Section 1

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Section 1.2 Chemistry and Matter

Practice Test

1.

Which of the following releases matter?a. engaging a car exhaust system

b. picking up iron filings with a magnet

c. turbines spinning in a power plant

d. turning on an electric hot plate

Hint

http://openhintpopup%28%27php-bin/displayHint.php?qi=81651&bk=167&qz=6224&wh=0%27)



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2.

Why can an object's weight vary with location but its mass cannot?

a. Determinations of mass account for varying gravitational pulls in different

locations.

b. Mass is a measure of the amount of matter in an object; weight is not.

c. Systems of measurement vary with location.

d. Weight is a measure of both mass and the effect of Earth's gravitational pull onthe matter. Mass measures only the amount of matter.

Hint

3. Anything that has mass and takes up space is _________.

a. matter

b. volume

c. pressure

d. weight

Hint

4. What is the measurement of matter whose value depends on the force of gravity?

a. . mass

b. volume

c. energy

d. weight

Hint










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5. Chemists whose specialty is determining the composition of chemicals work in the field

of ________.

a. analytical chemistry

b. organic chemistry

c. physical chemistry

d. inorganic chemistry

Hint

6. Chemists who study the chemistry of living organisms work in the field of

______________.


b. physical chemistry

c. organic chemistry

d. biochemistry

Hint

7. What branch of chemistry is most concerned with the study of carbon compounds?


b. inorganic chemistry

c. organic chemistry

d. physical chemistry

Hint

8. A branch of chemistry that is concerned with how and why chemicals interact is

______________.


b. biochemistry

c. theoretical chemistry


Hint

Section 1.3 Scientific Methods

Practice Test













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1.

A hypothesis is more likely to become a __________ because __________.

a. scientific law; fewer trials are needed during the experimental phase

b. scientific law; theories are conclusions reached by many scientists about

relationships in nature

c. theory, fewer experiments are needed in order to support it

d. theory, scientific laws are conclusions reached by many scientists about

relationships in nature

Hint

2. The general term for a systematic approach used in scientific study is _______________.

a. the scientific method

b. qualitative analysis

c. quantitative analysis

d. the scientific controversy

Hint

3. The type of data that is descriptive in nature is ___________.

a. quantitative data

b. qualitative data

c. random data

d. hypothetical data

Hint

4. What is the name given to a set of controlled observations that test a proposedexplanation?

a. hypothesis

b. experiment

c. theory










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d. conclusion

Hint

5. The variable that you plan to change during the course of an experiment is ___________.

a. the independent variable b. the dependent variable

c. a constant

d. a control

Hint

6. A judgment based upon the results of an experiment is a _____________.

a. hypothesis

b. theory

c. variable

d. conclusion

Hint

7. When an explanation has been supported by many experiments, the explanation is a __________.

a. hypothesis

b. theory

c. law

d. model

Hint

8. A tentative explanation for a series of observations is a __________.

a. hypothesis

b. theory

c. law

d. model

Hint

9. A ________ can be used to help visualize microscopic structures and events.

a. variable

b. hypothesis
















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c. theory

d. model

Hint

10. Suppose that you experimentally determine the mass of nylon formed as a result of eachof several similar chemical processes. What are the measurements of mass called?

a. qualitative data

b. independent variables

c. controls

d. quantitative data

Hint

Section 1.4 Scientific Research

Practice Test

1.

Which of the following materials was discovered by chance during research on artificial

fiber creation?

a. lysozyme

b. nylon







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c. ozone

d. silk

Hint

2. What is the name given to research that is undertaken to solve a specific problem?a. pure research

b. theoretical research

c. applied research

d. technology

Hint

3. Two items that must be worn during any laboratory experiment are ___________.

a. safety goggles and a lab apron

b. safety goggles and fire-proof mitts

c. gloves and a lab apron

d. rubber gloves and a face shield

Hint

4. In the laboratory, long hair can be a(n) ______________.

a. eye hazard

b. chemical hazard

c. fire hazard

d. distraction

Hint

5. When a laboratory session ends, you must always ___________.

a. wash your hands

b. wash your eyes in the eye-wash fountain

c. take a shower in the chemical hazard shower

d. eat or drink in the laboratory

Hint

6. The practical use of scientific information is ___________.

a. technology

b. pure research
















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c. applied research

d. engineering

Hint

7. Research that is conducted for the sake of increasing fundamental knowledge is __________.

a. technology

b. applied research

c. pure research

d. theoretical research

Hint

8. Research on solubility as a factor that affects the digestion of foods in humans would be

an example of ______________.

a. technology

b. fundamental research

c. applied research

d. theoretical research

Hint

Section 2.1 Units of Measurement

Practice Test

1. Which of the following conversion factors would be most useful in converting feet per

second to kilometers per hour?

a. 1 hour/60 seconds

b. 1 mile/5280 feet

c. 1 mile/1.6 kilometers

d. 5280 feet/1 mile

Hint

2. A sample of gold has a mass of 15.7 g and displaces 0.81 cm of water in a graduated

cylinder. What is the density of gold?

a. .05 g/cm

b. 0.81 g/cm













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c. 14.9 g/cm

d. 19.3 g/cm

Hint

3. In the SI system of measurement, there are seven ________ unitsa. metric

b. English

c. derived

d. base

Hint

4. Which of the following lists units in increasing order by volume?

a. microliter < kiloliter < centiliter < liter < milliliter

b. milliliter < centiliter < microliter < kiloliter < liter

c. centiliter < milliliter < liter < kiloliter < microliter

d. microliter < milliliter < centiliter < liter < kiloliter

Hint

5. Below is a list of common prefixes used in the SI and metric systems. Included with each

is an abbreviation and a definition. Which set contains an error?

a. mega m 10

b. deci d 10-

c. centi c 10-

d. micro μ 10-

Hint

6. Which of the following physical properties is not paired with the correct SI base unit or

measurement?

a. length: meter

b. time: secondc. mass: gram

d. temperature: Kelvin

Hint

7. A unit that is defined by a combination of base units is a(n) ________________.
















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a. metric unit

b. English unit

c. SI unit

d. derived unit

Hint

8. What scale provides the base unit for temperature in the SI system?

a. Celsius scale

b. Fahrenheit scale

c. Kelvin scale

d. Centigrade

Hint

9. What volume is occupied by 16.4 g of mercury? The density of mercury is 13.6 g/mL.

a. 1.21 mL

b. 0.829 mL

c. 223 mL

d. 30.0 mL

Hint

10. What is the density of a sample with a mass of 24.47 g and a volume of 13.2 mL?a. 1.8537 g/mL

b. 1.854 g/mL

c. 1.85 g/mL

d. 1.9 g.mL

Hint

11. What is the temperature 51°C expressed in kelvins?

a. 273.0 K

b. 222 K

c. 324° K

d. 324 K

Hint

12. How many kilograms of lead are present in 426 pounds of lead? Each kilogram contains
















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2.20 pounds.

a. 937 kg

b. 937.2 kg

c. 193.6 kg

d. 194 kg

Hint

13. A cube of metal is 1.42 centimeters on an edge. It has a mass of 16.3 grams. What is thedensity of this metal?

a. 4.68 g/cm

b. 4.30 g/cm

c. 5.69 g/cm

d. 6.14 g/cmHint

Section 2.2 Scientific Notation and Dimensional Analysis

Practice Test

1. Which of the following numbers is equal to 2.70 x 10-

?

a. 0.00027

b. 0.675

c. 27 x 10-

d. 270

Hint

2. How many milliliters are there in 1.0 microliter?

a. 1.0 x 10-

mL

b. 1.0 x 10-

mL

c. 1.0 x 10-

mL

d. 1.0 x 10-

mL

Hint

3. What is the area in square millimeters of a rectangle that is 7.431 cm long and 23.33 mmwide?













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a. 1734 mm

b. 173.4 mm

c. 17.34 mm

d. 1.734 mm

Hint

4. Dimensional analysis is a method of problem-solving that focuses on ________ .

a. units

b. error

c. accuracy

d. precision

Hint

Section 2.3 How reliable are measurements?

Practice Test

1. Why are plus and minus signs ignored in calculating percent error?

a. Accepted values for data do not include negative numbers.

b. An experimental value smaller than the accepted value is not considered an

error.

c. Only the size of the error matters, not whether the values in error are larger or

smaller than the accepted values.

d. The formula for percent error automatically cancels out the signs.

Hint

2. Which of the following numbers has 4 significant figures?

a. 0.05208

b. 0.052

c. 0.0521

d. 0.52089

Hint

3. The sum of 4.824 + 2.03 + 4.72319 + 123.4567 + 111.1 expressed to the proper number

of significant figures is ____________.

a. 246













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b. 246.1

c. 246.13

d. 246.134

Hint

4. Perform the indicated mathematical operations and report the answer to the proper number of significant figures. (21.55 x 4.12 ) / 42.42

a. 2.1

b. 2.09

c. 2.093

d. 2.093 0

Hint

5. The closeness of an experimental value to an accepted value is its ______.

a. accuracy

b. precision

c. percent error

d. error

Hint

6. A measure of the closeness of a series of measurements is their ___________ .

a. accuracy

b. precision

c. percent error

d. error

Hint

7. What is the percent error if a measured value is 24.59 g/mL and the accepted value is

25.49 g/mL?

a. 0.9647 b. 1.037

c. 0.0366

d. 3.5%

Hint
















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8. Which of the following is not a significant figure?

a. a value of 1 through 9

b. a zero between two digits that have values of 1 through 9

c. zeros to the right of a significant digit when a decimal point is present

d. zeros to the left of the first significant digit

Hint

9. How many significant figures will there be when the density value is calculated from thefollowing data? mass = 24.47g; volume =13.2 mL

a. 2

b. 3

c. 4

d. 5Hint

Section 2.4 Representing Data

Practice Test

1. The representation of data on graph that resembles a pizza is a _________.

a. circle graph

b. bar graph

c. line graph

d. inverse

Hint

2. How is the slope of a linear graph calculated?

a. slope = (y2 – y1) x 100

b. slope = (x2 – x1) / 100

c. slope = (x2 – x1) / (y2 – y1)

d. slope = (y2 – y1) / (x2 – x1)

Hint













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ection 3.1 Properties of Matter

Practice Test

1.

Which of the following is a vapor?

a. helium

b. hydrogen

c. oxygen

d. steam

Hint

2. Matter that has a uniform and unchanging composition is a ___________.

a. substance

b. solid

c. vapor

d. mixture

Hint

3. Seawater is not a substance because __________.

a. it is salty







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b. its composition may be different from sample to sample

c. it is a liquid

d. it has hydrogen as part of its composition

Hint

4. A characteristic that can be observed or measured without changing the sample’scomposition is ________________.

a. a chemical property

b. a physical property

c. a gaseous property

d. a crystalline property

Hint

5. Properties that are dependent on the amount of substance present are ________.

a. intensive properties

b. extensive properties

c. chemical properties

d. external properties

Hint

6. When one substance changes identity, it exhibits a(n) __________.

a. chemical property

b. physical property

c. extensive property

d. intensive property

Hint

7. Which of the following is not a chemical change?

a. freezing of water

b. rusting of iron

c. placing iron in hydrochloric acid and producing hydrogen gas

d. burning a piece of wood

Hint

8. What type of property is observed when milk sours?
















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a. an intensive property

b. an extensive property

c. a physical property

d. a chemical property

Hint

9. Which of the following is not a state of matter?

a. solid

b. liquid

c. gas

d. density

Hint

10. A form of matter that has a fixed shape and occupies a fixed volume is a _________.

a. plasma

b. gas

c. solid

d. liquid

Hint

11. A vapor is a _________.a. gas

b. solid

c. liquid

d. condensed state

Hint

12. Which of the following is a pure substance?

a. soda

b. gun powder

c. sugar water

d. steam

Hint
















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Section 3.2 Changes in Matter

Practice Test

1. How can the law of conservation of mass apply to a burning log, if all that remains of it isash?

a. The ash has the same mass as the log, although a large percentage of it blowsaway.

b. The law of conservation of mass applies to changes of state but not to chemical

reactions.

c. The law of conservation of mass applies to substitution and displacementreactions, but not to combustion reactions.

d. The masses of the gases and water vapor released into the air by the

combustion reaction plus the mass of ash equal the mass of the log before

burning.

Hint

2. The statement, "Mass can neither be created nor destroyed" is the ____________.

a. law of conservation of mass

b. law of conservation of energy

c. law of multiple proportions

d. law of gravity

Hint

3. In the following chemical reaction, how do you classify hydrogen and oxygen?

2H2 + O2 → 2H2O

a. products

b. reactantsc. physical changes

d. chemical properties

Hint

4. How would you read the following chemical reaction?










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2Cu + O2 → 2CuO

a. copper yields oxygen and copper(II) oxide

b. copper(II) oxide yields oxygen and copper

c. copper and oxygen yield copper(II) oxide

d. oxygen yields copper and copper(II) oxide

Hint

5. How can you tell when the following reaction is balanced?

2Cu + O2 → 2CuO

a. mass of the reactants = mass of products

b. mass of reactants = volume of products

c. volume of reactants = mass of products

d. volume of reactants = volume of products

Hint

Section 3.3 Mixtures of Matter

Practice Test







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1.

A heterogeneous mixture is poured through a piece of filter paper that is positioned over a

beaker. What is the substance that is collected in the beaker?

a. heterogeneous mixture

b. solid

c. solution

d. Not enough information is given.

Hint

2.




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Which of the following methods of separating mixtures is best if the solution cannot be

saturated and is temperature-sensitive?

a. chromatography

b. combustion

c. crystallization

d. distillation

Hint

3. Which of the following is a heterogeneous mixture?

a. sugar

b. sugar in water

c. samples of nitrogen and oxygen in the same container

d. samples of argon and iron in the same container

Hint

4. A solution of solids is a(n)___________.

a. filtration

b. alloy

c. pure metal

d. heterogeneous mixture

Hint

5. A technique that uses a porous barrier to separate heterogeneous mixtures is _______.

a. distillation

b. chromatography

c. filtration

d. crystallization

Hint

6. A technique that uses the differences in boiling point to separate homogeneous mixtures is

_________.

a. distillation

b. chromatography

c. filtration













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d. crystallization

Hint

Section 3.4 Elements and Compounds

Practice Test

1. Sucrose is composed of carbon, hydrogen, and oxygen. Which of the following data is notneeded in order to determine the percent by mass of the three elements?

Analysis Data of Two Copper Compounds

Compound % Cu % Cl

Mass copper (g)100.0 g

of compound

Mass chlorine (g)100.0 g

of compound

Mass ratio

mass Cu

————

mass Cl

I 64.20 35.80 64.20 35.801.793 g Cu/1 g

Cl

II 47.27 52.73 47.27 52.730.8964 g Cu/1 g

Cl

a. mass of carbon in sucrose sample

b. mass of hydrogen in sucrose sample

c. total mass of the sucrose sample

d. atomic numbers of the elements

Hint

2. Which of the following is a compound?

a. water

b. neon

c. steel

d. crude oil

Hint

3. A material that cannot be broken down further by chemical means is a(n) ________.










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a. substance

b. element

c. mixture

d. compound

Hint

4. Periods are ___________ on the periodic table.

a. horizontal rows

b. vertical columns

c. left to right diagonals

d. right to left diagonals

Hint

5. Elements on the left side of the periodic table are _______.

a. nonmetals

b. metalloids

c. metals

d. compounds

Hint

6. Fe2O3 is a(n) _____________.a. compound

b. element

c. heterogeneous mixture

d. homogeneous mixture

Hint

7. What is the percent by mass of the components of water?

a. 11% H, 89% O

b. 89% H, 11% O

c. 67% H, 33% O

d. 33% H, 67% O

Hint
















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Section 4.1 Early Theories of Matter

Practice Test

1. Which of the following is not a fundamental particle in an atom?

a. element

b. electron

c. proton

d. neutron

Hint

2. What is the smallest particle of an element that maintains the properties of the element?

a. molecule

b. mixture

c. cation

d. atom

Section 4.2 Subatomic Particles and the Nuclear Atom

Practice Test

1.

What can you conclude from the deflection of a cathode ray in a magnetic field?

a. The ray must be composed of charged particles.

b. The ray must be composed of iron.

c. The ray must have a positive charge.




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d. The ray must need to travel in a vacuum.

Hint

2.

Which of the results of Rutherford's gold foil experiments was not consistent with the

plum pudding atomic model?

a. Most of the alpha particles passed straight through the gold foil.

b. Some alpha particles were deflected straight back toward the particle source.

c. The paths of alpha particles were only slightly altered by collisions with

electrons.

d. The zinc sulfide coated screen produced a flash of light whenever it was struck

by an alpha particle.

Hint

3.

What properties did Rutherford use in the design of the gold foil experiment?







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a. alpha particle's negative charge and random distribution of protons

b. alpha particle's negative charge and gold foil's positive charge

c. alpha particle's positive charge and electron's negative charge

d. positively charged electrons distributed in a uniform negative charge

Hint

4. What is the negatively — charged particle in an atom?

a. proton

b. positron

c. neutron

d. electron

Hint

5. Which of the following particles has a mass that is almost the same as the mass of a

proton?

a. neutron

b. electron

c. positron

d. beta particle

Hint

6. Which of the following statements is correct?

a. An electron is about 2000 times more massive than a proton.

b. A proton is about 2000 times more massive than an electron.

c. A neutron will always be found orbiting the nucleus.

d. The nucleus is mostly empty space.

Hint

7. Which scientist determined that almost all of an atom’s mass of is located in its nucleus?

a. a. Dalton

b. Democritus

c. Rutherford

d. Thomson

Hint
















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Section 4.3 How Atoms Differ

Practice Test

1. What is the name for atoms of an element that have different masses?

a. isotopes

b. isomers

c. allotropes

d. alloforms

Hint

2. The atomic number of an element is defined by its number of ________.

a. protons

b. neutrons

c. electrons

d. nuclei

Hint

3. How many protons are present in an atom potassium-39?

a. 19

b. 20

c. 39d. 58

Hint

4. The sum of the protons and neutrons in a nucleus is __________.

a. the atomic number

b. the mass number

c. Avogadro’s number

d. the element number Hint

5. What is the name of the element that has an atomic number of 3?

a. lanthanum

b. lithium













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c. beryllium

d. helium

Hint

6. Which of the following is a correct statement about a neutral atom?a. Neutrons are present in the nucleus.

b. The atoms carry a positive or a negative charge.

c. The atom has the same number of proton and electrons.

d. The atom is radioactive.

Hint

7. Which of the following is true for any atom?

a. atomic number = number of protons = number of electrons

b. atomic number = number of neutrons = number of electrons

c. mass number = number of protons = number of electrons

d. mass number = number of protons = number of neutrons

Hint

8. Which of the following statements concerning hydrogen is false?

a. All hydrogen isotopes have one proton.

b. All hydrogen isotopes have one electron.

c. All hydrogen isotopes have one neutron.

d. Hydrogen’s electron is not found in the nucleus.

