Download Chemistry I Review - BarbaraElam-Rice

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MSL REVIEW
1) List the 5 postulates of Dalton’s Atomic Theory.
2) How was Dalton’s Atomic Theory proven to be wrong?
3) What was Democritus’s contribution to the early Atomic Theory?
4) Describe the results of the Cathode Ray, Oil Drop, and Gold Foil experiments.
5) As the wavelength gets longer, the frequency will ____________________. This is an example of an
_________________ relationship between wavelength and frequency.
6) The higher the energy, the greater the ____________________. This is an example of a
_______________________ relationship.
7) When an electron moves from a low energy to a high energy level energy is _________________.
8) When an electron moves from a high energy to a low energy level, energy is _____________ and a
_________________ is emitted.
9) Arrange the following from low to high energy: Infrared, Radio waves, Gamma rays, X-rays,
Microwaves.
10) Which type of light has the shortest wavelength and the highest frequency?
11) Which type of energy is produced when an electron moves from n=4 to n=2?
12) What are the 3 subatomic particles? Identify their location in the atom?
13) Define isotope. Identify the number of protons, neutrons and electrons for the following.
109
63
a. a. 107
Cu, 65Cu c. 112 Sn,115Sn
47 Ag , 47 Ag b.
14) When writing nuclear symbols, A represents the _________________ and Z represents the
__________________.
15) What is the difference between the atomic mass and mass number?
16) Write the hyphen notation and nuclear symbol for an atom that has 13 protons and 14 neutrons.
17) Identify the names for the following symbols: s, l, g, aq. Describe each phase. Identify the phase that
has the greatest amount of disorder.
18) What is the difference between a ternary and a binary compound?
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19) Write the formula for the following compounds. Identify the bond type and if the compound is ionic or
molecular.
a. a. Magnesium chloride b. sodium hydroxide
c. copper II sulfide
b. d. carbon monoxide
e. aluminum sulfate
f. ammonium phosphate
c. g. dinitrogen pentoxide
20) Write the name for the following compounds. Identify the bond type and if the compound is ionic or
molecular.
a. a. CuCl
b. ZnSO4
c. CCl4
b. d. CaCO3
e. NaOH
f. P2O5
c. g. PCl3
h. CuBr2
i. FeO
21) Write the name for the following acids.
a. a. HCl
b. HNO3
b. d. HC2H3O2
c. H2SO4
22) Write the formula for the following acids.
a. a. hydrobromic acid
b. phosphoric
b. d. carbonic acid
e. chloric acid
c. hydroiodic
23) A student collected the data shown below to determine experimentally the density of distilled water.
a. Mass of graduated cylinder + distilled H2O sample…163g
b. Mass of empty graduated cylinder….141g
c. Mass of distilled H2O sample….____g
d. Volume of distilled H2O sample….25.3mL
24) A cube has a volume of 8.0 cm3 and a mass of 21.6 grams. The density of the cube is…
25) A chemistry student is given the task of analyzing three unknown samples. Her data is listed in Data
Table 1. Use Data Table 1 to answer the questions below.
Data Table 1
Trial
Trial 1
Trial 2
Trial 3
Average
Sample A
Mass
Volume
(in
(in mL)
grams)
80.72
10.01
80.64
10.00
80.91
10.05
80.76
10.02
Sample B
Mass
(in
grams)
95.41
92.33
93.78
93.84
2
Volume
(in mL)
10.72
10.51
10.62
10.62
Sample C
Mass
(in
grams)
72.28
72.32
72.34
72.30
Volume
(in mL)
10.00
9.99
9.95
9.98
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The student compares her data to the following chart in the back of her textbook. Can she identify
Samples A, B, and C based on the data she recorded?
Table A Properties of Common Metals
Name
Color at room
temperature
Aluminum
Silver metal
Copper
Red metal
Iron
Silver metal
Nickel
Silver metal
Tin
White metal
Density
2.701
8.92
7.86
8.90
7.28
26) Review Handout on Solubility Curves
27) The Diagram below represents four 500-milliliter flasks. Each flask contains the gas represented by its
symbol. All gas samples are at STP. Each flask contains the same number of atoms/molecules/atoms
and molecules.
