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Chemistry Assessments

QUARTER ONEUnit 1

C2.5a Determine the age of materials using the ratio of stable and unstable isotopes of a particular type.

1. The isotope Np-238 has a half-life of 2.0 days. If 84 grams are produced on Monday, how much will be remaining 4.0 days later?

A. 18 gramsB. 21 gramsC. 35 gramsD. 42 gramsE. 84 grams

Answer: B

2. The isotope Np-238 has a half-life of 2.0 days. If 96 grams are produced on Monday, how much will be remaining 6.0 days later?

A. 8 gramsB. 12 gramsC. 24 gramsD. 48 gramsE. 96 grams

Answer: B

3. The half-life of a radioactive substance is 2.5 minutes. What fraction of the original radioactive substance remains after 10 minutes?

A. 1/2B. 1/4C. 1/8D. 1/16

Answer: D

Chemistry Assessments – May 2008 1

4. The half-life of 131I is 8.07 days. What fraction of a sample of 131I remains after 24.21 days?

A. 1/2B. 1/4C. 1/8D. 1/16

Answer: C

5. What does the half-life of an isotope mean?

A. It means that ½ of the atoms are isotopesB. It is the time it takes for ½ of the isotope atoms to decay into another materialC. It is ½ of the age of the isotopeD. It is ½ of the atomic mass of the atom

Answer: B

6. Which of the following best explains why a large amount of energy is released in this reaction?

A. The hydrogen converts the light into energy. B. Some of the reactant mass is converted into energy.C. All of the hydrogen isotopes undergo radioactive decay.D. The temperature of the products is lower than that of the reactants.

Answer: B

C3.5a Explain why matter is not conserved in nuclear reactions

1. Conservation of mass is not true for nuclear reactions because:

A. mass lost in a nuclear reaction is not destroyed, cut converted into energyB. alpha particles are converted into beta particlesC. energy lost in a nuclear reaction is converted into massD. the reactants of a nuclear reaction are neutrons

Answer: A


C4.7b Compare the density of pure water to that of a sugar solution.

1. The density of water is 1.0 g/cm3, the density of a 10% sugar solution is:

A. less than the density of waterB. greater than the density of waterC. equal to the density of waterD. none of the above

Answer: B

C4.8A Identify the location, relative mass, and charge for electrons, protons, and neutrons

1. What is the nuclear charge of an iron atom?

A. +26B. +30C. +56D. +82

Answer: A

Using the diagram above and a periodic table, answer the following questions:

2. What is the atomic number of this element?

A. 2B. 3C. 8D. 13E. 25

Answer: D


13P14N

e

e e

e e

e e

e

e

e ee

e

3. Which symbol represents this element?

A. HeB. LiC. OD. AlE. Mn

Answer: D

4. What is the mass number of this element?

A. 2B. 3C. 8D. 13E. 27

Answer: E

C4.8B Describe the atom as mostly empty space with an extremely small, dense nucleus consisting of the protons and neutrons and an electron cloud surrounding the nucleus

Using the diagram above, answer the following questions:

1. What can be inferred from the diagram about the structure of the atom?

A. the atom is very smallB. the electrons are moving very fastC. the atom is mainly empty spaceD. the electrons are very heavyE. the nucleus is composed of protons and neutrons

Answer: C


2. Experiments performed to reveal the structure of atoms led scientists to conclude that an atom’s

A. positive charge is evenly distributed throughout its volumeB. negative charge is mainly concentrated in its nucleusC. mass is evenly distributed throughout its volumeD. volume is mainly unoccupied

Answer: D

3. An experiment in which alpha particles were used to bombard thin sheets of gold foil led to the conclusion that an atom is composed mostly of

A. empty space and has a small, negatively charged nucleusB. empty space and has a small, positively charged nucleusC. a large, dense, positively charged nucleusD. a large, dense, negatively charged nucleus

Answer: B

4. Which of these conclusions can be drawn from Rutherford’s experiment?

A. Each atom contains electrons.B. The nucleus of an atom can be split.C. Each atom contains protons.D. Atoms are mostly empty space.

Answer: D


5. Ernest Rutherford performed an experiment in which he shot alpha particles through a thin layer of gold foil. He predicted that the alpha particles would travel straight through the gold atoms, as shown below

However, Rutherford observed that although most of the alpha particles passed straight through the foil, a few alpha particles were deflected, as shown below.

Which of the following statements about the atom did Rutherford’s experiment support?

A. An atom contains protons, neutrons, and electrons.B. An atom’s nucleus is small and has a positive charge.C. Electrons follow a predictable path around the nucleus.D. Different isotopes of an element have different masses.

Answer: B

C4.8C Recognize that protons repel each other and that a strong force needs to be present to keep the nucleus intact.

1. A nucleus consists of protons and neutrons that are tightly bound together by

A. the strong nuclear force B. repulsion of the protonsC. attraction of the electronsD. the positive charge of the protons

Answer: A


2. The neutrons in the nucleus of an atom produce a net attractive force to

A. repel the protonsB. produce a strong nuclear forceC. attract the electronsD. release energy

Answer: B

C4.8D Give the number of electrons and protons present if the fluoride ion has a -1 charge.

