Download name chemistry final review

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NAME _______________________
DATE _______________________
CHEMISTRY FINAL REVIEW
DIRECTIONS: Go through the review and try to answer as many questions as you can without having
to look anything up. Circle/highlight the questions you couldn’t answer right away. This will identify
areas that you need to concentrate on for the final. Go through your notes and worksheets to help you
answer the rest of the questions.
UNIT 1 - MEASUREMENTS
1.
What are significant figures?
The figures in a measurement that are certain plus one digit of uncertainty
that is estimated.
2.
How do you know how many sig figs should be in a measurement?
It is determined by the instrument used to make the measurement plus one
extra number that is estimated.
3.
Count the number of sig figs in the following measurements.
a. 2390
3
d. 120
2
b. 0.00987
3
e. 4,900
2
c. 0.03056
4
f. 300
1
4.
Express the following values in 4 sig figs.
a. 56
56.00
d. 0.087
b. 87,566
87,570
e. 1,999,999
c. 8.09449
8.094
f. 0.0063821
0.08700
2.000 x 106
0.006382
UNIT 2 – MATTER
5.
What is matter?
Anything that has mass and takes up space.
6.
What are the two forms of matter? Define.
Pure Substances and Mixtures
7.
What are the two types of mixtures? Define.
Homogeneous – Parts of the mixture look the same
Heterogeneous – Parts of the mixture are visibly different
8.
Give examples of the two types of mixtures.
Homogeneous – Iced Tea, Chocolate Milk, Gold Jewelry
Heterogeneous – Salad, Cereal, Sand
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9.
What are the two types of Pure Substances? Define.
Element – The simplest form of matter, only one type of atom
Compound – Two or more elements chemically combined
10.
Give examples of the two types of Pure Substances.
Element – Silver, Carbon, Lithium
Compound – H2O, C6H12O6, MgCl2
11.
What is an atom?
Simplest particle of matter.
12.
What is a molecule?
Two or more atoms chemically combined.
13.
What is a chemical property? Give two examples.
The ability of an object to change into a new substance; flammability,
corrosion.
14.
What is a physical property? Give two examples.
A physical description of a substance; color, hardness
15.
What is density?
The amount of mass per unit of volume.
16.
What is the density of a liquid that has a mass of 50.340 g and a volume of 300.00 mL?
0.16780 g/mL
17.
What is an intensive physical property? Give two examples.
A property that does not depend on how much of the substance there is;
density, color.
18.
What is an extensive physical property? Give two examples.
A property that does depend on how much of the substance there is; mass,
volume.
UNIT 3 – THE ATOM
19.
What are the three subatomic particles that make up an atom?
Proton, neutron, electron.
20.
Where are the three subatomic particles located in an atom?
Nucleus, nucleus, energy levels.
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21.
What are the charges of the three subatomic particles?
+1, 0, -1
22.
What are the masses of the three subatomic particles?
