Regents Chemistry
Review Packet for Midterm Exam
2012-2013
Name:
Significant Figures
How many sig figs are in each of the following measurements?
1. 6.25x10-6 g
_____3_____
6. 86270 mm
____4_________
2. 5.505x102 cm _____4_____
7. 86270.0 mm ____6_________
3. 0.0081 L
_____2_____
8. 140 m/hr
____2_________
4. 87.0˚C
_____3_____
9. 729 s
____3_________
10. 13.90 ml
____4_________
5. 0.00750 dm3 _____3_____
Vapor Pressure
1. What is vapor pressure?
pressure exerted by a vapor above a liquid
2. How is vapor pressure related to boiling point?
Normal boiling point of a liquid is the temp at which the vapor pressure of the liquid is
equal to the standard atmospheric pressure
3. If a substance has a high vapor pressure and a low boiling point what can you say about the IMF?
IMF are weak
4. What variable does vapor pressure depend on?
temperature
Nuclear Chemistry
Identify each of the following as natural transmutation (NT) or artificial transmutation (AT); if AT
identify if can be fission or fusion. Explain how to identify which type of reaction it is.
Nuclear reaction
1
238
92U
→
2
2
1H
3
1H
3
239
94Pu
4
23
11Na
5
14
7N
6
235
92U
+
7
212
→
+
+
83Bi
234 Th
90
→
1
+
0n
4
4
+
90
→
0n
38Sr
+
+
1
0n
→
24
11Na
4
2He
→
1
1H
1
→
92
36Kr
+
0n
212
84Po
+
explanation
NT
One reactant
+ energy
AT - fusion
147
56Ba
AT - fission
AT
2 small atoms join
together into 1 slightly
larger atom
1 large atom hit by
particle and splits into
2 smaller atoms
2 reactants
AT
2 reactants
AT - fission
1 large atom hit by
particle and splits into
2 smaller atoms
One reactant
2He
1
2He
type
+
0
+ 3 10n
17
8O
-1e
+
141
56Ba
+ 3 10n +
energy
NT
8. We use two laws to balance nuclear equations. These two laws are the conservation of ______
atomic number ___ and the conservation of ___atomic mass___.
9. In the symbol 31H what does the 3 represent? ____atomic mass______
10. In the symbol 126C what does the 6 represent? ___atomic number____
11. What nuclide is a radioisotope used to date rocks? ___U-238___
12. Find the missing term in the following nuclear equations:
a.
20
b.
180
9F
→
79Au
0
c. X →
d.
9
0
-1e
X = ___2010Ne____
+ X
→
176
77Ir
-1e
+
37
+ X
X = ___3718Ar____
19K
+ X →126C +
4Be
X = ___42He_____
1
0n
X = ____42He______
13. For fusion reactions to occur, high temperatures are required because both of the reacting
nuclei have
a. a large mass
b. many neutrons
c. a positive charge
d. a negative charge
14. Name one beneficial use and one destructive use of nuclear fission reactions?
Beneficial: large amounts of energy with no greenhouse effect
Destructive: exposure to harmful radiation and nuclear waste to dispose of
15. Explain why it is more difficult to cause an artificial transmutation reaction with an alpha
particle than with a neutron.
Alpha particles have a positive charge so need to be accelerated to high speeds in order
to overcome the repulsion forces of (+) to (+) particles for the nuclear reaction to occur
16. Where in the universe do fusion reactions take place?
The sun
17. Where do fission reactions take place in a controlled fashion? In an uncontrolled one?
Controlled: nuclear power plant
Uncontrolled: nuclear bomb
18. What is the total number of neutrons in an atom of Pb-207? __207-82 = 125 neutrons_
19. What is the nuclear charge of fluorine? ____+9 (9 protons) _______