1. To change F to F-, you need to:

A. Add one electronB. Remove one electronC. Remove one neutronD. Add one neutron

Answer: A

C4.10A List the number of protons, neutrons, and electrons of any given ion or isotope.

1. As a Ca atom undergoes oxidation to Ca2+, the number of neutrons in its nucleus:

A. decreasesB. increasesC. remains the same

Answer: C

2. An atom of carbon-12 and an atom of carbon-14 differ in:

A. atomic numberB. atomic massC. nuclear chargeD. number of electrons

Answer: B


3. All the isotopes of a given atom have:

A. the same atomic mass and the same atomic numberB. the same atomic mass but different atomic numbersC. different atomic mass but the same atomic numberD. different atomic mass and different atomic numbers

Answer: C

4. Atoms of the same element that have different numbers of neutrons are classified as:

A. charged atomsB. charged nucleiC. isomersD. isotopes

Answer: D

5. What are isotopes?

A. Atoms that only have neutrons in their nucleusB. Atoms that have lost or gained electronsC. Atoms with different numbers of neutrons for that elementD. Atoms that only have neutrons in their nucleus

Answer: C

6. What is the atomic number of an element that has six protons and eight neutrons?

A. 6B. 2C. 8D. 14

Answer: A

7. Why is one isotope of an element so much more common that the others?

A. The less common isotopes are usually unstableB. The heavier the isotope, the more common it isC. All isotopes are commonD. All isotopes are uncommon

Answer: A


C4.10B Recognize that an element always contains the same number of protons.

1. Is it possible for an element to have isotopes with different numbers of protons?

A. Yes, that is the definition of an isotopeB. It depends on the number of electrons the element hasC. No, because the element is defined by a specific number of protonsD. No, because the element is defined by a specific number of neutrons

Answer: C

2. Which statement about the mass of an electron is correct?A. The mass of an electron is equal to the mass of a proton.B. The mass of an electron is less than the mass of a proton.C. The mass of an electron is equal to the mass of a neutron.D. The mass of an electron is less than the mass of a neutron.

Answer: B

C4.10e Write the symbol for the isotope , where Z is the atomic number, A is the mass number and X is the symbol for the element.

1. What is the structure of a krypton-85 atom?

A. 49 electrons, 49 protons, 85 neutronsB. 49 electrons, 49 protons, 49 neutronsC. 36 electrons, 36 protons, 85 neutronsD. 36 electrons, 36 protons, 49 neutrons

Answer: D


C5.2C Draw pictures to distinguish the relationships between atoms in physical and chemical changes.

H2SO4 + 2NaOH Na2SO4 + 2H2O

1. The illustrations show a conservation-of-mass experiment. The solution in the beaker lost mass because:

A. materials have less mass at high temperaturesB. the mass of the reactants and products was less than 100 gC. sodium sulfate (Na2SO4) is lighter than airD. some of the water molecules turned into gas

Answer: D


Unit 2

C4.9A Identify elements with similar chemical and physical properties using the periodic table.

1. Elements in the Periodic Table are arranged according to their

A. atomic numberB. atomic massC. relative activityD. relative size

Answer: A

2. In which list are the elements arranged in order of increasing atomic mass?

A Cl, K, ArB. Fe, Co, NiC. Te, I, SbD. Ne, F, Na

Answer: B

3. Which element would you expect to have chemical properties similar to calcium (Ca)?

A. strontium, SrB. potassium, KC. rubidium, RbD. sodium, Na

Answer: A

4. Which of the following elements has characteristics of some metals and also of some nonmetals?

A. antimony ( 51Sb)B. calcium ( 20Ca)C. sulfur ( 16 S)D. zinc ( 30 Zn)

Answer: A


C4.9b Identify metals, non-metals, and metalloids using the periodic table.

1. Which elements are both classified as metalloids?

A. Ge and AsB. Bi and PoC. B and CD. Si and P

Answer: A

2. A characteristic of a nonmetal is

A. high melting pointsB. high electronegativityC. high electrical conductivityD. the ability to form positive ions

Answer: B

3. Determine from the periodic table which set of elements represents exclusively metals.

A. 0, N, Cl, BrB. Fe, Mn, Cl, HeC. Cu, Fe, Sn, LiD. Rn, 0, P, SE. Au, Ag, Na, Br

Answer: C

4. Determine from the periodic table which set of elements represents exclusively gases.

A. O, N, NeB. Al, Na, ClC. Sr, Ar, ZnD. C, P, OE. S, Fe, Ra

Answer: A


5. The table below contains a list of properties for an unidentified element, X.

Based on the properties in the table, to which of the following groups from the periodic table does element X most likely belong?

A. (1A)B. (2A)C. (4A)D. (6A)

Answer: A

C4.9c. Predict general trends in atomic radius, first ionization energy, and electronegativity of the elements using the periodic table.

