Download 1-Three states of matter . A: density, volume and weight B: solid

1-Three states of matter .
A: density, volume and weight
B: solid, liquid, and gas
C: water, metal and gases
2-Matter is something that take up space and has mass.
A: True
B: False
3-The temperature at which a substance changes from a liquid to a gas.
A: Freezing point
B: Melting point
C: Boiling point
D: Condensation point
4-The temperature at which a substance changes from a liquid to a solid
A: Freezing point
B: Boiling point
C: Melting point
D: Condensation point
5-How many grams of potassium chloride are produced if 25g of potassium chlorate
decompose? 2KClO3 --> 2KCl + 3O2
A: 74.6 g
B: 31.98 g
C: 15.2 g
6-How many grams of oxygen are produced if 25g of potassium chlorate decompose?
2KClO3 --> 2KCl + 3O2
A: 9.80 g
B: 4.90 g
C: 19.6 g
7-How many grams of hydrogen are necessary to react completely with
50.0g of nitrogen? 3H2 + N2 --> 2NH3
A: 25.2 g
B: 10.8 g
C: 9.75 g
8-How many grams of ammonia are produced from 50.0g of nitrogen
gas? 3H2 + N2 --> 2NH3
A: 121.4 g
B: 60.9 g
C: 15.2 g
9-How many grams of silver chloride are produced from 5.0g of silver nitrate
reacting with an excess of barium chloride? 2AgNO3 + BaCl2 --> Ba(NO3)2 +
2AgCl
A: 8.42 g
B: 71.6 g
C: 4.21 g
10-How many grams of potassium chloride are produced if 25g of potassium
chlorate decompose? 2KClO3 --> 2KCl + 3O2
A: 74.6 g
B: 31.98 g
C: 15.2 g
11-How many moles of hydrogen gas are needed to react completely
with two moles of nitrogen gas? 3H2 + N2 --> 2NH3
A: 3 mol
B: 6 mol
C: 9 mol
D: 12 mol
12-The type of matter of milk is ……….
a- Element
b-mixture
c-compound
d- all
13-The temperature 98,6 F° is equal to ……..in kelven
a-310
b-130
c-330
14-The unit of temperature in S.I. system is …………
d-350
a-Kelvin
b-Celsius
c-Fahrenheit
d-all
15-Which of the following is an intensive property…………
a-density
b-weight
c-mass
d-volume
16-The following number 1.002 has ……. Significant figure
a-
4
b- 3
c- 5
d- 2
17-The number of significant figure of 10500 is …………
a-3
b- 4
c- 5
d- all previous
18- The value of X come from the operation X= 15.3 + 0.335 is ………
a-15.635
b- 15.6
c- 15.63
d- 16
17 2-
19-The number of electrons in [8O ] is …………
a-8
b- 9
c- 10
d- 17
21-which of the following is Nobel gas?
a-H
b- Cl
c- Xe
d- O
22-Which quantity in each pair is larger?
a. 5 mL or 5 dL c. 5 cm or 5 mm
b. 10 mg or 10 μg d. 10 Ms or 10 ms
23-How many signifi cant fi gures does each number contain?
a. 16.00 c. 0.001 60 e. 1.06 g. 1.060 × 1010
b. 160 d. 1,600,000 f. 0.1600 h. 1.6 × 10–6
24- How many signifi cant fi gures does each number contain?
a. 160. c. 0.000 16 e. 1,600. g. 1.600 × 10–10
b. 160.0 d. 1.60 f. 1.060 h. 1.6 × 106
25- Round each number to three signifi cant fi gures.
a. 25,401 c. 0.001 265 982 e. 195.371
b. 1,248,486 d. 0.123 456 f. 196.814
26- Round each number in Problem 1.51 to four signifi cant
fi gures.
27- Carry out each calculation and report the answer using
the proper number of signifi cant fi gures.
a. 53.6 × 0.41 c. 65.2/12 e. 694.2 × 0.2
b. 25.825 – 3.86 d. 41.0 + 9.135 f. 1,045 – 1.26
1-Give the name of the elements in each group of three element symbols.
a. Au, At, Ag d. Ca, Cr, Cl
b. N, Na, Ni e. P, Pb, Pt
c. S, Si, Sn f. Ti, Ta, Tl
2- What element(s) are designated by each symbol or group
of symbols?
a. CU and Cu c. Ni and NI
b. Os and OS d. BIN, BiN, and BIn
3- Does each chemical formula represent an element or a
compound?
a. H2 b. H2O2 c. S8 d. Na2CO3 e. C60
4- Identify the elements in each chemical formula and tell how many atoms of each
are present.
