Download Chem 1 Worksheets WSHEET 1: Working with Numbers Practice

Chem 1 Worksheets
WSHEET 1: Working with Numbers Practice
1.
Rounding: round off the following numbers to the number of significant figures (sf) indicated.
a. 7.542 to 3 sf________
b. 16.365 to 3 sf________
c. 84.995 to 2 sf________
d. 6.02501 to 2 sf________
2.
Which of the following are exact numbers and which are measured numbers?
a. a baby weighs 1.75 kg________
b. a baker’s dozen________
c. a 5.0 L injection________
d. 78 students________
3.
Determine the number of significant figures in each of the following:
a. 0.00376________
b. 1.00376________
c. 43,000________
d. 14.05________
e. 14.00________
f. 3200.0________
4.
Do the calculations and then round off your answers to the correct number of significant
figures.
a. (5.321)(4.2)/(457) = ________ b. 24.31 – 3.5 = ________
c. 32.1 – 0.0035 = _______
d. 145.357 + 22.5 = ________
5.
Express the following in proper exponential notation:
a. 1512 = __________
b. 0.00529 = __________
c. 21.52x10-4 = __________
6.
Express the following as ordinary numbers:
a. 1.42x105 = ________
b. 2.069x10-3 = ________
7.
Come up to see me and then carry out the following without using your calculator: Get my
initials here:_______
a. 104x105 = _____
b. 104x10-5 = _____
c. 104/105 = _____
8.
Do the calculations, round off your answers to the correct number of significant figures, and
express the answer in proper exponential notation.
a. 2.23x107 * 3.0x10-4 = ________ b. 2.21x107/5.500x10-4 = ________
c. (2.25x107)1/2 = ________
d. 2.2x102 + 3.13x103 = ________
e. 7.63x10-2 – 3.15x10-4 = ________
9.
What would be the difference in the number of significant figures and in how the numbers were
expressed in exponential notation, given the following two pieces of data: 250 people vs. 250.
mL.
Apply the rules of significant figures in doing all these calculations.
10.
Convert each of the following to meters:
a. 2.30x1012 nm = __________ b. 3.6x10-4 km = __________
c. 100.0 yards = __________
11.
Convert each of the following to liters:
a. 1.60 mL = ________
b. 239.0 cm3 = ________
c. 1.00 gal = ________
12.
Convert the following temperatures as indicated:
a. 19.00oF to ______oC
b. 286.55 K to ________oC
o
c. –118 C to _______K
d. –40.00oC to ________oF
13.
In each case below, determine which is the lower temperature by putting them into the same
units of temperature.
a. 0oC or 0oF?
b. 0oC or 0 K?
14.
Calculate the volume of a sample of mercury with a density of 13.6 g/mL and a mass of 1.00 g.
________
A-1
Practice with multiple choice questions and Chp 2:WSHEET 2
Instructions: First work individually and show all work for the following, then select the best answer. Show
me your work and answers, then work with your group to see how well you all did and try to achieve agreement
on all the answers. The whole group should show me their work and answers when you have reached an
agreement.
1.
The density of a silver coin can be calculated from the following data:
Mass of silver coin
6.581 g
Volume of coin and water
23.7 mL
Volume of water alone
23.1 mL
The density of the coin should be reported as:
a: 10.847 g/mL
b. 10.85 g/mL
c. 10.8 g/mL d. 11 g/mL
e. 1x101 g/mL
2.
Copper makes up 1.1x10-4 percent by mass of a normal healthy human being. How many grams of
copper would be found in the body of a person weighing 150. Lb? (1.000 lb = 453.6 g)
a. 0.00075 g b. 0.075 g
c. 0.75 g
d. 7.5 g
e. 36 g
3.
How many TOTAL atoms are in 12 molecules of glucose, C6H12O6?
a. 24
b. 288
c. 2160
d. none of these
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A-2
Chem 1, Chp # 2 WSHEET 3( No Nomenculature).
1. Kaolinite, a clay mineral with the formula Al4Si4O10(OH)8, is used as a filler in slick-paper for magazines and
as a raw material for ceramics. Analysis shows that 14.35 g of kaolinite contains 8.009 g of oxygen. Calculate
the mass percent of oxygen in kaolinite.
A. 1.792 mass %
B. 24.80 mass %
C. 30.81 mass % D. 34.12 mass %
E. 55.81 mass % Work:
2. J. J. Thomson studied cathode ray particles (electrons) and was able to measure the mass/charge ratio. His
results showed that
A. the mass/charge ratio varied as the cathode material was changed.
B. the charge was always a whole-number multiple of some minimum charge.
C. matter included particles much smaller than the atom.
D. atoms contained dense areas of positive charge.
E. atoms are largely empty space.
3. Millikan's oil-drop experiment
A. established the charge on an electron.
B. showed that all oil drops carried the same charge.
C. provided support for the nuclear model of the atom.
D. suggested that some oil drops carried fractional numbers of electrons.
E. suggested the presence of a neutral particle in the atom.
4. Rutherford bombarded gold foil with alpha () particles and found that a small percentage of the particles
were deflected. Which of the following was not accounted for by the model he proposed for the structure of
atoms?
A. the small size of the nucleus
B. the charge on the nucleus
C. the total mass of the atom
D. the existence of protons
E. the presence of electrons outside the nucleus
5. Atoms X, Y, Z, and R have the following nuclear compositions:
Which two are isotopes?
A. X & Y
B. X & R
C. Y & R
D. Z & R
E. X & Z
6. Silicon, which makes up about 25% of Earth's crust by mass, is used widely in the modern electronics
industry. It has three naturally occurring isotopes, 28Si, 29Si, and 30Si. Calculate the atomic mass of silicon.