1. Arsenic and silicon are similar in that they both

A. have the same ionization energyB. have the same covalent radiusC. are transition metalsD. are metalloids

Answer: D

2. Which Group of the Periodic Table contains atoms with a stable outer electron configuration?

A. 1B. 8C. 16D. 18

Answer: D


3. The high electrical conductivity of metals is primarily due to:

A. high ionization energyB. filled energy levelsC. mobile electronsD. high electronegativities

Answer: C

4. According to the periodic table, which element most readily accepts electrons?

A. Fluorine B. Nitrogen C. Arsenic D. Aluminum

Answer: A

5. Which of the following trends in the periodic table should be expected as the atomic number of the halogens increases from fluorine (F) to iodine (I)?

A. Atomic radius decreases.B. Electronegativity decreases.C. Atomic mass decreases.D. Electron number decreases.

Answer: B

C4.10c Calculate the average atomic mass of an element given the percent abundance and mass of the individual isotopes.

1. The percent composition of Hydrogen is 99.98% found in nature. The percent abundance of Deuterium is .0156%. The mass of Hydrogen is 1.00794 amu and the mass of Deuterium is 2.01664amu. Calculate the average atomic mass of Hydrogen.

Answer: 1.00805


C4.10d Predict which isotope will have the greatest abundance given the possible isotopes for an element and the average atomic mass in the periodic table.

1. The average atomic mass of Chlorine is 35.453 amu. The isotopes of Chlorine are Chlorine-35 and Chlorine-37. Determine which isotope will be found in greatest abundance given the atomic mass.

Answer: Chlorine exists as two common isotopes. Chlorine-35 has an atomic mass of about 35 amu, Chlorine-37 has an atomic mass of about 37amu. The fractional abundance of Chlorine-35 in nature is 75 percent, while Chlorine-37 is found to be 25 percent. The average mass of Chlorine atoms is somewhere between 35 and 37, closer to 35, because Chlorine-35 is more abundant than Chlorine-37.

C5.2g Calculate the number of atoms present in a given mass of element.

1. How many atoms of Calcium are there in 80 grams?

A. 6.02 x 10 23

B. 12.04 x 10 23

C. 18.06 x 10 23

D. 24.08 x 10 23

Answer: D

2. How many atoms of carbon are there in 36 grams?

A. 6.02 x 10 23

B. 12.04 x 10 23

C. 18.06 x 10 23

D. 24.08 x 10 23

Answer: C

3. How many atoms of Oxygen are there in 64 grams?

A. 6.02 x 10 23

B. 12.04 x 10 23

C. 18.06 x 10 23

D. 24.08 x 10 23

Answer: D


C5.5A Predict if the bonding between two atoms of different elements will be primarily ionic or covalent.

1. Element X is in Group 2 and element Y is in Group 17. What happens when a compound is formed between these two atoms?

A. X loses electrons to Y to form an ionic bond.B. X loses electrons to Y to form a covalent bond.C. X gains electrons from Y to form an ionic bond.D. X gains electrons from Y to form a covalent bond.

Answer: A

2. Which atom will form an ionic bond with a Br atom?

A. NB. LiC. OD. C

Answer: B

3. Which elements combine by forming an ionic bond?

A. sodium and potassiumB. sodium and oxygenC. carbon and oxygenD. carbon and sulfur

Answer: A

4. Which type of bond is formed when an atom of potassium transfers an electron to a bromine atom?

A. metallicB. ionicC. nonpolar covalentD. polar covalent

Answer: B


C5.5B Predict the formula for binary compounds of main group elements.

Use a polyatomic ion table with oxidation numbers to answer the following questions.

1. What is the formula for the compound named sodium sulfate?

A. NaSOB. Na2SOC. Na2SO4

D. S2SO4

E. SSO4

Answer: B

2. What is the formula for the compound named potassium nitrate?

A. KNOB. KNO3

C. knoD. PNE. PNO

Answer: B

Rules for Naming Transition Metal Salts

1. The name of the metal is unchanged.2. For metals having more than one oxidation state, a Roman numeral in parenthesis indicates

the oxidation state of the metal.3. The anion is named by adding -ide to the stem of the element.

Using the rules above, answer the following questions:

1. What is the proper name for Cr2O3?

A. Chromium (II) oxideB. Chromium (III) oxideC. Chromium trioxideD. Dichromium trioxideE. Chromium oxide

Answer: A


2. What is the formula for Iron (II) sulfide?

A. FeSB. Fe2SC. FeS2

D. Fe2S2

E. Fe2S3

Answer: A

3. What is the proper name for ZnCl2?

A. Zinc chlorideB. Zinc (I) chlorideC. Zinc (II) chlorideD. Zinc (III) chlorideE. Zinc chlorine

Answer: A

C5.5c Draw Lewis structures for simple compounds

1. Which Lewis dot structure for NO2+ has the lowest energy?

Answer: A

2. Which of the following elements has the same Lewis dot structure as silicon?

A. germanium (Ge)B. aluminum (Al)C. arsenic (As)D. gallium (Ga)

Answer: A


3. Which of the following Lewis dot structures represents the compound methane (CH4)?

A.

B.

C.

D.

Answer: A

C5.5d Compare the relative melting point, electrical and thermal conductivity, and hardness for ionic, metallic, and covalent compounds.

