Download Math Review: Calculations and Scientific Notation

General Chemistry Practice Problems
Harvard University
Math Review: Calculations and Scientific Notation
1.
Use your calculator to determine the answers to each of the following calculations.
Please round your answers to three or four digits.
a)
423.1 + 100 + 0.256
b)
14.000 ÷ 6.1
c)
(6.11)(!)
d)
(4/3)!(2.16)3
e)
(14.3)(60)
f)
(6.0+9.7+0.61)(1.113)
g)
6.958 ! 10 8
5.91 ! 1012
h)
(1.173 " 10–3) + 3.6
i)
(1.1 " 106)(2.246 " 10–10)
1
General Chemistry Practice Problems
Harvard University
Dimensional Analysis I
Useful Information:
1 atm = 760 torr
1 torr = 0.01933 lb/in2
1 N = 0.225 lb
1 Pa = 1 N/m2
1 kg = 2.2 pounds
1 mL = 1 cm3
1 inch = 2.54 cm
1 foot = 12 inches
Use Dimensional Analysis to calculate the answers to each of the following problems.
1.
a)
The pressure inside a laboratory vessel is 35 lb/in2. Calculate the equivalent pressure in
atmospheres.
b)
A certain chemistry experiment takes place in a high-pressure chamber with a pressure
of 2.7 atmospheres. Calculate the pressure in torr.
c)
The atmospheric pressure is measured to be 754 torr. Calculate the pressure in Pa.
2.
The recommended adult dose of a drug used to treat asthma is 6.0 mg/kg of body mass.
Calculate the dose (in mg) for a 125-pound person.
3.
A standard sheet of laser printer paper is 0.09 mm thick. How many sheets of paper
must be stacked to reach the height of the Eiffel Tower at 1063.0 ft?
2
General Chemistry Practice Problems
Harvard University
Dimensional Analysis II
Useful Information:
1 gallon = 0.134 ft3
1 mL = 1 cm3
1 mile = 5280 feet
1 inch = 2.54 cm
Use Dimensional Analysis to calculate the answers to each of the following problems.
1.
A certain crystal of calcium chloride has a volume of 765.3 mm3.
a)
Calculate the volume of this crystal in cubic inches (in3).
b)
The density of CaCl2 is 2.51 g/cm3. Calculate the mass of this crystal.
2.
A reservoir has an area of 1.3 square miles with an average depth of 93.2 feet.
Calculate the volume of water (in gallons) held in the reservoir.
3
General Chemistry Practice Problems
Harvard University
Atoms, Molecules, and Ions
1.
For each of the following atoms or ions, provide the number of protons, neutrons, and
electrons:
a)
40Ar
b)
40Ca2+
c)
39K+
d)
39K
2.
Calcium chloride is an ionic substance. Determine the number of protons and electrons
in a calcium ion (as found in calcium chloride).
3.
The chloride ion exists in two common isotopes: 35Cl– and 37Cl–. Determine the
number of protons, neutrons, and electrons for both isotopes.
4.
a)
The chemical formulas for ammonia is NH3.
Calculate the total number of protons, neutrons, and electrons in one molecule of NH3.
(Assume that the only isotopes present are 14N and 1H.)
b)
In aqueous solution, some of the ammonia is present as the ammonium ion, NH4+.
Calculate the total number of protons, neutrons, and electrons in one ammonium ion,
NH4+. (Again, assume that the only isotopes present are 14N and 1H.)
4
General Chemistry Practice Problems
Harvard University
Naming Compounds
1.
2.
Write the chemical formula for each of the following species:
a)
Copper (I) phosphide
b)
Iron (III) sulfate
c)
Iodine
d)
Aluminum chloride
e)
Chromium (III) oxide
f)
Ammonia
g)
Lithium nitride
h)
Aluminum carbonate
i)
Cesium phosphate
j)
Zinc oxide
k)
Stannous chloride
l)
Dinitrogen pentoxide
m)
Silver Fluoride
n)
Phosphoric Acid
o)
Barium Phosphate
p)
Ammonium Sulfate
Write an acceptable chemical name for each of the following species.
a)
CCl4
b)
CuO
c)
Mg3N2
5
General Chemistry Practice Problems
Harvard University
Chemical Formulas: Percent Composition
1.
Arrange the following substances in order of increasing mass percent of carbon.
i) caffeine, C8H10N4O2
2.
a)
b)
ii) fructose, C6H22O6
iii) acetic acid, CH3COOH
The nerve gas Sarin, which was released in a Tokyo subway station in 1996, has a
molecular formula of C4H10PO2F.
Determine the percent composition by mass of Sarin.
An unknown compound is discovered in a raid on a terrorist organization; it is believed
that the compound is Sarin. When a 10.0-gram sample of this compound is completely
combusted, 15.6 g CO2 and 6.4 g H2O are produced, along with other combustion
products. Using numerical calculations, prove that this unknown compound cannot be
Sarin.
6
General Chemistry Practice Problems
Harvard University
Empirical and Molecular Formulas
1.
Tetrodotoxin is a neurotoxin found in several species including pufferfish and blueringed octopus. Tetrodotoxin is 41.38% C, 13.16% N, 5.37% H by mass, with the
remaining amount consisting of oxygen. It was found to have a molecular mass of
5.30 " 10–22 g. Determine the empirical formula and the molecular formula of
tetrodotoxin.
2.
Compound Z, which consists of only carbon, hydrogen, and oxygen, has just been
isolated from a tropical plant.
a)
When 5.467 grams of compound Z are burned in excess oxygen, 15.02 grams of CO2
and 2.458 grams of H2O are produced. Determine its empirical formula.
b)
Other experiments suggest that compound Z has a molar mass of approximately 250
g/mol. Calculate the true molar mass of compound Z.
3.
Hemoglobin, a critical protein found in red blood cells, is 0.3466 % Fe by mass. If the
hemoglobin molecule contains 4 Fe atoms, what is the molar mass of this protein?
7
General Chemistry Practice Problems
Harvard University
Writing and Balancing Equations
1.
a)
Balance the following equations using the simplest whole-number coefficients.
CO(NH2)2 (aq) + HOCl (aq) # NCl3 (aq) + CO2 (aq) + H2O (l)
(Hint: Balance the N first)
b)
Ca3(PO4)2 (s) + SiO2 (s) + C (s) # P4 (g) + CaSiO3 (l) + CO (g)
(Hint: Balance the P first)
2.