Isotope Isotopic Mass (amu)
Abundance %
28Si
27.976927
92.23
29Si
28.976495
4.67
30Si
29.973770
3.10
A-3
A. 29.2252 amu
B. 28.9757 amu
C. 28.7260 amu
D. 28.0855 amu
E. 27.9801 amu
Work:
7. Which of the following is a metalloid?
A. carbon, C, Z = 6
B. sulfur, S, Z = 16
C. germanium, Ge, Z = 32
D. iridium, Z = 77
E. bromine, Br, Z = 35
8. A row of the periodic table is called a
A. group B. period
C. isotopic mixture
D. family E. subshell
9. Which of the following ions occurs commonly?
A. P3+
B. Br7+
C. O6+
D. Ca2+
E. K10. Give the common name of the group in the periodic table to which each of the following elements belongs:
a. Rb ______________________ b. Br ________________________
c. Ba ________________________d. Ar ___________________________
Symbol
# protons
# neutrons
# electrons

17
18

Au

118



20
20
10. Fill in the blank spaces and write out all the symbols in the left hand column in full, in the form
include the appropriate values of Z and A as well as the correct symbol X).
A-4
(i.e.,
Chem 1: Chp # 3 WSHEET:4. Name: ______________________________
Show all formulas and working for problems given below.
1. Calcium fluoride, CaF2, is a source of fluorine and is used to fluoridate drinking water. Calculate its molar
mass.
2. Calculate the molar mass of Ca(BO2)2·6H2O.
3. Calculate the number of moles in 17.8 g of the antacid magnesium hydroxide, Mg(OH)2.
4. Calculate the number of oxygen atoms in 29.34 g of sodium sulfate, Na2SO4.
5. Calculate the mass in grams of 8.35  1022 molecules of CBr4.
6. Household sugar, sucrose, has the molecular formula C12H22O11. What is the % of carbon in sucrose, by
mass?
7. Terephthalic acid, used in the production of polyester fibers and films, is composed of carbon, hydrogen, and
oxygen. When 0.6943 g of terephthalic acid was subjected to combustion analysis it produced 1.471 g CO2 and
0.226 g H2O. What is its empirical formula?
8. Balance the following equation for the combustion of benzene:
C6H6(l) + O2(g)  H2O(g) + CO2(g)
9. Aluminum will react with bromine to form aluminum bromide (used as an acid catalyst in organic synthesis).
Al(s) + Br2(l)  Al2Br6(s) [unbalanced]
How many moles of Al are needed to form 2.43 mol of Al2Br6?
10. Magnesium reacts with iron(III) chloride to form magnesium chloride (which can be used in fireproofing
wood and in disinfectants) and iron.
3Mg(s) + 2FeCl3(s)  3MgCl2(s) + 2Fe(s)
A mixture of 41.0 g of magnesium (
= 24.31 g/mol) and 175 g of iron(III) chloride (
allowed to react. What mass of magnesium chloride = 95.21 g/mol) is formed?
A-5
= 162.2 g/mol) is
11. Tetraphosphorus hexaoxide (
= 219.9 g/mol) is formed by the reaction of phosphorus with oxygen gas.
P4(s) + 3O2(g)  P4O6(s)
If a mixture of 75.3 g of phosphorus and 38.7 g of oxygen produce 43.3 g of P4O6, what is the percent yield for
the reaction?
12. A 0.150 M sodium chloride solution is referred to as a physiological saline solution because it has the same
concentration of salts as normal human blood. Calculate the mass of solute needed to prepare 275.0 mL of a
physiological saline solution.
13. When 2.61 g of solid Na2CO3 is dissolved in sufficient water to make 250 mL of solution what is the
concentration of Na2CO3
14. How many milliliters of 1.58 M HCl are needed to react completely with 23.2 g of NaHCO3 (
g/mol)?
= 84.02
HCl(aq) + NaHCO3(s)  NaCl(s) + H2O(l) + CO2(g)
15. Analysis of a white solid produced in a reaction between chlorine and phosphorus showed that it contained
77.44% chlorine and 22.56% phosphorus. What is its empirical formula?
16. Consider the balanced equation:
Al2S3(s) + 6H2O(l)  2Al(OH)3(s) + 3H2S(g)
If 15.0g of aluminum sulfide and 10.0g of water are allowed to react as above, and assuming a complete
reaction
a. by calculation, find out which is the limiting reagent.
b. calculate the maximum mass of H2S which can be formed from these reagents.
c. calculate the mass of excess reagent remaining after the reaction is complete.
A-6
Chem 1 Wsheet 5: Chp # 5: Gas Laws
Name: ___________________________________________________________________________
1. Hydrogen gas exerts a pressure of 466 torr in a container. What is this pressure in atmospheres?
A. 0.217 atm
B. 0.466 atm
C. 0.613 atm
D. 1.63 atm
E. 4.60 atm
2. The air pressure in a volleyball is 75 psi. What is this pressure in torr?
A. 520 torr
B. 562 torr
C. 3900 torr D. 7600 torr E. 75,000 torr
3. "The pressure of an ideal gas is inversely proportional to its volume at constant temperature and number of
moles" is a statement of __________________ Law.
A. Charles' B. Boyle's
C. Amontons'
D. Avogadro's
E. Gay-Lussac's
4. "The rate of effusion of a gas is inversely proportional to the square root of its molar mass" is a statement of
______________________ Law.
A. Charles' B. Graham's
C. Dalton's
D. Avogadro's E. Boyle's
5. A sample of an ideal gas has its volume doubled while its temperature remains constant. If the original
pressure was 100 torr, what is the new pressure?
