Download Review # 3

HONORS CHEMISTRY - COMPREHENSIVE REVIEW FOR
MIDTERM EXAM
Topics to be covered:
I.
Introduction
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II.
Atoms, Molecules and Ions
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III.
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IV.
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V.
Scientific Method
Metric system
Density
Factor label method for solving problems
Measurements: calculations, significant figures, conversion factors, conversion between units, error analysis
Calculations: moles, molar mass, molar volume, molarity, stoichiometry
Classification of Matter: Elements, Compounds, Mixtures
Physical/Chemical properties/changes
History of Atomic Theory: key milestones, persons, and experiments
Atomic structure
Nuclear notations
Average atomic mass
Nuclear stability
Radioactive decay: ,  particles and  ray;  and  decays
Nuclear equations
Half-life, radioactive dating
Electron Configurations
Waves: main function is to transfer energy, frequency () and wavelength (),  = speed of the wave
Electromagnetic Radiation: no need for medium to travel and constant speed in vacuum,  = speed of light; E =
h
Atomic Emission Spectra and the flame test
Bohr’s Model
Duality of electrons, de Broglie
Heisenberg Uncertainty Principle
Schrodinger’s Quantum Mechanical Model: two principles (Pauli Exclusion and Aufbau) and a rule (Hund’s)
Ground state electron configurations, based on which valence shell electron configuration, Lewis dot structure are
established
The Periodic Table
History of development of the Periodic Table
Major features of the Periodic Table: groups/families (names for major groups), periods
Metals, nonmetals, metalloids
Correlations between ground state electron configurations and the Periodic Table: period number and principal
energy levels; valence electrons, Lewis dot structures
Octet Rule
Formation of Ions
Bonding: Ionic, Covalent and Metallic
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VI.
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Formation of ions
Ionic bond
Formulas and names of ionic compounds
Covalent bond
Formulas for covalent/molecular compounds
Metallic bond
Correlations between microstructures and macroscopic properties of pure substances: mainly physical properties of
ionic, covalent compounds and metals
Lab Skill
Bunsen Burner
Safety
Mass and Volume Measurement
Balance
Physical and Chemical Changes
% error
Working with crucibles, evaporating dishes
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Review 1
1.
A block of material measures 0.20 m by 150. cm by 6.0 mm and has a mass of 2.6 kg. What is its density in
g/cm3?
2.
24 g of Mg contain 6.02 x 1023 atoms. How many atoms are in 750 mg of Mg?
3.
Magnesium, Mg, has a density of 7.25 g/cm3. Given 12.0 cm3 of Mg
a. how many moles of Mg atoms does this represent?
b. how many atoms of Mg are contained in this volume?
( 3.63 mol)
(2.18 x 1024 atoms)
4.
How many molecules of aspirin are contained in a 100.0 mg tablet of aspirin, C 9H8O4?
(3.34 x 1020 molecules)
5.
What is the volume of 1.00x102 g of CH4 gas at STP?
6.
Calculate the density of an unknown object that has a mass of 55 g and occupies a volume of 17 mL.
(3.23 g/mL)
7
What is the mass of 2.47 cm3 of platinum? The density of platinum if 22.5 g/cm3.
8.
Write formulas for these compounds:
9.
10.
( 1.40x102 L)
a. sodium carbonate _____________________
c. iron(III) hydroxide _____________________
b. silicon tetrachloride __________________
d. ammonium iodide ______________________
Name these compounds:
a. KNO3 ___________________________
e. AgF ___________________________
b. CuO ___________________________
f. Al2(SO4)3 ___________________________
c. Mg3N2 ___________________________
g. H3PO4 ___________________________
d. CCl4 ___________________________
h. Ca(OH)2 ___________________________
State whether PE increases, decreases, or remains the same
a. falling ball
_________________
b. climbing rocket _________________
c. heating H2O at its boiling point _________
d. water going over a dam _________________
e. satellite circling the earth at constant altitude _________________
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Review # 2
Complete the work, whenever appropriate, on a separate sheet of paper SHOW all work
1.
2.