Photosynthesis (in plants) converts carbon dioxide and water into glucose (C6H12O6)
and oxygen. Write and balance the chemical equation for photosynthesis.
3.
Niobium metal (Nb) will react with solid iodine to produce solid triniobium octaiodide.
Write and balance the equation for this process.
4.
Write and balance the chemical equation for the complete combustion of liquid octane,
C8H18 (l).
8
General Chemistry Practice Problems
Harvard University
Stoichiometry of Reactions
1.
Sodium hypochlorite, the active ingredient in Clorox, can be made by the following
reaction:
2 NaOH (aq) + Cl2 (g) # NaCl (aq) + NaClO (aq) + H2O (l)
If chlorine gas is bubbled continuously through a solution containing 60.0 g of NaOH,
how many grams of NaClO can be produced, assuming the reaction goes to completion?
2.
For the following unbalanced chemical equation:
NiS + O2 + HCl # NiCl2 + H2SO4
a)
Write the balanced chemical equation for this reaction.
b)
What mass of NiCl2 will be produced if 0.458 g of NiS reacts?
3.
Ferrous oxalate, FeC2O4, decomposes upon heating:
FeC2O4 (s) # FeO (s) + CO2 (g) + CO (g)
Suppose the mixture of gaseous products is collected, what is the percent by mass of
CO2 in this mixture?
9
General Chemistry Practice Problems
Harvard University
Stoichiometry with Limiting Reagents
1.
Vanadium (V) oxide, V2O5, can be reduced by zinc to form vanadium (II) oxide, V2O2
and zinc oxide, ZnO.
a)
Write and balance the chemical equation for this process.
b)
What mass of vanadium (II) oxide can be produced from a mixture of 100.0 grams of
V2O5 and 100.0 grams of Zn?
c)
Assuming the reaction in part (b) proceeds to completion, what mass of V2O5 is left
over unreacted? What mass of Zn is left unreacted?
2.
a)
Many binary compounds of phosphorus and sulfur have been prepared.
Balance the following chemical equation for the preparation of P4S5.
(Hint: balance the S first.)
P4S3 + Br2 # P4S5 + PBr3
b)
What is the maximum mass of P4S5 that could be prepared from 100.0 g of P4S3 and
100.0 mL of liquid bromine? (Density of bromine is 3.12 g/mL)
10
General Chemistry Practice Problems
Harvard University
Stoichiometry of Mixtures I
1.
A certain mixture of CuS and Cu2S weighs 10.80 grams total. Complete reduction of
this mixture produces 8.06 grams of pure metallic Cu. Determine the amounts of CuS
and Cu2S in the original mixture.
2.
Barium carbonate, BaCO3, and calcium carbonate, CaCO3, will decompose upon
heating, releasing carbon dioxide gas and solid barium oxide or calcium oxide,
respectively. A 10.00 gram mixture of barium carbonate and calcium carbonate is
heated until both compounds have completely decomposed. The mass of the solid
products after heating is 7.01 grams. Determine the mass of barium carbonate and the
mass of calcium carbonate in the original mixture.
11
General Chemistry Practice Problems
Harvard University
Stoichiometry of Mixtures II
1.
You are given a steel vessel containing a mixture of methane gas (CH4 (g)) and propane
gas (C3H8 (g)). You completely combust this entire mixture in the presence of excess
oxygen gas, and you collect all of the carbon dioxide and water formed in the
combustion reaction. A total of 48.4 grams of CO2 and 32.4 grams of H2O are
collected. Determine the number of moles of CH4 (g) and C3H8 (g) that were present in
the initial mixture.
12
General Chemistry Practice Problems
Harvard University
Oxidation Numbers
1.
2.
Write the oxidation numbers of each atom in the following species:
BrO3–
H3AsO3
AsH3
CrCl3
KClO3
S2O32–
OF2
Na2O2
Fe3O4
Identify whether oxidation and reduction is taking place in each of the following
equations. If so, identify which species is being reduced and which is being oxidized.
N2 + 3 H2 # 2 NH3
2 FeCl3 + 3 KI # 2 FeCl2 + KI3 + 2 KCl
2 HCl + Na2CO3 # 2 NaCl + H2O + CO2
3 Cl2 + 6 OH– # 5 Cl– + ClO3– + 3 H2O
13
General Chemistry Practice Problems
Harvard University
Solutions: Molarity
1.
A 10.0 mL sample of concentrated sulfuric acid contains 17.7 g of H2SO4. Determine
the molarity of concentrated sulfuric acid.
2.
Pure acetic acid (C2H4O2) has a density of 1.049 g/mL. Calculate the molarity of pure
(anhydrous) acetic acid.
3.
A hydrochloric acid solution is prepared by dissolving 1.97g of hydrogen chloride gas
in 27.3 mL of water and then diluting that mixture to a total volume of 250.00 mL.
Calculate the molarity of the resulting solution.
4.
The concentration of NaClO in Clorox is 0.705 M. Calculate the mass of NaClO
present in 1.0 mL of Clorox.
5.
A solution is prepared by dissolving x grams of potassium nitrate in water and diluting
to a total volume of 100.0 mL. Another solution is prepared by dissolving y grams of
sodium chloride in water and diluting to a total volume of 500.0 mL. Both solutions are
then mixed together, giving a final concentration of KNO3 of 0.073 M and a final
concentration of NaCl of 0.128 M. Calculate x and y.
14
General Chemistry Practice Problems
Harvard University
Solution Stoichiometry I
1.
The toxic compound NCl3 can be formed from the reaction of bleach (NaClO) with
ammonia (NH3):
NaClO (aq) + NH3 (aq) # NCl3 (l) + NaOH (aq)
a)
Write a complete, balanced, net ionic equation for this process.
b)
You accidentally pour 1.0 mL of Clorox into a large bucket of ammonia solution. What
mass of NCl3 can be produced if the reaction goes to completion? (Clorox is a 0.705 M
solution of NaClO in water.)
2.
A sample of calcium carbonate weighing 6.35 grams is placed in 500.0 mL of 0.31 M
hydrochloric acid and allowed to react to form calcium chloride, water, and carbon
dioxide gas. Calculate the maximum mass of carbon dioxide gas that can be produced.
15
General Chemistry Practice Problems
Harvard University
Solution Stoichiometry II: Acid/Base Neutralizations
1.
A student has dissolved 87.5 g of sodium hydroxide in 1.53 L of water. This strongly
basic solution must be neutralized before disposal. What volume of 1.27 M
hydrochloric acid would be required to completely neutralize this solution?
