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 d-350

14-The unit of temperature in S.I. system is …………

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

19-The number of electrons in [8O17

]2-

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. 1s

2 2s

2 2p

6

b. [He]2s

2 2p

5

c. 1s

2 2s

2 2p

6 3s

2 3p

5

d. 1s

2 2s

2 2p

6 3s

2 3p

6 4s

1

Hint

2.

Which of these is the electron configuration of an atom most likely to gain an

electron?

a. 1s

2 2s

2 2p

6

b. 1s

2 2s

2 2p

6 3s

1

c. 1s

2 2s

2 2p

6 3s

2 3p

5

d. 1s

2 2s

2 2p

6 3s

2 3p

6 4s

1

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) Covalent bonds are strong intermolecular forces.

B) Covalent bonds are weak intramolecular forces.

C) Covalent bonds are weak intermolecular forces.

D) Covalent bonds are strong intramolecular forces.

27

What type of intermolecular force is responsible for the attraction between an ion and a polar molecule?

A) Dipole-dipole

B) Ion-induced dipole

C) Ion-dipole

D) Dipole-dipole

28

What type of intermolecular force is responsible for the attraction between a polar molecule that induces a charge on a non-polar molecule?

A) Dipole-dipole

B) Ion-dipole

C) Ion-induced dipole

D) Dipole-induced dipole

29

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)

C)

When there is a very strong dipole-dipole attraction between a hydrogen atom and a polar-molecule

D)

When an ion comes close enough to a non-polar molecule to change its electron density

30

Which of the following does not form hydrogen bonds?

A) H2S

B) H2O

C) NH3

D) HF

31

Which of the following is not true?

A) Hydrogen bonding helps explain why solid water floats on liquid water.

B) Hydrogen bonding is responsible for the relatively low boiling point of water.

C)

Water molecules are farther apart in solid water than they are in liquid water.

D) Hydrogen bonding is responsible for the relatively high boiling point of water.

32

Which of the following is not true?

A)

Deoxyribonucleic acid is stabilized by hydrogen bonds and London dispersion forces.

B)

van der Waals forces can exist between two different parts of the same large

molecule.

C)

The level of protein structure that is stabilized by London dispersion forces is called the secondary structure.

D)

The level of protein structure that is stabilized by London dispersion forces is called the tertiary structure.

33

Which of the following is true?

A) Hydrogen bonds are stronger than covalent bonds.

B)

A hydrogen bond is an electrostatic attraction between the nucleus of a hydrogen atom, bonded to fluorine, oxygen, or nitrogen, and the positive

end of a nearby dipole.

C)

In liquid water, each water molecule is hydrogen bonded to two other water molecules.

D) Hydrogen bonding is one type of dipole-dipole interaction.

34

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)

C)

Dipole-induced dipole forces have a greater energy than ion-induced dipole forces.

D) Dispersion forces have a greater energy than dipole-dipole forces.

35

Methanol, ethanol, ammonia, and methylamine are soluble in water because

A) they can form ion-induced dipoles

B) they can form hydrogen bonds

C) there are dipole-dipole forces

D) there are dispersion force

36

Which of the following has the highest boiling point?

A) MgO

B) Li

C) H2

D) H2O

37

Which of the following has the highest melting point?

A) Li

B) MgO

C) Cl2

D) H2O

38

Which of the following is true?

A)

Covalent network solids have very low boiling points and are insoluble in most liquids.

B) Metallic crystalline solids are formed by metals with low electronegativities.

C) Ionic crystalline solids have low electrical conductivity in liquid state.

D) Molecular crystalline solids have high electrical conductivity in liquid state.

39

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) Non-polar molecular crystals generally have low boiling points.

D) Non-polar molecular crystals generally have high boiling points.

40

An ionic crystalline solid forms between atoms with an electronegativity difference

A) less than 1.7

B) between 0.5 and 1.7

C) greater than 1.7

D) less than 0.5

41

Which of the following does not explain the malleability of metal solids?

A)

When stress is applied one layer can slide over another while the free electrons continue to bind the ions together.

B) Metallic bonds are non-directional.

C) When stress is applied, like charges become aligned and repel each other.

D)

Positive metal cations are layered as fixed arrays which can slide over one

another.

42

Which of the following substances is a good conductor in solid state?

A) P4

B) Zn

C) CO

D) NaCl

43

Which of the following is a good conductor when dissolved in water?

A) KCl

B) Br2

C) CO

D) CH3COOH

44

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) high boiling point, very high electrical conductivity, have a luster

D) high boiling point, low electrical conductivity, hard crystals

45

Which of the following are general properties of a molecular solid?

A) low boiling points and electrical conductivity, soluble in non-polar solvents

B) low boiling points, brittle, often soluble in water

C) high boiling points, hard crystals, insoluble in most liquids

D) 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.

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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. In a glass of sugar water, which substance is the solute?

a. water

b. sugar

c. glass

d. 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 Br ّ nsted–Lowry defi nition of

acids and bases.

2- Explain why NH3 is a Br ّ nsted–Lowry base but not an

Arrhenius base.

3- Which of the following species can be Br ّ nsted–Lowry

acids?

a. HBr b- AlCl3 c.- NO2 b.- Br2 d.- HCOOH f.- HNO2

4- Which of the following species can be Br ّ nsted–Lowry

acids?

a. H2O b. HOCl c. CH3CH2COOH b. I– e. FeBr3 f. CO2

5- Which of the following species can be Br ّ nsted–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.