a. K2Cr2O7
b. C5H8NNaO4 (MSG, fl avor enhancer)
c. C10H16N2O3S (vitamin B7)
5- Identify the element that fi ts each description.
a. an alkali metal in period 6
b. a transition metal in period 5, group 8
c. a main group element in period 3, group 7A
d. a main group element in period 2, group 2A
e. a halogen in period 2
f. an inner transition metal with one 4f electron
6- Identify the element that fi ts each description.
a. an alkaline earth element in period 3
b. a noble gas in period 6
c. a main group element in period 3 that has p orbitals
half-fi lled with electrons
d. a transition metal in period 4, group 11
e. an inner transition metal with its 5f orbitals completely
f. a transition metal in period 6, group 10 filled with electrons
7- Label each region on the periodic table.
a. noble gases e. alkaline earth elements
b. period 3 f. f block elements
c. group 4A g. transition metals
d. s block elements h. group 10
8- Identify each highlighted element in the periodic table and give its
[1] element name and symbol; [2] group number; [3] period; [4] classification (i.e.,
main group element, transition metal, or inner transition metal).
9-Write the element symbol that fi ts each description, using
a superscript for the mass number and a subscript for the
atomic number.
a. an element that contains 53 protons and 74 neutrons
b. an element with 35 electrons and a mass number of 79
c. an element with 47 protons and 60 neutrons
10- Which element in each pair is larger?
a. bromine and iodine c. silicon and potassium
b. carbon and nitrogen d. chlorine and selenium
11- Which element in each pair has its valence electrons
farther from the nucleus?
a. sodium and magnesium c. neon and krypton
b. carbon and fl uorine d. argon and bromine
12- Give the number of valence electrons in each element.
Write out the electronic confi guration for the valence
electrons.
a. sulfur b. chlorine c. barium d. titanium e. tin
13- Give the number of valence electrons in each element.
Write out the electronic confi guration for the valence
electrons.
a. neon c. aluminum e. zirconium
b. rubidium d. manganese
14-What is the maximum number of electrons that can be
contained in each shell, subshell, or orbital?
a. second shell c. 3p subshell e. fourth shell
b. 3s orbital d. 4f orbital f. 5p orbital
15-Write the element symbol that fi ts each description. Use a superscript for the mass
number and a subscript for the atomic number.
a. an element that contains 10 protons and 12 neutrons
b. an element with atomic number 24 and mass
number 52
16-How many protons, neutrons, and electrons are contained
in each element?
a. 27 13Al
b. 35
17Cl
c. 3416S
17-Give all of the terms that apply to each element:
[1] metal; [2] nonmetal; [3] metalloid; [4] alkali metal;
[5] alkaline earth element; [6] halogen; [7] noble gas;
[8] main group element; [9] transition metal; [10] inner
transition metal.
a. sodium c. xenon e. uranium
b. silver d. platinum f. tellurium
18- Give all of the terms that apply to each element:
[1] metal; [2] nonmetal; [3] metalloid; [4] alkali metal;
[5] alkaline earth element; [6] halogen; [7] noble gas;
[8] main group element; [9] transition metal; [10] inner
transition metal.
a. bromine c. cesium e. calcium
b. silicon d. gold f. chromium
19-For the given atomic number (Z) and mass number (A):
[1] identify the element; [2] give the element symbol;
[3] give the number of protons, neutrons, and electrons.
a. Z = 10, A = 20 d. Z = 55, A = 133
b. Z = 13, A = 27 e. Z = 28, A = 59
c. Z = 38, A = 88 f. Z = 79, A = 197
20-Give the name of the elements in each group of three
element symbols.
a. Au, At, Ag
d. Ca, Cr, Cl
b. N, Na, Ni
e. P, Pb, Pt
c. S, Si, Sn
f. Ti, Ta, Tl
1.
Which of these is the electron configuration of an atom most likely to lose an
electron?
a. 1s2 2s2 2p6
b. [He]2s2 2p5
c. 1s2 2s2 2p6 3s2 3p5
d. 1s2 2s2 2p6 3s2 3p6 4s1
Hint
2.