1 amu, 1 amu, 0 (1/2000thamu)
23.
What is the atomic number of an atom?
The number of protons in an atom.
24.
List the atomic number for the following elements.
a. Manganese 25
d. Barium
56
b. Silver
47
e. Uranium
92
c. Carbon
6
f. Rutherfordium 104
25.
What is the mass number of an atom?
The number of protons + the number of neutrons
26.
List the number of protons and neutrons for the following atoms.
a. Germanium – 73 32 p+ and 41 n
d. Oxygen – 17
+
b. Bromine – 80
35 p and 45 n
e. Nickel – 58
+
c. Calcium – 40
20 p and 20 n
f. Mercury – 200
8 p+ and 9 n
28 p+ and 30 n
80 p+ and 120 n
27.
What is an ion?
An atom that has a charge.
28.
How does an atom become an ion?
By gaining or losing electrons.
29.
What are the two types of ions and how are they different?
Cations are positively charged and anions are negatively charged.
30.
List the number of electrons in the following ions.
a. Fe+3
23 ed. K+1
18 eb. S-2
18 ee. I -1
54 ec. As-3
36 ef. Pb+4
78 e-
UNIT 4 – THE ELECTRON
31.
What is Quantum Mechanics?
The process that allows you to determine the most likely location of an
electron in an atom.
32.
What are the letters that designate the 4 Quantum Numbers?
n, l, m, s
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33.
What does each Quantum Number tell you?
n is the energy level, l is the sublevel, m is the orbital, s is the spin of the
electron.
34.
What is Pauli’s Exclusion Principle?
No two electrons can have the same four quantum numbers – basically, no two
electrons can be in the same place at the same time.
35.
What is Electron Configuration?
A way of indicating the energy level and sublevel of the electrons in an atom.
36.
Write out the full Electron Configuration of the following elements.
a. Chromium 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d4
b. Lead 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6, 6s2, 4f14, 5d10, 6p2
c. Europium 1s2, 2s2, 2p4
37.
Write the shortcut Electron Configuration for the following elements.
a. Molybdenum [Kr] 5s2, 4d4
b. Sulfur [Ne] 3s2, 3p4
c. Iridium [Xe] 6s2, 4f14, 5d7
38.
What are Orbital Diagrams?
A way of visually representing the energy level, sublevel, and spin of electrons
in the orbitals of an atom.
39.
Draw the Orbital Diagrams for the following elements. * Only the valence level *
a. Titanium () ( )( )( )( )( )
4s
3d
b. Tungsten ()
6s
()()()()()()() ( )( )( )( )( )
4f
5d
c. Rubidium ( )
5s
40.
What is a Valence Energy Level?
Highest energy level that contains electrons in an atom.
41.
What are Valence Electrons?
Electrons in the Valence energy level.
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42.
Determine the Valence Energy Level and the number of Valence Electrons for the
following elements.
a. Calcium VEL = 4; Ve-s = 2
b. Silicon VEL = 3; Ve-s = 4
c. Bismuth VEL = 6; Ve-s = 5
d. Cadmium VEL = 5; Ve-s = 2
e. Neon VEL = 2; Ve-s = 8
f. Carbon VEL = 2; Ve-s = 4
43.
What are Electron Dot Diagrams?
Diagrams that show the sublevel & number of valence electrons in an atom.
44.
Draw the Electron Dot Diagrams for the following elements.
a. Manganese
 