A chemist performs the same tests on two white crystalline solids, A and B. The results are show in the table below.

Solid A Solid B

Melting Point High, 801≈C Low, decomposesat 186≈C

Solubility in H2O(grams per 100.0 g in H2O at 0≈C)

35.7 3.2

Electrical Conductivity(in aqueous solution)

Good conductor Nonconductor

1. The results of these tests suggest that:

A. both solids contain only ionic bondsB. both solids contain only covalent bondsC. solid A contains only covalent bonds and solid B contains only ionic bondsD. solid A contains only ionic bonds and solid B contains only covalent bonds

Answer: D


2. What type of bond is used when hydrogen is bonded to oxygen in a water molecule?

A. covalentB. non-polarC. hydrogenD. ionicE. polar

Answer: A

3. Which are the properties of metals?

A. conductor, brittle low melting point, lustrousB. non conductor, brittle low melting point, lustrousC. malleable, conductor, high melting point, lustrousD. non conductor, brittle, high melting point, lustrousE. non conductor, brittle, high melting point, non lustrous

Answer: C

4. Several chemists examined a pure, unknown substance and observed and measured its physical properties. Their results are shown below.

Unknown Substance

Based on the data recorded in the table, answer the following.

a. What is the physical state of this substance at room temperature? Explain how the information in the table is used to make this classification of the substance’s state.

b. The substance is unreactive in water. What will happen if 10.00 g of this substance is added to 200.0 g of water at 20°C and standard pressure? Explain your response.

Answer: a. The unknown substance is a liquid at room temperature, since its melting point is –22.9oC. The substance is a solid below –22.9oC. The substance is a gas above 76.74oC. b. Since the solubility of the substance is .08g/100g water, then .16 g of the substance will dissolve in 200 g water.


Unit 3

C2.4a Describe energy changes in flame tests of common elements in terms of the (characteristic) electron transitions.

1. The color of light emitted by an atom is most related to

A. energy release by the electronB. mass of the electron C. potential energy of the electron in the ground stateD. size of the electron E. strength of charge on the electron

Answer: A

2. What causes the emission of radiant energy that produces characteristic spectral lines?

A. neutron absorption by the nucleusB. gamma ray emission from the nucleusC. movement of electrons to higher energy levelsD. return of electrons to lower energy levels

Answer: D

3. The characteristic bright-line spectrum of an element is produced when electrons

A. absorb quanta and return to lower energy levelsB. absorb quanta and move to higher energy levelsC. release quanta and return to lower energy levelsD. release quanta and move to higher energy levels.

Answer: C

4. During a flame test, ions of a specific metal are heated in the flame of a gas burner. A characteristic color of light is emitted by these ions in the flame when the electrons

A. gain energy as they return to lower energy levelsB. gain energy as they move to higher energy levelsC. emit energy as they return to lower energy levelsD. emit energy as they move to higher energy levels

Answer: C


5. Of the particles listed here, which is most involved with quantum theory?

A. electronB. protonC. neutronD. nucleusE. quark

Answer: A

C2.4b Contrast the mechanism of energy changes and the appearance of absorption and emission spectra.

1. The light produced by signs using neon gas results from electrons that are:

A. moving from a higher to a lower principal energy levelB. moving from a lower to a higher principal energy levelC. being lost by the Ne(g) atomsD. being gained by the Ne(g) atoms

Answer: A

C2.4c Explain why an atom can absorb only certain wavelengths of light.

Electromagnetic radiations are photons traveling in waves at the speed of light. The wavelength is the distance from the crest of one wave to the crest of another. Frequency is the number of waves that pass a given point in one second (waves/sec).

Answer the following questions using the above diagram and statement.


1. The electromagnetic spectrum shows the relationship between:

A. the mass of light and its forceB. the forces caused by light and the wavelengthC. the amount of energy and wavelengthD. the electrical power and strength of energyE. gamma and radio waves only

Answer: C

2. Which of the following occurs as energy of electromagnetic radiation increases?

A. wavelength increasesB. wavelength decreasesC. frequency increasesD. frequency decreasesE. both b and d are correct

Answer: E

3. Of the colors listed below, which one has the highest energy?

A. blueB. greenC. orangeD. redE. yellow

Answer: A

4. When an excited electron drops from the 4th orbital to the 2nd orbital, a photon that exhibits the color green. What would the color be of a photon emitted when an electron drops from the 3rd to the 2nd orbital?

A. blueB. greenC. purpleD. yellowE. white

Answer: D


Frequency

Energy Wavelength

Frequency

C2.4d Compare various wavelengths of light (visible and non-visible) in terms of frequency and relative energy.

1. From the above graphs, describe the relationship between energy, wavelength, and

frequency.

A. as the wavelength increases, the frequency and energy decreaseB. as the frequency increases, the wavelength and energy decreaseC. as the energy increases, the frequency and wavelength decreaseD. as the energy increases, the frequency and wavelength increase

Answer: A 2. From the above graphs, describe how you could increase the energy of light the most.