20. What happens to the half-life of a radioactive substance as the temperature of the substance
increases? ____ no change ___________
21. Based on Reference Table N, what fraction of a sample of gold-198 remains radioactive after
2.69 days?
HL Au-198 is 2.69 days
after I half-life there will be ½ of the sample left
22. How many days are required for 200.0 grams of radon-222 to decay to 50.0 grams?
HL Rn-222 = 3.82 days
200 → 100 → 50
2 half-lifes = (3.82 days)(2) = 7.64 days
23. How much time must elapse before 16 grams of potassium-42 decays, leaving 2 grams of the
original isotope?
HL K-42 = 12.4 h
16 → 8 → 4 → 2
(12.4 h)(3 HL) = 37.2 h
24. The half-life of iodine-131 is 8.07 days. What fraction of a sample of I-131 remains
after 24.21 days?
24.21 days
=
3 HL
1 → 1 → 1
8.07 days
2
4
8
25. What was the original mass of a radioactive sample that decayed to 25 grams in 4 half-lives?
1
2
3
4
400g → 200g → 100g → 50g → 25g
400 g was original mass
Atomic Structure
1. What is the definition of each of the following terms:
a. Atomic number = number of protons
b. Mass number = number of proton plus number of neutrons
c. Isotope = two (or more) atoms of the same element with different # neutrons (so
14
13
12
have different mass numbers)
examples:
6C and
6C and
6C
d. Nucleon = term used to collectively indicate the protons and neutrons in the nucleus
2. What is the charge, mass, location, and symbol of each of the following subatomic particles?
Particle
Charge
Mass
Location
symbol
proton
+1
1 amu
nucleus
1
1H
neutron
no charge
1 amu
nucleus
1
0n
electron
-1
negligible
electron cloud
0
-1e
or
1
1p
3. The number of electrons in a neutral atom is equal to the number of __protons____
4. Complete the following table:
Atom
# of protons
# of neutrons
12
# of electrons
Atomic #
Mass #
6C
6
6
6
6
12
7
3Li
3
4
3
3
7
23
11
12
11
11
23
3
1H
1
2
1
1
3
4
2He
2
2
2
2
4
11Na
5. What is the charge on the nucleus of an atom? ____positive____________.
6. Most of the volume of an atom is ____empty space_________.
7. Describe Rutherford’s famous experiment. What were the two results of the experiment and
what conclusion did Rutherford draw from each result?
1. positively charged part of atom is very small and is located in the center of the atom
2. majority of the atom is empty space
8. What is the difference between an isotope and an ion?
Isotopes are 2+ same atoms with different # neutrons
Ions are atoms that are charged due to losing/gaining electrons
9. Given that chlorine is 75% Cl-35 and 25% Cl-37, calculate the average atomic mass of chlorine.
(.75)(35) + (.25)(37) = 35.5 amu
10. What radioactive isotope is used to date organic material (wool, linen, wood, bones, etc)?
C-14
11. What is the absolute temperature scale?
Kelvin
12. In the Bohr model of the atom, describe the type of electron transition which will result in the
production of a spectral line. Is this transition endothermic or exothermic?
electrons transitioning from higher energy levels to lower energy levels emit wavelengths of
light (bright line spectra) - this is an exothermic process since energy is emitted
13. How many atoms are represented in the formula (NH4)2SO4?
2 atoms of N, 8 atoms of H, 1 atom of S, 4 atoms of O for a total of 15 atoms
14. What are two common substances that sublime?
CO2(s) [carbon dioxide]
and
I(s)
Periodic Table
1. What is nuclear charge? What is effective nuclear charge?