A. 10 torr
B. 50 torr
C. 100 torr
D. 200 torr
E. 1000 torr
6. A sample container of carbon monoxide occupies a volume of 435 mL at a pressure of 785 torr and a
temperature of 298 K. What would its temperature be if the volume were changed to 265 mL at a pressure of
785 torr?
A. 182 K B. 298 K
C. 387 K
D. 489 K
E. 538 K
7. A 750-mL sample of hydrogen exerts a pressure of 822 torr at 325 K. What pressure does it exert if the
temperature is raised to 475 K at constant volume?
A. 188 torr
B. 562 torr
C. 1.11  103 torr
D. 1.20  103 torr E. 1.90  103 torr
1
8. A sample of propane, a component of LP gas, has a volume of 35.3 L at 315 K and 922 torr. What is its
volume at STP?
A. 25.2 L
B. 30.6 L
C. 33.6 L
D. 37.1 L
E. 49.2 L
9. A carbon dioxide sample weighing 44.0 g occupies 32.68 L at 65C and 645 torr. What is its volume at STP?
A. 22.4 L
B. 31.1 L
C. 34.3 L
D. 35.2 L
E. 47.7 L
Work:
10. What is the density of carbon dioxide gas at -25.2C and 98.0 kPa?
A. 0.232 g/L
B. 0.279 g/L
C. 0.994 g/L
D. 1.74 g/L
E. 2.09 g/L
11. A 250.0-mL sample of ammonia, NH3(g), exerts a pressure of 833 torr at 42.4C. What mass of ammonia is
in the container?
A. 0.0787 g B. 0.180 g
C. 8.04 g
D. 17.0 g
E. 59.8 g
12. Magnesium metal (0.100 mol) and a volume of aqueous hydrochloric acid that contains 0.500 mol of HCl
are combined and react to completion. How many liters of hydrogen gas, measured at STP, are produced?
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g) Work:
A. 2.24 L of H2
B. 4.48 L of H2
C. 5.60 L of H2
D. 11.2 L of H2
E. 22.4 L of H2
13. Methane, CH4(g), reacts with steam to give synthesis gas, a mixture of carbon monoxide and hydrogen,
which is used as starting material for the synthesis of a number of organic and inorganic compounds.
CH4(g) + H2O(g)  CO(g) + H2(g) [unbalanced]
What mass of hydrogen is formed if 275 L of methane (measured at STP) is converted to synthesis gas?
A. 12.3 g
B. 24.7 g C. 37.1 g D. 49.4 g E. 74.2 g
2
14. Calculate the density in g/L of gaseous SF6 at 50.0C and 650. torr.
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Chem 1: Chp # 6 WSHEET 6
Name: __________________________________________________
1. A system that does no work but which transfers heat to the surroundings has
A. q < 0, E > 0 B. q < 0, E < 0 C. q > 0, E > 0
D. q > 0, E < 0 E. q < 0, E = 0
2. A system receives 575 J of heat and delivers 425 J of work. Calculate the change in the internal energy, E,
of the system.
A. -150 J B. 150 J
C. -l000 J
D. 1000 J
E. 575 J
3. A system initially has an internal energy E of 501 J. It undergoes a process during which it releases 111 J of
heat energy to the surroundings, and does work of 222 J. What is the final energy of the system, in J?
A. 168 J B. 390 J
C. 612 J
D. 834 J E. It cannot be calculated without more information
4. A system contracts from an initial volume of 15.0 L to a final volume of 10.0 L under a constant external
pressure of 0.800 atm. The value of w, in J, is
A. -4.0 J
B. 4.0 J
C. -405 J D. 405 J E. 4.05  103 J
5. Cold packs, whose temperatures are lowered when ammonium nitrate dissolves in water, are carried by
athletic trainers when transporting ice is not possible. Which of the following is true of this reaction?
A. H < 0, process is exothermic B. H > 0, process is exothermic
C. H < 0, process is endothermic D. H > 0, process is endothermic
E. H = 0, since cold packs are sealed
6. A Snickers® candy bar contains 280 Calories, of which the fat content accounts for 120 Calories. What is
the energy of the fat content, in kJ?
A. 5.0  10-1 kJ B. 29 kJ
C. 5.0  102 kJ
D. 1.2  103 kJ
E. 5.0  105 kJ
7. Natural gas, or methane, is an important fuel. Combustion of one mole of methane releases 802.3 kilojoules
of energy. How much energy does that represent in kilocalories?
A. 1.92  10-1 kcal
B. 1.92  102 kcal
C. 3.36  103 kcal
D. 1.92  105 kcal
6
E. 3.36  10 kcal
3
8. Calculate q when 28.6 g of water is heated from 22.0C to 78.3C.
A. 0.385 kJ
B. 1.61 kJ
C. 6.74 kJ
D. 9.37 kJ
E. 1.61  103 kJ
9. A 275-g sample of nickel at l00.0C is placed in 100.0 mL of water at 22.0C. What is the final temperature
of the water? Assume that no heat is lost to or gained from the surroundings. Specific heat capacity of nickel =
0.444 J/(g·K)
A. 39.6C B. 40.8C C. 61.0C D. 79.2C E. 82.4C
Work:
10. A piece of copper metal is initially at 100.0C. It is dropped into a coffee cup calorimeter containing 50.0 g
of water at a temperature of 20.0C. After stirring, the final temperature of both copper and water is
25.0C. Assuming no heat losses, and that the specific heat (capacity) of water is 4.18 J/(g·K), what is the heat
capacity of the copper in J/K?