Which of these is the electron configuration of an atom most likely to gain an
electron?
a. 1s2 2s2 2p6
b. 1s2 2s2 2p6 3s1
c. 1s2 2s2 2p6 3s2 3p5
d. 1s2 2s2 2p6 3s2 3p6 4s1
3. _____________ is the force that holds two atoms together.
a. A chemical bond
b. Glue
c. Nuclear force
d. Fission
4. What forms chemical
bonds?
a. atomic nuclei
b. valence electrons
c. inner-level electrons
d. noble gases
5. A positive ion forms when ___________.
a. an atom loses one or more valence electrons
b. an atom gains one or more valence electrons
c. electrons are pulled into the nucleus
d. electrons are pushed out of the nucleus
6.What is a negatively charged ion called?
a. nucleus
b. cation
c. anion
d. molecule
7. How many electrons are present in the valence level for all noble gases except
helium?
a. 6
b. 7
c. 8
d. 9
8. Why is the calcium ion (Ca2+) more stable than the calcium atom (Ca)?
a. Twenty electrons are more stable than
eighteen electrons.
b. Eighteen electrons are less stable than
twenty electrons.
c. The two electrons more than the noble gas
configuration is more stable.
d. The noble gas configuration is more
stable.
9. Which elements can either gain or
lose electrons to form stable octets?
a. metals
b. metalloids
c. nonmetals
d. transition metals
10.
Large differences in electronegativity result in __________ bonding between
atoms.
a. covalent
b. ionic
c. no
d. polar
11. The phosphorus pentachloride molecule is nonpolar and contains no unshared
electron pairs on the phosphorus atom. What are all the possible bond angles in
this molecule?
a. 120°
b. 180°
c. 90°, 120°, and 180°
d. 90° and 180°
12. Which element has the highest electronegativity?
a. N
b. O
c. F
d. Ne
13. Which compound consists of nonpolar molecules?
a. H2S
b. PH3
c. AsH3
d. SiH4
10- Which formulas represent ionic compounds and which
represent covalent compounds?
a. CO2 b. H2SO4 c. KF d. CH5N
11- Which formulas represent ionic compounds and which
represent covalent compounds?
a. C3H8 b. ClBr c. CuO d. CH4O
12- Which pairs of elements are likely to form ionic bonds
and which pairs are likely to form covalent bonds?
a. potassium and oxygen c. two bromine atoms
b. sulfur and carbon d. carbon and oxygen
13- Which pairs of elements are likely to form ionic bonds
and which pairs are likely to form covalent bonds?
a. carbon and hydrogen c. hydrogen and oxygen
b. sodium and sulfur d. magnesium and bromine
1. A(n) _________________ is the basic unit of matter.
A. electron
B. atom
C. proton
D. neutron
2. ________________ are negatively charged particles; located outside the atomic
nucleus.
A. Protons
B. Atoms
C. Electrons
D. Neutrons
3. A substance consisting entirely of one type of atom is known as _______________.
A. ion
B. isotope
C. atom
D. element
4. A(n) _______________ is an atom of an element that has a number of neutrons
different from that of other atoms of the same element.
A. element
B. electron
C. ion
D. isotope
5. A substance formed by the chemical combination of two or more elements in
definite proportions is a _________________.
A. atom
B. ion
C. compound
D. isotope
6. A(n) ________________ is a bond formed when one or more electrons are
transferred from one atom to another.
A. ionic bond
B. covalent bond
C. hydrogen bond
D. mixture
7. A _______________ is a bond formed by the sharing of electrons between atoms.
A. ionic bond
B. covalent bond
C. hydrogen bond
D. mixture
8. Material composed of two or more elements or compounds that are physically
mixed together but not chemically combined is a _______________.
A. solution
B. mixture
C. ionic bond
D. covalent bond
9. A ________________ is a substance that is dissolved in a solvent to make a
solution.
A. solute
B. solvent
C. solution
D. mixture
10. A ________________ is a substance in which a solute is dissolved to form a
solution.
A. solute
B. solvent
C. solution
D. mixture
11. The _______________ is a measurement system used to indicate the
concentration of hydrogen ions (H+) in solution; ranges from 0 to 14.
A. base
B. buffer
C. acid
D. pH scale
12. A(n) _______________is a compound that forms hydrogen ions (H+) in solution.
A. base
B. buffer
C. acid
D. pH scale
13. A(n) _______________is a compound that produces hydroxide ions (OH-) in
solution.
A. base
B. buffer
C. acid
D. pH scale
14. __________________ are small units that can join together with other small units
to form polymers.
A. Ions
B. Polymers
C. Monomers
D. Isotopes
15. Large compounds formed from combinations of many monomers are known as
_______________.
A. ions
B. polymers
C. monomers
D. isotopes
16. A compound made up of carbon, hydrogen, and oxygen atoms; major source of
energy for the human body is a _________________ .
A. nucleic acid
B. protein
C. carbohydrate
D. lipid
17. ________________ are macromolecules made mainly from carbon and hydrogen
atoms; includes fats, oils, and waxes.