Nuclear charge is the charge of the nucleus = positive sign with # protons in nucleus
- example: nuclear charge of K is +19

Effective nuclear charge: charge felt by valence electrons
o
Calculate:
# protons minus # kernel electrons
2. What is ionization energy?
amount of energy needed to remove the most loosely held electron (valence)from an atom
3. What is electro negativity?
ability of an atom to hold onto its own electrons & attract its neighbor’s electrons
4. What are the trends for atomic size, ionization energy, & electro negativity in the Periodic
Table?

Atomic size ↑ as go down a column from top to bottom and ↓ as go left to right across a row

electro negativity increases from bottom left (Fr) to upper right (F)

ionization energy ↓ as go down a column from top to bottom and ↑ as go left to right across a
row
Energy Changes: Identify the Q equation needed to solve each of the following energy word problems
below, then solve each problem.
1. How much energy is required to heat a 35.0 g sample of water from 35ºC to 75ºC?
Equation:____Q = mcΔT_______
Q = (35.0g)(4.18J/gºC)(40 ºC) = 5,852 J
2. An ice cube at 0˚C with a mass of 175.0g melts. How much energy does the ice absorb to melt?
Equation: ____Q = mHf_____________
Q = (175.0g)(334J/g) = 58,450 J
3. An 85.0g sample of water at 100.0˚C boils. How much energy is required to convert the water to
steam at 100.0˚C?
Equation: ____ Q = mHv ___________
Q = (85.0g)(2260J/g) = 192,100 J
4. A sample of water with a mass of 160.0g is heated from 15.0˚C to 70.0˚C. How much energy is
required to heat the water?
Equation: _____ Q = mcΔT ________
Q = (160.0g)( 4.18J/gºC)(55ºC) = 36,784 J
5. 450.0g of water at 0.0˚C is frozen to ice at 0.0˚C. How much energy is released to the
environment?
Equation: ______ Q = mHf ________
Q = (450.0g)(334J/g) = 150,300 J
6. 200.0g of steam at 100.0˚C condenses to water at 100.0˚C. How much energy is released to the
environment?
Equation: ____ Q = mHv ______
Q = (200.0g)(2260J/g) = 452,000 J
7. A 225g sample of water is cooled from 95.0˚C to 40.0˚C. How much energy is released to the
environment?
Equation: _____ Q = mcΔT ____
Q = (225g)(4.18J/gºC)(-55ºC) = (-) 51,727.5 J
8. A block of ice at 0.0˚C with a mass of 1.0kg melts. How much energy is absorbed from the
environment to convert it to water at 0.0˚C?
Equation: ____ Q = mHf_________
Q = (1000g)(334J/g) = 334,000 J
9. A 200.0g sample of water at 100.0˚C is boiled and converted to steam at 100.0˚C. How much
energy is required?
Equation: _____ Q = mHv_______
Q = (200.0g)(2260J/g) = 452,000 J
10. A 345.0g sample of steam at 100.0˚C condenses to water at 100.0˚C. How much energy is
released to the environment?
Equation: ____ Q = mHv_______
Q = (345.0g)(2260J/g) = 779,700 J
11. If you carry out a reaction in a Styrofoam cup calorimeter and
a. the temperature of the water increases, is the reaction endothermic or exothermic?
b. the temperature of the water decreases, is the reaction endothermic or exothermic?
c. How would you calculate the energy change? Q = mcΔT
Phase Changes
boiling
evaporation
increasing
temperature
endothermic
heat of fusion
heat of vaporization
solid
sublimation
decreasing
deposition
gas
phase change
constant
fusion
melting
condensation
exothermic
liquid
vaporization
1. _temperature_ is a measure of the average kinetic energy of the particles of a substance.
2. _vaporization_ is another word for boiling.
3. _heat of fusion_ is the amount of energy required to melt one gram of a solid at its melting
point.
4. _fusion__ is another word for melting.
5. One type of physical change is a _phase change___.
6. Solid to liquid or gas is an example of an _endothermic__ process.
7. __evaporation__ is the spontaneous change from liquid to gas at any temperature.
8. __sublimation__ is the change from solid phase to gas phase.
9. The average kinetic energy of the particles of a substance is increasing when the temperature
is ___increasing______.
10. __gas__ takes the shape and volume of its container.
11. __boiling__ is the change from liquid phase to gas phase at a constant temperature.
12. __condensation_ is a change from gas to liquid phase which is noticeable on glass.
13. __melting___ is the change from solid to liquid phase at a constant temperature.
14. __heat of vaporization__ is the amount of energy required to convert one gram of a liquid to
the gas phase at its boiling point.
15. __deposition_ is the change from gas phase to the solid phase.
16. __solid__ has a definite shape and a definite volume.