A. 2.79 J/K B. 3.33 J/K C. 13.9 J/K
D. 209 J/K E. None of these choices is correct.
11. Sand is converted to pure silicon in a three step process. The third step is
SiCl4(g) + 2Mg(s)  2MgCl2(s) + Si(s)
H = -625.6 kJ
What is the enthalpy change when 25.0 mol of silicon tetrachloride is converted to elemental silicon?
A. -25.0 kJ B. -7820 kJ C. -1.56  l04 kJ D. -3.13  104 kJ E. None of these choices is correct.
12. The highly exothermic thermite reaction, in which aluminum reduces iron(III) oxide to elemental iron, has
been used by railroad repair crews to weld rails together.
2Al(s) + Fe2O3(s)  2Fe(s) + Al2O3(s)
H = -850 kJ
What mass of iron is formed when 725 kJ of heat are released?
A. 47 g B. 65 g C. 95 g D. 112 g E. 130 g
13. Use Hess's Law to calculate the enthalpy change for the reaction
WO3(s) + 3H2(g)  W(s) + 3H2O(g)
from the following data:
2W(s) + 3O2(g)  2WO3(s)
2H2(g) + O2(g)  2H2O(g)
H = -1685.4 kJ
H = -477.84 kJ
A. 125.9 kJ B. 252.9 KJ C. 364.9 kJ D. 1207.6 kJ E. None of these choices is correct.
14. Use the following data to calculate the standard heat (enthalpy) of formation, Hf , of manganese(IV)
oxide, MnO2 (s).
4
2MnO2(s)  2MnO(s) + O2(g)
MnO2(s) + Mn(s)  2MnO(s)
A. -504 kJ
B. -372 kJ
C. -24 kJ
H = 264 kJ
H = -240 kJ
D. 24 kJ
E. 504 kJ
15. Calculate the Hrxn for the decomposition of calcium carbonate to calcium oxide and carbon dioxide.
Hf [CaCO3(s)] = -1206.9 kJ/mol; Hf [CaO(s)] = -635.1 kJ/mol; Hf [CO2(g)] = -393.5 kJ/mol
CaCO3(s)  CaO(s) + CO2(g)
A. -2235.5 kJ B. -1448.5 kJ C. -178.3 kJ D. 178.3 kJ E. 2235.5 kJ
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Chem 1, Chp # 7 wsheet. WSHEET 7 Name:_____________________
1. Who proposed a model that successfully explained the photoelectric effect? __________________
2. Who developed an empirical equation from which the wavelengths of lines in the spectrum of hydrogen
atoms can be calculated? ____________________________
3. Which scientist first proposed that particles of matter could have wave properties? _____________
4. Who proposed the principle which states that one cannot simultaneously know the exact position and
velocity of a particle? _________________
5. Select the arrangement of electromagnetic radiation which starts with the lowest energy and increases to
greatest energy.
A. radio, visible, infrared, ultraviolet B. infrared, visible, ultraviolet, microwave
C. visible, ultraviolet, infrared, gamma rays D. X-radiation, visible, infrared, microwave
E. microwave, infrared, visible, ultraviolet
6. Electromagnetic radiation of 500 nm wavelength lies in the ________ region of the spectrum.
A. infrared B. visible C. ultraviolet D. X-ray E. -ray
7. A radio wave has a frequency of 8.6  108 Hz. What is the energy of one photon of this radiation?
8. If the energy of a photon is 1.32  10-18 J, what is its wavelength in nm?
9. Use the Rydberg equation to calculate the frequency of a photon absorbed when the hydrogen atom
undergoes a transition from n1 = 2 to n2 = 4. (R = 1.096776  107 m-1)
5
10. An electron in the n = 6 level emits a photon with a wavelength of 410.2 nm. To what energy level does the
electron move?
11. The size of an atomic orbital is associated with
A. the principal quantum number (n). B. the angular momentum quantum number (l).
C. the magnetic quantum number (ml). D. the spin quantum number (ms).
E. the angular momentum and magnetic quantum numbers, together.
12. The orientation in space of an atomic orbital is associated with
A. the principal quantum number (n). B. the angular momentum quantum number (l).
C. the magnetic quantum number (ml). D. the spin quantum number (ms).
E. None of these choices is correct.
13. Which of the following is a correct set of quantum numbers for an electron in a 3d orbital?
A. n = 3, l = 0, ml = -1 B. n = 3, l = 1, ml = +3 C. n = 3, l = 2, ml = 3
D. n = 3, l = 3, ml = +2 E. n = 3, l = 2, ml = -2
14. The energy of a photon is directly proportional to the wavelength of the radiation.
True False
15. In the Bohr model of the hydrogen atom, the electron moves in a circular path which Bohr referred to as an
orbital.
True False
FILL IN THE BLANKS:
1.
The distance between identical points on a wave is called the ________________________.
2.
The unit “number of wavelengths per second” is called the_________________________.
3.
The frequency of light is inversely proportional to its _________________________.
4.
The lowest frequency of light that can eject electrons from the surface of a metallic element is called the
_______________________________________________.
5.
A particle of light that possesses one quantum of energy is called a _____________________.
6.
The lowest energy state of an atom is called the ____________________________________.
7.
When an electron moves from n=3 t n=2, energy is ______________________________.
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6
Chem 1, Ch 8 Worksheet WSHEET 8
Name : ______________________________________
1. "Each electron in an atom must have its own unique set of quantum numbers" is a statement of
A. the aufbau principle. B. the Pauli exclusion principle. C. Hund's rule.
D. the periodic law. E. Heisenberg's principle.
2. The ____________________ quantum numbers are associated with the energy of an electron in a manyelectron atom.
A. n and l B. n and ml C. l and ml D. n and ms E. n, l and ml
3. In a single atom, what is the maximum number of electrons which can have quantum number n = 4?
A. 16
B. 18
C. 32 D. 36 E. None of these choices is correct.
4. Select the correct set of quantum numbers (n, l, ml, ms) for the highest energy electron in the ground state of
potassium, K.