28) Gas Laws
a. A sample of a gas is at STP. As the pressure decreases and the temperature
b. increases, the volume of the gas _______________.
c. Under which condition will the volume of a given sample of a gas decrease
1. decreased pressure and decreased temperature.
2. decreased pressure and increased temperature.
3. increased pressure and decreased temperature.
4. increased pressure and increased temperature.
d.
e.
f.
g.
h.
i.
j.
If 10 liters of H2 (g) at STP is heated to a temperature of 546 K, pressure
remaining constant, the new volume of the gas will be 20 liters.
d. Draw the graph representing Boyle’s Law. Label the X & Y axis.
e. Draw the graph representing Charles’s Law. Label the X & Y axis.
f. Draw the graph representing Gay-Lussac’s Law. Label the X & Y axis.
g. At a temperature of 273 K, a 400 milliliter gas sample has a pressure of 769
millimeters of mercury. If the pressure is changed to 380 millimeters of mercury at
which temperature will this gas sample have a volume of 551 milliliters?
How many moles of gas does it take to occupy 120 liters at a pressure of 2.3 atmospheres and
temperature of 340K?
29) Describe the difference between cations and anions. How are they formed?
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30) An atom has 2 electrons in its valence shell. To what group does the atom belong? Will the atom for
an anion or a cation? What will be the oxidation number of the ion?
31) An intermolecular force that holds ionic compounds together is called electrostatic attraction.
32) Describe the 3 intermolecular forces? Which of these forces is the strongest? weakest?
33) How are intermolecular forces different from chemical bonds? Which is stronger?
34) If the electronegativity difference between atoms is greater than 1.7, what type of bond will form? If
the difference is less than 1.7, what type of bond will form?
35) What type of molecule has an electronegativity difference of 0?
36) Ionic compounds have high/low melting points, high/low boiling points, and can conduct electricity in
the _________________ state and in ____________________solutions.
37) Metallic bonds can be described as a _____________ of ______________ electrons.
38) List 6 properties of metals.
39) Atoms bond to become ________________________. _________________ energy lowers when
unstable atoms bond.
40) Which group on the table does not need to bond in order to be stable?
41) Which diatomic molecules have single covalent bond? double covalent bond? triple covalent bond?
Draw the Lewis Structure for each.
42) Draw Lewis Structures for the following. Identify the geometry and polarity of each molecule.
a. a. PH3
b.CCl4
c. CO2
d. SO3
e. H2O
f. H2S
b. g. SO2
h. NH3
43) Which bond type is the shortest…single, double or triple? Which has the greatest bond
energy….single, double or triple?
44) Draw the Lewis structures for the following molecules, I2, O2, N2. Which has the shortest bond length?
Which has the longest bond length? Which has the greatest bond energy?
45) The process that occurs when a solid turns into a gas is called ______________.
46) When looking at a heating/cooling curve, phase changes occur at __________________ temperature.
47) When held at constant pressure, what change will occur when the pressure is increased over boiling
water?
48) Identify the 6 phase changes that water can undergo.
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49) Identify the number of main groups on the periodic table and the name of each one.
50) How many periods are on the periodic table?
51) Groups are __________________ columns and periods are _________________ rows on the periodic
table.
52) Elements within the same _________________ have similar chemical properties, the same
_______________ number and the same number of _____________ ____________.
53) ______________ ________________ are responsible for the chemical reactivity of elements.
54) As you move _____________ the group, the reactivity of ____________ increases. As you move
_______ the group for _______________, the reactivity increases.
55) The ____________ metals are located in the ____________ of the periodic table.
56) _____________ numerals are used to identify the ____________________ number of transition
metals.