A. Nucleic acids
B. Amino Acids
C. Carbohydrates
D. Lipids
18. ________________ are macromolecules containing hydrogen, oxygen, nitrogen,
carbon, and phosphorous.
A. Nucleic acids
B. Amino Acids
C. Carbohydrates
D. Lipids
19. Compounds with an amino group (-NH2) on one end and a carboxyl group (COOH) on the other end are known as _______________ (Hint: They are the
building blocks of proteins).
A. nucleic acids
B. amino acids
C. carbohydrates
D. lipids
20. ________________ are macromolecules that contain carbon, hydrogen, oxygen,
and nitrogen; needed by the body for growth and repair and to make up enzymes.
A. Nucleic Acids
B. Proteins
C. Carbohydrates
D. Lipids
21. A process that changes one set of chemicals into another set of chemicals is
known as a ________________.
A. chemical reaction
B. Substrate
C. activation energy
D. catalyst
22. The _________________ is the energy needed to get a reaction started.
A. chemical reaction
B. substrate
C. activation energy
D. catalyst
23. A substance that speeds up the rate of a chemical reaction is known as a(n)
_______________.
A. enzyme
B. substrate
C. activation energy
D. catalyst
24. Proteins that act as biological catalysts are called _______________.
A. enzymes
B. substrates
C. activation energies
D. catalysts
25. A(n) _________________ is the reactant of an enzyme-catalyzed reaction.
A. enzyme
B. substrate
C. activation energy
D. catalyst
26
Which of the following is true?
A)
B)
C)
D)
27
B)
C)
D)
Covalent bonds are weak intermolecular forces.
Covalent bonds are strong intramolecular forces.
Dipole-dipole
Ion-induced dipole
Ion-dipole
Dipole-dipole
What type of intermolecular force is responsible for the attraction between a
polar molecule that induces a charge on a non-polar molecule?
A)
B)
C)
D)
29
Covalent bonds are weak intramolecular forces.
What type of intermolecular force is responsible for the attraction between an ion
and a polar molecule?
A)
28
Covalent bonds are strong intermolecular forces.
Dipole-dipole
Ion-dipole
Ion-induced dipole
Dipole-induced dipole
Which of the following can result in a dispersion force?
A)
When a non-polar molecule becomes slightly polar for an instant
When the oppositely charged ends of a polar molecule attract each other
B)
When there is a very strong dipole-dipole attraction between a hydrogen
C) atom and a polar-molecule
When an ion comes close enough to a non-polar molecule to change its
D) electron density
30
Which of the following does not form hydrogen bonds?
A)
B)
C)
D)
31
H2S
H2O
NH3
HF
Which of the following is not true?
A)
B)
Hydrogen bonding helps explain why solid water floats on liquid water.
Hydrogen bonding is responsible for the relatively low boiling point of water.
Water molecules are farther apart in solid water than they are in liquid
C) water.
D)
32
Hydrogen bonding is responsible for the relatively high boiling point of water.
Which of the following is not true?
Deoxyribonucleic acid is stabilized by hydrogen bonds and London dispersion
A) forces.
van der Waals forces can exist between two different parts of the same large
B) molecule.
The level of protein structure that is stabilized by London dispersion forces is
C) called the secondary structure.
The level of protein structure that is stabilized by London dispersion forces is
D) called the tertiary structure.
33
Which of the following is true?
A)
Hydrogen bonds are stronger than covalent bonds.
A hydrogen bond is an electrostatic attraction between the nucleus of a
hydrogen atom, bonded to fluorine, oxygen, or nitrogen, and the positive
B) end of a nearby dipole.
In liquid water, each water molecule is hydrogen bonded to two other water
C) molecules.
D)
34
Hydrogen bonding is one type of dipole-dipole interaction.
Which of the following statements is correct?
A)
Dipole-dipole forces have a greater energy than dipole-induced dipole forces.
Ion-induced dipole forces have a greater energy than dipole-dipole forces.
B)
Dipole-induced dipole forces have a greater energy than ion-induced dipole
C) forces.
D)
35
Methanol, ethanol, ammonia, and methylamine are soluble in water because
A)
B)
C)
36
D)
B)
C)
D)
they can form hydrogen bonds
there are dipole-dipole forces
there are dispersion force
MgO
Li
H2
H2O
Which of the following has the highest melting point?
A)
B)
C)
D)
38
they can form ion-induced dipoles
Which of the following has the highest boiling point?
A)
37
Dispersion forces have a greater energy than dipole-dipole forces.
Li
MgO
Cl2
H2O
Which of the following is true?