2.
What volume of 0.0843 M Ba(OH)2 would be required to completely neutralize a 1.00mL sample of 12.0 M acetic acid (CH3COOH)?
3.
Phosphoric acid can be neutralized by sodium hydroxide according to the equation:
H3PO4 + 3 NaOH # Na3PO4 + 3 H2O
What volume of 0.176 M NaOH would be required to neutralize 5.00 mL of
concentrated (14.8 M) phosphoric acid?
16
General Chemistry Practice Problems
Harvard University
Solution Stoichiometry III: Titrations
1.
You are given 2.00 grams of a white solid and told that it is a mixture of NaCl and
Na2CO3. You dissolve the sample in water and titrate with HCl; complete
neutralization requires 23.0 mL of 1.00 M HCl:
Na2CO3 + 2 HCl # H2CO3 + 2 NaCl
Calculate the mass of Na2CO3 in the unknown sample.
2.
A 10.0-gram sample of an unknown solid monoprotic acid is dissolved in water and
titrated with 0.789 M NaOH, requiring 40.6 mL to reach the endpoint. Determine the
molar mass of the unknown acid.
3.
HX is a monoprotic acid; it has only one acidic proton per molecule. When 1.736
grams of HX are titrated with 1.334 M NaOH, the endpoint is reached with 10.92 mL of
base. Calculate the molar mass of HX.
17
General Chemistry Practice Problems
Harvard University
Solution Stoichiometry IV: Precipitation Reactions Part 1
1.
To 250 mL of a 0.065 M solution of barium nitrate is added 100 mL of a 0.121 M
solution of sulfuric acid. Determine the concentration of aqueous barium ions
remaining in the solution after complete precipitation of barium sulfate.
2.
A 10.00-g mixture of calcium sulfate (CaSO4) and aluminum sulfate (Al2(SO4)3) is
completely dissolved in water. Excess barium nitrate is added, causing the sulfate ions
to precipitate completely as barium sulfate. If 17.50 g of barium sulfate were collected,
calculate the masses of CaSO4 and Al2(SO4)3 in the original mixture.
18
General Chemistry Practice Problems
Harvard University
Precipitation Reactions Part 2
1.
The following two solutions are mixed together:
200.0 mL of 0.300-molar lead (II) nitrate, Pb(NO3) 2
300.0 mL of 0.300-molar sodium fluoride, NaF
A white precipitate is formed in a double-displacement (exchange) reaction.
a)
Write a complete, balanced equation, and then a net ionic equation for this chemical
reaction. (Be sure to include state symbols such as (s), (aq), etc.)
b)
Calculate the mass of the lead (II) fluoride precipitate that would be formed, assuming
the reaction goes to completion, and determine the molar concentration of all ions in the
resulting solution.
19
General Chemistry Practice Problems
Harvard University
Putting It Together: Carminic Acid
1.
Carminic acid, a naturally occurring red pigment extracted from the cochineal insect,
contains only C, H, and O. It was commonly used as a dye in the first half of the
nineteenth century. Complete combustion of 10.00 g carminic acid yields 19.66 g CO2
and 3.658 g H2O.
a)
The empirical formula of carminic acid is CxHyO13. Determine the values of x and y.
Mass of C:
12.01 g C
19.66 g CO2 " 44.01 g CO
2
= 5.365 g C
Mass of H:
2.016 g H
3.658 g H2O " 18.016 g H O
= 0.4093 g H
Mass of O:
10.00 g – 5.365 g – 0.4093 g
= 4.2257 g O
2
Mass:
Moles:
Mole ratio:
Whole number ratio:
5.365 g C
0.4467
1.691
" 13 = 22 = x
0.4093 g H
0.4061
1.537
" 13 = 20 = y
4.2257 g O
0.2641
1.000
" 13 = 13
so the empirical formula is C22H20O13.
b)
The molecular formula of carminic acid is the same as the empirical formula. A
titration required 18.02 mL of 0.0406 M NaOH to neutralize 0.3602 g of carminic acid.
Determine the number of acidic hydrogen(s) in carminic acid.
20
General Chemistry Practice Problems
Harvard University
The Ideal Gas Law
1.
Chlorine (Cl2) has many uses in manufacturing and industry. You may be familiar with
its odor because it is sometimes used as a disinfectant in swimming pools. Humans can
detect the odor of chlorine when it is present at a pressure as low as 2.0 " 10–7 atm.
a)
Calculate the volume (in liters) of a classroom which is 20 meters long, 10 meters wide,
and 3.0 meters high. (Recall that 1 mL = 1 cm3).
b)
What mass of Cl2 must be released into the classroom described above at 25°C in order
for humans to detect the odor of chlorine?
2.
It requires 0.182 mol of O2 gas to exert a pressure of 1.50 atm in a rigid walled tank at
25°C. What mass of O2 would be required to exert a pressure of 17.2 atm in the same
tank at 100°C?
21
General Chemistry Practice Problems
Harvard University
Mixtures of Gases
1.
A 10.0-liter chamber contains a mixture of nitrogen and oxygen at a total pressure of
760 torr and a constant temperature of 25°C. The mole fraction of oxygen in the
mixture is 0.211.
a)
Calculate the number of moles of oxygen in the chamber.
b)
Gas is pumped out of the chamber until the total pressure is 0.100 torr. Calculate the
new partial pressure of oxygen in the chamber.
c)
Pure nitrogen gas is added to the chamber until the total pressure is again 760 torr, then
gas is pumped out of the chamber until the total pressure is 0.100 torr. Calculate the
new partial pressure of oxygen after this process.
2.
A sample of the gas butane (C4H10), of unknown mass, is contained in a vessel of
unknown volume, V, at 24.8°C and a pressure of 560.0 torr. To this vessel
8.6787 g of Ne are added in such a way that no butane escapes. The total pressure of
the vessel (at the same temperature) is 1420.0 torr. Calculate the volume of the vessel
and the mass of the butane.
22
General Chemistry Practice Problems
Harvard University
Reactions Involving Gases I
1.
A 1.25-gram sample of FeC2O4 is added to an evacuated 2.00-L steel vessel. The
vessel is heated to 400°C, at which point all the FeC2O4 is decomposed:
FeC2O4 (s) # FeO (s) + CO2 (g) + CO (g)
Calculate the total pressure inside the vessel at 400°C.
2.