57) Write electron configurations and noble gas configurations for the following elements/ions.
a. Ne
b. Ca
c. Au
d. Br
e. Mg
f. Potassium ion
g. Phosphide
h. Aluminum ion
j. Sulfide
k. O2l. Mg2+
58) Identify the following elements based on the electron configuration.
a. 1s22s22p63s23p1
b. 1s22s22p63s23p63d84s2 c. [Ne} 3s1
d. [Ar]3d34s2
59) S block elements include groups ________. P block elements include groups__________. D block
elements are groups ___________. F block elements are located _______________.
60) Identify the number of valence electrons from the following electron configurations in problems 58 a-l
and 59 a-d
61) How many electrons were lost/gained for letter 58 f-l? What is the oxidation number for each?
62) Define ionization energy. Write the group and period trend. Arrange the following elements in order
of decreasing ionization energy.
a. a. Cl, Br, F, I
b. C, B, N, O
c. Ca, Sr, Mg, Ba
63) Define atomic radius. Write the group and period trend. Arrange the following elements in order of
increasing atomic radius.
a. a. Cl, Br, F, I
b. C, B, N, O
c. Ca, Sr, Mg, Ba
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64) Define ionic radius. Write the group and period trend. Arrange the following elements in order of
decreasing ionic radius.
a. a. Cl-, Br-, F-, I-,
b. P3-, S2-, N3-, O2c. Ca2+, Sr2+, Mg2+, Ba2+
65) Define electronegativity. Write the group and period trend. Arrange the following elements in order of
increase electronegativity.
a. a. Cl, Br, F, I
b. C, B, N, O
c. Ca, Sr, Mg, Ba
66) Complete the following conversions.
a. a. 4.2 moles C to atoms
b. c. 1.6 x 1019 molecules H2 to moles
c. e. 4.7g C to atoms
d. g. 5.2 x 1012 molecules CH4 to g
e. 0.23 moles H2 to L
b. 2.94 moles CO2 to molecules
d. 6.1 x 1026 atoms S to moles
f. 2.73 g SO2 to molecules
h. 4.7L CH4 to moles
67) 1 mole of any gas at STP equals ______________.
68) How many grams of ammonium chloride are contained in 0.5 L of a 2 M solution?
69) What is the molarity of a solution on KNO3 that contains 404 grams of KNO3 in 2L of solution?
70) What is the molarity of a solution containing 20 grams of sodium hydroxide in .5 L of solution?
71) What is the total number of moles in 2 liters of 3 M of sodium hydroxide?
72) The empirical formula of a compound is CH2 and its molecular mass is 70. What is the molecular
formula of the compound?
73) If the empirical formula for an organic compound is CH2O, then the molecular mass of the compound
could be….135,45,60,15.
74) A compound consists of 40% sulfur and 60% oxygen by mass. What is the empirical formula of this
compound?
75) What is the empirical formula of a compound if a sample contains 8.52 grams of carbon and 1.42
grams of hydrogen?
76) What is the percent by mass on nitrogen in the compound NH4NO3?
77) Given the unbalanced equation N2(g) + H2(g)  NH3(g), when the equation is balanced, the ratio of
moles of hydrogen consumed to moles of ammonia produced is _______.3:2
78) Na + H2O NaOH + H2 What is the total number of moles hydrogen produced when 4 moles of
sodium react completely? 2
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79) C2H2(g) + O2(g)  CO2 + H2O(g) What is the total number of grams of O2(g) needed to react
completely with 0.50 mole of C2H2? 40g
80) CO2(g) + H2O(l)  CO2(s) + O2(g) What is the minimum number of liters of CO2(g) measured at STP,
needed to produce 32 grams of oxygen? 22.4L
81) H2 + O2  H2O The total number of grams of O2 needed to produce 54 grams of water is….48g
82) Write the balanced equation for the reaction of lead (II) nitrate with sodium iodide to form sodium
nitrate and lead (II) iodide.
a. If I start with 25.0 grams of lead (II) nitrate and 15.0 grams of sodium iodide, how many grams of
sodium nitrate can be formed?
b. What is the limiting reagent in the reaction?
c. How much of the excess reagent is used? How much remains?