Covalent network solids have very low boiling points and are insoluble in most
A) liquids.
B)
C)
D)
39
Metallic crystalline solids are formed by metals with low electronegativities.
Ionic crystalline solids have low electrical conductivity in liquid state.
Molecular crystalline solids have high electrical conductivity in liquid state.
Which of the following is not true?
A)
Non-polar molecular crystals are very soft and are soluble in non-polar solvents.
Non-polar molecular crystals are formed from symmetrical molecules with
covalent bonds between atoms with small electronegativity differences.
B)
C)
D)
40
Non-polar molecular crystals generally have high boiling points.
An ionic crystalline solid forms between atoms with an electronegativity difference
A)
B)
C)
D)
41
Non-polar molecular crystals generally have low boiling points.
less than 1.7
between 0.5 and 1.7
greater than 1.7
less than 0.5
Which of the following does not explain the malleability of metal solids?
When stress is applied one layer can slide over another while the free electrons
A) continue to bind the ions together.
B)
C)
Metallic bonds are non-directional.
When stress is applied, like charges become aligned and repel each other.
Positive metal cations are layered as fixed arrays which can slide over one
D) another.
42
Which of the following substances is a good conductor in solid state?
A)
B)
C)
D)
43
Zn
CO
NaCl
Which of the following is a good conductor when dissolved in water?
A)
B)
C)
D)
44
P4
KCl
Br2
CO
CH3COOH
Which of the following properties would indicate that a solid is ionic?
A)
low boiling point, low electrical conductivity, non-polar
high boiling point, high electrical conductivity, hard and brittle
B)
C)
D)
45
high boiling point, very high electrical conductivity, have a luster
high boiling point, low electrical conductivity, hard crystals
Which of the following are general properties of a molecular solid?
A)
B)
C)
D)
low boiling points and electrical conductivity, soluble in non-polar solvents
low boiling points, brittle, often soluble in water
high boiling points, hard crystals, insoluble in most liquids
high boiling points, hard and brittle, often soluble in water
1. The number of atoms in a mole of any pure substance is called
a) its atomic number
b) Avogadro's number
c) Its mass number
d) Its isotopic number
2. The atomic number of oxygen is 8. The atomic number of sulfur is 16. Compared with a
mole of oxygen, a mole of sulfur contains
a) twice as many atoms
b) half as many atoms
c) an equal number of atoms
d) 8 times as many atoms
3. To determine the molar mass of an element, one must know the element's
a) Avogadro constant
b) atomic number
c) number of isotopes
d) average atomic mass
4. Avogadro's number of atoms of any element is equivalent to
a) the atomic number of that element
b) the mass number of that element
c) 6.02 x 1023 particles
d) 100 g of that element
5. The mass of 1 mol of chromium is about
a) 12 g
b) 24 g
c) 52 g
d) 6.02 x 1023 g
6. A mass of 6.005 g of carbon contains
a) 1 mol of C
b) 2 atoms of C
c) 0.5000 mol of C
d) 1 atom of C
7. The mass of 2 moles of oxygen atoms is
a) 16 g
b) 32 g
c) 48 g
d) 64 g
8. What is the number of moles of atoms in 9.03 x 1024 atoms?
a) 1.50 mol
b) 9.03 mol
c) 10.0 mol
d) 15.0 mol
9. A sample of tin contains 3.01 x 1023 atoms. The mass of the sample is
a) 3.01 g
b) 59.3 g
c) 72.6 g
d) 11g
10. The mass of a sample of nickel is 11.74 g. It contains
a) 1.174 x 1023 atoms
b) 1.205 x 1023 atoms
c) 1.869 x 1023 atoms
d) 3.256 x 1023 atoms
11. Which of the following weighs more?
a) 1 mole of hydrogen
b) 0.25 moles of He
c) 0.1 mol of Ne
d) 0.2 mol of C
12. What is the molar mass of magnesium chloride, MgCl2?
a) 46g/mole
b) 59.763g/mole
c) 95.211g/mole
d) 106.354g/mole
13. What is the molar mass of (NH4)2SO4?
a) 114.09g/mole
b) 118.34g/mole
c) 128.06g/mole
d) 132.13g/mole
14. The molar mass of NO2 is 46.01 g/mole. How many moles of NO2 are present in
114.95g?
a) 0.4003mol
b) 1.000mol
c) 2.498mol
d) 114.95mol
15. The molar mass of CCl4 is 153.81g/mol. How many grams of CCl4 are needed to have
5.000 mol?