b. Berkelium
Mn
Na
c. Sodium
d. Nickel
 
 
e. Iodine
Ni
Bk
f. Aluminum

 

I 
 
 
Al 
45.
What makes an atom stable? What (general) rule tells us this?
The Octet Rule, 8 valence electrons make an atom/ion stable.
46.
How does the type of element determine the charge something will take?
Metals tend to lose electrons and nonmetals tend to gain electrons, metalloids
tend to act more like nonmetals in this aspect.
47.
Determine the charge that the following elements will take to become stable.
a. Li +1
b. N -3
c. Si +4 or -4
d. Ca +2
e. F -1
f. Al +3
UNIT 5 – THE PERIODIC TABLE
48.
What is the purpose of the Periodic Table?
The PT is used to organize all the elements in a useful way.
49.
What are the 3 types of elements?
Metals, nonmetals, and metalloids.
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50.
Where on the PT are these elements located?
Metalloids lie along the zig-zag line, metals are to the left of the line,
nonmetals are on the right except for Hydrogen.
51.
What are the vertical columns on the PT called?
Groups
52.
What are the horizontal rows on the PT called?
Periods
53.
What is a family?
A collection of elements with very similar chemical and physical properties.
54.
List the 5 metallic families and their location on the PT.
Alkali Metals – Group 1 (not H)
Alkaline Earth Metals – Group 2
Transition Metals – Groups 3 – 12
Lanthanides – Period 6 of the Inner Transition Metals
Actinides – Period 7 of the Inner Transition Metals
55.
List the 2 nonmetallic families on the PT.
Halogens – Group 17
Inert Gases – Group 18
56.
What is a trend on the PT? Why is it useful?
A trend is a pattern of properties that repeats on the PT. It can be used to
predict the properties of elements as well as make comparisons about
elements’ properties.
57.
What is Electronegativity?
EN is the ability of an atom to attract electrons.
58.
What is the trend of EN on the PT?
EN increases up and to the right along the PT.
59.
What is Ionization Energy?
IE is the amount of energy needed to remove an electron from an atom.
60.
What is the trend of IE on the PT?
IE increases up and to the right along the PT.
61.
What is atomic radius?
It is the radius of an atom – how big an atom is.
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62.
What is the trend of atomic radius on the PT?
Atomic radius increases down and to the left along the PT.
63.
How is ionic radius different from atomic radius? Be specific.
Cations are smaller than their parent atoms because they have lost electrons
and anions are larger than their parent atoms since they have gained
electrons.
UNIT 6 – FORMULAS AND NAMES
64.
What is an ionic compound?
An ionic compound is made up of two ions. Generally, it contains a metal and a
nonmetal or a polyatomic ion.
65.
What is a molecular compound?
A molecular compound is made up of two nonmetal elements.
66.
How do you determine the formula of a compound?
