Unit 1: Scientific Notation, Density, Etc.
0.1.1 Mass and Volume
1. For each of the following measurements, what quantity is being measured - density, mass,
length, or volume?
a. Millimeter-length
d. Liter-volume
3b. 4 cm volume
e. Kilogram-mass
c. Centigram-mass
f. 2 g/mL density
2. What does the Law of Conservation of Mass state?
Mass cannot be created or destroyed (this is why we must balance the number of atoms in
the reactants and products)
3. What is the difference between mass and volume? Mass is the amount of “atomic stuff”
Volume is the amount of ‘space’ the stuff takes up
4. You have a rectangular object with a length of 4 cm, a width of 9 cm and a height of 10 cm. What
is the volume of the object? 4 x 9 x 10 = 360 cm3
0.1.2 Density:
5. What is the volume of an object that has a mass of 5.1 grams and a density of 7.1 g/mL?
D = m/v
v = m/d = 5.1 / 7.1 = .72 L
6. What is the mass of an object that has a volume of 2.1 mL and a density of 4.5 g/mL?
D = m/v
m = d x v = 4.5 x 2.1 = 9.45 g
7. A student filled a graduated cylinder with water and read the meniscus at 26 mL. The student then
dropped a solid material into the graduated cylinder and the water level rose to 40 mL. If the solid
material had a density of 3.0 g/mL, determine the mass of the solid object.
40-26 = 14mL m = d x v = 3 x 14 = 42g
0.1.3 Scientific Notation:
8. Convert the following measurements into scientific notation:
a. 620,000 6.2 x 105
b.000323
3.23 x 10-4
9. Convert the following into standard notation:
a. 7.01 X 10-3 ..00701
10. Calculate:
a. 5.9 X 10-2 X 8.1 X 107 4.7 x 106
b. 9.2 X 104 92000
b.
5.9 X 105 8.1 X 10-2 7.3 x 106
0.1.4 Metric Conversions:
1.
2.
3.
4.
How many liters are in 402 mL? .402 L
How many centigrams is 3,000 grams? 300000 cg
How many milligrams are in 20 kilograms? 20000000 mg
How many kilometers are in 1,500 meters? 1.5 km
Modeling Chemistry
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Unit 2: Gas Laws
Combined Gas Law
V1 = initial volume
T1 = initial temperature
P1 = initial pressure
T2 = final temperature
V2 = final volume
P2 = final pressure
1. A 15,000 L sample of gas is initially at a pressure of 1.2 atm. If the pressure is increased to
250 atm, what is the new volume of the gas?
1.2 atm x 15000L = 250 atm x V2
cross-multiply: V2 = 72 L
2. If a scuba diver has 0.05 L of gas in his blood under a pressure of 250 atm, then rises very
fast where his blood has a pressure of 50.0 atm, what will the volume of gas in his blood be?
250 atm x 0.05 L = 50 atm x V2
cross-multiply: V2 = 0.25 L
3. If I inhale 2.2 liters of gas at a temperature of 180 C and it heats to a temperature of 380 C in
my lungs, what is the new volume of the gas?
18 + 273 = 291 K = T1
38 + 273 = 311 K = T2
2.2 L / 291 K = V2 / 311 K
cross-multiply: V2 = 2.3 L
Partial Pressures
1. A metal tank contains three gases: oxygen, helium, and nitrogen. If the partial pressures of
the three gases in the tank are 35 atm of oxygen, 5 atm of nitrogen, and 25 atm of Helium,
what is the total pressure inside of the tank?
35 + 5 + 25 = 65 atm
2. The total pressure in a closed system containing N2, O2, and CO2 is 60 atm. If the partial
pressure of N2 is 8 atm, and the partial pressure of O2 is 12 atm, what is the partial pressure of
CO2?
60 atm = 8 + 12 + P3
P3 = 40atm
Pressure Conversions
1 atm = 760 torr = 101.325 kPa = 760 mmHg = 14.7 psi
1. Convert 50.00 psi into torr 50 psi x (760 torr / 14.7 psi) = 2585 torr
2. Convert 987,654 torr into mmHg. 987654 torr x (760 mmHg / 760 torr) = 987654 torr
3. Convert 50 kPa to torr 50 kPa x (760 torr / 101.325 kPa) = 375 torr
Ideal Gas Law: PV = nRT
Use the ideal gas law to find one variable when everything is held constant
1. If I have 4 moles of a gas at a pressure of 5.6 atm and a volume of 12 liters, what is the
temperature?