a) 5.000g
b) 30.76g
c) 769.0g
d) 796.05g
16. How many Cl- ions are present in 2.00 mol of KCl?
a) 1.204 x 1024
b) 6.02 x 1024
c) 2.00
d) 0.5
17. How many OH- ions are present in 3.00 mol of Ca(OH)2?
a) 3.00
b) 6.00
c) 3.61 x 1024
d) 2.06 x 1023
18. What is the percent composition, by mass, of CO?
a) 50% C, 50% O
b) 12% C, 88% O
c) 25% C, 75% O
d) 43% C, 57% O
19. What is the percentage composition, by mass, of oxygen in H2O?
a) 15.99%
b) 33%
c) 88.8%
d) 99.8%
20. The empirical formula for a compound shows the symbols of the elements with
subscripts indicating the
a) actual numbers of atoms in a molecule
b) number of moles of the compound in 100 g.
c) smallest whole-number ratio of atoms
d) atomic masses of each element
21. A compound contains 259.2 g of F and 40.8 g of C. What is the empirical formula for this
compound?
a) CF4
b) C4F
c) CF
d) CF2
22. What is the empirical formula for a compound that is 53.3% O and 46.7% Si?
a) SiO
b) SiO2
c) Si2O
d) Si2O3
23. What is the empirical formula for a compound that is 31.9% potassium, 28.9% chlorine,
and 39.2% Oxygen?
a) KClO2
b) KClO3
c) K2Cl2O3
d) K2Cl2O5
24. What is the empirical formula for a compound that is 43.6% phosphorus and 56.4%
oxygen?
a) P3O7
b) PO3
c) P2O3
d) P2O5
25. To find the molecular formula from the empirical formula, one must determine the
compound's
a) density
b) molar mass
c) structural formula
d) shape
26. A compound's empirical formula is C2H5. If the molar mass is 58 g/mole, what is the
molecular formula?
a) C3H6
b) C4H10
c) C5H8
d) C5H15
27. A compound containing only hydrogen and oxygen is 5.9% hydrogen by mass. The molar
mass of the compound is 34 g/mole. What is the molecular formula of the compound?
a) H2O
b) H2O2
c) OH
d) H18O
28. The mass percentage of water in the hydrate CuSO4 ·5H2O is
a) 18%
b) 25%
c) 31%
d) 36%
e) 52%
29. The mass percent water in a hydrate of Na2CO3 is 62.98%. What is the formula for the
hydrate?
a) Na2CO3 · H2O
b) Na2CO3 · 3H2O
c) Na2CO3 · 5H2O
d) Na2CO3 · 10H2O
30-During chemical reactions, atoms are
a-broken down into smaller particles
b-never broken down into smaller particles
c-unchanged
31-According to the Law of Conservation of Mass
a-the mass of the reactants is greater than the mass of the products
b-the mass of the reactants is equal to the mass of the products
c-the mass of the products is greater than the mass of the reactants
32-You can tell a chemical reaction because it always produces
a-different substance .
b-reactants .
c change of state .
d-water
33-How many carbon atoms are there in C12H22O11?
a. 6
b. 12
c. 18
d. 45
34-Avogadro's number is
a-6.23 × 10–2
b-6.02 × 1023
c-6.23 × 102
35-What is the molar mass of silver nitrate? (AgNO3)
A: 107g/mol
B: 710 g/mol
C: 170 g/mol
36-What is the molar mass of carbon dioxide? (CO2)
A: 44g/mol
B: 40 g/mol
C: 28g/mo
37-Convert 3.57 moles of aluminum to grams.
A: 81 g
B: 96 g
C: 100 g
38-What is the mass of 4.26 moles of silicon?
A: 119g
B: 112g
C: 168g
39-How many moles are in 25.5g of silver?
A: .24 mol
B: .50 mol
C: 1 mol
C3H8 + 502 --> 3CO2 + 4H2O If there are two moles of oxygen available, which
mole ratio will tell you how much water is produced?
A: 4 mol H20/5 mol O2
B: 5 mol O2/4 mol H2O
1. The pressure of a sample of helium in a 1.0-L container is 0.857 atm. What is
the new pressure if the sample is placed in a 0.5-L container? (Assume the
temperature is constant.)
a. 0.143 atm
b. 0.429 atm
c. 1.38 atm
d. 1.71 atm
2. A sample of gas is held in a 10.0-L volume at 175 kPa. The temperature is kept
constant while the volume is decreased until the pressure is 350 kPa. What is
the new volume of the gas?
a. 1.0 L
b. 5.0 L
c. 10.0 L
d. 175 L
3. A 0.5-L container of nitrogen gas is heated under constant pressure to the
boiling point of water. What is its new volume?