If the cmpd is ionic then you must balance the charges of the ions that make
up the cmpd. If the cmpd is molecular then you must determine the formula
from the name listed (follow the prefixes).
67.
Which ion is written in the formula first?
Cation
68.
What is a polyatomic ion?
An ion made up of more than one atom.
69.
List the prefixes for numbers 1-10.
1 - mono
3 - tri
2 - di
4 - tetra
70.
5 - penta
6 - hexa
7 - hepta
8 - octa
9 - nona
10 - deca
Determine the formula of the following compounds. (Remember, you must first determine
if the compound is molecular or ionic!)
a. Magnesium Chloride
l. Hydrogen Nitrite
MgCl2
HNO2
b. Iron (III) Sulfide
m. Beryllium Chloride
Fe2S3
BeCl2
c. Barium Nitrate
n. Manganese (V) Selenide
Ba(NO3)2
Mn2Se5
d. Dihydrogen Monoxide
o. Lead (IV) Sulfite
H 2O
Pb(SO3)2
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e. Sodium Sulfate
Na2SO4
f. Aluminum Fluoride
AlF3
g. Lithium Permanganate
LiMnO4
h. Ammonium Bromide
NH4Br
i. Rubidium Oxide
Rb2O
j. Copper (II) Phosphate
Cu3(PO4)2
k. Calcium Chloride
CaCl2
71.
p. Tetracarbon Hexaiodide
C4I6
q. Silver Carbonate
Ag2CO3
r. Pentasulfur Octafluoride
S5F8
s. Chromium (III) Acetate
Cr(C2H3O2)3
t. Hydrogen Monobromide
HBr
u. Vanadium (I) Nitride
V3N
v. Gallium Oxide
Ga2O3
Write the name of the following compounds. (Remember, you must first determine if the
compound is molecular or ionic.)
a. SrCl2
l. CdO
Strontium Chloride
Cadmium (II) Oxide
b. NiS
m. K3PO3
Nickel (II) Sulfide
Potassium Phosphite
c. PBr5
n. NaF
Phosphorus Pentabromide
Sodium Fluoride
d. Ba(NO3)2
o. CrSe
Barium Nitrate
Chromium (II) Selenide
e. Mg3(PO4)2
p. FeCO3
Magnesium Phosphate
Iron (II) Carbonate
f. BI3
q. Al(NO2)3
Boron Tri-iodide
Aluminum Nitrite
g. Zr(C2H3O2)2
r. SrO
Zirconium (II) Acetate
Strontium Oxide
h. H2S
s. PtBr2
Dihydrogen Monosulfide
Platinum (II) Bromide
i. K2Se
t. Be3N2
Potassium Selenide
Beryllium Nitride
j. CaF2
u. ZnI2
Calcium Fluoride
Zinc Iodide*
k. NH4Cl
v. PbS2
Ammonium Chloride
Lead (IV) Sulfide
UNIT 7 – REACTIONS
72.
What is a reaction? What actually occurs in a reaction?
A reaction is a chemical change, atoms rearrange to form new substances.
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73.
What are the two parts in a reaction?
Reactants and products
74.
What is the name of the written form of a reaction?
Chemical Equation
75.
Why do chemical equations need to be balanced? What law dictates this?
Law of Conservation of Matter – matter cannot be created or destroyed. This
applies to atoms, mass, and energy.
76.
What is the name of the number that is used to balance chemical equations?
Coefficients
77.
What are the five types of reactions?
Synthesis, Decomposition, Single Replacement, Double Replacement,
Combustion.
78.
Label the type of reaction and then balance the following chemical equations.
a. Syn
4 Al +
3 O2
b. DR
3 CaBr2 + 2 H3PO4
c. Comb
C4H8 +
6 O2 
4 CO2
d. DR
3 KNO3
+ FeCl3