Ethylene, C2H4, will react with hydrogen gas under appropriate conditions to form
ethane, C2H6. A 10.0-liter vessel is charged with 1.5 atm of hydrogen and 1.0 atm of
ethylene at 25°C. The reaction is allowed to proceed to completion. Determine the
total pressure in the vessel at 25° at the completion of the reaction.
23
General Chemistry Practice Problems
Harvard University
Reactions Involving Gases II
1.
A 2.42 gram sample of PCl5 was placed into an evacuated 2.00-L flask and allowed to
partially decompose at 250°C according to the following equation:
PCl5 (g) # PCl3 (g) + Cl2 (g)
The total pressure in the flask after partial decomposition was 359 torr. Calculate the
mole fraction of each gas in the flask.
2.
Ozone (O3) can be prepared in the laboratory by passing an electrical discharge through
a quantity of oxygen gas (O2):
3 O2 (g) # 2 O3 (g)
A 10.00-L evacuated steel vessel is filled with 32.00 atm of pure O2 at 25°C. An
electric discharge is passed through the vessel, causing some of the oxygen to be
converted into ozone. As a result, the pressure inside the vessel drops to 30.64 atm at
25°C. Calculate the final percent by mass of ozone in the vessel.
24
General Chemistry Practice Problems
Harvard University
Collecting Gases Over Water
Useful Information: Vapor pressure of water at 25°C is 23.76 torr.
1.
In a sealed container at 25°C, 1.0 atm of A completely reacts with 1.0 atm of B to form
C and water according to the following equation:
A (g) + B (g) # C (g) + H2O (l)
Once the reaction is complete, we can conclude that:
i)
The total pressure inside the container is less than 1.0 atm.
ii)
The total pressure inside the container is greater than 1.0 atm.
iii) More information is needed to determine the total pressure inside the container.
2.
Hydrogen gas is generated in the laboratory and collected over water in a vessel at
25°C. The total pressure in the vessel is 1.00 atm, and the total volume of gas collected
is 1.37 L. Calculate the mass of hydrogen collected.
3.
A 2.00-liter steel cylinder of oxygen gas has a pressure of 17.2 atm at 25°C. You want
to reduce the pressure in the cylinder to 15.5 atm, so you allow oxygen to escape from
the cylinder and collect the escaping gas over water. What volume of gas should you
collect at 25°C and 1.00 atm in order to achieve the desired pressure in the cylinder?
25
General Chemistry Practice Problems
Harvard University
Kinetic-Molecular Theory of Gases
1.
A mixture of ammonium nitrate and butane are placed into a steel cylinder and
detonated. The product gases (CO2, H2O, and N2) reach a temperature of 725°C.
a)
Calculate the rms velocity of each gas at this temperature.
b)
Calculate the average kinetic energy per mole and per molecule for each gas.
26
General Chemistry Practice Problems
Harvard University
Putting It Together: Gases
1.
Consider the reaction used to produce NaAlCl4:
2 Al2O3 (s) + 6 Cl2 (g) + 4 NaCl (l)
#
4 NaAlCl4 (l) + 3 O2 (g)
A 10.0-L evacuated reaction vessel at 25°C is filled with some Al2O3 (s), NaCl (l), and
1.00-atm chlorine gas. Once the mixture is heated to 850°C some of the chlorine gas
will react. The final total pressure inside the vessel is 2.70 atm at 850°C.
a)
Determine the final partial pressures of Cl2 and O2 in the vessel.
b)
Calculate the mass of NaAlCl4 produced in the reaction.
c)
Calculate the RMS speed of Cl2 in the vessel at 850°C.
27
General Chemistry Practice Problems
Harvard University
Heat
1.
50 g of marble chips (heat capacity 0.94 J/g K) are heated from 25°C to 200°C.
a)
How much heat is consumed in this process?
b)
The hot marble chips are placed in 500 g of cold (10°C) water. Calculate the final
temperature of the system.
2.
A coffee-cup calorimeter contains 150 g of water at 24.6°C. A 110-g piece of
molybdenum metal is heated to 100°C and placed in the water in the calorimeter. The
system reaches equilibrium at a final temperature of 28.0°C. Calculate the specific heat
capacity of molybdenum metal. (The specific heat capacity of water is 4.184 J/g°C).
28
General Chemistry Practice Problems
Harvard University
Calorimetry I
1.
Consider the dissolution of sodium carbonate:
Na2CO3 (s) # 2Na+ (aq) + CO32– (aq)
When 15.0 grams of Na2CO3 are dissolved in 100 g of water, the temperature of the solution
increases from 25.0°C to 33.0°C. Calculate the change in enthalpy for the dissolving of
sodium carbonate. Assume no heat lost to the surroundings and the solution has a specific
heat capacity of 4.18 J/g°C.
2.
A 0.6037-g sample of solid naphthalene (C10H8) is combusted in a bomb calorimeter, and the
temperature rises by 2.27°C. The heat capacity of the calorimeter and its contents is 10.69
kJ/K. Calculate "H for the combustion of naphthalene. Assume that $H # $U.
3.
When a 0.235-g sample of benzoic acid is combusted in a bomb calorimeter, a 1.642°C rise in
temperature is measured. When a 0.265-g sample of caffeine (C8H10O2N4) is combusted, a
1.525°C rise in temperature is measured. Calculate the heat of combustion per mole of
caffeine. Assume the heat of combustion of benzoic acid is 26.38 kJ/g, and that $H # $U.
29
General Chemistry Practice Problems
Harvard University
Calorimetry II
1.
Ammonium nitrate (NH4NO3) is a particularly potent explosive when mixed with a
combustible liquid like fuel oil; the Oklahoma City bomb used such a mixture. The fuel will
explode along with the ammonium nitrate.
A mixture of 1.00 g NH4NO3 and 1.00 g C4H10 (a component of petroleum fuel) are burned
in a bomb calorimeter (Ccal = 10.43 kJ/°C). The heat of detonation of ammonium nitrate is
–118 kJ/mol, while the heat of combustion of butane is –2635 kJ/mol. The initial temperature
is 22.59°C. Predict the final temperature of the system. Assume that $H # $U.
2.
Given 200 g of hot tea at 90°C, what mass of ice at –15°C must be added in order to obtain a
final temperature of 10°C?
Useful Information: specific heat capacity of ice = 2.09 J/g°C
specific heat capacity of water = 4.18 J/g°C
specific heat capacity of tea = 4.18 J/g°C
enthalpy of fusion of ice $Hfus = +6.01 kJ/mol
30
General Chemistry Practice Problems
Harvard University
Enthalpy, Energy, Heat, and Work
Useful Information:
"H°f (CH3OH (g)) = –201.2 kJ/mol
"H°f (CH3OH (l)) = –238.6 kJ/mol
1.