a. 0.5 L
b. 0.64 L
c. 0.79 L
d. 0.86 L
4. How can gases be defined?
a. a physical state of matter that does not have a fixed shape or a fixed
volume
b. a physical state of matter that does not have a fixed shape but has a
fixed volume
c. a physical state of matter that has a fixed volume and a fixed shape
d. a chemical state of matter
5. Particles of matter that are in constant, random motion and that have a size that
is much smaller than the distance between them are _____________.
a. solids
b. liquids
c. gases
d. solutions
6. What is the name given to the relationship that shows that an increase in
pressure leads to a decrease in the volume of a gas?
a. Charles’s law
b. Boyle’s law
c. Avogadro’s number
d. Gay-Lussac’s law
7. How can the relationship between a gas at two sets of conditions be expressed
mathematically by Boyle’s law?
a. P1V1 = P2V2
b. P1/V1 = P2/V2
c. V1/T1 = V2/T2
d. V1T1 = V2T2
8. What relationship is demonstrated by the expansion of a gas—filled balloon
when it is heated?
a. Charles’s law
b. Boyle’s law
c. Avogadro’s number
d. Gay-Lussac’s law
9. A gas occupies a volume of 1.0 L at 1.0 atm pressure. What is the pressure
when the gas expands to fill 2.0 L?
a. 0.50 atm
b. 2.0 atm
c. 1.0 atm
d. 10 atm
10. A gas occupies a volume of 1.0 L at 25°C. What volume will the gas occupy at
100°C?
a. 1.0 L
b. 1.3 L
c. 0.80 L
d. 4.0 L
11. A gas occupies 2.0 L at STP. What volume will the gas occupy if the pressure is
increased to 2.0 atm, and the temperature is kept constant?
a. 1.0 L
b. 4.0 L
c. 0.50 L
d. 2.0 L
12. A sample of helium occupies 2.20 L at 1.0 atm. What
is the volume at 1.5 atm?
a. 1.5 L
b. 0.68 L
c. 2.20 L
d. 1.0 L
13. When will the molecules of all samples of ideal gases have the same average
kinetic energies?
a. at constant volume
b. at constant temperature
c. at constant amount
d. at constant pressure
14.
<.br>
What volume will 0.554 mol of gas occupy at STP?
a. 0.25 L
b. 3.34 L
c. 12.4 L
d. 40.4 L
15. Which of the following states that equal volumes of gases at the same
temperature and pressure contain the same number of particles?
a. Boyle’s law
b. Gay-Lussac’s law
c. Charles’s law
d. Avogadro’s principle
16. How can the molar volume of a gas be defined?
a. the volume that one mole occupies at STP
b. the volume that one gram occupies at STP
c. the volume that one mole occupies at 100°C and 1
atm pressure
d. the volume that one gram occupies at 100°C and 1
atm pressure
17. What is the volume of
2.0 moles of a gas at
STP?
a. 44.8 L
b. 22.4 L
c. 0.0223 L
d. 0.0446 L
18. If 1.00 L of a gas is 4.40 times as heavy as 1.00 L of O2 at the same
temperature and pressure, then what is the molar mass of the unknown gas?
a. 67.0 g/mol
b. 70.4 g/mol
c. 88.0 g/mol
d. 141 g/mol
1. What is the percent by mass of NaCl in a solution that contains 17.5 g NaCl per
500.0 g of water?
a. 3.38%
b. 3.50%
c. 3.61%
d. 14.80%
2. How much solvent is needed to make 200 ml of 50% rubbing alcohol?
a.
50 mL
b.
100 mL
c.
150 mL
d.
200 mL
3. Calculate the molarity of 0.75 L of a solution containing 0.83 g of dissolved
KCl.
a. 0.015 M
b. 0.75 M
c. 1.1 M
d. 6.2
4. How many grams of NaCl are dissolved in 500.0 mL of a 0.05M solution of
NaCl?
a. 0.05 g
b. 0.29 g
c. 1.46 g
d. 2.92 g
5. A solution that contains less solute per volume of solvent than another solution
made from the same components is said to be more ________.
a. dilute
b. concentrated
c. solvated
d. dissolved
6. Molarity is defined as the ____________.
a. mass of solute per mass of solution
b. volume of solute per volume of solution
c. moles of solute per liter of solution
d. moles of solute per kilograms of solvent
7. What volume of 12.6M HCl must be added to sufficient water to prepare 5.00
liters of 3.00MHCl?