17. Gas to liquid or solid is an example of an _exothermic__ process.
18. When the temperature is constant, the average kinetic energy of the particles in a substance
is ___constant____.
19. __liquid__ has definite volume and indefinite shape.
20. When temperature decreases, the average kinetic energy of the particles of a substance is
__decreasing___.
Mole Math
1. Draw the mole map in the space below
2. Calculate the formula mass of the following:
a. NH3__ 1(14) + 3(1) = 17g/mol__
d. CH4 _1(12) + 4(1) = 16g/mol__
b. NaCl __ 1(23) + 1(35) = 58g/mol__
e. C2H4 _2(12) + 4(1) = 28g/mol__
c. CO _1(12) + 1(16) = 28g/mol_
3. Perform the following conversions:
a. 2 moles NH3 to grams _34g_
f. 59.5g NH3 to moles _ 3.5mol_
b. 0.4 moles NaCl to grams __23.2g__
g. 3 moles CO to liters _67.2L__
c. 0.8 moles CO to grams __22.4mol__
h. 4.5 moles C2H4 to liters __100.8L__
d. 234.0g NaCl to moles __4.03mol__
i. 11.2 liters CO to moles _0.5mol___
e. 32.0g CH4 to moles __2mol__
j. 89.6 liters CH4 to moles _4mol__
4. What is the definition of percent? Part ÷ whole x 100 = %
5. Calculate the percent composition of the following:
a. N and H in NH3
N
FM= 17g.mol
14 x 100 = 82.35%
H
3 x 100 = 17.65%
17
17
b. C and H in CH4
C
12 x 100 = 75%
H
16
4 x 100 = 25%
16
c. Na and Cl in NaCl
Na
23 x 100 = 39.66%
58
Cl
35 x 100 = 60.34%
58
6. Calculate the percent by mass of water in the following:
a. BeSO4•2H2O
FM = 141g/mol
18 x 100 = 12.77%
141
b. CaO2•8H2O
FM = 216g/mol
144 x 100 = 66.67%
216
7. What is the empirical formula of a substance with the following molecular formulas:
a. P4O10 ___P2O5_____
b. N2O4
____NO2________
c. H2C2O4
____HCO2_____
8. What is the molecular formula of a substance with the empirical formula C4H5N2O and a molecular
mass of 194.0g/mol.
194.0
FM of empirical formula:
= 2
2(C4H5N2O)
=
4(12) + 5(1) + 2(14) + 16 = 97g/mol
C8H10N4O2
97
9. What is the molecular formula of a substance with the empirical formula CH2 and a molecular
mass of 84.0g/mol.
FM of empirical formula:
84 = 6
14
6(CH2) = C6H12
1(12) + 2(1) = 14g/mol
Gas Laws
1. What are the 4 assumptions of the KMT of gases?
1. Gas particles move in random, straight line motion
2. Gas particles undergo “elastic” collisions
3. Volume of gas particles is negligible
4. Gas particles do not attract or repel each other
2. What is an ideal gas?
Gas that behaves according to the KMT all the time
3. Under which two conditions do real gases behave ideally?
high temperature and low pressure
4. What is the Law of Conservation of Energy?
energy is neither gained or lost
5. What is the Law of Conservation of Matter?
matter is neither created nor destroyed
6. Given the reaction:
2X + 3Y → Q + 4Z
If 5g of X react completely with 30g of Y to produce 4g of Q, how much Z will be produced?
5g + 30g = 4g + Z
Z = 31g
7. What is Dalton’s Law of Partial Pressures?
The sum of the pressures of all the gases in a system is equal to the total pressure of the
system
8. What is Avogadro’s hypothesis?
One mole of any gas occupies 22.4 liters of space
9. What are the three equations for calculating an energy change? Describe when to use each
equation.
Q = mcΔT
Q = mHf
Q = mHv
10. What is the formula to calculate percent error?
% error =  actual – observed 
x
100
actual
11. What is the combined gas law?
P1V1 = P2V2
T1
T2
12. What temperature scale must you use in gas law word problems?
Kelvin
Intermolecular Forces
1. What are the three IMF? Which force is weakest? Which force is strongest?
Dispersion (weakest), diploe-dipole, H-bonds (strongest)
2. Describe the kind of molecule that exhibits each kind of intermolecular force (ex: polar, nonpolar, etc)
Dispersion:

monatomic elements, diatomic elements, non-polar molecules (symmetrical)
Ca, Ni, H2, O2, CH4, CO2
Dipole-dipole: polar molecules (asymmetrical)

CH3Cl, HCl
H bonds: H bonded to F, O, N

HF, H2O, NH3
3. Why does water have a high boiling point? has H bonds (strongest IMF)
4. What kind of IMF exists between each of the following molecules?
a. CH4
___dispersion____
b. HCl
__dipole-dipole____
c. CH3OH
__ H bonds___
d. CO2
___ dispersion ___
e. Xe
___ dispersion ____
f. NH3
___H bonds__
g. He
___ dispersion ____
h. N2
____ dispersion ___
i.
___H bonds____
HF
j. H2
____ dispersion ____
k. C5H12
___ dispersion ____
l.
___ dispersion ___
F2
m. HBr
____ dipole-dipole ____
5. If a substance has strong IMF what phase is it likely to be in at room temperature?
Solid or liquid
6. If a substance has weak IMF what phase is it likely to be in at room temperature?
Gas