Hint

9. An atom of an element contains eight electrons. What is the identity of this element?

a. nitrogen

b. oxygen

c. fluorine

d. carbon

Hint

10. How is the atomic mass unit (amu) defined?

a. 1/12 the mass of a carbon-12 atom

b. 1/14 the mass of a nitrogen-14 atom
















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c. 1/13 the mass of a carbon-13 atom

d. 1/16 the mass of an oxygen-16 atom

Hint

Section 4.4 Unstable Nuclei and Radioactive Decay

Practice Test

1. What is the name for the emission of rays and particles by a radioactive material?

a. radiation

b. nuclear reactivity

c. decay

d. radioactive seriesHint

2. What is the charge of an alpha particle?

a. 1+

b. 2+

c. 1 –

d. 0

Hint

3. What is the charge of a beta particle?

a. 1+

b. 2+

c. 1 –

d. 0

Hint

4. What is the charge of a gamma ray?a. 1

+

b. 2

c. 1 –

d. 0













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Hint

5. How do gamma rays differ from alpha particles and beta particles?

a. Alpha particle and beta particle emissions result in the formation of new atoms,whereas gamma ray emissions do not.

b. Gamma rays and beta particles result in the formation of new atoms, but alpha particles do not.

c. Gamma rays and alpha particles result in the formation of new atoms, but beta particles do not.

d. Gamma rays have mass, whereas alpha and beta particles do not.

Hint

6. What is the primary factor in determining an atom’s stability?

a. neutron to proton ratio

b. proton to electron ratio

c. neutron to electron ratio

d. alpha particle to beta particle ratio

Hint

7. What fundamental particle is identical to a beta particle?

a. the proton

b. the neutron

c. the electron

d. the positron

Hint

Section 5.1 Light and Quantized Energy

Practice Test

1.

Which of the following affects the amplitude of a wave?

a. amount of energy carried by the wave













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b. frequency

c. speed of light

d. wavelength

Hint

2.

What is the frequency of yellow light , which has a wavelength of 5.56 x 10-7

m?

a. 1.85 x 10 Hz

b. 1.85 x 10 m/s

c. 5.40 x 10 Hz

d. 5.40 x 10 m/s

Hint

3

.

What is the wavelength of a radio wave having a frequency of 3.75 x 107

Hz?







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a. .08 m

b. 0.8 m

c. 8.0 m

d. 80 m

Hint

4. The part of the electromagnetic spectrum that humans can see is the ________.

a. visible spectrum

b. infrared spectrum

c. ultraviolet spectrum

d. a, b, and c

Hint

5. One of mercury’s spectral lines has a wavelength of 6.234 x 10-

m. What is the line’s

frequency?

a. 4.81 x 10 s-

b. 1.87 x 10 s-

c. 4.81 x 10 s-

d. 1.87 x 10 s-

Hint

6. A wavelength of 500 nm is associated with the _______ portion of the electromagnetic

spectrum.

a. visible

b. infrared

c. ultraviolet

d. microwave

Hint

7. The minimum amount of energy that can be gained or lost by an electron is a _________.a. quart

b. queue

c. planck

d. quantum

Hint
















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8. Which of the following power sources makes use of the photoelectric effect?

a. solar power cells

b. nuclear power plants

c. alkaline batteries

d. diesel engines

Hint

9. When an atom loses energy by electron transition, a(n) _____________ is produced.

a. absorption

b. adsorption

c. ion

d. emissionHint

Section 5.2 Quantum Theory and the Atom

Practice Test

1. What is the maximum number of electrons that can be present in each principal energy

level of hydrogen?

Hydrogen’s First Four Principal Energy Levels

Principal

quantum number(n)

Sublevels (types of

orbitals) present

Number oforbitals

related to sublevel

Total number

of orbitals related to

principal energy level

(n

2

)1 s 1 1

2

s

p

1

34

3 s 135 9







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p

d

4

s

p

d

f

1

3

5

7

16

a. n

b. n

c. 2n

d. 2n

Hint

2. The Heisenberg uncertainty principle states that ____________.

a. no two electrons in the same atom can have the same set of four quantum

numbers

b. two atoms of the same element must have the same number of protons

c. it is impossible to simultaneously know the precise position and velocity of a

particle

d. electrons of atoms in their ground states enter energetically equivalent sets of

orbitals singly before they pair up in any orbital of the set

Hint

3. What is the lowest energy state of an atom called?

a. the ground state

b. the excited state

c. the solid state

d. the chaotic state

Hint

4. The ground state of hydrogen corresponds to the _________.

a. zeroeth energy level

b. first energy level










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c. second energy level

d. highest energy level

Hint

5. The concept that all moving particles have wave characteristics is attributed to ___________.

a. de Broglie

b. Thomson

c. Heisenberg

d. Bohr

Hint

6. In the periodic table, the principal quantum number corresponds to the ________.

a. group number

b. period number

c. lanthanides

d. transition metals

Hint

7. How many types of orbitals are present in atoms of elements in the third period of the

periodic table?

a. 1 b. 2

c. 3

d. 4

Hint

8. Which of the following statements is true?

a. Each set of d orbitals contains seven orbitals.

b. Each set of d orbitals can hold a maximum of 14 electrons.c. The first energy level contains only s and p orbitals.

d. All s orbitals are spherically shaped.

Hint
















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Section 5.3 Electron Configurations

Practice Test

1.

What is the wavelength of an object moving at 19.4 m/s and having a mass of 1000 kg?

a. 1.3 x 10-

m

b. 3.4 x 10-

m

c. 3.4 x 10-

m

d. 34 m

Hint

2. How many electrons can an orbital contain?

a. 1

b. 2

c. 3

d. 4

Hint

3. What is the symbol of the element that has the following electron configuration?1s

22s

22px

22py

22pz

23s

23px

23py

23pz

1

a. Si

b. P

c. S

d. Cl

Hint

4. The principle that states each electron occupies the lowest energy orbital available is the ______________.

a. aufbau principle

b. uncertainty principle

c. exclusion principle

d. photoelectric principle










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Hint

5. What is the electron configuration for tin (Sn)?

a. [Kr]5s 3d 3f 5p

b. [Kr]5s 3d 4d 5pc. [Kr]5s 3d 4f 5p

d. [Kr]5s 4d 5p

Hint

6. The valence orbitals in an atom are the ___________.

a. innermost orbitals

b. second energy level

c. d orbitals

d. outermost orbitals

Hint

7. What is an electron dot structure?

a. An element symbol surrounded by dots representing its valence electrons.

b. An element symbol surrounded by its innermost electrons.

c. An element symbol with a positive charge.

d. A filled noble gas in brackets plus the remaining electron configuration

expressed by filled orbitals.

Hint

8. An element with the outermost electron configuration, ns np , could be _______.

a. Pb

b. Zr

c. Mo

d. Se

Hint

9. An element whose atoms have four valence electrons is _______.

a. Nb

b. Cr

c. Sn
















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d. Ti

Hint

10. The total number of electrons in s orbitals in a germanium atom is ________.

a. 18 b. 15

c. 8

d. 6

Hint

11. How many electrons are in the highest energy p orbitals of a group 4A atom?

a. 1

b. 2

c. 3

d. 4

Hint

12. How many valence electrons does a group 1A metal atom have?

a. 1

b. 2

c. 3

d. 4

Hint

Section 6.1 Development of the Modern Periodic Table

Practice Test













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1.

Why was Mendeleev's periodic table widely accepted?

a. He organized the first 14 known elements.

b. He predicted the existence and properties of undiscovered elements.

c. He was the first to notice a pattern of similar properties among elements.

d. His periodic table listed all of the elements in the correct order.

Hint

2. In what part of the periodic table would you expect to find an element that emits light

when struck by electrons?

a. actinide series

b. lanthanide series

c. group 2A elements

d. group 3A elements

Hint

3. Columns on the periodic table are known as __________.

a. metalloids

b. periods

c. nonmetals







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d. groups

Hint

4. Which of the following are poor conductors of heat and electricity?

a. metals b. metalloids

c. nonmetals

d. alkaline earth elements

Hint

5. Which group on the periodic table is known as the halogens?

a. group 1A

b. group 2A

c. group 8A

d. group 7A

Hint

6. Which group on the periodic table is known as the alkaline earth metals?

a. group 1A

b. group 2A

c. group 8A

d. group 7A

Hint

7. Which of the following elements is a metalloid?

a. As

b. Na

c. W

d. F

Hint

8. Halogens are good disinfectants. Which of the following is a halogen?

a. N

b. O

c. Cl
















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d. Fe

Hint

Section 6.2 Classification of the Elements

Practice Test

1.

Why are both hydrogen and cesium s-block elements, when hydrogen has one electron

and cesium has 55?

a. All blocks contain at least one element from each period.

b. Blocks of elements on the periodic table are based only on an element's valenceelectrons.

c. The s-block includes only the most reactive elements.d. They have identical electron configurations.

Hint

2. What characteristic do atoms in the same group of elements share?

Period 1 hydrogen 1s

1s

Period 2 lithium 1s

2s

[He]2s

Period 3 sodium 1s

2s 2p

3s

[Ne]3s

Period 4 potassium 1s

2s 2p

3s

3p

4s

[Ar]4s

a. They have the same atomic mass.

b. The have the same number of electron orbitals.

c. They have the same number of valence electrons.

d. They have similar physical properties.







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Hint

3. Atoms of elements in group 1A have ____________.

a. one electron in their outermost energy level

b. two electrons in their outermost energy levelc. seven electrons in their outermost energy level

d. eight electrons in their outermost energy level

Hint

4. Which block consists of the elements in groups 3A through 8A?

a. s block

b. p block

c. d block

d. f block

Hint

5. Elements in the d block are also known as _______________.

a. alkali metals

b. alkaline earth metals

c. transition metals

d. lanthanide metals

Hint

6. Which of the blocks on the periodic table contains the most elements?

a. s block

b. p block

c. d block

d. f block

Hint

7. Which of the following is the general valence electron configuration for the alkaline earth

elements?

a. ns

b. ns

c. ns np
















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d. ns np

Hint

8. How would you classify an element that has an electron configuration of [Kr]5s 4d 5p ?

a. a representative element b. a transition metal

c. an alkali metal

d. a noble gas

Hint

Section 6.3 Periodic Trends

Practice Test

1.

Atoms with large ionization energy values are __________.

a. more likely to form positive ions

b. less likely to form positive ions

c. most likely to lose their outer electrons

d. lacking valence electrons

Hint










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2. The atomic radius generally _______________ on the periodic table.

a. increases from top to bottom and from right to left

b. increases from top to bottom and from left to right

c. increases bottom to top and from right to left

d. increases from bottom to top and from left to right

Hint

3. Atoms that lose electrons to form positive ions are __________.

a. nonmetals

b. metalloids

c. metals

d. noble gasesHint

4. What is the name for the amount of energy required to remove an electron from an atom?

a. ionization energy

b. electron affinity

c. ground state energy

d. dissociation energy

Hint

5. Which element has the lowest first ionization energy?

a. Li

b. Na

c. K

d. Cs

Hint

6. Which of the following have the same number of electrons? Cl –

, O –

, F, Ca , Fea. Ca

+and Fe

+

b. Ca+

and Cl –

c. F and Cl –

d. O –

and F

Hint
















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7. Which of the following elements is most likely to form a negatively-charged ion?

a. I

b. Br

c. Cl

d. F

Hint

8. In general, when comparing ions to their atoms, what trend is observed?

a. Positive ions are smaller and negative ions are larger than their corresponding

atoms.

b. Positive ions are smaller and negative ions are smaller than their corresponding

atoms.

c. Positive ions are larger and negative ions are larger than their correspondingatoms.

d. Positive ions are larger and negative ions are smaller than their corresponding

atoms.

Hint

9. Which of the following is an ion?

a. O-

b. C

c. HCl

d. lithium-3

Hint

10. How many protons and electrons are present in Al3+?

a. 13 protons, 13 electrons

b. 13 protons, 10 electrons

c. 10 protons, 13 electrons

d. 10 protons, 10 electrons

Hint

11. How many protons are present in Ba+?

a. 56

b. 54

c. 58













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d. 0

Hint

12. The ionic compound, sodium chloride, is formed from atoms of the elements sodium and

chlorine. What happens to the size of each atom when it forms an ion?

a. Sodium increases in size and chlorine increases in size.

b. Sodium increases in size and chlorine decreases in size.

c. Sodium decreases in size and chlorine decreases in size.

d. Sodium decreases in size and chlorine increases in size.

Hint

13. Which atom or ion has 16 protons and 18 electrons?

a. S-

b. Ne +

c. F-

d. N+

Hint

Section 7.1 Properties of s-Block Elements

Practice Test










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1.

Which of the following statements is true of alkali metals?

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a. They easily lose a valence electron.

b. They do not easily lose valence electrons.

c. They are highly stable.

d. They do not readily conduct electricity.

Hint

2.

Which of the following elements has the lowest first ionization energy?

a. B

b. C

c. F

d. Li

Hint

3. The elements in group 1A are grouped together because they __________.

a. are the lightest elements

b. are all metals

c. are all solids at room temperature

d. all have one valence electron

Hint

4. Elements in the s-block are found in groups __________.

a. 1A and 2A

b. 3A through 8A

c. 1B and 2B










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d. 1A through 8A

Hint

5. Magnesium has more properties in common with aluminum than with calcium. What kind

of relationship is demonstrated between these two elements?

a. vertical relationship

b. horizontal relationship

c. diagonal relationship

d. periodic relationship

Hint

6. Which element is the most abundant in the universe?

a. hydrogen

b. helium

c. lithium

d. oxygen

Hint

7. Alkali metals will lose electrons to form an ion that carries a charge of ________.

a. 1+

b. 2+

c. 1-

d. 2-

Hint

8. Which metal is the lightest?

a. aluminum

b. magnesium

c. lithium

d. boron

Hint

9. The most common sodium compound is ________.

a. sodium chloride

b. sodium nitrate
















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c. sodium hypochlorite

d. sodium phosphate

Hint

10. Which of the following alkali metals is an important human nutrient?a. lithium

b. potassium

c. rubidium

d. cesium

Hint

11. The major metallic element in the human skeleton is _________.

a. beryllium

b. calcium

c. strontium

d. barium

Hint

12. All of the isotopes of _____ are radioactive.

a. lithium

b. potassium

c. rubidium

d. francium

Hint

Section 7.2 Properties of p-Block Elements

Practice Test

1. Which group contains elements known as "salt formers"?

a. 1A

b. 3A

c. 7A













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d. 8A

Hint

2. Graphite and diamond are _________ of carbon.

a. ores b. allotropes

c. oxides

d. isotopes

Hint

3. Metals that are found to the right of the transition metals in the periodic table are located

in the _________.

a. s-block

b. p-block

c. d-block

d. f-block

Hint

4. Which element is the most abundant metal in Earth’s crust?

a. sodium

b. calcium

c. potassium

d. aluminum

Hint

5. An element or an inorganic compound that is found in nature as solid crystals is a __________.

a. mineral

b. trace element

c. solutiond. salt

Hint

6. Most of the elements in the p-block are _________.

a. metals
















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b. metalloids

c. nonmetals

d. liquids

Hint

7. Which one of the following elements will react most readily with chlorine?

a. lithium

b. sodium

c. rubidium

d. cesium

Hint

8. Elements in group 7A are __________.a. alkali metals

b. alkaline earth metals

c. halogens

d. noble gases

Hint

9. A lack of _______ can cause the thyroid gland to enlarge.

a. fluorine

b. chlorine

c. bromine

d. iodine

Hint

10. The most abundant noble gas on Earth is _________.

a. helium

b. neon

c. argon

d. krypton

Hint

Section 7.3 Properties of d-Block and f-Block Elements
















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Practice Test

1.

The fewer unpaired electrons in the d sublevel of a transition metal, the __________.

a. greater the hardness

b. higher the boiling point

c. higher the melting pointq

d. lower the boiling point

Hint

2. For which trace element in the table is the recommended daily allowance lowest?

Supplement Contents

Per Tablet

Element Amount % DV

Calcium 161 mg 16%

Chromium 25 µg 20%

Copper 2 mg 100%Iodine 150 µg 100%

Manganese 2.5 mg 125%

Magnesium 100 mg 25%

Molybdenum 25 µg 33%




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Phosphorus 09 mg 111%

Potassium 40 mg 1%

Selenium 20 µg 28%

Zinc 15 mg 100%

a. calcium

b. copper

c. molybdenum

d. selenium

Hint

3. Which of the following was the first trace element shown to be essential to humannutrition?

a. aluminum

b. zinc

c. iron

d. nickel

Hint

4. Which one of the following d-block elements would most likely have coloredcompounds?

a. titanium

b. zinc

c. vanadium

d. scandium

Hint

5. The chloride of ____________ is colorless.

a. Cr +

b. Co

c. Cu+

d. Ca+

Hint

6. Many transition metals are valued in manufacturing because these metals ________.













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a. are strong and malleable

b. are brittle

c. are non-conductors

d. are volatile

Hint

7. Which one of the following elements is unreactive enough to be found uncombined in

nature?

a. calcium

b. aluminum

c. gold

d. iron

Hint

8. The majority of actinide elements are known as _________.

a. halogen elements

b. coinage metals

c. lanthanide elements

d. transuranium elements

Hint

Section 8.1 Forming Chemical Bonds

Practice Test

1.

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










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a. 1s 2s 2p

b. [He]2s 2p

c. 1s 2s 2p 3s 3p

d. 1s 2s 2p 3s 3p 4s

Hint

2.

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

a. 1s 2s 2p

b. 1s 2s 2p 3s

c. 1s 2s 2p 3s 3p

d. 1s 2s 2p 3s 3p 4s

Hint

3. _____________ is the force that holds two atoms together.

a. A chemical bond

b. Glue

c. Nuclear force

d. Fission

Hint

4. What forms chemical bonds?

a. atomic nuclei

b. valence electrons

c. inner-level electrons

d. noble gases

Hint

5. A positive ion forms when ___________.

a. an atom loses one or more valence electrons

b. an atom gains one or more valence electrons













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c. electrons are pulled into the nucleus

d. electrons are pushed out of the nucleus

Hint

6. What is a negatively charged ion called?a. nucleus

b. cation

c. anion

d. molecule

Hint

7. How many electrons are present in the valence level for all noble gases except helium?

a. 6

b. 7

c. 8

d. 9

Hint

8. Why is the calcium ion (Ca+) 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.

Hint

9. Which elements can either gain or lose electrons to form stable octets?

a. metals

b. metalloids

c. nonmetals

d. transition metals

Hint

Section 8.2 The Formation and Nature of Ionic Bonds
















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Practice Test

1. Which of the following lattice energies represents the strongest force of attraction holding

the ions in place?