3 KCl
+ Fe(NO3)3
e. DR
MgS + 2 LiHCO3

Li2S
+ Mg(HCO3)2
f. DR
2 KMnO4
g. Syn
Ca
h. Decomp
2 KHCO3
i. SR
2 GaBr3
+
3 F2
j. DR
3 CuSO4
+
2 Fe(C2H3O2)3
k. DR
H2SO4
l. DR
KOH
+

+
→
6 HBr + Ca3(PO4)2

+ SrSO4 
F2
+
2 Al2O3
→
+ 4 H2O
K2SO4
+ Sr(MnO4)2
CaF2
H2O
+
→
FeCl2
→
AgNO3
→
CO2
+
K2CO3
2 GaF3
+
3 Br2
3 Cu(C2H3O2)2
→
2 HCl
AgOH
+
+ Fe2(SO4)3
FeSO4
+
KNO3
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m. DR
3 SrCO3
n. Comb
2 C3H6
2 H3PO4
+
+
9 O2
→
→
3 H2CO3
6 H2O
+
+
Sr3(PO4)2
6 CO2
UNIT 8 – THE MOLE
79.
What does the word “mole” mean?
Mole is a measurement of particles. It specifically refers to 6.02 x 1023
particles of a substance.
80.
What is Avogadro’s number?
6.02 x 1023
81.
What is molar mass? (* This is also called formula mass.)
Mass of one mole.
82.
What is a conversion factor?
Fraction in which the numerator is equal to the denominator – used to make
conversions between different units of measurement.
83.
Determine the number of moles that is equal to the following masses given. (Remember, to
first find the molar mass of the compound and “dump” the unit you don’t want anymore!)
*For extra names help – try to name the following compounds in the margin.
a. 100g of BaCl2
0.5 moles of BaCl2
b. 35g of (NH4)3PO4
0.23 moles of (NH4)3PO4
c. 750g of MgS
13 moles of MgS
d. 0.821g of BF3
0.0121 moles of BF3
e. 11.6g of Na2CO3
0.109 moles of Na2CO3
f. 68g of Al2(SO4)3
0.20 moles of Al2(SO4)3
g. 275g of NaCl
4.71 moles of NaCl
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h. 147g of HBr
1.82 moles of HBr
84.
Determine the mass that is equal to the following number of moles given. (Remember, to
first find the molar mass of the compound and “dump” the unit you don’t want anymore!)
*For extra names help – try to name the following compounds in the margin.
a. 2.16 moles of SrF2
271g of SrF2
b. 0.91 moles of C6H12O6
160g of C6H12O6
c. 4.33 moles of CuS
414g of CuS
d. 10 moles of AlI3
4000g of AlI3
e. 1.67 moles of KNO3
169g of KNO3
f. 5.00 moles of Fe(MnO4)3
2060g of Fe(MnO4)3
g. 0.075 moles of LiOH
1.8g of LiOH
h. 8.55 moles of HgCl2
2320g of HgCl2
85.
Convert the following quantities to volume.
a. 0.0750 moles Ar
24.1 L
b. 8.317 g SO2
2.908 L (via 0.1298 moles)
c. 3.86 x 1022 molecules H2O
1.40 L (via 0.0623 moles)
86.
What is percent composition?
Percent by mass of an element (part) in a compound (whole).
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87.
Determine the percent composition of each element in the following compounds.
a. (NH4)3PO4 28.19% N, 8.13% H, 20.77% P, 42.92% O
b. NiCl2 45.29% Ni, 54.71% Cl
c. BaCl2 · 3 H2O 52.36% Ba, 27.03% Cl, 2.31% H, 18.30% O
88.
What is a hydrate?
A hydrate is a compound that is surrounded by a certain number of water
molecules.
89.
Determine the percent composition of water in the following hydrates.
a. K3PO4 · 6 H2O 33.75%
b. Mg(ClO3)2 · 4 H2O 27.38%
c. Na2SO4 · 2 H2O 20.24%
90.
What is the purpose of heating a hydrate?
Heating a hydrate to 100C will boil the water off in the hydrate leaving
behind the anhydrous compound.
91.
What is an empirical formula?
The formula that shows the simplest whole number ratio of elements in the
compound.
92.
Determine the empirical formula of the unknown compound from the percent composition
given.
a. 62.6 % Tin and 37.4% Chlorine SnCl2
b.
15.27% Magnesium, 44.53% Chlorine, and 40.20% Oxygen Mg(ClO2)2 [MgCl2O4]
c. A sample of an unknown compound was determined to have 15.875 g of Silver, 2.063 g
of Nitrogen, and 7.063 g of Oxygen. AgNO3
93.
What is a molecular formula?
The formula that shows the actual whole number ratio of elements in the
compound.
94.
Determine the molecular formulas of the unknown compound based on the information
given.
a. An unknown compound with a molar mass of 166.22 g has a percent composition of
47.05% Potassium, 14.45% Carbon, and 38.50% Oxygen. K2C2O4
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b. An unknown compound was experimentally determined to contain 45.293g Carbon,
7.622 g Hydrogen, and 60.345g Oxygen. The molar mass of the compound is 180.18g.
C6H12O6
UNIT 9 – STOICHIOMETRY
95.
What does stoichiometry allow you to calculate?
Stoich allows you to calculate the amounts of other substances in a reaction.
96.
What is molar ratio?
Ratio of moles in a reaction determined by the balanced chemical equation.
97.
Determine the number of moles of ALL the other reactants and products in the following
chemical reactions using the number of moles of the one substance given in the equation.
(These are 1-step problems!)
a. 1.5 moles of HBr
3 CaBr2 + 2 H3PO4  6 HBr + Ca3(PO4)2
0.75 moles CaBr2, 0.50 moles H3PO4, 0.25 moles Ca3(PO4)2
b. 2.2 moles CO2
C4H8 + 6 O2  4 CO2
0.55 moles C4H8, 3.3 moles O2, 2.2 moles H2O
c. 0.75 moles of Al
4 Al + 3 O2
0.56 moles O2, 0.38 moles Al2O3