A. increase the wavelengthB. decrease the wavelengthC. decrease the frequencyD. increase the wavelength while decreasing the frequency

Answer: B

3. The proper relationship of wavelength (8) to light energy (E) is:

A. increased wavelength correlates to increased energyB. increased wavelength correlates to decreased energyC. wavelength is independent of energyD. wavelength is independent of frequency

Answer: B


C4.8e Write the complete electron configuration of elements in the first four rows of the periodic table.

1. 1s22s22p63s23p64s1 is the electron configuration for which element?

A. aluminum (Al)B. argon (Ar)C. potassium (K)D. sodium (Na)

Answer: C

C4.8f Write kernel structures for main group elements.

1. Sodium has an atomic number of 11 and a mass number of 23.

A. Identify the types of subatomic particles located in the nucleus of a sodium atom. Compare the properties of each type of particle.

B. Where is most of a sodium atom’s mass located? Explain your answer. C. Identify the subatomic particles that are found in the energy levels outside the

nucleus of a sodium atom. D. Describe the number and arrangement of these particles, including valence and

nonvalence (kernel) electrons.

Answers: a. Sodium has 12 neutrons and 11 protons. b. Most of the mass of a sodium atom is found in the nucleus. c. 11 electrons are found in the energy levels of the atom.d. The atom has 10 non-valance electrons and 1 valance electron. The configuration is 1s2 2s2

2p6 3s1.

C4.8g Predict oxidation states and bonding capacity for main group elements using their electron structure.

Use the diagram below to answer question 1.


1. Which element will gain only one electron during a chemical reaction?

A. silicon B. phosphorus C. sulfur D. chlorine

Answer: D

C4.8h Describe the shape and orientation of s and p orbitals.

1. What is the shape of the s orbital?

A. doughnutB. dumbbellC. sphericalD. elliptical

Answer: C

2. What is the shape of the p orbital?

A. doughnutB. dumbbellC. sphericalD. elliptical

Answer: B

C4.8i Describe the fact that the electron location cannot be exactly determined at any given time.

1. What is an electron cloud and why can’t an exact electron location be determined?

Answer: In atomic theory, an electron cloud is a visual prediction of where an electron may be located in respect to the nucleus, at a given time. Electrons do not travel in circular paths and it is impossible to predict where an electron is at a given time. Instead, scientists now describe electrons in terms of their probable location around the nucleus.


QUARTER TWOUnit 7

C2.2A Describe conduction in terms of molecules bumping into each other to transfer energy. Explain why there is better conduction in solids and liquids than gases.

1. Increasing the temperature increases the rate of a reaction by:

A. lowering the activation energyB. increasing the activation energyC. lowering the frequency of effective collisions between reacting moleculesD. increasing the frequency of effective collisions between reacting molecules

Answer: C

2. As the concentration of reacting particles increases, the rate of reaction generally

A. decreasesB. increasesC. remains the same

Answer: B

3. A 1.0-gram sample of powdered Zn reacts faster with HCl than a single 1.0-gram piece of Zn because the surface atoms in powered Zn have

A. higher average kinetic energyB. lower average kinetic energyC. more contact with the H+ ions in the acidD. less contact with the H+ ions in the acid

Answer: C

C2.2B Describe the various states of matter in terms of the motion and arrangement of the molecules (atoms) making up the substance.


C2.2c Explain changes in pressure, volume, and temperature for gases using the kinetic molecular model.

1. Four different gases are all observed to have the same temperature. Which of the following conclusions is supported by this observation?

A. All four gases must have the same mass. B. All four gases must have the same pressure. C. All four gases must have equal numbers of particles. D. All four gases must have equal average kinetic energies.

Answer: D

2. The illustration below shows four containers. Each container is full of helium gas at a different temperature.

If all of the containers are closed and have a pressure of 1 atm, which container has helium particles with the greatest average kinetic energy?

A. 1 *B. 2C. 3D. 4

Answer: A

C2.2f Compare the average kinetic energy of the molecules in a metal object and a wood object at room temperature.


1. What probably causes water to have the highest specific heat of the substances listed above?

A. Molecule sizeB. Molecular massC. Strong hydrogen bondsD. High density of ice

Answer: C

C3.3A Describe how heat is conducted in a solid.

1. Heat convection occurs in gases and liquids. Heat convection does not occur in solids because solids are unable to —

A. absorb heat by vibrating B. transfer heat by fluid motion C. emit radiation by reflecting light D. exchange heat by direct contact

Answer: A

C3.3B Describe melting on a molecular level.

1. Which of the following phase changes results in an overall increase in randomness of particles over the course of the change?

A. freezing B. condensationC. meltingD. deposition

Answer: D


Matter can exist in three phases. A solid has a definite shape and volume. A liquid has a definite volume, but takes on the shape of a container. A gas takes on both the volume and shape of a container. The differences in the physical properties of the three phases of matter can be explained by differences in how the particles (atoms or molecules) that make up matter interact, as in Figures A, B, and C.

2. A student took a small ice cube and heated it in a large, closed container. First, the ice cube melted and then the water boiled into steam. According to the information given, how would the volume of the water and the volume of the steam compare to the volume of the original ice cube?