A. 4, 1, -1, ½ B. 4, 1, 0, ½ C. 4, 0, 1, ½ D. 4, 0, 0, ½
E. 4, 1, 1, ½
5. Select the correct set of quantum numbers (n, l, ml, ms) for the first electron removed in the formation of a
cation for strontium, Sr.
A. 5, 1 , 0, -½ B. 5, 1, 0, ½
C. 5, 0, 1, ½ D. 5, 1, 1, ½ E. 5, 0, 0, -½
6. Write the correct electron configuration for Cu (Z = 29) ____________________________
7. An element with the electron configuration [noble gas]ns2(n - 1)d10np3 has ____________ valence electrons.
A. 2 B. 3 C. 5 D. 10 E. 15
8. Which of the following elements has the largest atomic size?
A. S B. Ca C. Ba
D. Po E. Rn
9. Which one of the following equations correctly represents the process relating to the ionization energy of X?
A. X(s)  X+(g) + e- B. X2(g)  X+(g) + X-(g) C. X(g) + e-  X-(g)
D. X-(g)  X(g) + eE. X(g)  X+(g) + e10. Which of the following elements has the smallest first ionization energy?
A. Rb
B. Mg C. I D. As E. F
7
11. Elements with the highest first ionization energies are found in the ___________ region of the periodic
table.
A. lower left B. upper left C. center D. lower right E. upper right
12. Elements with ________________ first ionization energies and ___________ electron affinities generally
form cations.
A. low, very negative B. high, positive or slightly negative C. low, positive or slightly negative
D. high, very negative E. None of these is generally correct.
13. Select the element with the least metallic character.
A. Sn
B. Sr C. Tl
D. Ge E. Ga
14. Which of the following atoms will be diamagnetic?
A. Cr
B. Ru
C. Fe
D. Pt
E. Cd
15. Select the diamagnetic ion.
A. Cu2+
B. Ni2+
C. Cr3+
D. Sc3+
E. Cr2+
16. Write down the maximum number of electrons in an atom which can have
a. quantum number n = 4.______________ b. orbital designation 3d.__________________
c. orbital designation 2pz. _________________
17. The maximum number of electrons in an atom with the same value of n is 2n2.
True False
18. Electron affinities of neutral atoms may be positive or negative.
True False
8
WSHEET 9: CHP 11:Chem 1 Worksheet for Hybrid and MO Theory
Work on separate paper and complete the following during lab time. Due at the end of lab.
1.
Describe the bonding in each of the following, using valence bond theory. Identify the hybridization on each atom besides
hydrogen and label all bonds as sigma or pi bonds. (Hint: draw Lewis structures first.)
a. SCN- (ion) b. CO2c. HOCN
2.
Draw the simplified molecular orbital diagram for each of the following. Based on that electron configuration, determine the number of
bonding pairs of electrons and the number of antibonding pairs of electrons, determine the bond order, and determine if the species is
paramagnetic or diamagnetic. Then identify which ones could NOT exist.
a. O2+
3.
d. H2C=CH-CH2-CH=CH2
b. O22- c. N2-
d. Ne2 e. Be2+
f. F2-
Consider the molecule below. For all atoms other than hydrogen, determine the electron-pair arrangement, hybridization, molecular
geometry, and identify each bond as sigma or pi. (Yes, the O has two lone pairs.)
:S:
|| ..
:N=C-C-O-H
..
9
Multiple Choice Questions:
1. A molecule with the formula AX3 uses __________ to form its bonds.
A. sp hybrid orbitals B. sp2 hybrid orbitals C. sp3 hybrid orbitals
D. sp3d hybrid orbitals E. sp3d2 hybrid orbitals
2. A molecule with the formula AX3E uses _________ to form its bonds.
A. s and p atomic orbitals B. sp3 hybrid orbitals C. sp2 hybrid orbitals
D. sp hybrid orbitals E. sp3d2 hybrid orbitals
3. A molecule with the formula AX4E2 uses _________ to form its bonds.
A. sp hybrid orbitals B. sp2 hybrid orbitals C. sp3 hybrid orbitals D. sp3d hybrid orbitals
E. sp3d2 hybrid orbitals
4. Valence bond theory predicts that carbon will use _____ hybrid orbitals in the carbonate anion, CO32-.
A. sp B. sp2 C. sp3 D. sp3d E. sp3d2
5. Valence bond theory predicts that iodine will use _____ hybrid orbitals in ICl2-.
A. sp2 B. sp3 C. sp3d D. sp3d 2 E. None of these choices is correct.
6. According to valence bond theory, the triple bond in ethyne (acetylene, C2H2) consists of
A. three  bonds and no  bonds. B. two  bonds and one  bond. C. one  bond and two  bonds.
D. no  bonds and three  bonds. E. None of these choices is correct.
7. Overlap of two sp2 hybrid orbitals produces a  bond.
A. True
B. False
8. In the valence bond treatment, a  bond is formed when two 2p orbitals overlap side to side.
A. True
B. False
10
WSHEET 10: LEWIS STRUCTURE HANDOUT
Molecule or ion
Lewis structure
Number of
LPs around
central
atom
EP Geometry
CCl4
H 2S
NF3
HCN
11
Molecular Geom.