Consider the evaporation of one mole of methanol, CH3OH, at its boiling point (64.4°C) at a
constant pressure of 1.00 atm. The density of liquid methanol is 0.791 g/mL.
a)
Calculate "H for the evaporation of one mole of methanol at its boiling point.
b)
Calculate the work (w) for the evaporation of one mole of methanol at it boiling point.
c)
Determine "U for the evaporation of one mole of methanol at its boiling point.
31
General Chemistry Practice Problems
Harvard University
Hess's Law
1.
Given the following enthalpies of reaction:
N2O4 (g) # 2 NO2 (g)
NO (g) + 1/2 O2 (g) # NO2 (g)
"H° = 57.20 kJ/mol
"H° = –57.07 kJ/mol
Calculate the enthalpy of reaction for:
2 NO (g) + O2 (g) # N2O4 (g)
2.
Given the following enthalpies of reaction:
NH3 (g) # NH3 (aq)
HNO3 (g) # HNO3 (aq)
NH4NO3 (s) # NH4NO3 (aq)
NH3 (aq) + HNO3 (aq) # NH4NO3 (aq)
Calculate the enthalpy of reaction for:
NH3 (g) + HNO3 (g) # NH4NO3 (s)
32
"H° = –34.10 kJ/mol
"H° = –72.3 kJ/mol
"H° = 26.4 kJ/mol
"H° = –52.3 kJ/mol
General Chemistry Practice Problems
Harvard University
Standard Enthalpies of Formation
1.
Given the following information:
"H°f (CaO (s)) = –635.5 kJ/mol
"H°f (H2O (l)) = –285.83 kJ/mol
"H°f (Ca(OH)2 (s)) = –986.2 kJ/mol
Calculate "H for the following reaction:
CaO (s) + H2O (l) # Ca(OH)2 (s)
2.
Given the following data for butane (C4H10):
"H°f (C4H10 (g)) = –124.7 kJ/mol
"Hvap(C4H10 (l)) = 22.9 kJ/mol
"H°f (H2O (g)) = –241.8 kJ/mol
Density of C4H10 (l) = 0.59 g/mL
"H°f (CO2 (g)) = –393.5 kJ/mol
and the balanced equation for the combustion of liquid butane:
C4H10 (l) + 13/2 O2 (g) # 4 CO2 (g) + 5 H2O (g)
Calculate the quantity of heat released by the combustion of 10.0 mL of liquid butane (the
amount of butane in a cigarette lighter).
33
General Chemistry Practice Problems
Harvard University
Putting It Together: Methanol Fuel
1.
Methanol (CH3OH) can be used as a substitute for gasoline in certain high-performance vehicles.
Although methanol produces comparable amount of CO2 to gasoline, its oxidation produces much
less toxic carbon monoxide or nitrogen oxides (from impurities in gasoline).
Useful Information:
Cm (molar heat capacity) of liquid methanol = 81.6 J/K mol $H°f (O2 (g)) = 0 kJ/mol
$Hvap of methanol = 38 kJ/mol
$H°f (CH3OH (g)) = –200.66 kJ/mol
Density of liquid methanol = 0.791 g/mL
$H°f (CO2 (g)) = –393.51 kJ/mol
Boiling point of methanol = 65°C
$H°f (H2O (g)) = –241.82 kJ/mol
3
1 L = 1000 cm
a)
The methanol in an automobile engine must be in the gas phase before it can react. Write an
expression for and calculate the heat (in kJ) that must be added to vaporize 3.79 L (1 gallon) of
liquid methanol starting at 25°C.
Total qadd
= qliquid + qvap = nCm$T + n$Hvap
Mole of CH3OH
= 93.54 mol
From CH3OH (l, 25°C) to CH3OH (l, 65°C), find the heat added (qliquid):
b)
Gaseous methanol can be combusted with oxygen in the air according to the following equation:
2 CH3OH (g) + 3 O2 (g)
#
2 CO2 (g) + 4 H2O (g)
Calculate q for the combustion of the quantity of methanol you vaporized in part (a) with excess
oxygen gas.
$H
= –1353 kJ (… per 2 mol of CH3OH (g))
q = n$H
= " (–1353 kJ) = –6.33 $ 104 kJ
34
General Chemistry Practice Problems
Harvard University
Putting It Together: Butane Heater
1.
A 200-L home heater tank contains butane gas (C4H10) at 26.0 °C. When some of the butane
is removed and burned in excess air, the pressure in the tank drops from 2.34 atm to 1.10 atm.
The energy generated in the combustion is then used to raise the temperature of 132.5 L of
water from 19.0°C to 71.3 °C.
Useful Information:
%H°f (CO2 (g)) = –393.5 kJ/mol
%H°f (H2O (l)) = –285.8 kJ/mol
Density of water = 1 g/mL
a)
Write a balanced chemical reaction for the complete combustion of butane.
b)
Calculate the enthalpy of formation (%H°f) of butane.
35
General Chemistry Practice Problems
Harvard University
Putting It Together: Nitrosyl Chloride
1.
Nitrosyl chloride, NOCl, is a yellow gas that is very toxic and irritating to the lungs, eyes and
skin. It can be prepared by reacting nitric oxide with chlorine according to the following
balanced chemical equation.
2NO (g) + Cl2 (g)
#
2NOCl (g)
0.10 moles of NO are allowed to react with 0.08 moles of Cl2 inside a steel reaction vessel
with a movable piston that allows the enclosed sample of gas to expand and contract against a
constant external pressure of 1.0 atm. The temperature is held constant at 25°C. Assuming
the reaction goes to completion, determine q, w, and %U (in kJ) for this reaction.
Useful Information:
1/
2
N2 (g) +
1/
2
O2 (g) #
N2 (g) + O2 (g) + Cl2 (g) #
36
NO (g)
"H = +90.3 kJ/mol
2 NOCl (g)
"H = +103.4 kJ/mol
General Chemistry Practice Problems
Harvard University
Energy, Particles, and Waves
Useful Information:
Speed of light c = 3.00 " 108 m/s
Planck's constant h = 6.63 " 10–34 J·s
Mass of electron me = 9.11 " 10–31 kg
1.
A photon produced by an X-ray machine has an energy of 4.70 " 10–16 J.
a)
What is the frequency of this photon?
b)
What is the wavelength of this photon?
2.