T = PV / nR = 5.6 x 12 / 4 x .0821 = 204.6 K
2. 0.275 moles of a gas occupy 0.97 L at a pressure of 1.75 atm. What is the temperature of the
gas?
T = PV / nR = 1.75 x .97 / .275 x 0.0821 = 75.2 K
3. If I have 21 moles of gas held at a pressure of 78 atm and a temperature of 900 K, what is the
volume of the gas?
V = nRT / P = 21 x 0.0821 x 900 / 78 = 19.9 L
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Combined Gas Law 3 variables
1. If I initially have a gas at a pressure of 12 atm, a volume of 23 liters, and a temperature of 200
K, and then I raise the pressure to 14 atm and increase the temperature to 300 K, what is the
new volume of the gas?
12 x 23 / 200 = 14 x V2 / 300
cross-multiply: 29.6 L
2. A gas takes up a volume of 17 liters, has a pressure of 2.3 atm, and a temperature of 299 K. If
I raise the temperature to 350 K and lower the pressure to 1.5 atm, what is the new volume of
the gas?
17 x 2.3 / 299 = 1.5 V2 / 350
cross-multiply: 30.5 L
3. A gas that has a volume of 28 liters, a temperature of 45 0C, and an unknown pressure has its
volume increased to 34 liters and its temperature decreased to 35 0C. If I measure the
pressure after the change to be 2.0 atm, what was the original pressure of the gas?
T1 = 45 + 273 = 318 K
T2 = 35 + 273 = 308 K
28 x P1 / 318 = 34 x 2 / 308
Modeling Chemistry
cross-multiply: 2.5 atm
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Review Fall 2012-2013
Unit 3: Heat and Energy
1) What is thermal energy? Energy to change temp What is phase energy? Energy to change
the phase (solid, liquid, gas)
2) What is the difference between heat and temperature? Temperature is the speed of
particles, Heat is the speed of particles AND the mass
3) Does temperature ever increase when the phase changes? NO!
4) A 100g sample of liquid water heats up from an initial temperature of 20oC to 78oC. How much
energy did it take in? 78-20 = 58 q = mCΔT = 100x 4.18 x 58 = 24244 J
5) During the winter, you accidentally forgot your 230g cup of iced tea outside and it froze. How
much energy was released? Q = mHf = 230 x 334 = 76820 J
6) How much energy is required to completely boil a 15 g sample of water at 100 oC?
Q = mHv = 15 x 2260 = 33900 J
Use the following heating curve at left for questions
7-12
7) At which sections is thermal energy increasing?
I, III, V
8) At which sections is phase energy increasing?
II, IV
9) Which section represents a liquid?
III
10) Which section represents melting?
II
11) Which section represents evaporating?
IV
12) At which sections is the phase energy constant?
I, III, V
Use the phase diagram at right for questions 13-16
13) What section represents a liquid? S
14) What sections represents a gas? K
15) What do we call point G? triple point
16) If a substance begins at 80°C at a pressure of 150 kPa
and the pressure decreased to 50 kPa what phase
change occurred? boiling
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Unit 4: The Atom
1.
2.
3.
4.
What is the atomic number of an element equal to? # of protons
What is the mass number equal to? # of protons & neutrons
How do you find the neutrons of an element? Atomic mass-atomic #
What is difference between a regular element and an ion of that same element? Ion has
gained or lost electrons
Use Bohr’s model of the atom for questions 5-8.
5.
6.
7.
8.