Make the following conversions:
a. 2.84 kilogram to gram
___________________________________________
b. 544 milliseconds to seconds
___________________________________________
c. 0.0656 gram to milligram
___________________________________________
d. 1102 cm to meters
___________________________________________
e. 3.6 nanometers to meters
___________________________________________
Use the factor label method to make each of the following conversions:
a. 5.235 x 10 3 ug to mg ___________________________________________
b. 63 dm3 to cm3
___________________________________________
c. 3.9 x 10 5 nm to m
___________________________________________
d. 0.128 L to mL
___________________________________________
e. 18.3 km to cm
___________________________________________
f. 33.4 mg to kg
___________________________________________
3.
Compute the density of clay if 42 g occupy 19.1 cm3.
4.
What is the number of significant figures in each of the following measurements?
a. 0.558 g ________
d. 0.0094 m ________
b. 7.3 m __________
e. 19.0000 g ________
c. 410 cm ________
f. 75.0 s _______
5.
Classify the following properties as physical or chemical.
a. flammability ____________
b. electrical conductivity _________________
c. ability to displace hydrogen from water _____________
d. ability to react with acids __________________
6.
Classify the following changes as physical or chemical.
a. distillation __________________
b. fermentation _______________
c. crystallization _______________
d. dissolving
7.
Element D has oxidation numbers 1+ and 2+. Element E has oxidation numbers 1- and 2-. List the
possible formulas of compounds of these elements.
8.
Density and speed are expressed in derived SI units. Give these units and tell why they are called derived units.
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9.
Make the following conversions.
a. 2.52 x 1021 formula units of ZrS2 to moles.
___________________________________________
b. 1.26 x 1025 formula units of Al(CH3COO)3 to grams. __________________________
c. 6.06 grams of iron(III) sulfate to moles. ___________________________________________
d. 88.4 grams MnI2 to moles. ___________________________________________
d. 0.00202 moles of nickel(II) hydroxide to grams. ________________________________________
10.
Tungsten is Swedish for “heavy stone”. Its symbol, W, comes from wolfram the German name for the
element. It has a density of 19.3 g/cm3. How many atoms are in a cubic centimeter of tungsten?
11.
Write names for the following compounds:
a. Sr(CH3COO)2
___________________________________________
12.
b. Mn(OH)2
___________________________________________
c. CdSO4
___________________________________________
d. Li3N
___________________________________________
e. Th3(PO4)4
___________________________________________
f. Ce2(CO3)3
___________________________________________
Write formulas for the following compounds:
a. sodium nitride
___________________________________________
b. sulfurous acid ___________________________________________
c. cerium(III) sulfide
___________________________________________
d. barium iodate ___________________________________________
e. hydrogen telluride
___________________________________________
f. silver sulfate
___________________________________________
g. hypochlorous acid
__________________________________________
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Review # 3
1.
The amount of mass per unit volume refers to the
a. density
b. specific weight
c. volume
2.
Which f the following is a physical property of sugar?
a. Its composition is carbon, hydrogen, and oxygen.
b. It turns black with concentrated sulfuric acid.
c. It can be decomposed with heat.
d. It is a white crystalline solid.
3.
A substance that can be farther simplified may be either
a. an element or a compound
c. a mixture or a compound
b. an element or a mixture
d. a mixture or an atom.
4.
A substance composed of two or more elements chemically combined is called
a. an isotope
b. a compound
c. an element
5.
An example of a chemical change is the
a. breaking of a glass bottle
b. rusting of iron
d. weight
d. a mixture
c. sawing a piece of wood
d. melting of an ice cube
6.
A substance which cannot be farther decomposed by ordinary chemical means is
a. water
b. air
c. sugar
d. silver
7.
An example of a physical change is the
a. fermenting of sugar to alcohol
b. burning of paper
c. rusting of iron
d. dissolving sugar in water
8.
Chemical action may involve one of the following except:
a. combining of atoms of elements to form a molecule
b. separation of the molecules in a mixture
c. breaking down compounds into elements
d. reacting a compound and an element to form a new compound and a new element.
9.
The energy of a system can be
a. easily changed to mass
b. transformed into a different form
10.