An electron is traveling with a velocity of 2.00 " 106 m/s
a)
What is the kinetic energy of this electron?
b)
What is the deBroglie wavelength of this electron?
37
General Chemistry Practice Problems
Harvard University
Light and Matter
Useful Information:
Speed of light c = 3.00 " 108 m/s
Planck's constant h = 6.63 " 10–34 J·s
1.
The Earth receives 3800 " 1021 J in solar energy per year. Assuming that the Sun radiates at
only one (highest intensity) wavelength at 487 nm, estimate the number of photons absorbed
by the Earth from the Sun in one year.
2.
One of the largest molecule so far to have its Debroglie wavelength measured is a
perfluoroalkylated nanosphere with molecular formula C180F250. Using an interferometer,
researchers were able to detect a Debroglie wavelength of 1 picometer (= 10–12 m) when
samples of the nanosphere were vaporized and passed through an interferometer. What is the
velocity that a single molecule of the nanosphere is traveling at to have this Debroglie
wavelength?
38
General Chemistry Practice Problems
Harvard University
The Photoelectric Effect
Useful Information:
Speed of light c = 3.00 " 108 m/s
Planck's constant h = 6.63 " 10–34 J·s
Mass of electron me = 9.11 " 10–31 kg
1.
a)
The photoelectric binding energy of chromium is 7.21 " 10–19 J.
Calculate the minimum frequency of light that could produce the photoelectric effect in
chromium.
b)
If 250 nm light strikes the surface of chromium in an evacuated glass tube, determine the
kinetic energy and the deBroglie wavelength of the ejected electron.
39
General Chemistry Practice Problems
Harvard University
Single Electron Atoms
Useful Information: RH = 2.18 " 10–18 J = 13.6 eV
1.
An excited hydrogen atom emits light with a wavelength of 397.2 nm to reach the energy
level for which n = 2. In which principal quantum level did the electron begin?
2.
At the Large Hadron Collider facility on the France/Switzerland border, protons are
accelerated to a very high velocity, and their collisions with other particles are studied.
Suppose a proton with a kinetic energy of 1.18 " 1012 eV is passed through a material
composed purely of He+ atoms. As the proton travels through this material it interacts with a
number of He+ atoms and transfers a portion of its energy to the helium electron. Suppose that
during each interaction the proton loses just enough energy to ionize the ground state electron
of He+. Calculate the number of collisions that would need to take place in order to stop the
proton.
40
General Chemistry Practice Problems
Harvard University
Orbitals and Quantum Numbers
1.
For each of the following subshells, identify the n and l quantum numbers, and state the total
number of electrons that could be found in each subshell.
Subshell
n
l
Total number of electrons
1s
3d
5p
2.
For each of the following orbitals, provide a perspective sketch of the orbital in the Cartesian
x, y, z coordinate system. Determine the number of angular nodes and the number of radial
nodes in the orbital, and show/describe where those nodes fall.
a)
3s
b)
4py
c)
4dx2-y2
41
General Chemistry Practice Problems
Harvard University
Electron Configurations of Neutral Atoms
1.
Write the electron configurations of the following neutral atoms. You may use the noble-gas
abbreviations (i.e. [Ar] . . . ).
C:
Mg:
Mn:
Se:
Cu:
Xe:
Ba:
Os:
Pb:
42
General Chemistry Practice Problems
Harvard University
Electron Configurations of Ions
1.
Write the ground state electron configurations for the following ions. Use the noble-gas
abbreviations if appropriate.
N– :
Al3+ :
O2– :
Zn2+ :
W6+ :
Cu2+ :
Gd3+ :
Se2– :
Fe2+:
Fe3+:
43
General Chemistry Practice Problems
Harvard University
Periodic Properties I
1.
Indicate “high” and “low” areas for each property on the mini periodic tables below.
Explain each trend briefly.
a)
Atomic Mass
b)
Atomic Radius
c)
Ionization Energy
d)
Electron Affinity
e)
Metallic Character
44
General Chemistry Practice Problems
Harvard University
Periodic Properties II
1.
Circle the element from each set with the largest atomic radius. Explain your choices.
a)
Ba
Ti
Ra
Li
b)
F
Al
In
As
2.
Circle the element from each set with the smallest ionization energy. Explain your choices.
a)
Tl
Po
Se
Ga
b)
Cs
Ga
Bi
Se
3.
Circle the element with the most negative electron affinity. Explain your choice.
Be
4.
N
O
F
Circle the ion with the largest radius. Explain your choice.
Se2–
F–
O2–
45
Rb+
General Chemistry Practice Problems
Harvard University
Periodic Properties III
1.
Provide brief explanations for the following observations:
a)
The first ionization energy of Se is less than the first ionization energy of As.
b)
More energy is released upon adding an electron to Br than upon adding one to Se.
c)
The reaction of Rb with water is much more violent that the reaction of Na with water.
d)
The electron affinity of fluorine is more negative than that of oxygen.
e)
In fact, the electron affinity of fluorine (–332 kJ/mol) is among the most negative of any of
the elements. This very negative electron affinity makes fluorine extremely reactive.
f)
The first ionization energy of xenon is 1170 kJ/mol. This substantial ionization energy helps
to make xenon very unreactive.
g)
The first ionization energy of iodine, xenon’s neighbor on the periodic table, is also quite high
(1020 kJ/mol). Iodine is quite reactive, however, unlike xenon.
46
General Chemistry Practice Problems
Harvard University
Putting It Together: Bismuth
1.
a)
Consider the electronic structure of the element bismuth (Bi).
The first ionization energy of bismuth is 703 kJ/mol. Calculate the longest possible
wavelength of light that could ionize an atom of bismuth.
b)
Write the electron configurations of neutral Bi and the Bi+ using the noble gas abbreviations.
Bi :
+
Bi :
c)
What are the n and l quantum numbers of the electron removed when Bi is ionized to
Bi+?
n=
l=
d)
Sketch an atomic orbital from which the electrons are removed from Bi to form Bi+. Clearly
show the shape and orientation of the orbital. Indicate the location of any radial and/or
angular nodes.
e)
Would you expect Element 113 to have an ionization energy that is greater than, equal to, or
less than that of bismuth?
(circle one):
greater than
equal to
less than
Explain briefly. (It is not enough to simply state a trend; you must provide an explanation
for the trend in terms of atomic structure.)
47
General Chemistry Practice Problems
Harvard University
Lewis Structures I: The Octet Rule
1.