What type of light is given off when an electron moves from n=5 to n=3? infrared
What type of light is given off when an electron moves from n=3 to n=2? 656-visible red
What color is a wavelength of 550 nm? green
What color is a wavelength of 621 nm?v orange
9. How much parent material is left after a substance goes through 3 half-lives? 1/16
10. If parent material is equal to 1/32, how much is daughter material? 1- (1/32) = 31/32
11. If you start with 100 g of a substance, how much is daughter material after 2 half-lives?
100 x (1/4) = 25 g
12. You start with 400g of Radon-14. Its half-life is 4 seconds. How much parent material (in grams) is
left after 16 seconds? 4 half lives > 1/16
400 x (1/16) = 25g
13. Write down the symbol for both an alpha particle and a beta particle.
14. Write out the formula for the alpha decay of Dysprosium (Dy) – 167. > Gd-163
15. Write out the formula for the alpha decay of Curium (Cm) – 257. > Pu-253
16. Write out the formula for the beta decay of Molybdenum (Mo) – 100. > Nb-100
17. Write out the formula for the beta decay of Mercury (Hg) – 214. > Au-214
*Alpha particle is a helium nucleus (atomic mass: 4, atomic #: 2)
*Beta particle (atomic mass: 0, atomic #: -1) proton is lost and converted to a neutron, so mass does not change
Substance Symbol Atomic
#
Hafnium
Hf
72
Cesium
Cs
55
Tungsten W
74
Arsenic
As-78
33
Iron
Fe-53
26
Copper
Cu-72
29
+3
Vanadium V
23
-2
Cadmium Cd
48
+5
Silicon
Si
14
Modeling Chemistry
Mass #
Protons Neutrons Electrons
178
133
150
78
53
72
51
101
28
72
55
74
33
26
29
23
48
14
5
106
78
76
45
27
43
28
53
14
72
55
74
33
26
29
20
50
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Review Fall 2012-2013
Unit 5: Periodicity
Patterns and Trends in the Periodic Table
1. What are the vertical columns of the periodic table called? groups
2. What are the horizontal rows called? periods
3. How does atomic mass change as you move down the periodic table? increases
4. How does atomic number change as you move across the periodic table? increases
5. Would you expect aluminum (Al) to be more like thallium (Tl) or sulfur (S)? Tl (same # of
valence electrons)
6. Is potassium (K) more similar to Lithium (Li) or Lead (Pb)? Why? Li (same # of valence
electrons)
7. Is helium (He) more similar to argon (Ar) or Cesium (Cs)? Give 2 reasons why. Ar
8. Which one of these four does not belong? Argon (Ar), Bromine (Br), Krypton (Kr), and Xenon
(Xe).
Bromine
9. Use your Periodic Table to arrange the following sets of elements in order of increasing atomic
size.
a. N, C, O, F __F, O, N, C__
b. F, Cl, Br, I ___F, Cl, Br, I__
c. Rb, P, As, Co _P, As, Co, Rb__
d. Mg, Si, Ar ___Ar, Si, Mg____
10. All about electronegativity:
a. What is Electronegativity? Ability to hold electrons to itself
b. As you go across a period in the Periodic Table does Electronegativity increase or
decrease? Why?
More protons pulling electrons tight to the nucleus
c. As you go down a group in the Periodic Table does Electronegativity increase or
decrease? Why?
Electrons are further away in larger orbitals
11. Use your Periodic Table to arrange the following in order of increasing electronegativity.
a. N, C, O, F __C, N, O, F___
b. F, Cl, Br, I ___I, Br, Cl, F__
c. Rb, P, As, Co __Rb, Co, As, P__
d. Ca, Se, Ni ___Ca, Ni, Se__
e. O, Po, S ____Po, S, O____
f. Cs, F, Ga ___Cs, Ga, F__
12. List 3 trends in the periodic table. Make sure to say how they change!
Electronegativity-increases across (left to right), decreases down
Ionization Energy-increases across, decreases down
Atomic Radius-decreases across, increases down
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Unit 6: Nomenclature
1. What is the correct name for BaO? Barium oxide (ionic)
2. What is the correct name for ClOz?
chloride dioxide (Covalent)
3. When writing the name of an ionic compound, are Greek prefixes used? NO!
4. When writing the name of a covalent compound, are Greek prefixes used? YES! (never put
-mono in front of the first element)
5. When the charge of a metal cation varies (i.e. is in the “crazy zone”) what is used to specify the
charge? Roman numeral (represents charge-not the subscript)
6. An ionic compound must have an overall charge equal to _____
Zero (overall neutral)
7. In order to determine the subscripts of ionic compounds, what do we need to do the charges?
Swap, drop, and chop the charges (use the table and group #s to determine charge)
8. Do we swap, drop, and chop anything when dealing with a covalent compound?
NO!-use greek prefixes
Properties of Ionic and Covalent Compounts
9. Which type of compound (ionic or covalent) has a high melting and boiling point?
Ionic
10. Can ionic compounds conduct electricity when melted or dissolved?
Yes!