A. water: about the same; steam: much greaterB. water: about the same; steam: about the sameC. water: much greater; steam: much greaterD. water: much greater; steam: about the same

Answer: A

3. If a liquid is heated, what happens to both its energy content and the speed of its particles?

A. The energy increases; the speed decreases.B. The energy increases; the speed increases.C. The energy remains the same; the speed remains the same.D. The energy remains the same; the speed increases.

Answer: B


4. You stir a glass of milk with a spoon. Which of the following best explains why the spoon keeps its shape while it passes through the milk?

A. In the spoon, the particles vibrate back and forth; in the milk, the particles maintain their relative positions.

B. In the spoon, the particles randomly move in all directions; in the milk, the particles are close together.

C. In the spoon, the particles maintain their relative positions; in the milk, the particles can only vibrate back and forth.

D. In the spoon, the particles maintain their relative positions; in the milk, the particles move in all directions.

Answer: D

C4.3A Recognize that substances that are solid at room temperature have stronger attractive forces than liquids at room temperature, which have stronger attractive forces than gases at room temperature.

1. If the attractive forces among solid particles are less than the attractive forces between the solid and a liquid, the solid will:

A. probably form a new precipitate as its crystal lattice is broken and re-formed.

B. be unaffected because attractive forces within the crystal lattice are too strong for the dissolution to occur.

C. begin the process of melting to form a liquid. D. dissolve as particles are pulled away from the crystal lattice by the liquid

molecules.

Answer: C

2. Compared with the particles in a solid, the particles in a liquid usually are:

A. Closer togetherB. Less fluidC. More massiveD. Higher in energy

Answer: D


3. The higher the temperature of a liquid, the greater the average speed of its molecules.What would happen to the temperature of an evaporating liquid if the molecules with higher speed are escaping the liquid and no energy is applied (like perspiration on skin)?

A. The temperature of the liquid will increase.B. The temperature of the liquid will remain unchanged.C. The temperature of the liquid will decrease.D. The temperature of the liquid will decrease briefly and then increase.

Answer: B

C4.3B Recognize that solids have a more ordered, regular arrangement of their particles than liquids and that liquids are more ordered than gases

The graph below compares three states of a substance

1. Which of the following choices is the best label for the y-axis?

A. molecular density B. molecular motion C. neutron density D. neutron motion

Answer: B


C4.5a Provide macroscopic examples, atomic and molecular explanations, and mathematical representations (graphs and equations) for the pressure-volume relationship in gases.

1. Which of the following best illustrates a graph of pressure versus volume for a gas at a constant temperature?

Answer: A

C4.5b Provide macroscopic examples, atomic and molecular explanations, and mathematical representations (graphs and equations) for the pressure-temperature relationship in gases.

1. Which of the following correctly describes molecules of two different gases if they are at the same temperature and pressure?

A. They must have the same mass.B. They must have the same velocity.C. They must have the same average kinetic energy.D. They must have the same average potential energy.

Answer: C

2. A cylinder of gas particles is shown below.

The cylinder is fitted with a movable piston that can be raised and lowered. Which of the following would result in an increase in the pressure of the gas below the piston?

A. increasing the volume of the cylinderB. removing some of the gas from the cylinderC. decreasing the volume of the cylinderD. decreasing the pressure outside the cylinder

Answer: C


3. Which of the following occurs when a rigid container of gas is heated?

A. The pressure inside the container increases. B. The pressure inside the container decreases. C. The pressure inside the container stays the same. D. The pressure inside the container changes the composition of the gas.

Answer: A

4. The graph shows the pressure of an ideal gas as a function of its volume. According to the graph, increasing the volume from 100 mL to 150 mL:

A. decreases the pressure by 160 kPaB. increases the pressure by 80 kPaC. increases the pressure by 160 kPaD. decreases the pressure by 80 kPa

Answer: D


C4.5c Provide macroscopic examples, atomic and molecular explanations, and mathematical representations (graphs and equations) for the temperature-volume relationship in gases

1. Assuming pressure is held constant, which of the following graphs shows how the volume of an ideal gas changes with temperature?

A. B.

C. D.

Answer: A

2. According to Charles’s law, what will happen to the piston when the gas is heated?

A. The piston will move up because the gas particles get larger.B. There will be no change because heat will not affect the system.C. The piston will move up because the gas particles move faster and get

farther apart.D. The piston will move down because the gas particles move slower and

get closer together.

Answer: C


Unit 4INTRODUCTION TO BONDING

C2.1a Explain the changes in potential energy (due to electrostatic interactions) as a chemical bond forms and use this to explain why bond breaking always requires energy.

1. Because we can observe two types of electrostatic interactions, attractive and repulsive, we postulate that there are two kinds of charge. Do you think that there are also two types of mass? Why do you think this? Do you think that there are two types of magnetic poles? Why do you think this?

Answer: Mass is the quantity responsible for gravitational interactions. Since all gravitational interactions are attractive, we can assume that there is only one type of mass. Magnetic poles are responsible for magnetic interactions. We find that magnetic interactions can be attractive or repulsive. Therefore, there must be two types of magnetic poles (called North and South).