83) Which 2 solutions when mixed together will undergo a double replacement reaction and form a white
solid substance? a. NaCl(aq) and LiNO3(aq)
b. KCl(aq) and AgNO3(aq) c. KCl(aq) and LiCl(aq)
d.
NaNO3 and AgNO3(aq)
84) Be able to distinguish between single replacement, double replacement, synthesis, and decomposition
reactions. Refer to reference tables
85) Predict the products of the following reactions. Use the reference table for assistance.
a. ____ Ca(OH)2 + ____ HF 
b. NaC2H3O2 + ____ H2SO4 
c. Aluminum bromide and chlorine gas react to form …
86) What must be present for combustion to take place? What are the products of a combustion
reaction?
87) Neutralization is a type of _________________ reaction. Neutralization occurs between
_____________and _______________. The products of this reaction are _____________ and
____________.
88) An atom in an _______________ state is above its ground state because of the absorption of
___________ resulting _____________ moving to a ___________ energy level. When the atoms
returns to its ___________ state, the electron releases energy in the form of a ___________ that
resembles energy from the electromagnetic spectrum.
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89)
a.
b.
c.
d.
e.
f.
What is the boiling point of the substance?
Solid and liquid phases can exist in equilibrium between points….
Which segments represent potential energy?
Which segments represent kinetic energy?
Is this an example of a heating or cooling curve?
Which segment(s) represent an endothermic process?
90) Endothermic means heat energy is going ________; Exothermic means heat energy is going _______.
91)
a.
b.
c.
d.
e.
At point A, what phase change will take place as the pressure decreases and
temperature remains constant?
At what point do all three phases exist in equilibrium?
At point J, what phase change will take place as the temperature decreases and
pressure remains constant?
92) . The universe tends towards __________. This is called ____________. Arrange the following in order
of increasing entropy—g, l, s; ionic solutions, ionic compound
93) Write the symbols for: a. alpha particle_______________; b. beta particle______________; c. gamma
particle _________________.
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94) In the equation X
95) In the equation
234
91
226
88
Ra  24He , X represents…
Pa  X + 10 e , X represents…
96) Define half-life. Review the half-life calculations.
97) What is the different between fusion and fission?
98) In order to satisfy the Law of Conservation of Mass, chemical equations must be ________________.
99) In a balanced chemical equation, the mass of the ____________ = the mass of the _____________.
100)
Describe the difference between a total/complete ionic equation and a net ionic equation.
a. NaCl(aq) +
Pb(NO3)2(aq) 
PbCl2(s)
+
NaNO3(aq)
101)
If ∆H+=endothermic; ∆H-=exothermic
102)
List properties of acids/bases.
103)
What is the formula for concentration?
104)
What is the difference between strength and concentration as related to an acid?
105)
Identify the 6 strong acids and bases.
106)
Write the equation for dilution.
107)
If the pH is below 7, the substance is a(n) ____________, above 7 is a(n) ___________, and
equal to 7 is ______________.
108)
Litmus is __________in an acid and ____________in a base.
109)
110)
111)
112)
Write the equation used to calculate pH.
Review the pH and pOH calculations.
Review the dilution calculations.
Name the following acids:
a. HCl
b. H2NO3
c. H3PO4
d. HBr
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113)
a.
b.
c.
d.
e.
f.
g.
h.
Which letter represents the activation energy?
Is this a diagram of an endothermic or exothermic reaction?
Which letter represents the activation energy for the reverse reaction?
Is the reverse of the reaction endothermic or exothermic?
Which letter represents the potential energy?
Which letter represents the heat of reaction?
What will change as a result of the addition of a catalyst.
Review the handouts on potential energy. You should be able to do the calculations.
114) Given the equation 12.6 kcal + H2(g) + I2(g)  2HI(g), identify the equilibrium shift based on the
changes.
a. Add H2
b. Remove I2
c. Increase temperature
d. decrease pressure
e. Increase pressure
115) Given the equation NaOH(s) Na+(aq) + OH- (aq) + 10.6 kcal HINT: pure solids and liquids do
not affect equilibrium.
a. Add NaOH
b. Increase temperature
c. decrease pressure
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