11. Can covalent compounds conduct electricity when melted or dissolved?
No!
12. Which type of compound (ionic or covalent) has low melting and boiling points?
Covalent
13. Write the correct formula for cobalt (II) chloride
CoCl2
14. Write the correct formula for aluminum bromide
AlBr3
15. Write the correct formula for diphosphorouspentoxide
P 2 O5
16. What kind of bonds transfer electrons? Ionic (stealing electrons)
17. ______Covalent____ bonds share electrons
18. A positive ion is also called a ___cation__.
19. A negative ion is also called a(n) anion
Ionic or Covalent?
20. sodium oxide – ionic or covalent? ionic
21. Trinitrogenpentasulfide – ionic or covalent? covalent
22. sodium sulfate – ionic or covalent? ionic
23. lead (II) chloride– ionic or covalent? ionic
24. disulfur trioxide – ionic or covalent? covalent
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25. iron(II) hydroxide– ionic or covalent? ionic
26. N2O4– ionic or covalent? Covalent
Unit 7: Chemical Reactions
For the following equations, identify the types of reaction as synthesis, combustion, decomposition,
single replacement, or double replacement.
1) Na3PO4 + 3 KOH  3 NaOH + K3PO4 ____double replacement____
2) MgCl2 + Li2CO3 MgCO3 + 2 LiCl ____ double replacement ____
3) C6H12 + 9 O2 6 CO2 + 6 H2O _______combustion___
4) Pb + FeSO4 PbSO4 + Fe ______ single replacement __
5) CaCO3CaO + CO2 ____decomposition___
Directions: Predict the products of the following reactions and balance the reaction.
Single Replacement
____Ba + ____AuSO4 -----> BaSO4 + Au
____Na + ____MgCl2 -----> 2 NaCl + Mg
Double Replacement
____BaCl2 + ____Na2CO3-------> BaCO3 + 2 NaCl
____Al2 (SO4) 3 + ____H3PO4 ------> 2 AlPO4 + 3 H2SO4
Balancing
1. __2_ SO2 + __1_ O2 __2_ SO3
2. __2_CH4 + _3__ O2 __2_CO + _4__H2O
3. _2__ P + _3__Cl2 __2_PCl3
Changes
1. What is the definition of a physical change? Change without a new substance
2. How do you know a physical change has occurred? (What are three indicators?)
size, shape, texture
3. What is the definition of a chemical change? A new substance is made
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4. How do you know a chemical change has occurred? (What are 4 indicators?)
Change in temperature, color, smell, gas or solid (precipitate) is produced
5. HNO3 + Ba(OH)2 ––> H2O + Ba(NO3)2
In the chemical reaction above, which substances are reactants? Which substances are
products?
Reactants = HNO3 + Ba(OH)2
Products = H2O + Ba(NO3)2
6. An endothermic reaction is (hot/cold) because it (releases/absorbs) energy.
7. An exothermic reaction is (hot/cold) because it (releases/absorbs) energy.
Unit 8: The Mole
Two-Step Mole Conversions
1. How many molecules are there in 450 grams of Na2SO4?
450 g x (1 mol / 142g) x (6.022 x 1023 / 1 mol) = 19 x 1023 = 1.9 x 1024 particles