Lattice Energies of Some Ionic Compounds

Compound Lattice energy Compound Lattice energy

KI – 632 KF – 808

KBr – 671 AgCl – 910

RbF – 774 NaF – 910

NaI – 682 LiF – 1030

NaBr – 732 SrCl2 – 2142

NaCl – 769 MgO – 3795

a. -632

b. -769

c. -910

d. -2142

Hint

2. Which ion has the electron configuration 1s 2s 2p 3s 3p ?

a. H+

b. F-

c. Cl-

d. Br -

Hint

3. How are ionic bonds formed?

a. sharing of electrons between atoms

b. electrostatic forces between ions

c. loss of electrons by atoms

d. gain of electrons by atoms

Hint

4. An ionic compound that will conduct an electric current when it forms an aqueous










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solution is a(n) _________.

a. electrolyte

b. nonelectrolyte

c. molecular compound

d. crystal lattice

Hint

5. What is the name given to the energy required to separate one mole of the ions of an ioniccompound?

a. heat energy

b. kinetic energy

c. potential energy

d. lattice energyHint

Section 8.3 Names and Formulas for Ionic Compounds

Practice Test

1. How many ions are produced when one unit of sodium phosphate (Na3PO4) dissolves in

water?

a. 2 b. 3

c. 4

d. 8

Hint

2. What is the name given to the simplest ratio of ions in an ionic compound?

a. empirical formula

b. formula unit

c. mole

d. monatomic ion

Hint

3. What kind of ion is iodide (I-)?













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a. monatomic anion

b. polyatomic anion

c. monatomic cation

d. polyatomic cation

Hint

4. What is the charge on a monatomic ion?

a. reduction number

b. oxidation number

c. mass number

d. Avogadro’s number

Hint

5. Sulfate (SO4-) is a ____________.

a. monatomic anion

b. polyatomic anion

c. monatomic cation

d. polyatomic cation

Hint

6. What is the correct name of Fe2O3?a. iron oxide

b. iron(I) oxide

c. iron(II) oxide

d. iron(III) oxide

Hint

7. What is the formula for ammonium bromide?

a. NH4Br

b. NH3Br

c. (NH4)2Br

d. NH4Br 2

Hint

8. Name the pair that contains an incorrect formula.
















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a. aluminum phosphate AlPO4

b. iron(II) sulfite FeSO3

c. silver carbonate AgCO3

d. magnesium hydroxide Mg(OH)

Hint

9. Which one of the following classifications of ions in the table contains an error?

Hydroxide OH-

monatomic anion

carbonate CO3-

polyatomic anion

ammonium NH+

polyatomic cation

magnesium Mg+

monatomic cation

a. hydroxide

b. carbonate

c. ammonium

d. magnesium

Hint

Section 8.4 Metallic Bonds and Properties of Metals

Practice Test

1. If tungsten were not available for use in a lightbulb filament, which of the elements in thetable would be the best choice for a substitute?

Melting Points and Boiling Points of Some Metals

Element Melting point (°C) Boiling point (°C)

Lithium 180 1347

Tin 232 2623

Aluminum 660 2467

Barium 727 1850

Silver 961 2155

Copper 1083 2570







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a. aluminum

b. barium

c. copper

d. lithium

Hint

2. Why are metals described as having a "sea of electrons"?

a. The electrons are wet.

b. The electrons move in waves.

c. The electrons are free to move between the atoms in a metal.

d. The electrons are fixed to a particular metal atom.

Hint

3. Which of the following is a correct description of a metallic property?

a. Metals are brittle.

b. Metals have a dull appearance.

c. Metals gain electrons to form anions.

d. Metals are malleable.

Hint

4. Whichof the following is a mixture of elements that has metallic properties?a. an alloy

b. a suspension

c. a gas

d. a pure metal

Hint

Section 9.1 The Covalent Bond

Practice Test













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1.

Which of the following bond length represents the strongest bond?

a. 0.95 x 10-

m

b. 1.10 x 10-

m

c. 1.21 x 10-

m

d. 1.43 x 10-

m

Hint

2. When two or more atoms join together by sharing their outermost electrons, ________.

a. an ionic compound is formed

b. a molecular compound is formed

c. a mixture is formed

d. a solution is formed

Hint

3. Single covalent bonds are also referred to as ___________.

a. sigma bonds

b. pi bonds

c. delta bonds

d. hydrogen bonds

Hint

4. What is the charge on the simple monatomic ion that sulfur forms?

a. 1

b. 2+










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c. 1 –

d. 2-

Hint

5. How many sigma and how many pi bonds are present in the ethene molecule (C2H4)?a. 4 sigma and 2 pi

b. 5 sigma and 2 pi

c. 4 sigma and 1 pi

d. 5 sigma and 1 pi

Hint

6. Side — on overlaps of atomic orbitals are also referred to as ___________.

a. sigma bonds

b. pi bonds

c. delta bonds

d. hydrogen bonds

Hint

Section 9.2 Naming Molecules

Practice Test

1. What is the formula for the binary compound of potassium and chlorine?

Binary Molecular Compound Binary Acid

Contains only two different

elementsContains hydrogen and one other element

Second element uses name + -ideUses hydro + root of second element + -ic and

acid

a. KCl b. KCl2

c. K 2ClO3

d. KHCl

Hint













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2. Which of the following is the formula for nitric acid?

Formula Acid Name

H2SO3 Sulfurous acid

H2SO4 Sulfuric acid

a. HN

b. HNO

c. HNO2

d. HNO3

Hint

3. What is the molecular compound name for hydrazine, N2H4?

Number of atoms Prefix

1 mono-

2 di-

3 tri-

a. dinitrogen hydride

b. dinitrogen tetrahydride

c. trinitrogen pentahydride

d. trinitrogen quatrohydride

Hint

4. A binary compound is a ___________.

a. compound made of one type of atom

b. compound made of polyatomic ions

c. compound made of two types of atoms

d. compound made of three types of atoms

Hint

5. What is the name of the aqueous solution of H2SO4?

a. hydrogen persulfide

b. sulfuric acid










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c. sulfurous acid

d. hydrogen persulfate

Hint

6. What is the formula for the binary compound of aluminum and nitrogen?a. Al2 N2

b. Al2 N

c. AlN2

d. AlN

Hint

7. The correct molecular compound name for H2O2 is ________.

a. hydrogen oxide

b. dihydrogen dioxide

c. dihydrogen monoxide

d. hydrogen dioxide

Hint

8. What is the correct name for HBr in an aqueous solution?

a. hydrobromic acid

b. hydrogen bromide

c. bromic acid

d. bromous acid

Hint

9. What is the correct name for H2SO3 in an aqueous solution?

a. sulfuric acid

b. sulfurous acid

c. hydrosulfuric acid

d. hydrosulfic acid

Hint

10. What is the chemical formula for dintrogen hexoxide?

a. NO3

b. NO2
















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c. N2O5

d. N2O6

Hint

Section 9.3 Molecular Structures

Practice Test

1. When a hydrogen atom is part of a molecular structure, it is always a ________ atom.

a. terminal

b. ionic

c. centrald. 1

–

Hint

2. In the compound H2SO4, which element will be the central atom in the molecular structure?

a. H

b. S

c. O

d. both S and O

Hint

3. In forming a molecular structure, each atom should have ________ electrons around it

unless there is some specific reason why this cannot be achieved.

a. 2

b. 4

c. 6

d. 8Hint

4. Which element can share only one pair of electrons?

a. H

b. C













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c. N

d. O

Hint

5. To determine the number of bonding electron pairs, you must divide the number of electrons available for bonding by _______.

a. 1

b. 2

c. 3

d. 4

Hint

6. How many lone pairs of electrons are present in a molecule of water?

a. 1

b. 2

c. 3

d. 4

Hint

7. When two or more Lewis structures can be used to represent a single molecule, these

multiple structures are ______________.

a. resonance structures b. isomeric structures

c. polar structures

d. liquid structures

Hint

8. Which of the following atoms will be an exception to the octet rule in a molecule?

a. O

b. Nc. C

d. B

Hint
















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Section 9.4 Molecular Shape

Practice Test

1. Many simple molecules contain two lone pairs around the central atom. Which type of

electron — to — electron repulsions are smallest in these molecules?a. bonding pair to bonding pair

b. bonding pair to lone pair

c. lone pair to lone pair

d. All repulsions are equal.

Hint

2. What hybridization is present in an octahedral molecular geometry?

a. sp

b. sp

c. sp

d. sp d

Hint

Section 9.5 Electronegativity and Polarity

Practice Test

1.







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Large differences in electronegativity result in __________ bonding between atoms.

a. covalent

b. ionic

c. no

d. polar

Hint

2. 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°

Hint

3. Which element has the highest electronegativity?

a. N

b. O

c. F

d. Ne

Hint

4. Which compound consists of nonpolar molecules?

a. H2S

b. PH3

c. AsH3

d. SiH4

Hint

Section 10.1 Reactions and Equations

Practice Test













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1.

What is the correct word equation for the chemical equation 2Fe(s) + 3Cl2(g) →

2FeCl3(s)?

a. Aqueous iron and chlorine gas react to produce solid iron chlorate.

b. Aqueous iron and chlorine gas react to produce aqueous iron chlorate.

c. Solid iron and chlorine gas react to produce solid iron chlorate.

d. Solid iron and chlorine gas react to produce solid iron chloride.

Hint

2. Which of the following observations could be an indication of a chemical reaction of

antimony?

a. It is a solid at room temperature.

b. It melts at 631°C.

c. It expands upon freezing.

d. It burns in an atmosphere of chlorine.Hint

3. In the following unbalanced chemical reaction, which components are classified as the

reactants?

C7H16 + O2 → CO2 + H2O

a. C7H16 and O2

b. C7H16 and CO2

c. C7H16 and H2O

d. CO2 and H2O

Hint

4. The _________ symbol separates the reactants from the products and means "to yield".

a. +










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b. (s)

c. →

d. (g)

Hint

5. In a chemical equation, the number written in front of the chemical formula is the _________.

a. reactant

b. product

c. equation

d. coefficient

Hint

6. When there are exactly the same number and type of atoms on both sides of a chemicalequation, the equation is ____________.

a. balanced

b. ionic

c. synthetic

d. decomposed

Hint

7. A balanced chemical equation is a direct presentation of the ___________.a. law of conservation of momentum

b. law of conservation of energy

c. law of conservation of mass

d. ideal gas law

Hint

8. Balance the following equation. What is the coefficient for hydrogen fluoride?

SiO2 + HF → SiF4 + H2O

a. 1

b. 2

c. 3

d. 4













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Hint

9. Consider the following balanced equation. Which one of the following statements is

false?

a. One molecule of O2 will react with 2 molecules of H2.

b. One molecule of O2 will produce 2 molecules of H2O.

c. Two molecules of H2 will produce two molecules of H2O.

d. Two molecules of H2 will produce one molecule of H2O.

Hint

Section 10.2 Classifying Chemical Reactions

Practice Test

1. Will the double displacement reaction NaOH + CaBr 2 occur? If so, what is the product of

the reaction?

a. No

b. Yes, Br(OH)2

c. Yes, CaNa

d. Yes, Ca(OH)2

Hint










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2.

Which of the following single-replacement reactions will occur?

a. Br 2(g) + HF(aq) →

b. Cl2(g) + HBr(aq) →

c. Au(s) + Cu(NO3)2 (aq) →

d. Cu(s) + ZnCl2(aq) →

Hint

3. What type of reaction occurs when potassium and chlorine gas produce potassium

chloride?

Predicting Products of Chemical Reactions

Class of reaction Reactants Probable products

Synthesis Two or more substances One compound

Combustion A metal and oxygen The oxide of the metal




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A nonmetal and oxygen

A compound and oxygen

The oxide of the nonmetal

Two or more oxides

Decomposition One compound Two or more elements and/or compounds

Single-replacement

A metal and a compound

A nonmetal and

acompound

new compound and the replaced metal

A new compound and the replaced

nonmetal

Double-

replacementTwo compounds

Two different compounds,

one of which is often a solid, water, or a gas

a. combustion

b. decomposition

c. replacement

d. synthesis

Hint

4. What are the products of the reaction between barium hydroxide and hydrochloric acid?

a. barium chloride and water

b. barium hydroxide and water

c. barium chloride and barium hydride

d. chloric acid and barium

Hint

5. The following chemical equation is classified as a ______________.

CaCO3 → CaO + CO2

a. synthesis reaction

b. decomposition reaction

c. single-replacement reaction

d. double-replacement reaction

Hint

6. In the combustion reaction of propane, C3H8, what are the products?










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a. C3H8 and O2

b. C3H8 and CO2

c. O2 and H2O

d. H2O and CO2

Hint

7. The major physical indicator of a combustion reaction is _______.

a. heat

b. a color change in solution

c. the formation of a precipitate

d. cold

Hint

8. The chemical equation 2H2 + O2 → 2H2O represents a __________.

a. synthesis reaction




Hint

9. The chemical equation Zn + 2HCl → H2 + ZnCl2 represents a _________.a. synthesis reaction




Hint

Section 10.3 Reactions in Aqueous Solutions

Practice Test













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1.

Which of these is a spectator ion in the reaction 2NaOH(aq) + CuCl2 → 2NaCl(aq) +

Cu(OH)2(s)?

a. Cu

b. Cl and Na

c. Na

d. OH

Hint

2. What is the complete ionic equation for the following chemical equation?

Cl2(aq) + Na2SO4(aq) → BaSO4(s) + 2NaCl(aq)

a. Ba+(aq) + Cl

-(aq) + Na+(aq) + SO4

-(aq) → Na+(aq) + Cl

-(aq) + BaSO4(s)

b. Ba+(aq) + 2Cl

-(aq) + 2Na

+(aq) + SO4

-(aq) → 2Na

+(aq) + 2Cl

-(aq) + BaSO4(s)

c. BaCl2(aq) + 2Na (aq) + SO4-(aq) → 2Na (aq) + 2Cl

-(aq) + BaSO4(s)

d. Ba (aq) + SO4-(aq) → BaSO4(s)

Hint

3. In the chemical equation below, which of the components is a precipitate?







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HCl(aq) + AgNO3(aq) → AgCl(s) + HNO3(aq)

a. HCl(aq)

b. AgNO3(aq)

c. AgCl(s)

d. HNO3(aq)

Hint

4. The components of a chemical equation that are left after the spectator ions are removed

are the ____________.

a. spectator ions

b. reactants

c. products

d. participating ions

Hint

5. In the reaction below, which of the components will precipitate?

NaCl(aq) + KBr(aq) → KCl(aq) + NaBr(aq)

a. KCl(aq)

b. NaBr(aq)

c. KBr(aq)

d. none of the components precipitate

Hint

6. Which of the following components will always be produced during a chemical reaction

between an acid and a base?

a. an acid

b. a base

c. a gas

d. water

Hint

7. Na+(aq) is a _________________ in the following complete ionic equation.

H+(aq) + Cl

-(aq) + Na

+(aq) + OH

-(aq) → Na

+(aq) + Cl

-(aq) + H2O(l)













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a. participating ion

b. reacting compound

c. reacting molecule

d. spectator ion

Hint

Section 11.1 Measuring Matter

Practice Test

1.

Determine the number of atoms in 3.54 mol S.

a. 1.70 x 10

b. 1.70 x 10

c. 2.13 x 10

d. 2.13 x 10

Hint

2. What is the SI base unit used to measure the amount of a substance?

a. kilogram

b. mole

c. kelvin

d. meter

Hint

3. How many particles are present in one mole of particles?

a. 6.02 x 10










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b. 3.0 x 10

c. 1

d. 12

Hint

Section 11.2 Mass and the Mole

Practice Test

1. How many moles of chlorine are in 100 g chlorine (Cl)?

Element Molar Mass (g/mol)

Hydrogen 1.01

Carbon 12.01

Chlorine 35.45

a. 0.355

b. 2.82

c. 64.6

d. 100

Hint

2. How many atoms of carbon are in a pure sample of carbon having a mass of 72.0 g?

Element Molar Mass (g/mol)

Hydrogen 1.01

Carbon 12.01

Chlorine 35.45

a. 3.6 x 10

b. 36 x 10

c. 3.6 x 10

d. 4.33 x 10

Hint










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3. What is the mass of one mole of carbon-12 atoms?

a. 12 amu

b. 12 grams

c. 6 amu

d. 6 grams

Hint

4. How many moles of hydrogen are present in 2.0 grams of hydrogen gas?

a. 2.0 mol

b. 0.5 mol

c. 1.0 mol

d. 4.0 molHint

5. What is the mass of 5.0 x 10 molecules of water?

a. 8.3 x 10-

g

b. 0.15 g

c. 5.0 g

d. 6.02 x 10 g

Hint

6. How many moles of atoms are present in one mole of water?

a. 3 mol

b. 6 mol

c. 1 mol

d. 2 mol

Hint

7. How many moles of hydrogen atoms are present in 2.0 moles of ammonia NH(3)?a. 3.0 mol

b. 6.0 mol

c. 8.0 mol

d. 2.0 mol

Hint
















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8. What is the molar mass of H2CrO4?

a. 118 g

b. 52 g

c. 64 g

d. 2 g

Hint

9. How many molecules of ammonia are present in 34.0 g of ammonia NH(3)?

a. 1.20 x 10

b. 6.02 x 10

c. 1.77 x 10

d. 34Hint

10. How is one gram defined?

a. one mole of amu

b. one mole of protons

c. 10 moles of amu

d. 0.1 mole of amu

Hint

11. A sample of a compound is found to contain 0.404 grams of Fe and 0.174 grams of O.

What is the empirical formula for this compound?

a. FeO

b. FeO2

c. Fe2O3

d. Fe3O4

Hint

12. What is the mass of 1.0 x 10 molecules of O2?

a. 1.9 x 10 g

b. 6.0 x 10 g

c. 2.7 x 10-

g

d. 5.3 x 10-

g













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Hint

13. How many moles of C3H8 are present in 451 g C3H8?

a. 3.84 mol

b. 0.879 molc. 10.2 mol

d. 1.44 mol

Hint

14. Calculate the number of atoms in 4.00 x 10-

g sodium.

a. 1.05 x 10 atoms

b. 3.92 x 10 atoms

c. 3.24 x 10 atoms

d. 5.54 x 10 atoms

Hint

Section 11.3 Moles of Compounds

Practice Test

1.