For each of the following molecules, draw the best possible Lewis structure. Include all nonzero formal charges, and indicate resonance if appropriate.
CF4
OF2
HCN
CO
NSF
HNO3
SF3+
C22–
NO2–
PH2–
CO32–
BH4–
48
General Chemistry Practice Problems
Harvard University
Lewis Structures II: Less Than an Octet
1.
For each of the following molecules or ions, draw the best possible Lewis structure. Include
all non-zero formal charges, and indicate resonance if appropriate.
BBr3
BeH2
AlF3
NO
NO2
49
General Chemistry Practice Problems
Harvard University
Lewis Structures III: More Than an Octet
1.
For each of the following molecules or ions, draw the best possible Lewis structure. Include
all non-zero formal charges, and indicate resonance if appropriate.
SeF4
XeF2
AsCl5
AsF6–
XeF3+
BrF2–
50
General Chemistry Practice Problems
Harvard University
Bond Enthalpies
Average bond enthalpies (in kJ/mol):
C–N
C=N
C%N
O–O
293
615
891
146
C–O
C=O
C%O
O=O
358
799
1072
495
C–C 348
C–Cl 328
Cl–Cl 243
N–N 163
N=N 418
N%N 941
1.
a)
Cyanogen, NCCN, is a highly toxic gas.
Draw a correct Lewis structure for cyanogen
b)
Estimate the change in enthalpy for the complete combustion of cyanogen:
NCCN (g) + 2 O2 (g) # N2 (g) + 2 CO2 (g)
2.
Phosgene (COCl2), a substance used in poisonous gas warfare in WWI, is so named because
it was first prepared by the action of sunlight on a mixture of carbon monoxide and chlorine
gases. Its name comes from the Greek words phos (light) and genes (born of).
a)
Draw the best possible Lewis structures of CO, Cl2, and COCl2. Include any non-zero formal
charge and resonance structures if appropriate.
b)
Estimate the enthalpy of formation ("H°f) for gaseous phosgene.
"Hsub (C (s)) = 718 kJ/mol
51
General Chemistry Practice Problems
Harvard University
Lattice Energy: The Born-Haber Cycle
1.
Construct a Born-Haber cycle for FeCl3 and determine its lattice energy.
IEFe (1,2,3) = 758, 1558, and 2952 kJ/mol
"Hsub (Fe) = 415.5 kJ/mol
EACl = –349 kJ/mol
DCl–Cl = 243.4 kJ/mol
"H°f (FeCl3) = –401 kJ/mol
2.
Construct a Born-Haber cycle for the ionic solid barium chlorate Ba(ClO3)2 and determine its
lattice energy. Assume that the bonds between Cl and O in this compound are single bonds.
"H°f (Ba(ClO3)2 (s)) = –772 kJ/mol
"Hsub (Ba (s)) = 180 kJ/mol
DCl–Cl = 243 kJ/mol
DCl–O = 272 kJ/mol
"H°f
IE1 (Ba) = 503 kJ/mol
IE2 (Ba) = 965 kJ/mol
EA (ClO3) = –252 kJ/mol
DO=O = 498 kJ/mol
= "Hsub(Ba(s)) + DCl–Cl + 3 DO=O – 6 DCl–O + IE1 + IE2 + 2EA(ClO3) – LE
52
General Chemistry Practice Problems
Harvard University
Molecular Geometry I: Neutral Molecules
1.
For each of these molecules, draw the best possible Lewis structure. Based on the
Lewis structure, provide the electron-pair geometry, the molecular geometry, and the
hybridization of the central atom. Lastly, indicate whether the molecule is polar.
CO2
ONF
e– pair geometry:
e– pair geometry:
molecular geometry:
molecular geometry:
hybridization:
hybridization:
polar?
yes
no
polar?
yes
BF3
ICl3
e– pair geometry:
e– pair geometry:
molecular geometry:
molecular geometry:
hybridization:
hybridization:
polar?
yes
no
polar?
53
yes
no
no
General Chemistry Practice Problems
Harvard University
Molecular Geometry II: Ions
1.
For each of the following ions, draw the best possible Lewis structure. Based on the
Lewis structure, provide the electron-pair geometry, the molecular geometry, and the
hybridization of the central atom.
IF4–
PCl4+
e– pair geometry:
e– pair geometry:
molecular geometry:
molecular geometry:
hybridization:
hybridization:
SeO32–
I3–
e– pair geometry:
e– pair geometry:
molecular geometry:
molecular geometry:
hybridization:
hybridization:
54
General Chemistry Practice Problems
Harvard University
Molecular Geometry III: Polycentric Molecules and Ions
1.
For each of the following molecules, draw the best possible Lewis structure. On the
basis of that Lewis structure, predict the electron-pair geometry, the molecular
geometry, and the hybridization of every non-terminal atom.
a)
F3S—S—F
b)
CH3—CH=CH—CO2–
c)
I5– (shaped like a big “V”)
55
General Chemistry Practice Problems
Harvard University
Sigma and Pi Bonding I
1.
a)
Draw the best possible Lewis structure for formaldehyde (H2CO) and determine the
geometry and hybridization of the carbon atom.
Describe each of the &- and !-bonds in the molecule. (e.g. “There are two &-bonds
between H s orbitals and C sp2 hybrid orbitals”)
c)
Imagine that formaldehyde is positioned in the y,z plane as shown below. On this
framework, draw a perspective sketch of the formaldehyde molecule that indicate the
location and orientation of the sigma bonding and the pi bonding. Label them as sigma
(&) or pi (!) bonds, and indicate the orbitals involved in forming each bond (e.g. s, p,
sp2, etc.).
x-axis
b)
H
C
O
z-axis
H
xis
y-a
2.
a)
Draw the best possible Lewis structure of CO2. Determine the molecular geometry and
the hybridization for the central atom.
b)
What carbon hybrid orbital(s) are participating in &-bonding in CO2?
c)
What carbon orbital(s) are participating in !-bonding in CO2?
56
General Chemistry Practice Problems
Harvard University
Sigma and Pi Bonding II
1.
a)
b)
The molecule allene, H2C=C=CH2, is similar to CO2 in several respects. Draw a
complete Lewis structure for allene and determine the hybridization of each carbon
atom.
Allene is not planar; the two ends of the molecule are in different perpendicular planes.
On the provided framework, draw a perspective sketch of the allene molecule that
indicate the location and orientation of the pi bonding orbitals, showing clearly why the
two ends of the molecule must be perpendicular to one another.