2. How many grams are there in 2.3 x 1024 atoms of silver?
2.3 x 1024 x (1 mol / 6.022 x 1023 ) x (108 g / 1 mol) = 41 x 101 = 410 g
3. What is the volume of 32 grams of carbon dioxide gas at STP?
32g x (1mol / 44 g) x (22.4 L / 1 mol) = 16.3 L
4. What is the mass of 2.00 L of nitrogen gas (N2) at STP?
2L x (1mol / 22.4 L) x (28g / 1mol) = 2.5 g
Percent composition
1. What is the percent by mass of chlorine in FeCl3?
Cl3 = 3 x 35 = 105g
FeCl3 = (56 + 105) = 161g
105/161 = .65 x 100 = 65%
2. What is the percent by mass of silver in Ag2CrO4?
Ag2 = 108 x 2 = 216g
Ag2CrO4 = 216 + 52 + (4x16) = 332g
216/332 = .65 x 100 = 65%
3. What is the percent by mass of oxygen in NO2?
O2 = 16 x 2 = 32g
NO2 =14 + 32 + 46g
32/46 = .70 x 100 = 70%
4. What is the percent by mass of iodine in NaI?
I = 127 g
NaI = 23 + 127 = 150 g
127/150 = .85 x 100 = 85%
One-Step Mole Conversions
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1. How many moles are in 15 grams of lithium? 15g x (1mol / 7g) = 2.1 mol
2. How many moles are in 1.91 x 1024molecules of beryllium iodide, BeI2?
1.91 x 1024 x (1 mol / 6.022 x 1023 ) = .3 x 101 = 3mol
3. How many grams are in 3.3 moles of potassium sulfide, K2S?
3.3 mol x (110g / 1mol) = 363 g
4. How many moles are in 1.2 liters of ammonia, NH3?
1.2 L x (1.mol / 22.4 L) = .05 mol
Calculate Molar Mass of each compound.
1. Cl2
35x2 = 70g
2. BeCl2
9+(2 x 35) = 79g
3. Mg(OH)2
24+(2x16)+2 = 58g
4. H3PO4
3+31+(4x16) = 98g
5. Pb(NO3)2 207+(2x14)+(6x16)=331g
Unit 9: Stoichiometry
1. For the reaction CaO + SO3 CaSO4, how many moles of calcium sulfate are produced from 50 g
of sulfur trioxide?
50g x (1 mol SO3 / 80 g SO3) x (1 mol CaSO4/ 1 mol SO3) = .625 mol CaSO4
/)=
2. For the reaction Co + F2 CoF2, how many moles of fluorine are required to produce 175 g of
cobalt(II) fluoride?
175g CoF2 = (1 mol CoF2 / 97 g CoF2) x (1 mol F2/ 1 mol CoF2) = .01 mol F2
3. In the reaction between excess K(s) and 5.21 g of O2(g), potassium oxide is formed . What mass
of potassium oxide would you expect to form? Equation: 4 K(s) +
1 O2(g)
2 K2O(s)
5.21g O2 x (1 mol O2/ 32 g O2 ) x (2 mol O2 / 1 mol K2O) x (94 g K2O / 1 mol K2O) = 30.6 g K2O
4. Determine the mass of carbon dioxide that should be produced in the reaction between 8.98 g of
carbon and excess O2.
Equation: 1 C(s) +
1 O2(g)
1 CO2(s)
8.98g C x (1 mol C/ 12 g C) x (1 mol CO2 / 1 mol C) x (44 g CO2 / 1 mol CO2) = 33 g CO2
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For problems #5 - #8, use solubility graph at right.
5. Which substance is most soluble at 45º C ? potassium iodide
6. Which two substances have the same solubility at 23º C ?
potassium nitrate & sodium chloride
7. Which substance’s solubility changes the most from 0º C to 100º C ? potassium nitrate
8. What is the solubility of potassium nitrate at 65º C ? 130 g/ mL of H20
9. A 32.1 g sample of potassium nitrate is dissolved in 225 mL of solution. What is the molarity of the
solution?
32.1 g x (1 mol / 85g) = .38mol
225 mL x (1 L / 1000 mL) = .225 mL
M = .38 mol / .225 L = 1.69 M
10. What is the molarity of a solution made from 200
g of CuSO4 dissolved in 4.25 L of water?
200g x (1mol / 160g) = 1.25mol
M = 1.25 mol / 4.25 L = .29 M
11. 4.15 g of I2 are dissolved in 305 mL of solution.
Calculate the molarity of iodine in the solution.
4.15g x (1 mol / 127g) = .03 mol
305 mL x (1 L / 1000 mL) = .305 L
.03 mol / .305 L = .11 M
12. What volume would be needed in order to
prepare a 0.75 M NaCl solution from 7.6 g of NaCl ?
7.6 g x (1 mol / 58g) = .13 mol
M = mol / L
L = mol / M = .13 mol / .75 L = .17L
Identify the following compoundsas either acids or bases:
1. HCl ____acid__
4. HClO3 ____ acid ____
2. HF
____ acid ___
5. CH3OH ____ base ___
3. NaOH ____base____
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