Calculate the number of moles of hydrogen atoms in 4.75 mol sulfuric acid.

a. 2.38 mol

b. 4.75 mol

c. 9.50 mol

d. 33.2 molHint

Section 11.4 Empirical and Molecular Formulas

Practice Test













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1. What relationship is described by the formula below? (mass of the element/mass of the

compound) x 100

a. percent composition

b. mole ratio


d. empirical formula

Hint

2. What is the percent composition of carbon dioxide CO2?

a. 73% C and 27% O

b. 27% C and 73% O

c. 33% C and 66% O

d. 50% C and 50% O

Hint

3. How can the actual molecular formula be determined?

a. from the empirical formula alone

b. from the molar mass alone

c. from the molar mass divided by the mass of the empirical formula

d. from the molar mass multiplied by the mass of the empirical formula

Hint

4. An empirical formula mass is found to be 17 g/mol. The molar mass is found to be 34g/mol. What is the number of empirical formula units in this molar mass?

a. 2 units

b. 0.5 unit

c. 17 units

d. 1 unit

Hint

5. Analysis of a covalent compound showed that it contained 14.4% hydrogen and 85.6%carbon by mass. What is the empirical formula for this compound?

a. CH

b. CH2













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c. CH3

d. C2H3

Hint

6. A compound contains sulfur, oxygen, and chlorine. Analysis of a sample showed that itcontained by mass 26.95% sulfur and 59.61% chlorine. What is the simplest formula for this compound?

a. SOCl

b. SOCl2

c. SO2Cl2

d. SO2Cl

Hint

7. A 4.628-g sample of an iron oxide was found to contain 3.348 g of iron and 1.280 g of oxygen. What is the simplest formula for this compound?

a. FeO

b. Fe2O3

c. Fe3O4

d. FeO2

Hint

8. What mass of iron is contained in 62.8 grams of pyrite FeS2?a. 40.3 g

b. 29.2 g

c. 31.7 g

d. 58.5 g

Hint

9. What mass of calcium metal could be obtained from 1 kg of limestone that is 50.0% pure

CaCO3?

a. 0.05 kg

b. 0.2 kg

c. 0.4 kg

d. 0.5 kg

Hint
















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Section 11.5 The Formula for the Hydrate

Practice Test

1. What is the correct chemical formula for barium hydroxide octahydrate?

Prefix Molecules H2O

Mono- 1

Di- 2

Tri- 3

Tetra- 4

Penta- 5

Hexa- 6

Hepta- 7

Octa- 8

Deca- 10

a. Ba(OH)2 . 8H2O

b. Ba(OH)2 . H2O

c. 8Ba(OH)2 . H2O

d. Ba . 8(OH)2 Hint

2. A compound that has a specific number of water molecules bound to the structure is

______________?

a. a hydrate

b. an anhydrate

c. aqueous

d. waterlogged

Hint

3. Which word indicates that two water molecules are associated with a substance in a

chemical compound?

a. mono-

b. di-







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c. tri-

d. an-

Hint

Section 12.1 What is stoichiometry?

Practice Test

1. What is the scientific law that states that matter is not created or destroyed but only

transformed in a chemical reaction?

a. law of conservation of energy

b. law of conservation of mass

c. law of conservation of momentum

d. law of gravity

Hint

2. What is the first thing you must do to solve a stoichiometric problem?

a. Find the limiting reactant.

b. Find the excess reactant.

c. Write a balanced chemical equation.

d. Find the empirical formula.

Hint

3. What do the coefficients in the following balanced chemical equation mean?

2H2 + O2 → 2H2O

a. 2 moles of hydrogen and 1 mole of oxygen in the reactants

b. 2 moles of oxygen and 1 mole of hydrogen in the reactants

c. 2 moles of oxygen and 1 mole of hydrogen in the products

d. 2 moles of hydrogen and 1 mole of oxygen in the products

Hint

4. What is the ratio between the coefficients of any two substances in a balanced equation?

a. molar mass balanced equation

b. mole ratio













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c. quadratic equation

d. chemical formula

Hint

5. What is the balanced chemical equation for the reaction between hydrochloric acid andcalcium carbonate when the products of the reaction are calcium chloride, carbon dioxide,and water?

a. HCl + CaCl2 → CaCO3 + CO2 + H2O

b. HCl + CaCO3 → CaCl2 + CO2 + H2O

c. 2HCl + CaCO3 → CaCl2 + CO2 + H2O

d. 2HCl + CaCl2 → CaCO3 + CO2 + 2H2O

Hint

6. How many moles of CO2 would be produced from 56 moles of O2 according to thefollowing balanced equation?

2C2H6 + 7O2 → 4CO2 + 6H2O

a. 16 mol

b. 32 mol

c. 224 mol

d. 48 mol

Hint

7. How many grams of Fe2O3 are present in 0.82 mol Fe2O3?

a. 130 g

b. 65 g

c. 195 g

d. 260 g

Hint

8. How many grams of O2 are required to burn 18 g C5H12?

C5H12 + 8O2 → 5CO2 + 6H2O

a. 16 g

b. 32 g

c. 64 g













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d. 80 g

Hint

9. Balance the following equation with the smallest whole-number coefficients. What is the

coefficient for H2O in this equation?

PBr 3 + H2O → H3PO3 + HBr

a. 2

b. 3

c. 4

d. 1

Hint

10. Balance the following equation with the smallest whole — number coefficients. How manymoles of Ag2S can be prepared from 3 moles of Ag?

Ag + H2S + O2 → Ag2S + H2O

a. 1.0 mol

b. 1.5 mol

c. 2.0 mol

d. 4.0 mol

Hint

11. Balance the following equation with the smallest whole — number coefficients. What is

the coefficient for O2 in the balanced equation?

C5H10 + O2 → CO2 + H2O

a. 5

b. 10

c. 15

d. 20

Hint

Section 12.2 Stoichiometric Calculations

Practice Test













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1.

Based on the following equation, how many moles of hydrochloric acid are needed to

react with 0.64 moles of potassium permanganate?

2KMnO4 + 8 HCl → 3 Cl2 + 2 MnO2 + 4 H2O + 2KCl

a. 0.21 mol HCl

b. 0.64 mol HCl

c. 2.7 mol HCl

d. 5.1 mol HCl

Hint

2. If sufficient hydrochloric acid is used to react completely with 48.6 g of magnesium, howmuch hydrogen will be produced?

2HCl + Mg → MgCl2 + H2

a. 3 g

b. 2 mol

c. 1 mol

d. 6 g

Hint

3. What mass of SrF2 can be prepared from the reaction of 10.0 g Sr(OH)2 with excess HF?

Sr(OH)2 + 2HF → SrF2 + 2H2O

a. 9.67 g

b. 9.82 g

c. 10.0 g

d. 10.3 g

Hint

4. How many moles of sulfur are present in 5 moles of H2SO4?

a. 1 mol

b. 2 mol

c. 5 mol










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d. 10 mol

Hint

5. Balance the following equation with the smallest whole — number coefficients. How many

grams of O2 will be produced if 23.2 g of XeF2 reacts with excess water?

XeF2 + H2O → Xe + HF + O2

a. 2.19 g

b. 1.10 g

c. 3.31 g

d. 4.42 g

Hint

Section 12.3 Limiting Reactants

Practice Test

1. The substance that limits the extent of a chemical reaction has a special name. What isthis substance called?

a. limiting reactant

b. limiting product

c. excess reactant

d. excess product

Hint

2. What mass of SrF2 can be prepared from the reaction of 8.05 g of Sr(OH)2 with 3.88 g of HF?

Sr(OH)2 + 2HF → SrF2 + 2H2O

a. 11.7 g

b. 12.2 g

c. 10.5 g

d. 8.32 g

Hint

3. How many moles of HCl will just react with 0.424 g Ba(OH)2?













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2HCl + Ba(OH)2 → BaCl2 + 2H2O

a. 4.94 x 10-

mol

b. 9.90 x 10-

mol

c. 2.48 x 10-

mol

d. 1.24 x10- mol

Hint

4. In the chemical reaction below, 3.27 grams of Zn are reacted with 3.30 grams of HCl.Which component will limit the reaction?

Zn + 2HCl → ZnCl2 + H2

a. Zn

b. HCl

c. ZnCl2

d. H2

Hint

5. Which reactant in the following reaction is in excess when 9.8 grams of Ca(OH)2 is

reacted with 9.8 grams of H3PO4?

3Ca(OH)2 + 2H3PO4 → Ca3(PO4)2 + 6H2O

a. Ca(OH)2

b. H3PO4

c. Ca3(PO4)2

d. H2O

Hint

6. What is the maximum amount of Ca3(PO4)2 that can be prepared from 9.8 grams of

Ca(OH)2 and 9.8 grams of H3PO4?

3Ca(OH)2 + 2H3PO4 → Ca3(PO4)2 + 6H2O

a. 6.8 g b. 8.6 g

c. 10.3 g

d. 13 g

Hint













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Section 12.4 Percent Yield

Practice Test

1. What is the maximum amount of product that can be produced from a given amount of

reactant?

a. percent yield

b. theoretical yield

c. actual yield

d. mole ratio

Hint

2. What is the percent yield of CO2 if a reaction using 10.0 g CO with excess O2 produces12.8 g CO2?

2CO + O2 → 2CO2

a. 76.40%

b. 78.10%

c. 81.50%

d. 84.40%

Hint

3. The reaction of 5.0 grams of fluorine with excess chlorine produced 5.6 grams of ClF3.What percent yield of ClF3 was obtained?

Cl2 + 3F2 → 2ClF3

a. 58%

b. 69%

c. 76%

d. 86%

Hint

4. What is the percent yield for the reaction between 9.8 grams of Ca(OH)2 and 9.8 grams of

H3PO4 when 2.5 grams of

Ca3(PO4)2 are actually obtained?










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3Ca(OH)2 + 2H3PO4 → Ca3(PO4)2 + 6H2O

a. 19%

b. 9%

c. 27%

d. 36%

Hint

Section 13.1 Gases

Practice Test

1.

According to the kinetic-molecular theory, which of these describes a gas?

a. large particles in constant, random motion

b. large particles far apart in uniform motion

c. small particles in constant, random motion

d. small particles far apart in uniform motion

Hint







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2.

Which of the following illustrates effusion of a gas?

a. Cooking aromas from the kitchen can be detected by a person on the front

porch.

b. A dog follows a scent during a search and rescue mission.

c. Fragrance of cologne is present in a room in which none has been used.

d. A tire deflates after being punctured by a nail.

Hint

3.

A mixture of oxygen, carbon dioxide and nitrogen has a total pressure of 0.78 atm. What




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is the partial pressure of O2, if the partial pressure of CO2 is 0.46 atm and the partial

pressure of N2 is 0.08 atm?

a. 0.15 atm

b. 0.24 atm

c. 1.69 atm

d. 2.40 atm

Hint

4. The theory that describes the behavior of gases in terms of the motion of gas particles is

_____________.

a. the theory of intermolecular forces

b. Graham’s law of effusion

c. Dalton’s law of parhal pressures

d. the kinetic-molecular theory

Hint

5. The measure of the average kinetic energy of the particles in a sample of matter is _____.

a. temperature

b. speed

c. velocity

d. heat

Hint

6. Which of the following is the general definition of a gas?

a. Matter with a fixed volume and a fixed shape.

b. Matter with a fixed volume but no fixed shape.

c. Matter with no fixed volume and no fixed shape.

d. Matter with a fixed shape but no fixed volume.

Hint

7. Which of the following gases will diffuse the fastest at room temperature?

a. Ne

b. CO2

c. N2

d. H2













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Hint

8. At the same ______, the particles of different gases have the same average kinetic energy.

a. volume

b. pressurec. density

d. temperature

Hint

9. A device that is used to measure atmospheric pressure is a ___________.

a. barometer

b. thermometer

c. kilometer

d. micrometer

Hint

10. One atmosphere is equal to ____________.

a. 1 mm Hg

b. 760 cm Hg

c. 760 mm Hg

d. 1 cm Hg

Hint

11. In a balloon filled with air (about 80% nitrogen and 20% oxygen), the pressure in the

balloon is primarily ___________________________.

a. the pressure exerted by nitrogen

b. the pressure exerted by oxygen

c. the sum of the pressures exerted by nitrogen and oxygen

d. zero

Hint

Section 13.2 Forces of Attraction

Practice Test
















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1.

Which of the following molecules can form hydrogen bonds?

a. BaH3

b. CH3

c. NaH3

d. NH3

Hint

2. The boiling points of the halogens increase in the order F2 < Cl2 < Br 2 < I2 due to an

increase in ____________.

a. permanent dipoles

b. hydrogen bonding

c. ionic interactions

d. dispersion forces

Hint

3. In a polar molecule, which atom will have the greatest partial negative charge?

a. the largest atom

b. the smallest atom

c. the most electronegative atom

d. the least electronegative atom

Hint

4. Which molecule will not undergo hydrogen bonding?

a. HF

b. H2O










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c. NH3

d. CH4

Hint

Section 13.3 Liquids and Solids

Practice Test

1. Which of the following materials has the highest density?

a. water vapor

b. steam

c. liquid water

d. iceHint

2. Which of the following is the general definition of a liquid?

a. Matter that has a fixed volume and a fixed shape.

b. Matter that has a fixed volume but no fixed shape.

c. Matter that has neither a fixed volume nor a fixed shape.

d. Matter that has a fixed shape but no fixed volume.

Hint

3. Which of the following is the general definition of a solid?

a. Matter that has a fixed volume and a fixed shape.

b. Matter that has a fixed volume but no fixed shape.

c. Matter that has neither a fixed volume nor a fixed shape.

d. Matter that has a fixed shape but no fixed volume.

Hint

4. Which of the following materials will have the greatest viscosity?a. air

b. cooking oil

c. water

d. vinegar













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Hint

5. The curved shape of water in a glass pipette is an example of _________.

a. adhesion

b. cohesionc. viscosity

d. compression

Hint

6. A solid whose individual particles are arranged in an orderly, geometric, three-

dimensional structure is a ___________.

a. crystalline solid

b. amorphous solid

c. unit cell

d. liquid crystal

Hint

7. Which of the following is a covalent network solid?

a. ice

b. table salt

c. table sugar

d. quartz

Hint

8. Which of the following solids is a metallic solid?

a. graphite

b. sulfur

c. nickel

d. iodine

Hint

9. Which of the following is an ionic solid?

a. graphite

b. nickel

c. ammonium chloride
















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d. dry ice (solid carbon dioxide)

Hint

10. Which of the following is a molecular solid?

a. potassium sulfate b. lead

c. dry ice (solid carbon dioxide)

d. ammonium chloride

Hint

11. Glass is a(n) _____________.

a. liquid

b. crystalline solid

c. metallic solid

d. amorphous solid.

Hint

Section 13.4 Phase Changes

Practice Test

1.

According to the phase diagram for CO2, what is the critical point for carbon dioxide?

a. -100










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b. -78

c. -45

d. 31

Hint

2. The phase change that describes the direct conversion of a solid into a gas is _______.

a. vaporization

b. sublimation

c. melting

d. deposition

Hint

3. Which of the following phase changes will release energy during the transition?a. vaporization

b. condensation

c. sublimation

d. melting

Hint

4. The point on a phase diagram where the solid state, the liquid state, and the gas vapor

state can coexist is _________.

a. the pressure point

b. the absolute zero point

c. the critical point

d. the triple point

Hint

Section 14.1 The Gas Laws

Practice Test

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 isconstant.)













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a. 0.143 atm

b. 0.429 atm

c. 1.38 atm

d. 1.71 atm

Hint

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 LHint

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

Hint

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

Hint

5. Particles of matter that are in constant, random motion and that have a size that is muchsmaller than the distance between them are _____________.

a. solids

b. liquids

c. gases













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d. solutions

Hint

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


d. Gay-Lussac’s law

Hint

7. How can the relationship between a gas at two sets of conditions be expressed

mathematically by Boyle’s law?

a. P 1V 1 = P 2V 2

b. P 1/V 1 = P 2/V 2

c. V 1/T 1 = V 2/T 2

d. V 1T 1 = V 2T 2

Hint

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


d. Gay-Lussac’s law

Hint

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

Hint

10. A gas occupies a volume of 1.0 L at 25°C. What volume will the gas occupy at 100°C?
















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a. 1.0 L

b. 1.3 L

c. 0.80 L

d. 4.0 L

Hint

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

Hint

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

Hint

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

Hint

Section 14.2 The Combined Gas Law and Avogadro's Principle

Practice Test













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1.

<.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

Hint

2. 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

Hint

3. 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

Hint

4. 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










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d. 0.0446 L

Hint

5. 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

Hint

Section 14.3 The Ideal Gas Law

Practice Test

1. What is the name given to a gas whose particles do not take up space, do not have

intermolecular attractive forces, and follow the gas laws under all conditions of

temperature and pressure?

a. noble gas

b. diatomic gas

c. ideal gas

d. real gasHint

2. What will happen to a real gas when the temperature is lowered and the pressure is

raised?

a. The gas eventually condenses to become a liquid.

b. The gas does not change.

c. The gas expands.

d. The number of gas particles doubles.

Hint

3. If the pressure exerted by a gas at 100°C in a volume of 0.044 L is 3.81 atm, how many

moles of gas are present?

a. 5.5 x 10-

mol

b. 1.8 x 10 mol













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c. 1.0 mol

d. 100 mol

Hint

4. How many grams of ammonia (NH3) are present in a sample that occupies 2.0 L at a pressure at 2.0 atm and 25°C?

a. 0.16 grams

b. 2.7 grams

c. 6.3 grams

d. 0.36 grams

Hint

5. Determine the Celsius temperature of 1.50 moles of ammonia contained in a 10.0-L vessel

under a pressure of 2.0 atm.

a. -111°C

b. 162°C

c. -50°C

d. 0.0°C

Hint

6. What is the molar mass of a gas if 0.104 g of the gas occupies 48.7 mL at STP?

a. 28.2 g/mol b. 34.5 g/mol

c. 40.0 g/mol

d. 47.9 g/mol

Hint

7. How many moles of an ideal gas are contained in 8.21 L at 73°C and 380 torr?

a. 0.25

b. 1.5 x 10c. 0.14

d. 7.5 x 10

Hint
















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Section 14.4 Gas Stoichiometry

Practice Test

1. Calculate the number of moles of gas contained in a 16.0-L vessel at 373 K with a

pressure of 2.50 atm.a. 2.8 x 10

-mol

b. 1.31 x 10-

mol

c. 0.766 mol

d. 1.31 mol

Hint

2. Determine the Celsius temperature of 1.50 moles of ammonia contained in a 10.0-L vessel

under a pressure of 2.0 atm.

a. -111

b. 162°C

c. -50°C

d. 0.0°C

Hint

3. A gaseous compound is 30.6% nitrogen and 69.4% oxygen by mass. A 5.25-g sample of

the gas occupies a volume of 1.00 L and exerts a pressure of 1.26 atm at – 4.0°C. What is

the molecular formula for the gas?a. NO

b. NO2

c. N3O6

d. N2O4

Hint

4. A 10.0-L vessel contains gas A at a pressure of 300.0 torr. A 3.00-L vessel contains gas B

at a pressure of 400.0 torr. Gas A is forced into the second vessel. Calculate the resulting

pressure in torr. Assume the temperature remains constant.a. 1000 torr

b. 1400 torr

c. 1800 torr

d. 2000 torr










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Section 15.1 What are Solutions?