11.
c. measured only as potential energy
d. measured only as kinetic energy
What is the approximate formula mass of Ca(NO3)2?
a. 70
b. 82
c. 102
d. 150
Which of the following is classified as a diatomic element?
a. water
b. milk
c. chlorine
12.
In this reaction XClO3 + A → XCl
a. X
b. XClO3
13.
How many atoms are represented in the formula Ca3(PO4)2?
a. 5
b. 8
c. 9
d. 12
14.
15.
+ O2 +
c. A
The oxidation number of sulfur in H2SO4 is
a. +2
b. +3
c. +4
e. 164
d. ozone
A which substance is the catalyst?
d. XCl
e. O2
d. +6
e. 13
e. +8
The present scale of atomic mass is based on 1 amu being equal to the mass of:
a. 1 hydrogen atom
c. 1/10 of an oxygen-16 atom
b. 1/12 of carbon-12 atom
d. 1/32 of an oxygen molecule
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16.
If the density of a diatomic molecule of gas is 1.43 g/L, what is its gram-molecular mass?
a. 16 g
b. 32 g
c. 48 g
d. 64 g
e. 14.3 g
17.
The correct formula for calcium hydrogen sulfate is
a. CaH2SO4
b. CaHSO4
c. Ca(HSO4)2
d. Ca2HSO4
18.
Which of the following involves a physical change?
a. The formation of HCl and H2 from H2 and Cl2
b. The color change when NO is exposed to air
c. The formation of steam from burning H2 and O2
d. The solidification of corn oil at low temperatures
e. the odor of NH3 when NH4Cl is rubbed together with Ca(OH)2 powder?
19.
How many atoms are in one mole of water?
a. 3
b. 6.02 x 1023
c. 2 (6.02 x 1023) d. 54
e. Ca2H2SO4
e. 3 (6.02 x 1023 )
20
The Law of Definite Composition is based on definite composition by
a. mass
b. volume
c. density
d. specific mass
e. freezing point and boiling point
21.
Which of the following atoms normally forms monatomic molecules?
a. Cl
b. H
c. N
d. He
23.
What is the mass of 1 mole of KClO3 . 12 H2O?
a. 132
b. 180
c. 339
24.
d. 516
e. 474
How many atoms are present on the formula KAl(SO4)2?
a. 7
b. 9
c. 11
d. 12
e. 13
25.
If a crystal of CuSO4 . 5H2O is heated, what will the result be?
a. a blue crystal
b. a pink crystal
c. a blue powder
d. a white crystal
e. a white powder
26.
Which of the following is not a compound?
a. copper sulfate
b. carbon dioxide
d. air
e. lime
27.
34.
35.
c. sugar
What volume would 1.5 moles of hydrogen gas occupy at STP?
a. 11.2 L
b. 22.4 L
c. 33.6 L
d. 44.8 L
One mole of water contains
a. 18 grams
b. 6.02 x 10-23 atoms
d. 6.02 x 10 -23 molecules
c. 6 x 1023 ions
A gas at STP which contains 6.02 x 1023 atoms and forms diatomic molecules will occupy
a. 11.2 L
b. 22.4 L
c. 33.6 L
d. 67.2 L
e. 67.2 L
e. 1.06 quarts
36.
A compound whose molecular mass is 90 contains 40.0% carbon, 6.67% hydrogen, and 53.33 % oxygen. What is
the molecular formula of the compound?
a. CH2O
b. C3H6O3
c. C2H2O2
d. C2H2O3
37.
The simplest unit of water that retains its property is called
a. an atom
b. an element
c. a proton
38.
d. a hydroxide
The number of atoms of nitrogen represented in the formula NH4NO3 is
a. 1
b. 2
c. 3
4. 4
5.
e. a molecule
5
For questions 39-40 : What is the apparent oxidation number of the underlined element in each of the following
compounds?
a. +2
b. -2
c. +3
d. -3
e. +5
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39.
NaNO3
40.
CaSO4
41.
NH3
42.
How many grams of sulfur are present in 1 mole of H2SO4?
a. 2
b. 32
c. 49
d. 64
43.
45.
46.
47.