2. Imagine having a collection of electrically charged marbles which retain their charge even if they touch other objects. Red marbles are positively charged, blue marbles are negatively charged.

What happens if you place (a) a bunch of red marbles, (b) a bunch of blue marbles, or (c) an equal mixture of red and blue marbles close together on a flat horizontal surface? (d)

What happens in (c) if you have few more red marbles than blue ones? (e) As a whole, is the collection of marbles in (d) positively charged, negatively charged, or neither?

ANSWER: (a) The red marbles, all carrying a positive charge, exert repelling forces on each other and so they fly apart so as to get as far away as possible from each other, see Figure S26.10a. (b) The blue marbles, all carrying a negative charge, also exert repelling forces on each


other and so they, too, spread out as illustrated in Figure S26.10b. (c) Red and blue marbles are attracted to each other and so they form red-blue pairs. A red-blue pair carries no net charge and so it does not interact with other red-blue pairs of marbles. Consequently the marbles organize themselves in pairs; because the pairs don’t interact they do not spread out, see Figure S26.10c. (As we will see later, because the positive and negative charges do not overlap completely — a red and a blue marble cannot be at the same place — some residual interaction is left). (d) Each blue marble forms a red-blue pair; the left-over red marbles repel each other and spread out. (e) Because of the surplus of positively-charged red marbles, the entire collection of marbles carries a positive charge. Remember it is not the type of charge present in an object that determines the net charge, but the type of charge causing the surplus.

C2.1b Describe energy changes associated with chemical reactions in terms of bonds broken and formed (including intermolecular forces).

1. Given the reaction:Cl(g) + Cl(g) Cl2(g) + energy

Which statement best describes the reaction?

A. A bond is formed and energy is absorbed.B. A bond is formed and energy is released.C. A bond is broken and energy is absorbed.D. A bond is broken and energy is released.

Answer: B Use the following equation to answer the question:

Energy + 2C + 3H2 C2H6

carbon hydrogen ethane 2. On the right hand side of the equation are two carbons and six hydrogens. On the left hand

side are two carbons and six hydrogens. This is because:

A. Only two carbons can react with six hydrogens.B. The conservation of matter is reflected in this model.C. The conservation of volume is reflected in this model.D. The formation of ethane requires small amounts of energy.E. The formation of ethane requires large amounts of energy.

Answer: B


3. Which of the statements given concerning chemical bonding is FALSE?

A. Energy is required to break any chemical bond.B. Energy must be released if two atoms are to form a chemical bond.C. A chemical bond occurs only if the potential energy is lowered.D. All chemical reactions require a transfer of electrons in forming new bonds.E. Chemical bonds are due to electrostatic forces of attraction and repulsion.

Answer: D

C3.2b Describe the relative strength of single, double, and triple covalent bonds between nitrogen atoms.

1. What is the formal charge on a triply-bonded oxygen?

A. +1B. 0C. -1 D. Cannot be determined

Answer: A

2. Which of the following statements is correct?

A. the C(triple bond)C bond is three times as strong as a C-C bond. B. the C(triple bond)C bond is composed of two sigma bonds and one pi bond. C. the C(triple bond)C bond is composed of one sigma bond and two pi bonds. D. the pi bond is stronger than the sigma bond in an alkyne.

Answer: C

3. Which of the following is a correct explanation of why it is difficult to draw the Lewis structure for NO.

A. As an odd electron molecule, it is impossible to complete the octet on both atomsB. One can only draw structures with high formal charges. C. It is impossible to tell where to place the odd electron. D. None of the above

Answer: A


4. The Lewis structure of NO is difficult to draw due to its odd number of electrons. Use molecular orbital theory to predict which atom has the odd electron.

A. OB. NC. Neither D. Cannot be determined

Answer: B

C3.3c Explain why it is necessary for a molecule to absorb energy in order to break a chemical bond.

1. The molar entropies of NO(g), N(g) and O(g) at 298oK are 210.6, 153.2, and 161.0 J mole-1 K-1 respectively. Using the Go given below calculate the bond energy of NO.

NO(g) N(g) + O(g)Go = 600.1 kJ

A. 609.0B. 620.0C. 631.0D. 639.0

Answer: C

2. The entropies of CH4(g), C(g) and H(g) are 186.2, 158.0, and 114.6 J mole-1 K-1 respectively at 298oK. Using the Go given below calculate the bond energy of CH4.

CH4(g) C(g) + 4 H(g)Go = 1535.2 kJ

A. 440.0B. 422.6C. 415.8D. 406.9

Answer: C


3. The entropies of I2(g) and I(g) are 260.6 and 180.7 J mole-1 K-1 respectively at 298oK. Using the data given below calculate the bond energy of I2.

I2(g) 2 I(g)Go = 121.2 kJ

A. 145.2B. 148.3C. 151.2D. 153.5

Answer: C

C4.4a Explain why at room temperature different compounds can exist in different phases.

The table below shows some characteristics of four substances at 1 atm pressure.

1. Which of the following substances is a liquid at temperatures ranging from -50°C to 0°C?

A. bromineB. chlorineC. ethanolD. mercury

Answer: C

C4.4b Identify if a molecule is polar or nonpolar given a structural formula for the compound.