Polarity:
molecule
Bond angles
SO3
CO2
PCl5
XeF4
SF6
PO43-
Also try OCN12
Chem 1 WSHEET 11: Chp # 12
Name:___________________________________________________
1. Examine the phase diagram for the substance Bogusium (Bo) and select the correct statement.
A. Bo(s) has a lower density than Bo(l).
B. The triple point for Bo is at a higher temperature than the melting point for Bo.
C. Bo changes from a solid to a liquid as one follows the line from C to D.
D. Bo changes from a liquid to a gas as one follows the line from C to D.
E. Point B represents the critical temperature and pressure for Bo.
2. Examine the following phase diagram and determine what phase exists at point F.
A. vapor + liquid
B. vapor
C. liquid
D. solid
13
E. supercritical fluid
3. Ammonia's unusually high melting point is the result of
A. dipole-dipole forces.
B. London dispersion forces.
D. covalent bonding.
E. ionic bonding.
C. hydrogen bonding.
4. In hydrogen iodide __________________ are the most important intermolecular forces.
A. dipole-dipole forces
B. London dispersion forces
C. hydrogen bonding
D. covalent bonds
E. polar covalent bonds
5. Which of the following atoms should have the greatest polarizability?
A. F
B. Br
C. Po
D. Pb
E. He
6. The strongest intermolecular interactions between ethyl alcohol (CH3CH2OH) molecules arise
from
A. dipole-dipole forces.
B. London dispersion forces.
C. hydrogen bonding.
D. ion-dipole interactions.
E. carbon-oxygen bonds.
7. The strongest intermolecular interactions between hydrogen fluoride (HF) molecules arise
from
A. dipole-dipole forces. B. London dispersion forces.
C. hydrogen bonding.
D. ion-dipole interactions. E. ionic bonds.
8. Which of the following should have the highest boiling point?
A. CF4
B. CCl4
C. CBr4
D. CI4
E. CH4
9. Which of the following should have the highest surface tension at a given temperature?
A. CH4
B. CF4
C. CCl4
D. CBr4
E. CI4
10. A metal such as chromium in the body-centered cubic lattice will have _______________
atom(s) per unit cell.
A. 1
B. 2
C. 3
D. 4
E. 9
11. Iron crystallizes in the body-centered cubic lattice. What is the coordination number for Fe?
A. 4
B. 6
C. 8
D. 10
E. 12
12. Of the five major types of crystalline solid, which would you expect each of the following to
form? (e.g., H2O: molecular)
a. Sn
b. Si
c. KCl
d. Xe
e. F2
13. The energy of a hydrogen bond is greater than that of a typical covalent bond.
True False
14. A single water molecule can participate in at most two hydrogen bonds at any instant.
True False
14
15. Consider the phase diagram shown below.
a. What phase(s) is/are present at point A?
b. What phase(s) is/are present at point B?
c. Name point C and explain its significance.
d. Starting at D, if the pressure is lowered while the temperature remains constant, describe what
will happen.
15
Chem 1 WSHEET 13 , Chp # 13: Name:
___________________________________________________
1. For a given solution, which of the following concentration values will change as temperature
changes?
A. mass percent B. molality C. mole fraction D. molarity E. None of these choices is
correct.
2. Potassium fluoride is used for frosting glass. Calculate the molarity of a solution prepared by
dissolving 78.6 g of KF in enough water to produce 225 mL of solution.
A. 0.304 M B. 0.349 M C. 1.35 M D. 3.29 M E. 6.01 M
3. Calculate the molarity of a solution prepared by diluting 1.85 L of 6.5 M KOH to 11.0 L.
A. 0.28 M B. 0.91 M C. 1.1 M D. 3.1 M E. 3.9 M
4. Saccharin, one of the first non-nutritive sweeteners used in soft-drinks, is 500 times sweeter
than sugar in dilute aqueous solutions. The solubility of saccharin is 1.00 gram per 290 mL of
solution. What is the molarity of a saturated saccharin solution?
saccharin= 183.2 g/mol
A. 0.0188 M B. 0.632 M C. 1.58 M D. 3.45 M E. None of these choices is correct.
5. Copper(II) bromide is used as a wood preservative. What mass of CuBr2 is needed to prepare
750.0 mL of a 1.25 M solution?
A. 134 g B. 209 g C. 372 g D. 938 g E. > 1 kg
6. What is the molality of a solution prepared by dissolving 86.9 g of diethyl ether, C4H10O, in
425 g of benzene, C6H6?
A. 0.362 m B. 0.498 m C. 2.01 m D. 2.76 m E. None of these choices is correct.
7. The solubility of the oxidizing agent potassium permanganate is 7.1 g per 100.0 g of water at
25C. What is the mole fraction of potassium permanganate in this solution?
A. 0.0080 B. 0.0086 C. 0.066 D. 0.45 E. 0.48
8. Sodium hydroxide is a common ingredient in drain cleaners such as Drano. The mole fraction
of sodium hydroxide in a saturated aqueous solution is 0.310. What is the molality of the
solution?
A. 0.310 m B. 0.690 m C. 1.24 m D. 12.4 m E. 25.0 m
9. The mole fraction of potassium nitrate in an aqueous solution is 0.0194. The solution's density
is 1.0627 g/mL. Calculate the molarity of the solution.
A. 0.0194 M B. 0.981 M C. 1.05 M D. 1.96 M E. 19.4 M
10. Colligative properties depend on
A. the chemical properties of the solute. B. the chemical properties of the solvent.
C. the masses of the individual ions.
D. the molar mass of the solute.
E. the number of particles dissolved.
16
11. From the following list of aqueous solutions and water, select the one with the lowest
freezing point.