50.
e. 98
What is the approximate mass of 1 liter of dinitrogen monoxide, N2O, at STP?
a. 1 g
b. 2 g
c. 11.2 g
d. 22 g
What is the gram-formula mass of calcium carbonate?
a. 68 g
b. 75 g
c. 82 g
d. 100 g
What is the mass in grams of 2.0 moles of NO2?
a. 92
b. 60
c. 46
e. 44 g
e. 116 g
d. 30
Which sample contains a total of 3.00 x 1023 molecules?
a. 14 g N2
b. 14 g of Li
c. 4.0 g of H2
d. 4.0 g of He
A Ca (+2) ion differs from a Ca atom in that the Ca ion has
a. more protons
c. more electron
b. fewer protons
d. fewer electrons
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Review # 4
1.
What is the total number of nucleons (protons and neutrons) in an atom of selenium ( at. # = 34; mass # = 79)
a. 34
b. 45
c. 79
d. 113
2.
Isotopes of an element have a different
a. number of electrons
c. number of protons
b. atomic number
d. mass number
3.
A neutral oxygen atom, O, differs from an oxide ion in that the atom has
a. more electrons
c. more protons
b. fewer electrons
d. fewer protons
4.
What is the total number of neutrons in an atom of K, whose mass is 39, and atomic number is 19?
a. 19
b. 39
c. 20
d. 58
5.
The amount of hydrogen chloride that the formula HCl represents is one
a. atom
b. gram
c. liter
d. molecule
6.
The mass number of an atom is equal to the total number of its
a. electrons only
c. protons only
b. electrons and protons
d. protons and neutrons
7.
The number of protons in an atom of Cl - 36 is
a. 17
b. 35
c. 18
d. 36
8.
When a chlorine atom reacts with a sodium atom to form an ion, the chlorine atom will
a. lose one electron
c. lose two electrons
b. gain one electron
d. gain two electrons
9.
Element X exists in three isotopic forms. The isotopic mixture consists of 10.0 % X-10, 20.0 % X-11, and
70.0 % X-12. What is the average atomic mass of this element?
a. 11.0 amu
b. 12.0 amu
c. 11.6 amu
d. 12.4 amu
10.
The nucleus of a fluorine atom has a charge of
a. 1+
b. 19+
c. 9+
11.
The correct formula for nickel (III) sulfate is
a. Ni2S3
b. Ni2(SO4)3
c. Ni3S2
d. 0
d. Ni3(SO4)2
12.
What is the total number of neutrons in an atom of fluorine, whose atomic number is 9 and whose mass is 19?
a. 9
b. 19
c. 10
d. 28
13.
Which particle has exactly a mass of approximately 1 mass unit and a unit positive charge?
a. neutron
b. electron
c. proton
d. alpha particle
14.
When an atom of bromine becomes a bromide ion, its size
a. decreases
b. increases
c remains the same
15.
What is the total number of electrons in a magnesium (+2) ion?
a. 10
b. 12
c. 2
d. 24
16.
Which species has a negative charge?
a. a lithium ion
b. an aluminum ion
17.
Which particle is electrically neutral?
a. proton
b. positron
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c. an electron
c. neutron
d. a sodium atom
d. electron
9
18.
An atom that contains 35 protons, 45 neutrons, and 35 electrons has an atomic number of
a. 35
b. 45
c. 80
d. 115
19.
Two isotopes of the same element will have the same number of
a. neutrons and electrons
c. neutrons and nucleons
b. protons and nucleons
d. protons and electrons
20.,
Which sample contains the same number of atoms as 24 grams of carbon?
a. 80. g Ar
b. 10. g He
c. 24 g Mg
d. 4.0 g He
21.
How many neutrons are in the nucleus of an atom that has an atomic number of 17 and mass number of 35?
a. 17
b. 18
c. 35
d. 52
22.
Isotopes are atoms which have different
a. atomic masses
b. atomic numbers
c. atomic radii
d. number of electrons
23.
What is the total charge on anion that contains 10 electrons, 13 protons, and 15 neutrons?
a. -1
b. +1
c. -3
d. +3
24.
The element whose properties are most similar to those of tellurium is ______.
a. Be
b. O
c. S
d. Po
25.