Practice Test

1.

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 in1.0 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

Hint

2. In a solution, the substance that does the dissolving is ____________.a. the solvent

b. the solute

c. saturated

d. miscible

Hint

3. Solutions can be mixtures of _______________.

a. solids

b. liquids

c. gases

d. all of the above

Hint










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

a. water

b. sugar

c. glass

d. none of the above

Hint

5. The process of a solvent and a solute completely mixing and forming a solution is ________.

a. solvation

b. salvation

c. crystallization

d. ionizationHint

6. 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

Hint

7. 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

Hint

8. What is a common means of identifying a supersaturated solution?

a. precipitation

b. dissolution

c. solvation

d. hydration













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Hint

9. 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

Hint

Section 15.2 Solution Concentration

Practice Test

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%

Hint

2.

How much solvent is needed to make 200 ml of 50% rubbing alcohol?










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a. 50 mL

b. 100 mL

c. 150 mL

d. 200 mL

Hint

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

Hint

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

Hint










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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

Hint

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

Hint

7. What volume of 12.6 M HCl must be added to sufficient water to prepare 5.00 liters of

3.00 M HCl?

a. 1.19 L

b. 21.0 L

c. 0.840 L

d. 7.56 L

Hint

8. What mass of Ca(OH)2 is contained in 1500 mL of 0.0250 M Ca(OH)2 solution?

a. 3.17 g

b. 2.78 g

c. 1.85 g

d. 2.34 g

Hint

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













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c. 3.2m

d. 6.3m

Hint

10. Calculate the molality of 10% H3PO4 solution in water.a. 0.380m

b. 0.760m

c. 1.13m

d. 1.51m

Hint

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

Hint

12. What is the molarity of 2500 mL of a solution that contains 160 grams of ammonium

nitrate (NH4 NO3)?

a. 0.333 M b. 0.450 M

c. 0.600 M

d. 0.800 M

Hint

Section 15.3 Colligative Properties of Solutions

Practice Test

1. What is the name for a substance that dissolves in water but does not form ions or conduct

an electric current?

a. nonelectrolyte













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b. electrolyte

c. insoluble

d. saturated

Hint

2. Which of the following is not a colligative property?

a. freezing point depression

b. boiling point elevation

c. vapor pressure lowering

d. increasing solubility

Hint

3. If 4.27 g sucrose (C12H22O11) are dissolved in 15.2 g water, what is the boiling point of the resulting solution? K b for water = 0.512°C/m.

a. 101.64°C

b. 100.42°C

c. 99.626°C

d. 100.73°C

Hint

4. Calculate the freezing point of a solution that contains 8.0 g of sucrose (C12H22O11) in 100

g of H2O. K f for H2O = 1.86°C/m a. -0.044°C

b. -0.22°C

c. -0.39°C

d. -0.44°C

Hint

5. A 4.305-g sample of a nonelectrolyte is dissolved in 105 g of water. The solution freezes

at -1.23°C. Calculate the molar mass of the solute. K f for water = 1.86°C/m.

a. 39.7 g/mol

b. 58.4 g/mol

c. 46.2 g/mol

d. 62.0 g/mol

Hint
















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Section 15.4 Heterogeneous Mixtures

Practice Test

1. A suspension is _________.

a. a heterogeneous mixture

b. a homogeneous mixture

c. a solution

d. impossible to prepare

Hint

2. A colloid contains particles that ____________________.

a. are smaller than atoms

b. are between 1 nm and 1000 nm in diameter

c. settle out if left undisturbed

d. are atomic-size in scale

Hint

3. The Tyndall effect describes ______________.

a. precipitation of colloidal particles using electrically charged plates

b. the adsorption of positive ions onto the surface of a hydrophilic solid

c. hydrophobic interactions between nonpolar molecules

d. the scattering of light by colloidal particles

Hint

4. Which one of the following is an example of an emulsion?

a. shaving cream

b. fog

c. mayonnaise

d. styrofoam

Section 16.1 Energy

Practice Test










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1. One calorie is equal to 4.184J. How much energy in joules is supplied by a breakfast bar containing

nutritional calories?

a. 170 J

b. 711 J

c. 1.7 x 10 J

d. 7.11 x 10 J

Hint

2. The temperature of a 25.0 g sample of aluminum changed from 25°C to37°C when heated. How mu

energy was absorbed by the aluminum?

Specific Heats of Common Substances at 298 K (25°C)

Specific heatSubstance J/(g·°C)

Water(l)

(liquid)

4.184

Water(s)

(ice)

2.03

Water(g)

(steam)

2.01

Ethanol(l)

(grain alcohol)2.44

Aluminum(s) 0.897

Granite(s) 0.803

Iron(s) 0.449

Lead(s) 0.129

Silver(s) 0.235

Gold(s) 0.129

a. 10.8 J

b. 22.4 J




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c. 269 J

d. 334 J

Hint

3. The temperature of a 2.5 kg sample of silver was heated from 25°C to 45°C. How much energy waabsorbed by the silver?


Specific heat

Substance J/(g·°C)

Water(l)

(liquid)

4.184

Water(s)

(ice)

2.03

Water(g)

(steam)

2.01

Ethanol(l)

(grain alcohol)2.44

Aluminum(s) 0.897

Granite(s) 0.803

Iron(s) 0.449

Lead(s) 0.129

Silver(s) 0.235

Gold(s) 0.129

a. 12 J

b. 118 J

c. 1.2 x 10 J

d. 2.6 x 10 J

Hint

4. A body in motion possesses kinetic energy because of its __________.







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a. position

b. composition

c. motion

d. inertia

Hint

5. If 5.0 grams of copper cools from 35.0°C to 22.6°C and loses 23.6 joules of heat, what is the specif

heat of copper?

a. 0.076 J/(g·°C)

b. 3.8 ´ 102 J/(g·°C)

c. 0.38 J/(g·°C)

d. 0.62 J/(g·°C)

Hint

6. The specific heat of aluminum is 0.900 J/(g·°C). How many joules of heat are absorbed by 30.0 g o

aluminum if it is heated from 20.0°C to 40.0°C?

a. 540 J

b. 270 J

c. 812 J

d. 1.14 x 104 J

Hint

7. The same amount of heat is added to a 10-g sample of each of the following metals. If each metal i

initially at 20.0°C, which metal will reach the highest temperature?

a. beryllium 1.82 J/(g·°C)

b. calcium 0.653 J/(g·°C)

c. copper 0.385 J/(g·°C)

d. gold 0.129 J/(g·°C)

Hint

8. What is the final temperature of a mixture of chromium metal and water when 50.0 grams of chrom

at 15°C (specific heat = 0.448 J/(g·°C)) is added to 25 mL of water (specific heat = 4.18 J/(g·°C)) a45°C? The density of water is 1 g/mL.

a. 25°C

b. 30°C













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c. 35°C

d. 40°C

Hint

Reset

Section 16.2 Heat in Chemical Reactions and Processes

Practice Test

1. If 500 g of water at 100°C loses 27,000 J of heat, what is the final temperature of the

water?


Specific heat Substance J/(g·°C)

Water(l)

(liquid)

4.184

Water(s)

(ice)

2.03

Water(g)

(steam)2.01

Ethanol(l)

(grain alcohol)

2.44

Aluminum(s) 0.897



http://www.mcgraw-hill.com/

http://www.glencoe.com/






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Granite(s) 0.803

Iron(s) 0.449

Lead(s) 0.129

Silver(s) 0.235

Gold(s) 0.129

a. 6.5°C

b. 12.9°C

c. 87.1°C

d. 93.5°C

Hint

2. Chemical reactions and physical changes that absorb energy from their surroundings are __________.

a. endothermic

b. exothermic

c. isothermic

d. mesothermic

Hint

3. In the chemical reaction between aqueous solutions of hydrochloric acid (HCl) and

sodium hydroxide (NaOH), what is the chemical system?

a. HCl and water

b. NaOH and water

c. HCl + NaOH → NaCl + H2O

d. aqueous HCl and NaOH

Hint

4. The heat content of a system at constant pressure is defined as the ___________.

a. enthalpy

b. entropy

c. work

d. heat

Hint













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5. For the reaction below, the enthalpy change is +624.7 kJ. How would you classify this

reaction?

SiO2(g) + 3C(s) → SiC(s) + 2CO(g)

a. endothermic reaction, heat is lost from the system

b. endothermic reaction, heat is gained by the system

c. exothermic reaction, heat is lost from the system

d. exothermic reaction, heat is gained by the system

Hint

6. How much heat is absorbed in the complete reaction of 3.00 grams of SiO2 with excess

carbon in the reaction below? ΔH° for the reaction is +624.7 kJ.

SiO2(g) + 3C(s) → SiC(s) + 2CO(g)

a. 366 kJ

b. 1.13 ´ 105 kJ

c. 5.06 kJ

d. 31.2 kJ

Hint

7. How much heat energy is liberated when 11.0 grams of manganese is used in the

formation of Mn2O3?

ΔHf ° for Mn2O3 is -962.3 kJ/mol

a. 96.3 kJ

b. 192 kJ

c. 289 kJ

d. 460 kJ

Hint

8. Which of the following substances has an enthalpy value of 0 at 298 K and 1 atmosphere

of pressure?a. HCl(aq)

b. Na(s)

c. NaOH(s)

d. CO2(g)

Hint













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Section 16.3 Thermochemical Equations

Practice Test

1. How much heat is evolved when 27.0 g of glucose is burned according to this equation?

C6H12O6 + 6O2 → 6CO2 + 6H2O; Δ H comb. = -2808kJ


Substance Specific heat

J/(g·°C)

Water(l)

(liquid)

4.184

Water(s)

(ice)2.03

Water(g)

(steam)

2.01

Ethanol(l)

(grain alcohol)

2.44

Aluminum(s) 0.897

Granite(s) 0.803

Iron(s) 0.449

Lead(s) 0.129

Silver(s) 0.235

Gold(s) 0.129

a. 136 kJ

b. 280 kJ

c. 421 kJ

d. 421 J

Hint




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2. How much heat is required to melt 200 grams of ice? The heat of fusion is 333 J/g.

a. 66.6 kJ

b. 666 J

c. c. 1.66 J

d. d. 6.66 kJ

Hint

3. Calculate the amount of heat absorbed by 10.0 grams of ice at -15.0°C as it is convertedto liquid water at 50.0°C. The specific heat of H2O(s) = 2.09 J/(g·°C). The specific heat of

H2O(l) = 4.18 J/(g·°C). The heat of fusion is 333 J/g.

a. 5.73 kJ

b. 0.676 kJ

c. 0.170 kJ

d. 2.83 kJ

Hint

4. From the following data,

H2 (g) + Cl2 (g) → 2HCl(g) Δ H °v = -185 kJ

2H2 (g) + O2 (g) → 2H2O(g) Δ H ° = -483.7 kJ

calculate Δ H ° for the following reaction.

4HCl(g) + O2 (g) → 2Cl2 (g) + 2H2O(g)

a. 299 kJ

b. -114 kJ

c. -299 kJ

d. 114 kJ

Hint

5. Given the following information,

SO3 (g) + H2O(l) → H2SO4 (l) Δ H ° = -133 kJ

Pb(s) + PbO2 (s) + 2H2SO4(l) → 2PbSO4(s) + 2H2O(l) Δ H ° = -509 kJ

calculate the Δ H ° for the reaction below.










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Pb(s) + PbO2 (s) + 2SO3 (g) → 2PbSO4 (s)

a. 376 kJ

b. b. -376 kJ

c. -642 kJ

d. -775 kJ

Hint

Section 16.4 Calculating Enthalpy Changes

Practice Test

1. Calculate ΔH° for the following reaction.

Fe3O4(s) + CO(g) → 3FeO(s) + CO2(g) ΔHf °(kJ/mol) -1118 -110.5 -272 -393.5

a. -263 kJ

b. 54 kJ

c. 19 kJ

d. neg. 50 kJ

Hint

2. Calculate the enthalpy for the reaction:

2ZnS(s) + 3O2(g) → 2ZnO (s) + 2SO2(g) ΔHf ° (kJ/mol) -206.0 0 -348.3 – 296.8

a. -270.6 kJ

b. -878.2 kJ

c. +270.6 kJ

d. +878.2 kJ

Section 16.5 Reaction Spontaneity

Practice Test

1. Which of the following processes occurs with a decrease in entropy?

a. freezing of water

b. boiling water







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c. sublimation of dry ice

d. dissolving salt in water

Hint

2. At one atmosphere pressure and the melting point of a substance, which statement is trueabout this general equation? solid 3 liquid

a. ΔH = 0 for the process

b. ΔS = 0 for the process

c. ΔG = 0 for the process

d. Both ΔH and ΔS = 0 for the process

Hint

3. Calculate the ΔG° for the reaction:

2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(l)ΔG° (kJ/mol)209.2, 0, -394.4, -237.2

a. -1409 kJ

b. -2599 kJ

c. -1643 kJ

d. -2470 kJ

Hint

4. A process can never be spontaneous when it is _________.a. exothermic, and there is an increase in disorder

b. endothermic, and there is an increase in disorder

c. exothermic, and there is a decrease in disorder

d. endothermic, and there is a decrease in disorder

Hint

Section 17.1 A Model for Reaction Rates

Practice Test

1. Use the data in the table to determine how long it will take to completely consume 1L of

1M butyl chloride [C4H9Cl], assuming the same reaction rate.













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Molar Concentration of [C4H9Cl]

[C4H9Cl]

at t= 0.00 s

[C4H9Cl]

at t= 4.00 s

0.220M 0.100M

a. 0.0300 s

b. 0.300 s

c. 3.00 s

d. 33.3 s

Hint

2. Use the data in the table to calculate the average reaction rate expressed in moles H2

consumed per liter per second.

Time (s) [H2] (M) [Cl2] (M) [HCl] (M)

0 0.03 0.05 0.55

10 0.005 0.025 0.05

a. 0.002 mol/(L s)

b. 0.0025 mol/(L s)

c. 0.003 mol/(L s)

d. 0.030 mol/(L s)

Hint

3. Use the data in the table to calculate the average reaction rate expressed in moles Cl2

consumed per liter per second.

Time (s) [H2] (M) [Cl2] (M) [HCl] (M)

0 0.03 0.05 0.55

10 0.005 0.025 0.05a. 0.002 mol/(L s)

b. 0.0025 mol/(L s)

c. 0.003 mol/(L s)

d. 0.030 mol/(L s)







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Hint

4. Use the data in the table to calculate the average reaction rate expressed in moles HCl

produced per liter per second.

Time (s) [H2] (M) [Cl2] (M) [HCl] (M)

0 0.03 0.05 0.55

10 0.005 0.025 0.05

a. 0.002 mol/(L s)

b. 0.0025 mol/(L s)

c. 0.005 mol/(L s)

d. 0.0500 mol/(L s)

Hint

5. What Greek symbol is used to represent change?

a. →

b. Δ

c. ∈

d. Σ

Hint

6. The term that describes the change in the concentration of a reactant divided by the

change in time of the reaction is the ____________.

a. rate constant

b. average reaction rate

c. activated complex

d. chemical reaction

Hint

7. What is the term given to a temporary arrangement of atoms that may produce either products or reactants?

a. transition state

b. collision

c. reactant

d. activation energy













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Hint

8. What is the term used to describe the minimum energy needed to form an activated

complex?

a. temperature

b. kinetic energy

c. activation energy

d. potential energy

Hint

9. How much time has passed when 5 mol/L have reacted at an average rate of 0.5 mol/L·s?

a. 1.0 s

b. 10 s

c. 100 s

d. 0.1 s

Hint

10. Which of the following is the theory that states that atoms, molecules, and ions must

collide for a chemical reaction to occur?

a. collision theory

b. accident theory

c. reaction theoryd. mechanism

Hint

Section 17.2 Factors Affecting Reaction Rates

Practice Test

1. When the concentration of a reactant is increased, the _______________.

a. reaction speeds up

b. reaction slows down

c. reaction stops

d. reaction rate does not change













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Hint

2. What will happen when you decrease the surface area of a reactant?

a. The reaction speeds up.

b. The reaction slows down.c. The reaction stops.

d. The reaction rate does not change.

Hint

3. When you increase the temperature of a chemical reaction, the ____________.

a. reaction speeds up

b. reaction slows down

c. reaction stops

d. reaction rate does not change

Hint

4. A substance that speeds up the reaction rate but is not consumed in the reaction is a(n) ____________.

a. reactant

b. product

c. catalyst

d. inhibitor

Hint

5. What happens to the activation energy of the reaction when a catalyst is present?

a. The activation energy does not change.

b. The activation energy decreases.

c. The activation energy increases.

d. The reaction stops.

Hint

6. What is the name given to a catalyst that exists in the same physical state as the reaction

being catalyzed?

a. homogeneous catalyst

b. heterogeneous catalyst
















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c. solid-state catalyst

d. ephemiral catalyst

Hint

7. Automobiles use a catalytic converter made of metal to ensure complete combustion of gasoline. What type of catalyst is present in automobiles?

a. homogeneous catalyst

b. heterogeneous catalyst

c. solid-state catalyst

d. ephemiral catalyst

Hint

Section 17.3 Reaction Rate Laws

Practice Test

1.

The rate law for the reaction between nitrogen monoxide and hydrogen is Rate =k[NO2]2[H2]. What happens to the rate if the concentration of nitrogen doubles?

a. The rate doubles.

b. The rate triples.

c. The rate quadruples.

d. The rate is reduced by half.

Hint

2. In the rate-law expression, Rate = k [A], what does the symbol k represent?

a. instantaneous rate

b. concentration

c. specific rate constant

d. reaction order

Hint













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3. The exponents in a rate-law expression define the __________, which describes how the

rate is affected by the concentration of the reactant.

a. instantaneous rate

b. concentration

c. specific rate constant

d. reaction order

Hint

4. What is the reaction order for each reactant in the following rate-law expression? Rate =

k[A][B]2[C]

3

a. first order for A, first order for B, first order for C

b. first order for A, second order for B, first order for C

c. first order for A, second order for B, third order for C

d. first order for A, second order for B, second order for C

Hint

5. What is the overall reaction order for the following rate-law expression? Rate =

k[A][B]2[C]

3

a. 4

b. 5

c. 6

d. 7

Hint

6. What is the most common experimental method for evaluating a reaction order called?

a. method of final rates

b. method of instantaneous rates

c. method of initial rates

d. method of experimental rates

Hint

7. How are the rate law and the order for a complex reaction determined?

a. by calculation

b. by experimentation













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c. by trial and error

d. by estimation

Hint

Section 17.4 Instantaneous Reaction Rates and Reaction Mechanisms

Practice Test

1. What quantity is defined by the change in concentration of a component in a chemical

reaction divided by the change in time of the reaction?

a. specific rate constant

b. rate order

c. instantaneous rate

d. catalyst

Hint

2. What is the instantaneous rate for a reaction that is second order in A and first order in B

when [A] = 2 M and [B] = 3 M ? The specific rate constant for the reaction is 0.1 M -2s-1.

a. 12 M /s

b. 1.2 M /s

c. 0.12 M /s

d. 0.012 M /s

Hint

3. What is the slowest of the elementary steps in a complex reaction called?

a. mechanism

b. propagating step

c. rate-limiting step

d. fast stepHint

4. What are the units that generally express a reaction rate?

a. liters per mole per second

b. moles per liter per second













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c. seconds per liter per mole

d. seconds per mole per liter

Hint

5. What is the name given to a chemical produced in one elementary step and used up in asubsequent elementary step?

a. complex step

b. elementary step

c. reaction mechanism

d. intermediate

Hint

6. What are the individual steps in a complex reaction known as?

a. complex steps

b. elementary steps

c. intermediate steps

d. rate-determining steps

Hint

Section 18.1 Equilibrium: A State of Dynamic Balance

Practice Test

1.