A. 0.75 M (NH4)3PO4 B. l.0 M CaSO4 C. l 0 M LiClO4
D. 1.5 M CH3OH, methyl alcohol
E. pure water
12. Select the strongest electrolyte from the following set.
A. CH3CH2OH, ethanol
B. LiNO3
C. C6H12O6, glucose
13. Select the weakest electrolyte from the following set.
A. Na2SO4 B. KCl C. CH3CH2COOH, propionic acid
D. CCl4
D. CaCl2
E. HF
E. LiOH
14. How many moles of bromide ions are present in 750.0 mL of 1.35 M MgBr2?
A. 0.506 mol B. 1.01 mol C. 2.03 mol
D. 3.04 mol
E. None of these choices is correct.
15. Calculate the vapor pressure of a solution prepared by dissolving 0.500 mol of a non-volatile
solute in 275 g of hexane (
= 86.18 g/mol) at 49.6C. Phexane= 400.0 torr at 49.6C.
A. 54 torr B. 154 torr C. 246 torr D. 346 torr E. 400. Torr
16. Determine the freezing point of a solution which contains 0.31 mol of sucrose in 175 g of
water. Kf = 1.86C/m
A. 3.3C B. 1.1C
C. 0.0C D. -1.1C E. -3.3C
17. Dimethylglyoxime, DMG, is an organic compound used to test for aqueous nickel(II) ions.
A solution prepared by dissolving 65.0 g of DMG in 375 g of ethanol boils at 80.3C. What is
the molar mass of DMG?
Kb = 1.22C/m, boiling point of pure ethanol = 78.5C
A. 44.1 g/mol B. 65.8 g/mol C. 117 g/mol D. 131.6 g/mol E. 553 g/mol
*****************************************************************************
WSHEET 14
Chem 1 : Equations Packet!!!!!
WRITING CHEMICAL EQUATIONS
Know the meaning of reactant, product, yield, stoichiometric coefficient, Law of conservation of
mass.
AIDS IN BALANCING
1. Remember the seven diatomic elements.
2. In general, if a polyatomic ion is NOT changed during a chemical reaction, you may
carry it over and count it unaltered.
3. Start balancing with the most complex formula in the equation and finish with the
simplest.
17
4. NEVER change a formula’s subscripts to balance elements.
5. Check carefully to make sure that every element is balanced.
TYPES OF CHEMICAL EQUATIONS
I. COMBINATION OR SYNTHESIS A + B  C. The chemical union of two or more
elements or compounds to form a more complex substance.
2 Na + Cl2  2 NaCl
NH3 + H2ONH4OH
II. DECOMPOSITION C  A + B. The reaction in which a compound is broken up into
its elements or into simpler compounds.
2 KClO3 2 KCl + 3 O2
2 H2O  2 H2 + O2
III. SINGLE REPLACEMENT A + BC  B + AC. The reaction in which a free element
replaces another element in a compound. Requires the use of the Activity Series to
predict if one element will replace another.
Zn + Pb(NO2)2  Pb + Zn(NO2)2Cl2 + 2 KBr  Br2 + 2 KCl
Fe + H2SO4  H2 + FeSO4
Cu + HCl  No reaction
IV. DOUBLE REPLACEMENT AB + CD  AD + CB. The reaction in which two
compounds exchange ions to form two new compounds. Three sub-categories:
AgNO3 + NaCl  AgCl + NaNO3 precipitation
HCl + NaOH  H2O + NaCl neutralization
Na2S + 2 HCl  2 NaCl + H2S gas-forming
V. COMBUSTION __X + __O2  __XO?. The rapid reaction of oxygen with elements
or compounds to produce the more oxidized forms. Hydrocarbons’ combustion
reactions with oxygen always from CO2 and H2O if complete combustion occurs.
2 Mg + O2  2 MgO
S + O2  SO2
CH4 + O2  CO2 + 2 H2O
SUPPLEMENT TO SILBERBERG: BALANCING REDOX REACTIONS
Keep this handout to use with Chapter 4 and beyond. See the Dry Lab in your lab manual also.
Steps for balancing by the half-reaction (ion-electron) method in acidic solution:
1. Determine oxidation numbers for all atoms in the reaction. Identify what has been oxidized and what
has been reduced and connect them with arrows. Determine how many electrons have been lost or
gained per atom. Write the start of a half-reaction for what is oxidized and another for what is reduced.
2. For each half-reaction:
a. Balance all atoms other than hydrogen or oxygen.
b. Balance oxygen by adding water.
c. Balance hydrogen by adding H+.
d. Balance electrical charge by adding electrons to the appropriate side. (You have already
determined how many electrons in step one; check to see that this number works.)
3. Use least common denominator factors to multiply each half-reaction to make the electrons lost by the
oxidation half-reaction equal to the electrons gained by the reduction half-reaction.
18
4. Add the two half-reactions to make a whole reaction, canceling all species that are common to both
sides.
In base:
Complete the above four steps, then:
5. Looking at the whole reaction, add OH- to both sides equal to the number of H+.
6. On the side of the reaction where H+ = OH-, combine them to make water.
7. Cancel water common to both sides.
8. Rewrite the whole reaction neatly.
********************************************************************************************************************
WSHHET 15 : NET IONIC EQUATIONS
NAME_______________________________
Only reactions involving aqueous solutions can be written as net ionic equations, i.e.,
equations showing only those ions which react to form an insoluble product (precipitate), a gas
or a liquid. This section will lead you through the steps in writing net ionic equations, then you
will practice.
Step 1: Does the reaction include reactants or products in aqueous solutions?
Example 1: 4 Fe(s) + 3 O2(g)  2 Fe2O3(s)
There is no (aq) symbol anywhere, therefore there will be no net ionic equation.
Example 2: 2 Al(s) + 6 HCl(aq)  3 H2(g) + 2 AlCl3(aq)
There are two (aq) symbols, therefore one can determine the net ionic equation for this
reaction.