How many electrons are in a neutral atom of Li?
a. 4
b. 7
c. 10
d. 3
26.
The number of protons in the nucleus of carbon-13 is
a. 19
b. 13
c. 7
d. 6
27.
The mass number of an atom is equal to the total number of its
a. electrons, only
c. protons, only
b. electrons and protons
d. protons and neutrons
28.
Which pair of atoms are isotopes?
12
a.
29.
6
12
C and
7
39
N
b.
19
K and
18
38
Ar
226
c.
88
Ra and
222
86
Rn
Which set of particles are arranged in order of increasing mass?
a. H2, H, H+
b. H+, H, H2
c. H2, H+, H
39
d.
19
40
K and 19K
d. H, H+, H2
30.
The existence of fractional atomic masses is best explained by the
a. mass of electrons
c. existence of isotopes
b. inaccuracies in determining masses
d. varying number of protons in the nucleus of the atom
31.
Which pair of nuclei contain the same number of neutrons?
a. Li-7 and Be-9 b. K - 40 and Ar - 40
c. Na -23 and Na-22
32.
33.
34.
An atomic mass unit is equal to ____.
a. the mass of a hydrogen atom
b. the mass of a C-12 atom
d. Na - 23 and Mg -24
c. 1/1836 the mass of hydrogen atom
d. 1/12 the mass of a C-12 atom
The mass number of an atom is equal to the total number of its ____.
a. electrons, only
b. protons, only
c. electrons and protons
d. protons and neutrons
Which atom has a mass of approximately two atomic mass units?
3
a.
H
1
2
b.
H
1
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c.
He
2
4
d.
He+2
2
10
35
35.
The number of protons in an atom of 17Cl is ____.
a. 17
b. 18
c. 35
d. 36
36.
Which nucleus contains a total of 2 protons and 1 neutron?
1
a.
H
1
3
b.
3
H
c.
1
4
He
d.
2
He
2
1
37.
38.
What is the mass number of an
a. 1
b. 2
H atom?
c. 3
1
d. 4
The neutral atoms in a given sample of an element could have different
a. mass number
b. number of protons
c. atomic numbers
d. number of electrons
39.
The accepted value for the boiling point of a substance is 120 oC. A student performs an experiment and reports
the boiling point to be 110 oC. The percent error of the student’s observation is
a. 83%
b. 120%
c. 20 %
d. 80%
e. 8.0%
40
Which of the following does not have a net charge?
a. a proton
b. a neutron
c. an electron
d. an ion
e. the nucleus
41.
Which of the following pairs show the principle of the Law of Multiple Proportions?
a. H2O(s) and H2O(aq)
c. NaCl and KCl
b. N2O and NO2
d. C2H5Cl and C2H5Br
42.
What is the total number of subatomic particles in an atom having an atomic number of 17 and atomic mass of 35?
a. 17
b. 18
c. 34
d. 35
e. 52
43.
Which of the following is not part of Dalton’s Atomic Theory?
a. Atoms of a given element have the same mass.
b. All elements are made of indivisible, indestructible atoms.
c. Atoms of one element can be converted into a different element.
d. Compounds are the result of the combination of atoms of different elements.
e. All atoms of a given element are identical.
44.
The term nucleon refers to
a. protons, only
b. neutrons, only
d. both protons and neutrons
45.
c. electrons, only
e. both protons and electrons
What is the net charge on an atom with 20 protons, 20 neutrons, and 18 electrons?
a. 0
b. 2+
c. 2d. 18e. 20+
Note:
For practice questions regarding electron configurations, the Periodic Table and
bonding, review those interactive multiple-choice questions posted on Schoolwire
site. Homework assignments are good resource in preparing for the stoichiometry
calculations.