What is the equilibrium constant expression for the following reaction?

ClNO2 + NO ‹–› NO

2+ ClNO

a. K eq = [ClNO2][NO]/[NO2][ClNO]

b. K eq = [ClNO2][ClNO]/[NO][NO2]

c. K eq = [NO] [ClNO]/[ClNO] [NO]

d. K eq = [NO2][ClNO]/[ClNO2][NO]

Hint













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2.


H2 + I2 ‹–› 2HI

a. K eq = [HI] /[H2][I2]

b. K eq = [H2][I2]/ [HI]

c. K eq = [HI] /[H] [I]

d. K eq = [2HI]/ /[H2][I2]

Hint

3.


PCl5 ‹–› PCl3 + Cl2 a. K eq = [PCl] /[PCl] [Cl]

b. K eq = [PCl5]/[PCl3][Cl2]

c. K eq = [PCl] [Cl] /[PCl]

d. K eq = [PCl3][Cl2]/ [PCl5]

Hint

4.


2SO3 ‹–› 2SO2 + O2







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a. K eq = [SO2] [O2]/[SO3]

b. K eq = [2SO2][2O]/[2SO3]

c. K eq = 2[SO2] [O2]/2[SO3]

d. K eq = [SO3] /[SO2] [O2]

Hint

5.


N2 + 3H2 ‹–› 2NH3

a. K eq = [N2][H2] /[NH3]

b. K eq = [N2] [H2]/2[NH3]

c. K eq = [NH3] /[N2][H2]

d. K eq = 2[NH3]/ [N2] [H2]

Hint

6. When Δ G is zero, ______________.

a. the reaction is spontaneous

b. the reaction is not spontaneous

c. the reaction is at equilibrium

d. the reaction has no energy

Hint

7. When the system A + B ↔ C + D is at equilibrium, ___________.

a. the forward reaction has stopped

b. the reverse reaction has stoppedc. both the forward and reverse reactions have stopped

d. neither the forward nor the reverse reaction has stopped

Hint

8. A molecular system that is a dynamic process with two opposing reactions balancing each

other is a(n) ___________.













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a. chemical equilibrium

b. spontaneous reaction

c. nonspontaneous reaction

d. unreacted set of chemicals.

Hint

9. What is the equilibrium constant expression for the following reaction?

4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g)

a. K eq = [NH3] [O2]/[NO] [H2O]

b. K eq = [NH3] [O2] /[NO] [H2O]

c. K eq = [NO] /[NH3][O2]

d. K eq = [NO] [H4O] /[NH3] [O2]Hint

10. What is the equilibrium constant expression for the following reaction?

2NH4 NO3(s) → 2NH3(g) + 2NO(g) + H2(g) + 2O2(g)

a. K eq = [NH3][NO][H2][O2]

b. K eq = [NH3] [NO] [H2][O2]

c. K eq = ([NH3][NO][H2][O2])/[NH4 NO3]

d. K eq = ([NH3] [NO] [H2][O2] )/[NH4 NO3]

Hint

11. For the reaction, 2A + B → C + 2D at 35°C, the value of the forward rate constant, K f, is

3.0 × 10-3 M

-1s

-1and the value of the reverse rate constant, K r, is 1.5 × 10

-2 M

-2s

-1.

Calculate the value of K eq for this reaction.

a. 2

b. 0.5

c. 0.2

d. 5

Hint

12. Consider the following reaction in which all reactants and products are gases. 1.00 mole

of A and 2.00 moles of B are placed in a 5.0 — L container. After equilibrium has been

established, 0.50 mole of D is present. What is the equilibrium constant for this reaction:A + 2B → 2C + D?













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a. 1

b. 0.15

c. 0.33

d. 3

Hint

13. When will the least time be required for a reaction to reach equilibrium?

a. K eq is very small

b. K eq is very large

c. K eq is approximately 1.00

d. cannot tell because the time to reach equilibrium does not depend on K eq

Hint

Section 18.2 Factors Affecting Chemical Equilibrium

Practice Test

1. Which of the following states that an equilibrium system that is disturbed will respond in

a manner to restore equilibrium?

a. Le Châtelier’s principle

b. Boyle’s law

c. Charles’s law

d. Dalton’s law

Hint

2. For the system, H2(g) + CO2(g) → H2O(g) + CO(g) at equilibrium, the addition of H2(g)

would cause ____________.

a. only more H2O to form

b. only more CO to formc. more H2O and CO to form

d. only more CO2 to form

Hint

3. For the reaction, 2SO2(g) + O2(g) → 2SO3(g) at equilibrium, the removal of O2 would













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cause _______________.

a. the concentration of SO2 to increase, and the concentration of SO3 to increase

b. the concentration of SO2 to increase, and the concentration of SO3 to decrease

c. the concentration of SO2 to decrease, and the concentration of SO3 to increase

d. the concentration of SO2 to decrease, and the concentration of SO3 to decrease

Hint

4. For the reaction, 2SO2(g) + O2(g) → 2SO3(g) at equilibrium, what will be the effect onthe net amount of SO3 present if the volume of the container is increased?

a. The concentration of SO3 will increase.

b. The concentration of SO3 will decrease.

c. The concentration of SO3 will remain the same.

d. This question cannot be answered without knowing the value of Keq.Hint

5. For the reaction, 2SO2(g) + O2(g) → 2SO3(g) + heat, at equilibrium, what will be theeffect on the net amount of SO3 present if the temperature of the container is increased?

a. The concentration of SO3 increases.

b. The concentration of SO3 decreases.

c. The concentration of SO3 remains the same.

d. This question cannot be answered without knowing the value of Keq.

Hint

6. For the gas phase reaction, SO2(g) + ½O2(g) → SO3(g), ΔH° = -1.6 × 102 kJ for theforward reaction. To increase the yield of SO3, the reaction should be run at ________.

a. high pressure and high temperature

b. high pressure and low temperature

c. low pressure and high temperature

d. low pressure and low temperature

Hint

7. For the reaction, H2(g) + Cl2(g) → 2HCl(g) + heat, what will be the effect on theequilibrium constant, K eq, if the pressure of the vessel is decreased at constant

temperature?

a. K eq increases

b. K eq decreases













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c. K eq does not change

d. The question cannot be answered without knowing the initial value of K eq.

Hint

8. For the reaction, 2NOCl(g) +75 kJ → 2NO(g) + Cl2(g), which of the following will shiftthe reaction to the left?

a. add a catalyst

b. heat the reaction vessel

c. decrease the volume of the container

d. add more NOCl

Hint

9. For the reaction, 2Cl2(g) + 2H2O(g) + heat → 4HCl(g) + O2(g), which of the following

will shift the equilibrium to produce more HCl?

a. add more Cl2

b. lower the temperature of the reaction

c. decrease the volume of the reaction vessel

d. add more O2

Hint

10. Which of the following would force the forward reaction to completion?

CaCO3(s) + 2H3O+(aq) → Ca

2+(aq) + 3H2O(l) + CO2(g)

a. add base to neutralize the H3O+

b. add more Ca+

to the mixture

c. conduct the experiment in an open container

d. add CO2 to the mixture

Hint

Section 18.3 Using Equilibrium Constants

Practice Test

1. Which of the solubility product expressions is incorrect?

a. K sp(Ag2S) = [Ag+] [S

-]

b. K sp (Sb2S3) = [Sb+] [S

-]













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c. K sp (CaF2) = [Ca+][F

-]

d. K sp(CuS) = [Cu+][S

-]

Hint

2. The molar solubility for BaCO3 is 9.0 x 10-

M at 25°C. What is the solubility productconstant, K sp, for BaCO3?

a. 1.2 x 10-

b. 8.1 x 10-

c. 5.3 x 10-

d. 4.0 x 10-

Hint

3. The value of the K sp for SrSO4 is 2.8 x 10-

. What is the molar solubility of SrSO4?

a. 7.6 x 10- M

b. 5.8 x 10- M

c. 5.3 x 10- M

d. 5.7 x 10- M

Hint

4. If NaCl is added to a 0.010M solution of AgNO3 in water at 25°C, what will be [Cl-]

when precipitation of AgCl begins? The Ksp for AgCl is 1.8 x 10-10

.

a. 1.0 x 10-

M b. 1.3 x 10

- M

c. 1.8 x 10- M

d. 1.8 x 10- M

Hint

5. How many grams of MgF2 will dissolve in 150 mL of 0.100 M NaF solution? Ksp for

MgF2 is 6.4 x 10-9

.

a. 6.4 x 10-

g

b. 4.1 x 10-

g

c. 1.0 x 10-

g

d. 6.0 x 10-

g

Hint
















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Section 19.1 Acids and Bases: An Introduction

Practice Test

1.

Identify a conjugate acid-base pair in the reaction NH3 + H2O ‹–› NH4+

+ OH-

a. H2O and OH-

b. NH3 and H2O

c. NH3 and OH-

d. H2O and NH4 Hint

2.

Which of these is a product of the first step in the ionization of phosphoric acid, H3PO4?

a. H3PO4

b. H2PO4

c. HPO4

d. PO4

Hint

3. What will a nonmetallic anhydride produce in aqueous solution?

a. acid

b. base

c. salt

d. solid precipitate







7/16/2019 Section 1


Hint

4. A substance that contains hydrogen and produces H+

ions in aqueous solution is a(n)

________.

a. acid

b. base

c. salt

d. water

Hint

5. According to the BrÆnsted-Lowry theory, a base is a(n) ___________.

a. electron pair acceptor

b. hydrogen ion donor

c. hydrogen ion acceptor

d. electron pair donor

Hint

6. Which one of the following could not be a BrÆnsted-Lowry acid?

a. H2O

b. H3O+

c. NH4+

d. BF3

Hint

7. When HClO4 ionizes in water, ClO4-is the ______________.

a. acid

b. base

c. conjugate acid

d. conjugate base

Hint

8. Which of the following is not a conjugate acid-base pair?

a. H2O and OH-

b. H3O+

and OH-

c. HCl and Cl-
















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d. HNO3 and NO3-

Hint

9. H2SO4 is a ____________.

a. monoprotic acid b. monoprotic base

c. diprotic acid

d. diprotic base

Hint

10. Which of the following is a weak acid?

a. HCl

b. HNO3

c. HClO4

d. HF

Hint

11. Which of the following is a strong base?

a. NH3

b. HCO3-

c. NaOH

d. HCOOH

Hint

12. Which of the following is the strongest base according to the Bronsted-Lowry theory?

a. Cl-

b. NO3-

c. F-

d. I-

Hint

13. Under what condition is the OH-ion concentration in water expected to be zero?

a. in a solution of strong acid

b. in a solution of strong base

c. in a solution of weak base
















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d. never

Hint

14. Which of the following compounds is not a salt?

a. K 2SO4 b. BaCrO4

c. FeCl2

d. Mg(OH)2

Section 19.2 Strengths of Acids and Bases

Practice Test

1.

Compared to strong acids, weak acids produce __________ ions and conduct electricity __________ efficiently.

a. fewer, more

b. fewer, less

c. more, more

d. more, less

Hint







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2.

The acid ionization constant, K a, is __________ for __________ acids.

a. largest, anhydride

b. largest, weak

c. smallest, strong

d. smallest, weak

Hint

3. A 0.001 M solution of HCl is ____________.

a. a weak acid solution

b. a concentrated acid solution

c. a dilute acid solution

d. neutral

Hint

Q82217A287209 6293 188 2 1 0

3

Section 19.3 What is pH?

Practice Test

1.

Why is it possible to derive the ion product constant for water, but not for aqueous acids?

a. At a given temperature, aqueous acids have constant concentrations.







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b. Aqueous acids never reach equilibrium.

c. The concentration of pure water is constant.

d. Water does not ionize completely.

Hint

2. pH is defined as __________.

a. log [H+]

b. -log [H+]

c. log [OH-]

d. -log [OH-]

Hint

3. A solution of pH = 2.1 would be described as __________.a. distinctly basic

b. slightly basic

c. slightly acidic

d. distinctly acidic

Hint

4. When a solution has a pH of 4, what is the pOH of that solution?

a. 10

b. 4

c. 18

d. 7

Hint

5. At 298 K, pure water has a pH of __________.

a. 0

b. 14

c. -14

d. 7

Hint

6. What is the concentration of hydrogen ions in a solution that has a pH of 4.32?
















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a. 4.8 x 10- M

b. 6.2 x 10- M

c. 5.1 x 10- M

d. 8.6 x 10- M

Hint

7. Calculate the pH of 0.075 M KOH.

a. 10.4

b. 11.12

c. 11.46

d. 12.88

Hint

8. Calculate the pH of a solution that has a [OH-] = 2.50 x 10

- M .

a. 0.4

b. 3.6

c. -3.6

d. 10.4

Hint

Section 19.4 Neutralization

Practice Test

1. Neutralization is the chemical process in which ____________.

a. sodium ions react with chloride ions to form sodium chloride

b. hydrogen ions react with chloride ions to form hydrogen chloride

c. sodium ions react with hydroxide ions to form sodium hydroxide

d. hydrogen ions react with hydroxide ions to form water

Hint

2. What is the net ionic equation for the neutralization reaction between HF and KOH?

a. H+

+ OH- → H2O

b. K +

+ F- → KF













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c. H+ + KOH → H2O + K

+

d. HF + OH- → H2O + F

-

Hint

3. Which of the following cannot be a buffer?a. a mixture of a weak acid and a weak base

b. a mixture of a weak acid and a strong base

c. a mixture of a strong acid and a weak base

d. a mixture of a strong acid and a strong base

Hint

4. Which of the following salts will produce a basic solution?

a. NaCl

b. KBr

c. KCN

d. Ca(NO3)2

Hint

5. Calculate the H ion and OH-ion concentrations in a 0.50 M solution of HBr.

a. [H+] = 0.50 M and [OH

-] = 0.50 M

b. [H+] = 0.50 M and [OH

-] = 2.0 M

c. [H+] = 0.50 M and [OH

-] = 2.0 ´ 10

- M

d. [H+] = 1.0 ´ 10

- M and [OH

-] = 1.0 ´ 10

- M

Hint

Q82228A287253 6295 188 2 1 0

5

Section 20.1 Oxidation and Reduction

Practice Test













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1.

What is the oxidation number of Group 17 nonmetals?a. -1

b. 0

c. 1

d. 3

Hint

2. Which of the following describes oxidation?

a. gain of electrons b. gain of positive charge

c. loss of electrons

d. loss of positive charge

Hint

3. When is the oxidation number of oxygen not -2?

a. in water

b. in oxyacids

c. in peroxide

d. in sulfuric acid

Hint

4. When an atom undergoes oxidation, the ___________ changes.










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a. mass

b. identity

c. isotope

d. charge

Hint

5. Oxidation is __________.

a. the loss of electrons and an increase in charge

b. the loss of electrons and a decrease in charge

c. the gain of electrons and an increase in charge

d. the gain of electrons and a decrease in charge

Hint

6. What is the oxidation number of N in KNO3?

a. 6

b. 5

c. 3

d. -3

Hint

7. What is the oxidation number of O in Na2SO4?a. 6

b. 1

c. +2

d. -2

Hint

8. The highest possible oxidation number for carbon is ________.

a. +8

b. 6

c. 4

d. 2

Hint
















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9. Select the compound in which chlorine has the highest possible oxidation number.

a. HCl

b. HClO

c. HClO3

d. HClO4

Hint

10. Which of the following reactions is a redox reaction?

a. H+(aq) + OH

-(aq) → H2O(l)

b. 2KBr(aq) + Pb(NO3)2(aq) → 2KNO3(aq) + PbBr 2(s)

c. CaBr 2(aq) + H2SO4(aq) → CaSO4(s) + 2HBr(aq)

d. 2Al(s) + 3H2SO4(aq) → Al2(SO4)3(aq) + 3H2(g)

Hint

11. In the following reaction, which chemical species is the oxidizing agent?

5H2O2 + 2MnO4-+ 6H

+ → 2Mn

2++ 8H2O + 5O2

a. H2O2

b. MnO4-

c. H+

d. Mn

Hint

12. In a redox reaction, the number of electrons lost by the reducing agent ____________.

a. equals the number of protons in the reducing agent

b. equals the number of protons in the oxidizing agent

c. equals the number of electrons lost by the oxidizing agent

d. equals the number of electrons gained by the oxidizing agent

Hint

13. The most electronegative atom in a compound has a charge that is __________.

a. positive

b. negative

c. neutral

d. zero













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Hint

14. Which of the following types of reactions do not usually involve redox?

a. synthesis

b. decompositionc. single-replacement

d. double-replacement

Hint

15. What element is oxidized in the following reaction?

3Cu + 8HNO3 → Cu(NO3)2 + 2NO + 4H2O

a. Cu

b. H

c. N

d. O

Hint

16. Predict the products of the following single — replacement reaction.

Fe(s) + CuSO4(aq) → ?

a. CuS(s) + Fe2SO4(aq)

b. Fe(s) + Cu(s) + SO4(aq)

c. Cu(s) + FeSO4(aq)

d. FeCuSO4

Hint

17. Group 1A metal ions have an oxidation number of _______.

a. 1

b. 2

c. 3

d. -1

Hint

18. What is the oxidizing agent in the following reaction?

6KOH(aq) + 3Cl2(g) → KClO3(aq) + 5KCl(aq) + 3H2O(l)
















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a. KOH

b. Cl2

c. KClO3

d. KCl

Hint

Section 20.2 Balancing Redox Equations

Practice Test

1. Name the spectator ion in the following equation, including the correct coefficient:

Cu(s) + 4HNO3 (aq) → Cu(NO3)2 (aq) + 2NO2 + 2H2O (l)

a. 2NO2

b. 3NO2

c. 2NO3

d. 3NO3

Hint

2. Determine the net change in oxidation number of chlorine in the following reaction:

ClO4-(aq) + Br

- (aq) → Cl

-(aq) + Br 2(g) (in acid solution)

a. -8

b. -7

c. 7

d. 8

Hint

3. Complete and balance the following redox equation. When properly balanced with

whole — number coefficients, the coefficient of S is _________.