Step 2: If there is an aqueous solution present, we can write the ions present in solution as
separate ions dissolved in water, EXCEPT for weak acids and weak bases and poorly soluble
ionic compounds.* Remember that each ion is surrounded by water molecules, and that they are
called hydrated ions. In Example 2 above, we can write the following:
6 HCl(aq) as 6 H+(aq) + 6 Cl-(aq) and 2 AlCl3(aq) as 2 Al3+(aq) + 6 Cl-(aq).
(We have to keep the same total number of ions present, per the Law of Conservation of Mass,
therefore look at both formula subscripts and stoichiometric coefficients.)
Step 3: Now we rewrite the normal chemical equation as ionic equations, i.e., showing all
dissolved ions separately. (NOTE that I do not call the normal chemical equation a molecular
equation as many textbooks do. Are ionic compounds molecular? Nope.)
For Example 2 above, we write:
2 Al(s) + 6 H+(aq) + 6 Cl-(aq)  3 H2(g) + 2 Al3+(aq) + 6 Cl-(aq)
19
Note: nothing that was solid, liquid or gas was changed in any way – no charges added, no
separation. If something is not ionic and dissolved in water, don’t change it!
Step 4: In a normal chemical equation, the products have to be different from the reactants, i.e,
we would not put the same thing on both sides of the equation and say that a reaction had
occurred. In an ionic equation, we can see that some ions have not changed from the reactant
side to the product side. These are called SPECTATOR IONS. In Example 2 above, the
chloride ions are spectator ions. We need to be able to identify these.
Step 5: After we identify the spectator ions, we can rewrite the ionic equation without them;
after all, they aren’t contributing anything or reacting with anything. If we do that, we have a
NET IONIC EQUATION. It does not show only ions, but it does show only those “things”
which have reacted or formed. In Example 2 above, the net ionic equation would be:
2 Al(s) + 6 H+(aq)  3 H2(g) + 2 Al3+(aq)
Notice that the chloride ions have been left out. Now there are changes on both sides of the
equation. It is important that the charges add up to the same number on both sides of the
equation. Notice that 6 x (+1) = 2 x (+3) so our net ionic equation is fine.
*Actually, not everything that is written with (aq) is in ionic form in water. We will learn about three solutions that
become a gas and water in open containers: NH4OH, H2CO3 and H2SO3. Take H2O from each of these and what’s
left is a gas. Also a salt or base may be designated (aq) but be only poorly soluble, in which case you leave the
formula whole. The third exception is weak acids and weak bases, which should be kept written as molecules. So,
only soluble salts and soluble strong bases and strong acids can be separated into ions.
PRACTICE: Balance, then write the following as total ionic equations with the spectator ions
circled. Then rewrite them as net ionic equations.
1. ___Zn(s) + ___Pb(NO3)2(aq)  ___Pb(s) + ___Zn(NO3)2(aq)
Total ionic:
Net ionic:
2. ___(NH4)2S(aq) + ___Cu(NO3)2(aq)  ___NH4NO3(aq) + ___CuS(s)
Total ionic:
Net ionic:
3. ___Pb(NO3)2(aq) + ___HCl(aq)  PbCl2(s) + ___HNO3(aq)
Total ionic:
Net ionic:
4. ___HNO3(aq) + ___Cu(s)  ___Cu(NO3)2(aq) + ___NO2(g) + ___H2O(l)
20
Total ionic:
Net ionic:
5. ___Na2S(aq) + ___O2(g) + ___H2O(l)  ___Na2S2O3(aq) + ___NaOH(aq)
Total ionic:
Net ionic:
(Hint: S2O32-(aq) is an anion called thiosulfate; don’t separate it any further.)
WSHEET 16: BALANCING OXIDATION-REDUCTION REACTIONS PRACTICE
The following unbalanced reactions need to be balanced using redox methods, in acidic solution:
1. ___H2S(g) + ___NO3-(aq)  ___NO(g) + ___S(s)
2. ___MnO4-(aq) + ___Cl-(aq)  ___Cl2(g) + ___Mn2+(aq)
3. ___Cl-(aq) + ___Cr2O72-(aq)  ___Cr3+(aq) + ___Cl2(g)
4. ___HgS(s) + ___NO3-(aq)  ___HgCl4-(aq) + ___NO(g) + ___S(s)
5. ___Fe2+(aq) + ___MnO4-(aq)  ___Fe3+(aq) + ___Mn2+(aq)
6. ___BiO3-(aq) + ___Mn2+(aq)  ___MnO4-(aq) + ___Bi3+(aq)
7. ___WO3(s) + ___Sn2+(aq) + ___Cl-(aq)  ___W3O8(s) +___ SnCl62-(aq)
8. ___NiO2(s) + ___Ag(s)  ___Ni2+(aq) + ___Ag+(aq)
9. ___Fe3+(aq) + ___H2S(g)  ___Fe2+(aq) + ___S(s)
10. ___PbO2(s) + ___Mn2+(aq)  ___Pb2+(aq) + ___MnO4-(aq)
11. ___C2H4(g) + ___MnO4-(aq)  ___Mn2+(aq) + ___CO2(g)
The following unbalanced reactions need to be balanced using redox methods, in basic solution:
12. ___Cr3+(aq) + ___MnO2(s)  ___Mn2+(aq) + ___CrO42-(aq)
13. ___Bi(OH)3(s) + ___SnO22-(aq)  ___Bi(s) + ___SnO32-(aq)
14. ___Cr3+(aq) + ___H2O2(aq)  ___CrO42-(aq)
15. ___S2O82-(aq) + ___Ni(OH)2(s)  ___SO42-(aq) + ___NiO2(s)
21