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Midterm Review: Answer Key
Review 1:
1. 10 g/cm3 (1.4 g/cm3)
7. 55.6 g Pt
8. Na2CO3
b. SiCl4
9. a. potassium nitrate
d. carbon tetrachloride
g. phosphoric acid
10. a-D
b-I
c-I
2. 1.9 x 1022 atoms
c. Fe(OH)3
d. NH4I
b. copper(II) oxide
e. silver fluoride
h. caclium hydroxide
d-D
e-constant
c. magnesium nitride
f. aluminum sulfate
Review #2:
1.a. 2840 g
b. 5.44 x 10-1 sec
2. a. 5.235 mg
b. 6.3 x 104 cm3
6
e. 1.83 x 10 cm
f. 3.34 x 10-5 kg
3. 2.2 g/cm3
4. a. 3 b. 2
c. 2
5. a. ch
b. ph c. ch d. ch
7. DE; D2E; DE2
8. g/cm3; m/sec
d. 0.287 mol
e. .187 g
10. 6.32 x 1022 atoms
11.a. strontium acetate
d. lithium nitride e. thorium(IV) phosphate
12. a. Na3N
b. H2SO3
c. Ce2S3
c. 65.6 mg
c. 3.9 x 10-4 m
d. 11.02 m
e. 3.6 x 10-9 m
d. 128 mL
d. 2
e. 6
f. 3
6. A. ph b. ch c. ph
9. A. 4.19 x 10-3 mol
d. ph
b. 4.27 x 103 g
c. 0.0152 mol
b. manganese(II) hydroxide
c. cadmium sulfate
f. cerium(III) carbonate
d. Ba(IO3)2
e. H2Te
f. Ag2SO4
g. HClO
Review #3:
1-a
14-d
27-c
39-e
2-d
15-b
3-c
16-b
5-b
18-d
35-a
43-b
6-d
7-d
19-e
20-a
36-C3H6O2
44-d
45-d
8-b
21-d
37-e
46-a
9-b
41-d
4-b
17-c
34-a
42-b
40-a
3-b
16-c
29-b
42-e
4-c
17-c
30-c
43-c
5-d
18-a
31-d
6-d
19-d
32-d
45-b
8-b
21-b
34-a
9-c
22-a
35-a
10-e
23-c
38-b
47-a
11-c
24-d
12-c
25-d
13-e
26-d
12-c
25-d
38-a
13-c
26-d
39-e
50-d
Review #4
1-c
14-b
27-d
40-b
2-d
15-a
28-d
41-b
7-a
20-a
33-d
10-c
23-d
36-c
11-b
24-d
37-a
22.
An analysis of gas gave: C = 85.7%, and H = 14.3%. if the formula mass of this gas is 42 amu, what is the
empirical formula and the molecular formula?
a. CH; C4H4
b. CH2; C3H6
c. CH3; C3H9
d. C2H2; C3H6
e. C2H4; C3H6
22-b
28.
The empirical formula of a compound is CH. its molecular mass could be:
a. 21
b. 40
c. 51
d. 78
29.
The percent by mass of oxygen in MgO is closest to
a. 16%
b. 24%
c. 40%
d. 60%
30.
A 20 g of hydrate is heated until all water of hydration is driven off. The mass of the anhydrous compound
remaining is 15%. What is the percent of water in the hydrate?
a. 15 %
b. 85%
c. 1.33%
d. 75 %
31.
What is the mass in grams of 1 mole of (NH4)2S?
a. 68 g
b. 50 g
c. 64 g
midc_rev.doc: review for chemistry
6/28/2017
d. 54 g
12
32.
33.
How many molecules are contained in 127 grams of iodine, I2?
a. 1.50 x 1023
b. 3.01 x 1023
c. 9.03 x 1023
d. 12.4 x 1023
If the empirical formula for an organic compound is CH2O, then the molecular mass of the compound could be
a. 135
b. 60
c. 45
d. 15
44.
If the simplest formula of a substance is CH2 and its molecular mass is 56, what is its true formula
( molecular formula)?
a. CH2
b. C2H4
c. C3H4 d. C4H8 e. C5H10
48.
A compound contains 0.5 mole of carbon for each mole of hydrogen. The empirical formula of this compound is
a. CH
b. CH2
c. C2H
d. C2H2
49.
A compound has an empirical formula NO2. Its molecular formula could be:
a. NO2
b. N2O
c. N4O2
d. N4O4
28-d
29-c
30-85% 31-68g 32-b 33-b 44-d 48-b
midc_rev.doc: review for chemistry
6/28/2017
49-a
13