a. 1.19 L
b. 21.0 L
c. 0.840 L
d. 7.56 L
8. What mass of Ca(OH)2 is contained in 1500 mL of 0.0250M Ca(OH)2 solution?
a. 3.17 g
b. 2.78 g
c. 1.85 g
d. 2.34 g
9. Calculate the molality of a solution that contains 25 g of H2SO4 dissolved in
80 g of water.
a. 1.6m
b. 2.2m
c. 3.2m
d. 6.3m
10. Calculate the molality of 10% H3PO4 solution in water.
a. 0.380m
b. 0.760m
c. 1.13m
d. 1.51m
11. What is the mole fraction of ethanol (C2H5OH) in a solution of 47.5 g of
ethanol in 850 g of water?
a. 0.021
b. 0.18
c. 0.032
d. 0.98
12. What is the molarity of 2500 mL of a solution that contains 160 grams of
ammonium nitrate (NH4NO3)?
a. 0.333M
b. 0.450M
c. 0.600M
d. 0.800M
13.
If 12.0 g of a gas at 2.5 atm dissolve in 1.0 L of water at 25°C, how much will
dissolve in 1.0 L of water at STP?
a. 0.21 g/L
b. 2.1 g/L
c. 4.8 g/L
d. 12.0 g/L
14. In a solution, the substance that does the dissolving is ____________.
a. the solvent
b. the solute
c. saturated
d. miscible
15. Solutions can be mixtures of _______________.
a. solids
b. liquids
c. gases
d. all of the above
16.
a.
b.
c.
d.
In a glass of sugar water, which substance is the solute?
water
sugar
glass
none of the above
A solution is said to be ________ when more solute can be dissolved in the solvent at
a given temperature.
a. solvation
b. salvation
c. crystallization
d. ionization
18. Which of the following will not increase the rate of solvation?
a. agitating the mixture
b. increasing the surface area
c. increasing the temperature
d. formation of a precipitate
19. A solution is said to be ________ when more solute can be dissolved in
the solvent at a given temperature.
a. supersaturated
b. saturated
c. unsaturated
d. solvated
20. What is a common means of identifying a supersaturated
solution?
a. precipitation
b. dissolution
c. solvation
d. hydration
21. The decrease in solubility of a gas in a solution when the pressure is reduced is
described by _________.
a. Boyle’s law
b. Henry’s law
c. Charles’s law
d. the ideal gas law
1- Explain the difference between the Arrhenius defi nition
of acids and bases and the Brّnsted–Lowry defi nition of
acids and bases.
2- Explain why NH3 is a Brّnsted–Lowry base but not an
Arrhenius base.
3- Which of the following species can be Brّnsted–Lowry
acids?
a. HBr b- AlCl3 c.- NO2
b.- Br2
d.- HCOOH
f.- HNO2
4- Which of the following species can be Brّnsted–Lowry
acids?
a. H2O b. HOCl c. CH3CH2COOH b. I– e. FeBr3 f. CO2
5- Which of the following species can be Brّnsted–Lowry
bases?
a. –OH b. C2H6 c. –OCl
b. Ca2+ d.
PO43– f. MgCO3
6- Calculate the value of [–OH] from the given [H3O+] and label the solution as
acidic or basic.
a. 10–8 M c. 3.0 × 10–4 M
b. 10–10 M d. 2.5 × 10–11 M
7- Calculate the value of [–OH] from the given [H3O+] and
label the solution as acidic or basic.
a. 10–1 M c. 2.6 × 10–7 M
b. 10–13 M d. 1.2 × 10–12 M
8- Calculate the value of [H3O+] from the given [–OH] and
label the solution as acidic or basic.
a. 10–2 M c. 6.2 × 10–7 M
b. 4.0 × 10–8 M d. 8.5 × 10–13 M
9- Calculate the value of [H3O+] from the given [–OH] and
label the solution as acidic or basic.
a. 10–12 M c. 6.0 × 10–4 M
b. 5.0 × 10–10 M d. 8.9 × 10–11 M
10- Calculate the pH from each H3O+ concentration
calculated in Problem 9.77.
11- Calculate the pH from each H3O+ concentration
calculated in Problem 9.78.
12- Calculate the H3O+ concentration from each pH: (a) 12;
(b) 1; (c) 1.80; (d) 8.90.
13- Calculate the H3O+ concentration from each pH: (a) 4;
(b) 8; (c) 2.60; (d) 11.30.
14- If a urine sample has a pH of 5.90, calculate the concentrations of H3O+ and –
OH in the sample.
15- If pancreatic fl uids have a pH of 8.2, calculate the concentrations of
H3O+ and –OH in the pancreas.