 J, b. 1000 nm, c. infra-red, d. absorbed
65.
px, py, and pz orbitals, see p. 290, figure 7.25 in your text
66.
[H2] = 0.068 M, [I2] = 0.018 M, [HI] = 0.264 M
67.
177 kJ/mol
68.
1.73 x 10–3
69.
10.20
70.
3.8 x 10–16 mol/L
71.
a. +81.1 kJ, b. 2.82 x 10–14, c. 6.77
72.
–2486 kJ
73.
–908 kJ
74.
6.88 kJ released
75.
1.82 x 104 g/mol
76.
a. pentane, b. 343 Torr, c. 0.796
77.
a. NH3, b. I2, c. I, d. He < C2H4 < H2CO < CH3OH < LiCl, e. Cl–Cl < O–Cl < Na–Cl
78.
a. 760 Torr, b. 109 Torr
170
Pt  α + 166Os
79.
a.
N2+ 9 e– total; refer to notes, or see me with questions about diagram
b.
bond order = 2.5
c.
paramagnetic
d.
(σ2s)2(σ2s*)2(π2p)4(σ2p)1
e.
N2+
f.
C2
g.
C2
80.
a. T, b. F, c. F, d. T, e. T
81.
i. b, ii. f, iii. i, iv. e, v. g
82.
a.
silicon
b.
electron configuration of neon: 1s2 2s2 2p6
c.
2
d.
n = 3, l = 0, ml = 0, ms = +1/2 or – 1/2
e.
valence
f.
3p
83.
a. ionization energy, b. Al+ (g)  Al2+ (g) + e–, c. greater than
84.
a. 8, b. S2–, Cl–, Ar, K+, Ca2+, Sc3+
85.
a.
ν = 4.5 x 1012 Hz
b.
microwaves
c.
neither – both are the same
d.
λ = 2 x 10–2 nm
86.
a. –1, 0, or +1, b. no, c. 6
87.
a. 175 K, b. l  g, c. l  s