+
2 Al2O3
d. 3.9 moles of FeCl3 3 KNO3 + FeCl3  3 KCl
12 moles KNO3, 12 moles KCl, 3.9 moles Fe(NO3)3
98.
4 H2 O
+
Fe(NO3)3
Determine the masses of ALL the other reactants and products in the following chemical
reactions using the mass of the one substance given in the equation. (These are tricky
problems! Think about why you were given two amounts of reactants.)
a. 200.0 g C3H6 and 200.0 g of O2
2 C3H6 + 9 O2 → 6 H2O + 6 CO2
O2 is the LR, C3H6 is in excess. There is 141.5g of C3H6 left over and 75.08g
H2O and 183.4g CO2 produced.
b. 45.9 g CuSO4 and 67.3 g of Fe(C2H3O2)3
3 CuSO4 + 2 Fe(C2H3O2)3 → 3 Cu(C2H3O2)2 + Fe2(SO4)3
CuSO4 is the LR, Fe(C2H3O2)3 is in excess. There is 22.6 g of Fe(C2H3O2)3 left
over and 52.2g Cu(C2H3O2)2 and 38.3g Fe2(SO4)3 produced.
c. 0.82 g GaBr3 and 1.0 g of F2
2 GaBr3 + 3 F2 → 2 GaF3 + 3 Br2
GaBr3 is the LR, F2 is in excess. There is 0.85g of F2 left over and 0.34g GaF3
and 0.64g Br2 produced.
d. 25 g KHCO3
2 KHCO3 → H2O + CO2
The reaction will produce 2.2g H2O, 5.5g CO2, and 17g K2CO3.
+
K2CO3
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UNIT 10 – BONDING
1.
What is a bond?
A bond is a connection between two atoms or ions using electrons that hold the two
particles together.
2.
What are the 3 main categories of bonds? What is the difference?
Metallic, ionic, and covalent. Metallic is the connections between metallic atoms, ionic
bonds are the attractions between two oppositely charged ions, and covalent bonds are the
attractions between two nuclei and a pair of electrons.
3.
What are the 2 types of covalent bonds? What is the difference?
Polar covalent and nonpolar covalent. Polar covalent is unequal sharing of electrons
and nonpolar covalent bonds occur when electrons are shared equally.
4.
What is a double bond? Triple bond?
Double bond is a covalent bond that shares two pair of electrons. Triple bond is a covalent
bond that shares three pair of electrons.
5.
What is the VSEPR Theory? Explain.
Valence Shell Electron Pair Repulsion Theory – Valence electron pairs around a central
atom will repel and move as far apart as they can to minimize repulsions. It is used to
predict molecular shapes.
6.
What are the 5 basic shapes formed by molecules with only bonds surrounding the center
atom? (No unbonded electrons)
Linear, Trigonal Planar, Tetrahedron, Trigonal Bipyramid, Octahedron.
7.
What are the shapes formed when there are unbonded pairs of electrons around the center
atom?
Bent, 117º (Trigonal Planar)
Trigonal Pyramid, Bent 104.5 º (Tetrahedron)
See-saw, T-Shaped, Linear (Trigonal Bipyramid)
Square Pyramid, Square Planar (Octahedron)
8.
Determine the shape formed by the following molecules.
a. SeF6 Octahedron
b. ClF3 T-Shaped
c. AsF5 Trigonal Bipyramid
d. NH3 Trigonal Pyramid
e. BF3 Trigonal planar
f. H2O Bent 104.5 º
g. SiBr4 Tetrahedron
9.
What is resonance?
Phenomenon that occurs in a molecule with single/double bonds between the same
elements. The extra bond will constantly switch between the atoms.
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UNIT 11 – GASES
Henry’s Law – The amount of a gas that can dissolve in a liquid varies directly with the partial
pressure of the gas. So if pressure increases, so does solubility.
10.
What does Boyle’s Law state? Its formula? What 2 factors must be held constant for this
law to be true?
Pressure and volume are inversely related to each other when temperature and the number
of moles remain constant. P1V1 = P2V2
11.
What does Charles’ Law state? Its formula? What 2 factors must be held constant for this
law to be true?
Volume and temperature are directly related to each other when pressure and number
of moles are held constant. V1/T1 = V2/T2
12.
What does the Combined Gas Law state? Its formula? What factor must be held constant
for this law to be true?
The combined gas law relates pressure, volume, and temperature of a gas when the number
of moles is kept constant. (P1V1)/T1 = (P2V2)/T2
13.
What is the Ideal Gas Law? Its formula?
The Ideal Gas Law relates all factors of one gas in one situation. PV = nRT
14.
What is R? What is its value?
R is the Universal Gas Law Constant, its value is 8.31 kPa∙L/mol∙K or 0.0821 atm∙L/mol∙K
15.
What unit does temperature have to be measured in to solve any gas law problem?
Kelvin
16.
How do you convert degrees Celsius into the unit from #15?
Celsius + 273
17.
Solve the following problems using the best Gas Law. Remember that standard
temperature is 0C and standard pressure is 1 atm, 101 kPa, 760 mm Hg or 760 torr.
a.
Oxygen gas kept at 45.00C has a pressure of 105.0 kPa. The pressure is decreased