H2S + HNO3 → S + NO + H2O

a. 1

b. 2

c. 3

d. 4










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Hint

4. When the following equation is balanced what is the coefficient of Sn?

Sn + HNO3 → SnO2 + NO2 + H2O

a. 1

b. 2

c. 3

d. 4

Hint

5. Balance the following ionic equation. What is the coefficient of the reducing agent in this

reaction?

Fe2+(aq) + MnO4-(aq) + H+(aq) → Fe3+(aq) + Mn2+(aq) + H2O(l)

a. 5

b. 1

c. 8

d. 4

Hint

Section 20.3 Half-Reactions

Practice Test

1. A balanced reaction that shows only the oxidation process is a ___________.

a. balanced chemical equation

b. single-replacement reaction

c. half-reaction

d. synthesis reaction

Hint

2. Balance the following equation. How many electrons must be transferred between thereducing agent and the oxidizing agent in this reaction?

H2S + HNO3 → S + NO + H20

a. 2













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b. 3

c. 4

d. 6

Hint

3. What is the net change in oxidation number of iodine in the following ionic reaction?

2MnO4-+ I

-+ H2O → 2MnO2 + IO3

-+ 2OH

-

a. 1

b. 2

c. 3

d. 6

Hint

4. Which of the following half-reactions represents oxidation?

a. O2 + 4e- → 2O

-

b. Fe + 3e- → Fe

c. Fe → Fe+

+ 2e-

d. Cu+

+ 2e- → Cu

Hint

5. Complete and balance the following chemical reaction with the smallest set of coefficients. What is the coefficient for iodine in this reaction?

HI + HNO3 → I2 + NO + H2O

a. 1

b. 2

c. 3

d. 4

Hint

6. Balance the following equation with the smallest whole — number coefficients. How many

moles of zinc will react with 6 moles of cobalt(III) chloride?

Zn(s) + CoCl3(aq) → ZnCl2(aq) + Co(s)

a. 2













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b. 3

c. 6

d. 9

Hint

Q82256A287365 6298 188 2 1 0

6

Section 21.1 Voltaic Cells

Practice Test

1.

Which of the following must be true of a salt bridge?

a. Ions can move through the salt bridge.

b. The salt bridge must prevent the flow of charged particles.

c. The solutions joined by the salt bridge can mix freely.

d. The tube of the salt bridge must remain open.

Hint

2. In any electrochemical cell, the cathode is the ________.







7/16/2019 Section 1


a. positive electrode

b. negative electrode

c. electrode at which some species gains electrons

d. electrode at which some species loses electrons

Hint

3. In a voltaic cell, the salt bridge __________.

a. is not necessary for the cell to work

b. acts as a mechanism to allow mechanical mixing of the solutions

c. allows charge balance to be maintained in the cell

d. drives electrons from one half-cell to another

Hint

4. A cell is constructed by immersing a strip of lead in a 1.0 M Pb(NO3)2 solution and a strip

of silver in a 1.0 M AgNO3 solution. The circuit is completed by a wire and a salt bridge.As the cell operates, the strip of silver gains mass (only silver), the strip of lead loses

mass, and the concentration of lead ions increases in the solution around the lead strip.

Which of the following represents the reaction that occurs at the cathode in this cell?

a. Pb+

+ 2e- → Pb

b. Pb → Pb+

+ 2e-

c. Ag+

+ e- → Ag

d. Ag → Ag+ + e-

Hint

5. Using a table of standard reduction potentials, which of the following is the strongest

oxidizing agent?

a. Al+

b. Al

c. F2

d. F-

Hint

6. Which of the following is the weakest reducing agent?

a. Al+

b. Al













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c. F2

d. F-

Hint

7. A voltaic cell consists of a standard hydrogen electrode connected by a salt bridge and awire to an electrode consisting of a strip of Cd metal dipping into a 1 M solution of Cd(NO3)2. When the cell produces a current, the electrons flow through the wire from the

_________ electrode to the ________ electrode. In this cell, the _______ electrode acts as

the cathode.

a. Cd, H2, Cd

b. H2, Cd, Cd

c. H2, Cd, H2

d. Cd, H2, H2

Hint

8. Under standard conditions, what is the standard cell potential for the cell? Cd|Cu+

||

Cu2+

|Cu?

a. +0.74 V

b. -0.74 V

c. 0.06 V

d. -0.06 V

Hint

9. Under standard conditions, which of the following is the net reaction that occurs in thecell? Cd|Cd

2+|| Cu

2+|Cu

a. Cu + Cd+ → Cd + Cu

+

b. Cu + Cd → Cu+

+ Cd+

c. Cu+

+ Cd+ → Cu + Cd

d. Cu+ + Cd → Cu + Cd

+

Hint

10. What is the numerical value for the standard cell potential for the following reaction?2Cr

3+(aq) + 3Cu(s) → 2Cr(s) + 3Cu

2+(aq)

a. -1.08 V

b. -0.40 V

c. 0.40 V













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d. 1.08 V

Hint

11. How many moles of electrons have been transferred between the species in the following

reaction? 2Cr 3+

(aq) + 3Cu(s) → 2Cr(s) + 3Cu2+

(aq)

a. 2

b. 3

c. 4

d. 6

Hint

12. A voltaic cell in which the reaction involves the combustion of one reactant with oxygen

to produce electric energy is a(n) _____________.

a. primary battery

b. secondary battery

c. fuel cell

d. electrolytic cell

Hint

Section 21.2 Types of Batteries

Practice Test

1.

Why are alkaline dry cell batteries smaller in size than zinc-carbon dry cells?

a. Alkaline batteries do not need the carbon rod cathode.

b. Alkaline batteries do not require steel cases.










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c. No zinc is reduced in the alkaline batteries.

d. Zinc paste takes up less space than zinc powder.

Hint

2.

What is the anode reaction when a NiCad battery is used to generate electric current?

a. oxidation of cadmium in the presence of an acid

b. oxidation or cadmium in the presence of a base

c. reduction of cadmium in the presence of an acid

d. reduction of cadmium in the presence of a base

Hint







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3.

How does paint protect steel surfaces from corrosion?

a. Paint acts as a sacrificial anode.

b. Paint corrodes leaving steel intact.

c. Paint galvanizes steel.

d. Paint prevents the oxidation-reduction reaction of corrosion from occurring.

Hint

4. A redox reaction that is not easily reversed is the source of electric energy in a(n)

_______.

a. primary battery


c. fuel cell


Hint

5. A redox reaction that is easily reversed, can produce electric energy, and can be recharged

is a(n) __________.

a. primary battery


c. fuel cell


Hint

6. The pieces of metal that are placed on the outside of ships to help prevent corrosion are

_________.

a. galvanized metals

b. sacrificial cathodes










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c. sacrificial anodes

d. electrolyzed metals

Hint

7. A chemical reaction generates electric energy spontaneously in ___________.a. a galvanic cell

b. corrosion

c. electrolysis

d. spontaneous combustion

Hint

Section 21.3 Electrolysis

Practice Test

1. During the electrolysis of molten sodium iodide, sodium ions move ________.

a. to the anode, which is positively charged

b. to the anode, which is negatively charged

c. to the cathode, which is positively charged

d. to the cathode, which is negatively charged

Hint

2. The use of electric energy to bring about a chemical reaction is ____________.

a. a galvanic cell

b. corrosion

c. electrolysis

d. spontaneous combustion

Hint

3. Why is electrolysis needed to produce pure metals?a. Metals would rather be cations.

b. Metals would rather be anions.

c. Metals are easily reduced.

d. Metals are always neutral.













7/16/2019 Section 1


Hint

4. The Hall-Héroult process, which is used to make aluminum metal, electrolyzes a molten

mixture of aluminum oxide and cryolite (Na3AlF6) to reduce aluminum cations to

aluminum metal. Why is the cryolite used in this mixture?

a. The cryolite lowers the cell potential for the reduction of aluminum cations.

b. The cryolite raises the cell potential for the reduction of aluminum cations.

c. The cryolite increases the conductivity of the mixture.

d. The cryolite lowers the melting point of the mixture and thereby lowers the costof manufacture by lowering the amount of heat energy needed to melt the

mixture.

Hint

5. Which of the following is not obtained from the electrolysis of an aqueous solution of

sodium chloride?a. Na

b. NaOH

c. Cl2

d. H2

Hint

6. In an electrolytic cell, the electrode that acts as a source of electrons to the solution is

called the __________; the chemical change that occurs at this electrode is called

________.

a. anode, oxidation

b. anode, reduction

c. cathode, oxidation

d. cathode, reduction

Hint

7. The electrolysis of an aqueous sodium chloride solution using inert electrodes produces

gaseous chlorine at one electrode. At the other electrode, gaseous hydrogen is produced,and the solution around the electrode becomes basic. Which of the following equations is

the correct equation for the cathode half-reaction in this electrolytic cell?

a. 2Cl- → Cl2 + 2e

-

b. 2H2O + 2e- → H2 + 2OH

-

c. Cl2 + 2e- → 2Cl

-













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d. H2 + 2OH- → 2H2O + 2e

-

Hint

8. Consider the electrolysis of an aqueous solution of aluminum fluoride. Which one of the

following statements describes what will be observed?

a. Al metal is produced at one electrode, and O2 and H are produced at the other.

b. O2 and H are produced at one electrode, and H2 and OH-are produced at the

other.

c. Al metal is produced at one electrode, and F2 is produced at the other.

d. Al metal is produced at one electrode, and O2 and H+

are produced at the other.

Hint

9. Electroplating is an example of __________.

a. a galvanic cell b. a primary battery

c. a fuel cell

d. electrolysis

Hint

Section 22.1 Alkanes

Practice Test

1. The branch of chemistry that is devoted to the study of carbon compounds is _________.

a. inorganic chemistry

b. organic chemistry

c. analytical chemistry


Hint

2. What element is combined with carbon to form the simplest organic compounds, alsoknown as hydrocarbons?

a. hydrogen

b. alcohols

c. nucleic acids

d. lipids













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Hint

3. What is the class of hydrocarbons that only have single bonds?

a. alkynes

b. alkenesc. alkanes

d. alkylhalides

Hint

4. What is the prefix name that indicates when six carbon atoms are present in the alkane

chain structure?

a. pent

b. hex

c. hept

d. oct

Hint

5. What is the name given to a series of compounds that differ from one another by a

repeating unit?

a. homologous series

b. homogeneous series

c. heterogeneous seriesd. hydrocarbon series

Hint

6. The side chain on an alkane that appears to replace a hydrogen on the straight chain is a

_________________.

a. parent chain

b. substituent chain

c. longest chain

d. cross-linked chain

Hint

Q82294A287517 6302 188 2 1 0

6
















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Section 22.2 Cyclic Alkanes and Alkane Properties

Practice Test

1. The smallest cycloalkane has __________ carbon atoms.

a. one

b. two

c. three

d. six

Hint

2. What is the name given to a hydrocarbon that contains a ring?

a. straight-chain hydrocarbon

b. branched-chain hydrocarbon

c. cyclic hydrocarbon

d. unsaturated hydrocarbon

Hint

3. What type of organic compound has at least one double or triple bond?

a. alkane

b. saturated hydrocarbon

c. cyclic hydrocarbon

d. unsaturated hydrocarbon

Hint

4. What is the principle use of alkanes?

a. as polar solvents

b. as fuels

c. as antiseptics

d. as flavoringsHint

Section 22.3 Alkenes and Alkynes

Practice Test













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1. The structural formula -C=C- indicates a ___________ bond between the carbon atoms.

a. single covalent

b. single ionic

c. double covalent

d. double ionic

Hint

2. What is the identifying structural feature of alkenes?

a. single bonds

b. attachment of a halogen

c. double bond

d. triple bondHint

3. What organic compound is characterized by the presence of a triple bond?

a. alkene

b. alkyne

c. alkane

d. alkyl halide

Hint

Section 22.4 Isomers

Practice Test

1.

How do structural isomers differ from stereoisomers?

a. Structural isomers have different formulas; stereoisomers have the same

formula.










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b. Stereoisomers have different chemical and physical properties; structural

isomers have the same chemical and physical propeties.

c. Structural isomers' atoms are bonded in the same order; stereoisomers' atoms

are bonded in different orders.

d. Structural isomers' atoms are bonded in different orders; stereoisomers' atoms

are bonded in the same order.

Hint

2. Compounds that are composed of the same number and type of atoms but have themarranged in different ways are ___________.

a. isomers

b. isotopes

c. polymers

d. alkanes

Hint

3. What physical structure must be present to have geometric isomers, also known as cis — and trans — isomers?

a. double bond

b. single bond

c. triple bond

d. cyclic structure

Hint

4. What property can be assigned to a carbon atom that has four different atoms or structuralgroups attached to it?

a. isomerity

b. chirality

c. chivalry

d. geometry

Hint

5. What term is used to describe isomers that have all the atoms bonded in the same order but arranged differently in space?

a. structural isomers

b. nonisomers

c. polymers













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d. stereoisomers

Hint

6. Chiral carbons will produce _____________.

a. structural isomers b. optical isomers

c. polymers

d. monomers

Section 22.5 Aromatic Hydrocarbons and Petroleum

Practice Test

1.

Hydrocarbons with __________ carbon atoms will be collected at the highest point on afractionating tower.

a. 4

b. 8

c. 12




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d. 20

Hint

2. What kind of organic compounds contain benzene rings as part of their structure?

a. aliphatic compounds b. alkalinic compounds

c. aromatic compounds

d. adiabatic compounds

Hint

3. The process used to separate fractions from petroleum is ___________.

a. fractional distillation

b. isomerization

c. chirality

d. drilling

Hint

4. Organic compounds that are composed of carbon chains are referred to by a special name.

What kind of compounds are these?

a. aliphatic compounds

b. aromatic compounds

c. allylic compounds

d. alkalinic compounds

Hint

5. Substances that cause cancer are ______________.

a. substituted

b. unsaturated

c. carcinogenic

d. aliphatic

Hint

6. What is the main source of hydrocarbons?

a. petroleum

b. grains
















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c. waste products

d. combustion

Hint

7. In fractional distillation the hydrocarbons composed of fewer carbon atoms rise up thecolam in the form of a _______________.

a. solid

b. liquid

c. distillate

d. gas

Hint

Section 23.1 Functional Groups

Practice Test

1.

What is the difference between an alkyl halide and and aryl halide?

a. An alkyl halide contains a halogen atom bonded to an aromatic group; an arylhalide contains a halogen bonded to an aliphatic carbon atom.

b. An alkyl halide contains a halogen atom bonded to an aliphatic carbon atom; anaryl halide contains a halogen bonded to an aromatic group.

c. Alkyl halides contain double bonds; aryl halides do not.

d. Alkyl halides contain functinal groups; aryl halides do not.

Hint

2.

What is a possible product of a reaction between chloromethane and sodium hydroxidesolution?

a. ethanol










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b. ethyl amine

c. methanal

d. methanol

Hint

3. An organic compound that contains a halogen is known as a(n) ______________.

a. hydrocarbon

b. halocarbon

c. oxycarbon

d. alkalicarbon

Hint

4. In a molecule of 1-

fluoropropane, there is one fluorine atom attached to carbon number 1in the chain. What is fluorine in this molecule?

a. a dimeric element

b. a metal

c. a functional group

d. an unsaturation

Hint

5. In a substituted alkane that has a bromine atom and an iodine atom as functional groups,

which functional group is named first?a. Bromine is named before iodine.

b. Iodine is named before bromine.

c. Bromine is the only functional group named.

d. Iodine is the only functional group named.

Hint

6. When chlorine is added to an alkane, what name is used to identify that chlorine is presentin the molecule?

a. chlorine

b. chloride

c. chloro

d. Chlorine is never named in organic molecules.

Hint
















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7. An organic compound that contains a halogen bonded to a benzene ring is known as an

_________.

a. alkyl halide

b. aryl halide

c. alcohol

d. alkane

Hint

8. In general, what can you conclude about the boiling points and densities of alkyl halides

when compared to the corresponding unsubstituted alkanes?

a. The boiling point of the alkyl halide is lower, and the density of the alkyl halide

is lower.

b. The boiling point of the alkyl halide is lower, and the density of the alkyl halideis higher.

c. The boiling point of the alkyl halide is higher, and the density of the alkyl

halide is lower.

d. The boiling point of the alkyl halide is higher, and the density of the alkylhalide is higher.

Hint

9. The process in which a halogen atom replaces a hydrogen atom of an alkane moleculethrough substitution is known as ___________.

a. halogenation

b. hydrogenation

c. hydroxylation

d. hydrocarbonation

Hint

10. A reaction of an alkyl halide with a hydroxyl group will produce an ____________.

a. alkane

b. alcohol

c. alkene

d. amine

Hint

11. A reaction of an alkyl halide with ammonia will produce an ____________.

a. alkane













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b. alcohol

c. alkene

d. amine

Hint

Section 23.2 Alcohols, Ethers, and Amines

Practice Test

1.

The general formula for an amine is __________.

a. RNH2

b. ROR'

c. ROH

d. RX

Hint

2. The functional group for an alcohol is ________________.

a. -OH

b. -Cl

c. -NH2

d. -COOH

Hint










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3. What is the name of the organic compound that is formed when a hydroxyl group is added

to ethane?

a. ethane

b. ethanol

c. ethyl ether

d. ethene

Hint

4. When ethanol is made unfit to drink by adding a noxious chemical, the ethanol is said to

be _________________.

a. dehydrated

b. denatured

c. dehydrogenated

d. dehalogenated

Hint

5. What is the name of the two-carbon, substituted alkane containing two hydroxyl groups?

a. ethanol

b. ethanediol

c. ethanetriol

d. cycloethanol

Hint

6. An organic compound that contains an oxygen atom bonded to two carbon atoms is a(n)

______________.

a. alkane

b. alcohol

c. ether

d. carboxylic acid

Hint

7. The organic compound that has a nitrogen atom bonded to a carbon atom in an aliphatic

chain or an aromatic ring is known as a(n) _______________.

a. amine

b. alkane













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c. alcohol

d. carboxylic acid

Hint

Section 23.3 Carbonyl Compounds

Practice Test

1.

What is the difference between an aldehyde and a ketone?a. Aldehydes have carbonyl groups; ketones have carboxyl groups.

b. Aldehydes have carbonyl groups at the end of a carbon chain; ketones have

carbonyl groups within carbon chains.

c. Aldehydes have carboxyl groups; ketones have carbonyl groups.

d. Aldehydes have carboxyl groups at the end of a carbon chain; ketones havecarboxyl groups within carbon chains.