H
H
C
C
C
H
H
2.
a)
b)
Draw a complete Lewis structure for tetrolate ion (shown below) and determine the
hybridization of each carbon atom.
CH3–C%C–CO2–
Provide a sketch of the tetrolate molecule that indicate the location and orientation of
the sigma bonding and the pi bonding. Label them as sigma (&) or pi (!) bonds, and
indicate the orbitals involved in forming each bond (e.g. s, p, sp2, etc.).
57
General Chemistry Practice Problems
Harvard University
Molecular Orbital Theory: Orbital Overlap
1.
For each of the following pairs of atomic orbitals, draw an image and provide the name
(e.g. “!*2p”) of the molecular orbital that would result from the combination of those
two orbitals. Be sure to clearly indicate the nodes where appropriate.
Note: The shaded and unshaded atomic orbitals represent the positive wavefunction and the negative
wavefunction, respectively.
Atomic orbitals
Molecular orbitals
a)
+
1s
1s
&*1s
b)
+
2p
2p
&2p
c)
+
2p
2p
!2p
d)
+
2p
2p
!'2p
58
General Chemistry Practice Problems
Harvard University
MO Diagrams: Homonuclear Diatomic Molecules
1.
a)
Consider two forms of oxygen: diatomic oxygen (O2) and the peroxide anion (O22–).
Draw an energy-level diagram for the valence molecular orbitals of O2. Include the
relevant atomic orbitals in your diagram. Clearly label each molecular orbital (e.g &2s,
!2p, etc.) and fill in the appropriate number of electrons.
b)
Write the molecular orbital electron configuration (i.e. (&2s)2 etc.) for O2 and O22–.
c)
Will O2 be attracted by a magnetic field?
d)
Do you expect O22– to have a stronger or weaker bond than the bond in O2? Briefly
explain your reasoning.
YES
NO
(circle one)
O22– will have a weaker bond because there are two more electrons in the antibonding
orbital.
e)
Based on your diagram in part (a), draw the geometry of the lowest unoccupied
molecular orbital (LUMO) and the highest occupied molecular orbital (HOMO) for O2.
59
General Chemistry Practice Problems
Harvard University
MO Diagrams: Heteronuclear Diatomic Molecules
1.
a)
b)
Consider the bonding in NO from the viewpoint of molecular orbital theory.
Draw an energy-level diagram for the valence molecular orbitals of NO. Include the
relevant atomic orbitals in your diagram. Clearly label each molecular orbital (e.g &2s,
!2p, etc.) and fill in the appropriate number of electrons. Use the same energy order as
the neutral O2 molecule.
The NO molecule can lose or gain an electron to form NO+ or NO–, respectively. Circle
the single best choice for each of the following questions.
Has a double bond
NO
NO+
NO–
Is diamagnetic
NO
NO+
NO–
Has the longest bond of the three species
NO
NO+
NO–
Is isoelectronic with CO
NO
NO+
NO–
60
General Chemistry Practice Problems
Harvard University
Coordination Complexes: Introduction
1.
Provide the transition metal oxidation state, the number of d electrons, and the correct
name for each of the following:
a)
[Pt(NH3)4](ClO4)2
b)
[ReH9]2–
(Hint: What is the ox. state of H?)
c)
Na2[Fe(CO)4]
d)
[IrF6]3–
e)
K3[Fe(CN)6]
f)
[PtCl2(NH3)2]
61
General Chemistry Practice Problems
Harvard University
Geometry and Isomerism I: Monodentate Ligands
1.
Each of the following complexes contains a metal atom in an octahedral geometry. For
each complex, draw all the possible stereoisomers. Indicate if any of the isomers are
chiral, and be sure to draw the enantiomer of any chiral isomers.
a)
[Co(NH3)3(H2O)3]3+
b)
[Pt(NH3)2Cl2Br2]
2.
When the octahedral complex [Co(NO2)3(NH3)3] is treated with HCl, one can isolate a
complex [CoCl2(NH3)3(H2O)]+ in which the two chloride ligands are trans to one
another.
a)
Draw the two possible stereoisomers of the starting material [Co(NO2)3(NH3)3].
(The NO2– ligands are all bound through the nitrogen atom.)
b)
Assuming that the NH3 groups remain in place, which of the two starting isomers could
give rise to the observed product?
62
General Chemistry Practice Problems
Harvard University
Geometry and Isomerism II: Chelates
1.
Consider the octahedral complex [Co(Cl)2(H2O)2(en)]+,
where en = ethylenediammine = H2NCH2CH2NH2.
Draw all isomers of this complex. For each isomer, indicate if it is chiral or achiral, and
draw the enantiomer of any chiral complexes. There are 6 or fewer isomers in total.
(You can represent ethylenediammine in your drawings as: N
63
N)
General Chemistry Practice Problems
Harvard University
Electronic Structure of Coordination Compounds
1.
For each of the following, give the electronic configuration of the d-orbitals (for
example t2g3eg1), indicate whether the complex is high-spin or low-spin (or neither),
and give the number of unpaired electrons.
a)
[Fe(H2O)6]2+
b)
[Ni(H2O)6]2+
c)
[Fe(CN)6]4 –
d)
W(CO)6
(The CO ligand is similar to CN–)
2.
The three complexes [Co(NH3)6]3+, [Co(NH3)5(H2O)]3+, and [CoF6]3– are colored red,
yellow, and blue (but not necessarily in that order). Match each complex with its color,
and explain your reasoning.
3.
Explain why [FeF6]3– is basically colorless whereas [CoF6]3– is colored.
64
General Chemistry Practice Problems
Harvard University
Putting It Together: Coordination Compounds
1.
The drug Nipride, Na2[Fe(CN)5NO], is an inorganic complex that is used as a source of
NO to lower blood pressure during surgery.
a)
Provide a systematic name for the compound Na2[Fe(CN)5NO]. Note: The NO ligand
in this complex is neutral and is named “nitrosyl.”
b)
For the complex [Fe(CN)5NO]2–, give the electronic configuration of the d orbitals (for
example t2g3eg1), indicate whether the complex is high-spin or low-spin (or neither),
and give the number of unpaired electrons.
c)
The [Fe(CN)5NO]2– complex is red-violet in color (“ruby red”), while the similar
complex [Fe(CN)6]3– is red-orange in color (“bright red”).
The CN– ligand is a strong-field ligand. By comparison, the NO ligand is: (circle one)
much weaker
slightly weaker
slightly stronger
Briefly explain your choice using the language of crystal field theory.
65
much stronger