to 90.00 kPa. What temperature will allow this to happen?
272.6 K
b.
What was the original pressure of Sulfur gas at 35.00C and 590.0 cm3 when it is now
in a 500.0 cm3 container with a pressure of 680.0 mm Hg and at 25.00C?
595.6 mm Hg
c.
What was the original volume of Hydrogen gas at 15C if it is now at 23C and
in a 2.0 L container?
1.9 L
d.
What is the original pressure of a gas in a 1.10 L container if its new volume is
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1.35 L with a pressure of 1.05 atm?
1.29 atm
e.
What is the new volume of Chlorine gas at STP when it was in a 2.5 L container
with a pressure of 720 torr at 33C?
2.1 L
UNIT 12 – SOLIDS & LIQUIDS
18.
Explain the difference between solids, liquids and gases in terms of molecular spacing?
Solid particles are tightly packed and in position, while liquid particles are close together
but are able to move past one another, and gas particles are very far apart.
19.
Explain the difference between solids, liquids and gases in terms of Kinetic Energy and
IMF?
Solids have a high IMF compared to their low KE, they do move – but it I sin position.
Liquids have more KE and less IMF than a solid which allows them to move around but
not separate from each other. Gases are very far apart and are constantly moving in a
straight line. They have high KE and low IMF.
20.
What is Heat of Fusion? Heat of Vaporization?
Heat of Fusion is the amount of heat that is required to completely melt one gram of a solid
at the melting point. Heat of Vaporization is the amount of heat required to completely
vaporize one gram of a liquid at the boiling point.
21.
Why does the temperature of a substance remain the same during a phase change?
The energy is being used to break the IMF of the substance rather than increase the
temperature.
22.
What is a crystal?
A rigid solid made up of a repeating pattern of atoms.
23.
What is a unit cell?
The simplest repeating pattern of atoms in a crystal.
24.
What is evaporation? How does evaporation occur?
Evaporation is the process by which a liquid becomes a gas below the boiling point. It
occurs when liquid particles at the surface collide and enough energy is transferred in the
collision to overcome the IMF and allows that particle to become a gas.
UNIT 13 – IMF
25.
What is IMF?
Intermolecular Forces are the attractive forces that occur between molecules.
26.
What is the importance of the IMF?
IMF allows for solids and liquids to occur.
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27.
What are the three types of IMF?
Dispersion forces, Dipole-dipole forces, and Hydrogen bonding.
28.
Which is the strongest form of IMF? Why?
Hydrogen bonding is the strongest because it occurs between two dipoles so the attraction
is constant and it occurs between molecules that contain a polar covalent bond between
Hydrogen and Nitrogen, Oxygen, or Fluorine. This bond is extremely strong since the
atoms are very small and allows the bond to be very close causing a greater partial charge.
29.
Which is the weakest form of IMF? Why?
Dispersion forces is the weakest form of IMF since it occurs between two momentary
dipoles. The force is not constant.
30.
What is a dipole?
A dipole is a molecule that has two opposite areas of partial charges due to polar covalent
bonds.
31.
What is a momentary dipole?
A momentary dipole is a dipole that was formed from a nonpolar molecule due to the
constant movement of electrons. At a given moment, the electrons may have arranged
themselves closer to one atom causing the partial charge to occur.
UNIT 14 – SOLUTIONS
32.
What is a solution?
A solution is another name for a homogeneous mixture.
33.
What are the 2 parts to a solution?
Solute and solvent.
34.
What is the difference between the 2 parts in #33?
Solute dissolves into the solvent.
35.
What is solubility?
Solubility is the ability of a solute to dissolve in a solvent at a given temperature.
36.
What are the three descriptions of solutions with respect to solubility?
Saturated, supersaturated, unsaturated.
37.
How do you know if a solution is saturated?
You know the solution is saturated if you add more solute and it doesn’t dissolve.
38.
What types of substances dissolve in each other?
Miscible substances completely dissolve in one another. Polar substances dissolve in polar
substances and nonpolar substances dissolve in nonpolar substances.
39.
What is a solution with water as the solvent?
Aqueous
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40.
What is molarity? What is the formula for molarity?
Molarity is a measure of the number of moles of solute per liter of solution.
41.