Hint

2. The carbonyl group is a(n) _______________.a. carbon single bonded to an oxygen

b. carbon double bonded to an oxygen

c. carbon triple bonded to an oxygen

d. oxygen single bonded to another oxygen

Hint

3. An organic compound that has a carbonyl group bonded to a hydroxyl group is known as

a(n) __________________.

a. alkane

b. alcohol

c. aldehyde

d. carboxylic acid

Hint













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4. An organic compound that contains a carboxyl group with the hydrogen replaced with an

alkyl group is known as a(n) _______________.

a. ether

b. ester

c. aldehyde

d. carboxylic acid

Hint

5. What is the name of the process in which a carboxylic acid reacts with an alcohol

producing an ester and releasing water?

a. hydration

b. halogenation

c. hydroxylation

d. condensation

Hint

Section 23.4 Other Reactions of Organic Compounds

Practice Test

1. The formation of an alkene from an alkane is known as a(n) _____________ reaction.

a. addition b. elimination

c. condensation

d. halogenation

Hint

2. The formation of an alkane from an alkene is known as a(n) ____________ reaction.

a. addition

b. elimination

c. condensation

d. halogenation

Hint

3. The oxidation of 2 — propanol will produce ______________.













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a. 2 — propanol

b. 2 — propanal

c. 2 — propanone

d. 2 — propanoic acid

Hint

4. The formation of polyethylene from ethene is an example of a(n) ____________ polymerization reaction.

a. addition

b. elimination

c. condensation

d. halogenation

Hint

5. The formation of a polyester from a dicarboxylic acid and a diol is a example of a(n) ____________ polymerization reaction.

a. addition

b. elimination

c. condensation

d. halogenation

Hint

Section 23.5 Polymers

Practice Test










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1.

What is the product of the addition polymerization of ethene?a. PVC

b. polyethylene

c. nylon

d. Kevlar

Hint

Section 24.1 Proteins

Practice Test

1.

A chain of 15 amino acids joined by peptide bonds is called a(n) __________.

a. enzyme

b. peptide

c. polypeptide

d. protein




7/16/2019 Section 1


Hint

2. What is the name given to a series of amino acids that are linked by peptide bonds?

a. protein

b. lipidc. carbohydrate

d. ester

Hint

3. What two functional groups are contained in all amino acids?

a. carbonyl and halogen

b. carboxyl and amino

c. amido and ether

d. amino and amido

Hint

4. What type of chemical functional group is made when a peptide bond is formed?

a. amine

b. ester

c. amide

d. carbonyl

Hint

5. A polypeptide of 50 or more amino acids is called a protein. What kind of molecule is the polypeptide?

a. isomer

b. monomer

c. stereomer

d. polymer

Hint

6. How many different amino acids are commonly found in living things?

a. 20

b. 10

c. 15
















7/16/2019 Section 1


d. 30

Hint

7. Proteins can form two three-dimensional structures. Changes in temperature, pH, and

other factors will disrupt these structures. By what name do we know this disruption?

a. denaturation

b. helix

c. sheet

d. polypeptide

Hint

8. Most enzymes are ____________.

a. lipids

b. carbohydrates

c. proteins

d. denatured

Hint

9. Enzymes act as __________ in your body.

a. catalysts

b. carbohydrates

c. lipids

d. substrates

Hint

10. To function properly, the reactive site of an enzyme must have the same shape as the

substrate to which it binds. This fit is commonly known as _____________.

a. tongue and groove

b. lock and key

c. projection and enclosured. heme and globin

Hint

11. What are the four primary functions and uses of proteins in living cells?

a. energy storage, digestion, denaturation, transport
















7/16/2019 Section 1


b. recycling, digestion, polymerization, transport

c. catalysis, structure, transport, carrying signals

d. denaturation, structure, transport, carrying signals

Hint

Section 24.2 Carbohydrates

Practice Test

1.

By what type of reaction are monosaccharides linked together to form disaccharides?

a. addition

b. combustion

c. condensation

d. dehydration

Hint

2. What functional groups characterize carbohydrates?

a. multiple carboxyl and one hydroxyl

b. multiple hydroxyl and one carbonyl

c. multiple carbonyl and multiple hydroxyl

d. multiple carbonyl and one hydroxyl

Hint

3. What is the name of the carbohydrate that is composed of two monosaccharide units?

a. monosaccharide

b. disaccharide

c. oligosaccharide

d. polysaccharide

Hint













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4. What six-carbon sugar is found in the blood and provides immediate energy for the body?

a. sucrose

b. fructose

c. galactose

d. glucose

Hint

5. What is the name of the polysaccharide that cannot be digested by humans and provides

dietary fiber?

a. starch

b. amylose

c. glycogen

d. celluloseHint

6. The primary function of these two polysaccharides is to store energy.

a. fructose and sucrose

b. glucose and cellulose

c. starch and glycogen

d. glucose and starch

Hint

Section 24.3 Lipids

Practice Test

1.

What functional group do fatty acids and amino acids have in common?

a. amide

b. amino










7/16/2019 Section 1


c. carbonyl

d. carboxyl

Hint

2. Large biological molecules that are nonpolar belong to a class of molecules called _________.

a. lipids

b. proteins

c. carbohydrates

d. enzymes

Hint

3. ________________ are the building blocks of lipids.

a. Polysaccharides

b. Polymers

c. Carbohydrates

d. Fatty acids

Hint

4. ____________ is the common name given to a triglyceride that is a solid at room

temperature.

a. An oil b. A fat

c. A wax

d. Cholesterol

Hint

5. What is the process in which triglycerides react with a strong, inorganic base such assodium hydroxide to produce carboxylate salts and glycerol?

a. lipogenesis

b. saponification

c. peptide linkage

d. polymerization

Hint

6. Steroids are a type of lipid. What is the structure of a steroid?
















7/16/2019 Section 1


a. The basic structure is composed of four rings.

b. The structure has many fatty acids.

c. The structure uses peptide bonds.

d. The structure is composed entirely of water — soluble molecules.

Hint

Q82376A287821 6314 188 2 1 0

6

Section 24.4 Nucleic Acids

Practice Test

1.

If a molecule contains a 5-carbon sugar, the nitrogen base uracil, and can be used in

protein synthesis, then the molecule is __________.

a. DNA

b. RNA

c. thymine

d. cytosine

Hint

2. . Nucleic acids are made of nucleotides. There are three components to a nucleotide.

Which of the following correctly identifies those three parts?

a. nitrogen base, phosphate, and a five-carbon sugar

b. hydroxyl base, phosphate, and a five-carbon sugar







7/16/2019 Section 1


c. nitrogen base, phosphate, and a six-carbon sugar

d. nitrogen base, a protein, and a five-carbon sugar

Hint

3. DNA is a molecule made of two long chains of nucleotides that are arranged in adistinctive, three-dimensional physical structure. What is this three-dimensional structure?

a. a spherical macromolecule

b. a double helix

c. a single helix

d. a pleated sheet

Hint

4. What is the principle function of DNA?

a. stores genetic information

b. causes molecules to transfer from one place in the body to another

c. sends messages to different parts of the body

d. stores energy for later use

Hint

5. What sugar is used to make DNA?

a. ribose

b. ribulose

c. dextrose

d. deoxyribose

Hint

Section 24.5 Metabolism

Practice Test

1. What is the product of an endothermic reaction in which an inorganic phosphate group is

added to adenosine diphosphate?

a. ADP

b. ATP

c. glucose













7/16/2019 Section 1


d. lactic acid

Hint

2. The myriad set of chemical reactions that take place in the body so that cells can function

is __________________.

a. cannibalism

b. catabolism

c. metabolism

d. anabolism

Hint

Section 25.1 Nuclear Radiation

Practice Test

1.

Unstable atomic nuclei emit radiation to __________.

a. attain more stable atomic configurations

b. gain electrons







7/16/2019 Section 1


c. gain neutrons

d. lose protons

Hint

2. What are isotopes of atoms with unstable nuclei called?a. radioactivity

b. radioisotopes

c. radiation

d. radioactive decay

Hint

3. What particle emitted during radioactive decay has a mass of 4 amu?

a. alpha particle

b. beta particle

c. gamma particle

d. delta particle

Hint

4. Which particle emitted during radioactive decay is indistinguishable from an electron?

a. alpha particle

b. beta particle

c. gamma particle

d. delta particle

Hint

5. The type of radiation that has the greatest penetrating ability is ___________.

a. alpha radiation

b. beta radiation

c. gamma radiation

d. delta radiation

Hint

Section 25.2 Radioactive Decay
















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Practice Test

1. If a thorium-230 atom undergoes alpha decay, what are the products of the reaction?

a. actinium and an alpha particle

b. actinium and a positronc. radium and an alpha particle

d. radium and a positron

Hint

2. What ratio is used to evaluate the stability of a nucleus?

a. the electron to proton ratio

b. the neutron to proton ratio

c. the electron to neutron ratio

d. the atomic number to mass number ratio

Hint

3. What conditions in the nucleus usually result in beta decay?

a. The nucleus has too many neutrons.

b. The nucleus has too many protons.

c. The valence electrons are lost.

d. The atomic mass is greater than 300 amu.

Hint

4. What happens during positron emission?

a. A proton is converted into a neutron and a positron.

b. A neutron is converted into a proton and a positron.

c. A proton is converted into a neutron and an electron.

d. An atomic explosion occurs.

Hint

5. What happens to the mass of an atom when it undergoes alpha decay?

a. There is no change in mass.

b. The mass decreases by 1.

c. The mass increases by 4.













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d. The mass decreases by 4.

Hint

6. The atomic number of an atom ________ when it undergoes positron emission.

a. decreases by 2 b. increases by 2

c. decreases by 1

d. increases by 1

Hint

7. A series of nuclear reactions that begins with an unstable nucleus and results in the

formation of a stable nucleus is _______________.

a. a radioactive decay series

b. an isotope series

c. a band of stability

d. the Balmer series

Hint

Section 25.3 Transmutation

Practice Test

1. The half-life for tritium is 12.32 years. How long will it take for a 10.00-g sample of

tritium to decay until 1.875 g remain?

Data Table 1

Parent and daughter nuclei data

Number of half-lives Parent fraction Daughter fraction Daughter-to-parent ratio

0 1

1 1/2

2 1/4

3 1/8

4 1/16

a. 0.6594 years










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b. 5.333 years

c. 24.64 years

d. 30.80 years

Hint

2. The conversion of an atom of one element into an atom of a different element is _______________.

a. isotopic abundance

b. radioactivity

c. transuranium element

d. transmutation

Hint

3. The elements with atomic numbers greater than 92 are _______________.

a. transuranium elements

b. lanthanides

c. actinides

d. halogens

Hint

4. The goal of ancient alchemy was to change lead (atomic number 82) into gold (atomic

number 79). Which of the following could result in the transformation of lead into gold?a. an induced transmutation that removes an alpha particle and a positron from a

lead atom

b. an induced transmutation that adds an alpha particle and removes a positronfrom a lead atom

c. an induced transmutation that removes an alpha particle and a beta particle

from a lead atom

d. an induced transmutation that adds an alpha particle and removes a beta particlefrom a lead atom

Hint

5. The time required for one-half of a radioactive isotope to decay into its products is

_______________.

a. half-time

b. half-life

c. transmutation













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d. isomerization

Hint

6. Carbon-14 is used to date archaeological artifacts. If carbon-14 decays by loss of a beta

particle, what new element is formed?

a. nitrogen-13

b. nitrogen-14

c. carbon-13

d. boron-14

Hint

Section 25.4 Fission and Fusion of Atomic Nuclei

Practice Test

1.

If one fission reaction of a uranium-235 atom produced two neutrons, how many neutrons

would be released if the chain reaction occurred three more times?

a. 2







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b. 4

c. 8

d. 16

Hint

2. To hold the nucleons together in the nucleus, energy is required. What is the name of thisenergy?

a. binding energy

b. kinetic energy

c. thermal energy

d. free energy

Hint

3. What is the name of the process in which a large radioactive isotope is broken intosmaller isotopes?

a. transmutation

b. nuclear fission

c. nuclear fusion

d. beta emission

Hint

4. What is the name given to the amount of a radioactive substance that is massive enough tosustain a chain reaction?

a. critical mass

b. critical condition

c. mole

d. mass defect

Hint

5. In a nuclear reactor, what is the heat generated by the nuclear fission reaction used for?

a. fusing other unstable isotopes together

b. generating steam

c. freezing water

d. making building products

Hint
















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6. The energy produced by the Sun is a result of ___________.

a. nuclear fission

b. alpha emission

c. nuclear fusion

d. radiocarbon dating

Hint

7. The reaction products of nuclear fusion are __________ than the reactants.

a. more massive

b. less massive

c. less radioactive

d. cooler Hint

8. What is the major problem associated with the development of fusion as a controlled

energy source?

a. The containment of the radioactive decay products.

b. The low energy yield of the fuel.

c. The containment of the extremely high-temperature plasma.

d. The resulting air pollution.

Hint

Section 25.5 Applications and Effects of Nuclear Reactions

Practice Test










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1.

Which of the following would be the best choice for use in detecting diseased tissue as

part of medical diagnostics?

a. gamma ray

b. ionizing radiation

c. radiotracer

d. X-ray

Hint

2. Radiation that is energetic enough to ionize matter is called ionizing radiation. Which of

the following devices could be used to detect ionizing radiation?

a. film badge

b. scintillation counter

c. Geiger counter

d. all of the above

Hint

Section 26.1 Earth's Atmosphere

Practice Test

1.







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The photodissociation of diatomic oxygen in the atmosphere produces ___________.

a. O

b. 2O

c. ozone

d. oxygen and electrons

Hint

2. Which part of the atmosphere is closest to Earth’s surface?

a. troposphere

b. stratosphere

c. mesosphere

d. thermosphereHint

3. Which gas is present in the highest concentration in the atmosphere?

a. oxygen

b. carbon dioxide

c. nitrogen

d. argon

Hint

4. Which part of the atmosphere contains the ozone layer?

a. troposphere

b. stratosphere

c. mesosphere

d. thermosphere

Hint

5. What atoms make up ozone?a. three oxygen atoms

b. two oxygen atoms

c. two oxygen atoms and one nitrogen atom

d. two nitrogen atoms













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Hint

6. Which of following is the blanket of pollution that is formed when sunlight reacts with

chemicals in the air?

a. ozone

b. chlorofluorocarbons

c. smog

d. rain

Hint

7. What are the two most common acid components of acid rain?

a. nitric acid and carbonic acid

b. nitric acid and phosphoric acid

c. nitric acid and sulfuric acid

d. sulfuric acid and carbonic acid

Hint

8. Acid rain chemically reacts with calcium carbonate, the major component of limestone

and marble, which are used to make buildings and statues. What are the chemical

products formed when calcium carbonate reacts with nitric acid?

a. carbon dioxide, water, and calcium nitrate

b. carbonic acid and calcium hydroxide

c. calcium carbonate and nitric acid

d. nitrous oxide and calcium hydroxide

Hint

9. Photodissociation in the upper atmosphere prevents some harmful radiation from reaching

Earth’s surface. What type of radiation is absorbed during photodissociation?

a. visible radiation

b. infrared radiation

c. high-energy ultraviolet radiation

d. radio waves

Hint

Section 26.2 Earth's Water
















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Practice Test

1.

__________ use the greatest percentage of freshwater in the United States.

a. Farms

b. Governments

c. Industries

d. Residents

Hint

2.

__________ is difficult to remove from sewage and wastewater and contributes to

bacterial and algal overgrowth in freshwater.

a. carbon

b. lead

c. nitrogen

d. oxygen

Hint

3. What term is used to describe all the water on Earth’s surface?

a. troposphere

b. oceans

c. aquifer

d. hydrosphere







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Hint

4. What is the source of power that drives the water cycle?

a. nuclear power plants

b. the Sunc. gravity

d. tides

Hint

5. What is the measure of salt concentration in seawater called?

a. salinity

b. molarity

c. density

d. aquation

Hint

6. What process is used to remove salt from seawater so that it can be used by humans?

a. precipitation

b. desalination

c. denaturation

d. aeration

Hint

7. What percentage of the hydrosphere is actually liquid freshwater?

a. 0.60%

b. 72%

c. 97%

d. 2.10%

Hint

8. During the municipal treatment of water, a process is used to combine air with dissolved

harmful, organic substances and convert them to harmless compounds. What is this process called?

a. sedimentation

b. aeration
















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c. coagulation

d. filtering

Hint

9. What chemical is used to kill any microorganisms that may have remained in the water after the initial cleanup?

a. chlorine

b. lime

c. alum

d. silicon dioxide

Hint

10. Seawater is desalinated commercially in some areas of the world by _____________.

a. reverse osmosis

b. precipitation

c. distillation

d. adsorption

Hint

Section 26.3 Earth's Crust

Practice Test










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1.

Which of the following resulted in Earth's core becoming more dense than its crust?

a. gravity

b. high electronegativity

c. oxygenation

d. rapid heating

Hint

2. What is the name given to a solid inorganic, compound found in nature?

a. solution

b. precipitation

c. mineral

d. vitamin

Hint

3. What is the solid part of Earth’s crust called? a. hydrosphere

b. troposphere

c. atmosphere

d. lithosphere

Hint










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4. What term is used to describe a metal-containing mineral that can be processed for a

reasonable profit?

a. gangue

b. slag

c. ore

d. crust

Hint

5. What is the most abundant element in Earth’s crust?

a. nitrogen

b. oxygen

c. carbon

d. sodium

Hint

Q82448AC28810 6323 188 2 1 0

5

Section 26.4 Cycles in the Environment

Practice Test







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1.

The two routes for nitrogen fixation in nature are __________.

a. bacteria and the water cycle

b. bacteria and lightning

c. lightning and ultraviolet radiation

d. lightning and ozone formation

Hint

2. What chemicals are used by plants to make carbohydrates during photosynthesis?

a. carbon dioxide and water

b. carbon dioxide and calcium carbonate

c. water and calcium carbonate

d. carbon dioxide and nitrogen

Hint

3. What is the natural warming of Earth’s surface that occurs when certain gases in the

atmosphere absorb heat?

a. vulcanism

b. carbon cycle

c. greenhouse effect

d. photosynthesis







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Hint

4. When an increase in greenhouse gases leads to an increase in the temperature over all of

Earth, a specific condition occurs. What is the common name for this condition?

a. global warming

b. ozone hole

c. photosynthesis

d. carbon cycle

Hint

5. Although nitrogen gas cannot be used by plants, it can be converted into forms that plants

can use. What process converts nitrogen gas into this more useful form?

a. global warming

b. nitrogen fixation

c. photosynthesis

d. precipitation

Hint

6. Nitrogen is oxidized in the atmosphere and in the soil by nitrogen-fixing bacteria. What is

the most oxidized form of nitrogen?

a. nitrogen monoxide

b. nitrogen dioxide

c. nitrate

d. ammonia

Hint

http://highered.mcgraw-hill.com/sites/0078664187/student_view0/self-check_quizzes.html

http://highered.mcgraw-hill.com/sites/0078664187/student_view0/chapter1/web_links.html













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Section 1

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