What is the molarity of a 5.0 L solution that contains 2.0 moles of K2CO3?
0.40 M
42.
How many moles of solute are there in 0.500 L of a 1.28 M NaNO3 solution?
0.640 moles NaNO3
43.
How many liters of solution will contain 1.5 moles of CaCl2 if the solution is 6.0 M CaCl2?
0.25 L solution
UNIT 15 – ACIDS & BASES
44.
What is an acid?
An acid is a substance that has Hydrogen ions (or hydronium Ions) to donate in reactions.
45.
What is a base?
A base is a substance that accepts Hydrogen ions (or Hydronium Ions) in reactions.
46.
What are the properties of acids? Bases?
Acids are electrolytes, react with metals to produce Hydrogen gas, sour, turn litmus paper
red, is colorless in phenolphthalein, and has pH values less than 7.
Bases are electrolytes, do not react with metals, bitter, slippery, turn litmus paper blue, is
pink in phenolphthalein, and has pH values greater than 7.
47.
What is a neutral substance?
A neutral substance has equal amounts of acidic and basic character and has a pH value of
exactly 7.
48.
What is the difference between weak acids and bases and strong acids and bases?
Strong acids and bases dissociate 100% while weak acids and bases dissociate less than
10%.
49.
List the strong acids.
Hydrochloric Acid – HCl, Hydrobromic Acid – HBr, Hydroiodic Acid – HI, Nitric Acid –
HNO3, Perchloric Acid – HClO4, Sulfuric Acid – H2SO4
50.
List the strong bases.
Lithium Hydroxide – LiOH, Sodium Hydroxide – NaOH, Potassium Hydroxide – KOH,
Rubidium Hydroxide – RbOH, Cesium Hydroxide – CsOH, Calcium Hydroxide –
Ca(OH)2, Strontium Hydroxide – Sr(OH)2, Barium Hydroxide – Ba(OH)2
51.
What is a conjugate acid/base pair?
A conjugate acid/base pair is the beginning and ending substance of a substance acting as
an acid and the result of that change and vice versa.
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52.
What is a neutralization reaction? What are the products?
A neutralization reaction is any reaction between an acid and a base. If the reaction is
between a strong acid and a strong base, the products will be salt and water.
53.
If the pH of a solution is 3.55, what is the pOH, [H3O+1], and [OH-1] of the solution?
pOH = 10.45, [H3O+1] = 2.82 x 10-4, [OH-1] = 3.55 x 10-11
54.
What is the Keq (or Ka) expression for HC2H3O2?
Ka = [H+1][C2H3O2-1]
[HC2H3O2]
UNIT 16 – REACTION RATES
Some reactions occur in several steps. Each step may occur at a different rate. The reaction can
only go as fast as its slowest step. This step is called the “rate-determining step”.
55. What is the rate of reaction?
The rate at which the products disappear or the rate at which the products appear. It is the
speed at which a reaction occurs.
56. How does a reaction occur?
A reaction occurs when two molecules collide with one another and then if they have enough
energy (activation energy) they will form an activated complex. This activated complex may
break apart to form new molecules (products).
57. What is an activated complex?
An activated complex is one large molecule that is formed when two reactant particles
collide and allow for the possibility of new particles to form.
58. What can increase the rate of reaction?
Rate of reaction can be increased by increasing the concentration of a reaction,
increasing the temperature, determining the nature of the reactants, and adding a catalyst.
59. What is used to slow down the rate of reaction?
An inhibitor slows down a reaction.
60. How does a catalyst speed up a reaction?
A catalyst speeds up a reaction by lowering the activation energy needed to form an activated
complex. This increases the chances that the activated complex will form more often in
order to produce products more frequently.
61. What is equilibrium?
Equilibrium occurs in a reversible reaction where the rate of the forward reaction is
equal to the rate of the reverse reaction.
62. State LeChatelier’s Principle.
Le Chatelier’s Principle states that a system at equilibrium will shift that equilibrium in order
to relieve an added stress.
63. What factors affect equilibrium?
Concentration of any substance in the system, temperature, and pressure affects the
equilibrium of the system.
64.
What is the equilibrium expression for the following reaction?
2 Mg + O2  2 MgO
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Keq =
[MgO]2 .
[Mg]2[O2]
65. Fill in the following table for the following reaction.
CH4 (g) + 2 O2 (g) + 35 kcal ↔ CO2 (g) + 2 H2O (l)
If…
Equilibrium
[CH4]
[O2]
Then…
Shift
[CO2]
[H2O]
[O2] ↑

↓
X
↑
↑
[CH4] ↑

X
↓
↑
↑
[CO2] ↑

↑
↑
X
↓
[H2O] ↓

↓
↓
↑
X
Pressure ↓

↑
↑
↓
↓
Temp ↓

↑
↑
↓
↓
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