1. Consider the following molecules.Which compound has the most polar bonds?

A. CH4B. NCl3 C. BF3D. CS2

Answer: C


2. Which compound has the most polar bonds?

A. SiCl4

B. SSl2

C. PCl3

Answer: C

3. What type of bond is formed when a highly electronegative element and an electropositive element form a bond?

A. Nonpolar covalent B. Polar covalentC. Nonpolar ionic D. Ionic

Answer: B

C5.8A Draw structural formulas for up to ten carbon chains of simple hydrocarbons.

1. The diagram shows the structural formula of benzene. The empirical and molecular formulas of benzene are, respectively –

A. CH, C2H2

B. CH, C3H3

C. C3H3, C6H6

D. CH, C6H6

Answer: D

C5.8B Draw isomers for simple hydrocarbons.

1. How would the bond strength of the double bond in an alkene compare to that of a single bond in the corresponding alkane?

A. The double bond would have the same strength as the single bond.B. The double bond would have less than twice the strength than the corresponding

single bond.C. The double bond would have exactly twice the strength of the single bond.D. The double bond would have more than twice the strength of the single bond.

Answer: B


2. The enolate anion of ethyl acetate has the following structure:

Which resonance structure has the lowest energy?

A.

B.

Answer: A

C5.8C Recognize that proteins, starches, and other large biological molecules are polymers.

1. Which of the following best describes polymers?A. large molecules made up of smaller organic molecules, called isomers,

that are linked together to form new bondsB. hydrocarbons in which some of the carbon atoms form double or triple

covalent bonds with other carbon atomsC. proteins, carbohydrates, and lipidsD. large molecules made up of smaller organic molecules, called monomers,

that are linked together to form new bonds

Answer: D

2. Both lipids and carbohydrates are hydrocarbon polymers that contain carbon, hydrogen, and oxygen. What quality distinguishes lipids from carbohydrates?

A. Lipids are a type of protein while carbohydrates are a type of sugar.B. For a balanced diet, you should eat more lipids than carbohydrates.C. Lipids are a more concentrated source of energy for the body than

carbohydrates.D. Your body can break down lipids but cannot break down carbohydrates.

Answer: C


Unit 8ADVANCED BONDING CONCEPTS

C4.3c Compare the relative strengths of forces between molecules based on the melting point and boiling point of the substances.

1. The temperature at which a liquid boils is the boiling point of the liquid. The boiling point is an indication of the submicroscopic forces that hold the matter in the liquid state. Water, H2O, boils at 100oC. Ethanol, C2H6O boils at 78oC. Ammonia, NH3, boils at -33oC. Which one of the following ranks the submicroscopic forces in these liquids from the strongest to the weakest?

A. Water > Ammonia > EthanolB. Ammonia > Ethanol > WaterC. Ethanol > Water > AmmoniaD. Water > Ethanol > Ammonia

Answer: D

C4.3d Compare the strength of the forces of attraction between molecules of different elements. (For example, at room temperature, chlorine is a gas and iodine is a solid.)

1. At room temperature, chlorine exists as a gas, bromine exists as a liquid, and iodine exists as a solid. The physical states of these elements indicate that melting point —

A. decreases from top to bottom with group17 elementsB. is independent of periodic positionC. increases from top to bottom within group 17 elements _D. is constant within group 17 elements

Answer: C

C4.3e Predict whether the forces of attraction in a solid are primarily metallic, covalent, network covalent or ionic based upon the elements’ location on the periodic table.

1. Which of the pairs of elements listed will NOT form an ionic solid?

A. barium and iodineB. calcium and oxygenC. lithium and chlorineD. oxygen and hydrogenE. cesium and fluorine

Answer: D


2. Which combination(s) is/are most likely to form predominantly covalent bonds? A. Potassium-ChlorineB. Oxygen-OxygenC. Hydrogen-CarbonD. More than one of the aboveE. None of the above

Answer: D

3. Which combination(s) is/are most likely to form predominantly ionic bonds?

A. Potassium-ChlorineB. Oxygen-OxygenC. Hydrogen-CarbonD. More than one of the above E. None of the above

Answer: A

4. In reactions to form ionic compounds, metals generally:

A. lose electronsB. gain electronsC. become non-metalsD. do not react

Answer: A

C4.3f Identify the elements necessary for hydrogen bonding (N, O, F).

1. What type of bond is used when hydrogen is bonded to oxygen in a water molecule?


Answer: A


2. What type of bond forms between carbon and hydrogen in a molecule of methane?


Answer: A

C4.3g Given the structural formula of a compound, indicate all the intermolecular forces present (dispersion, dipolar, hydrogen bonding).

1. What type of bond is formed when a highly electronegative element and an electropositive element form a bond?

A. Nonpolar covalent B. Polar covalent C. Nonpolar ionic D. Ionic

Answer: D

2. Which of the following molecules has a net dipole?

A. NH3 B. SF6 C. CCl4 D. SiH4

Answer: A


C4.3h Explain properties of various solids such as malleability, conductivity, and melting point in terms of the solid’s structure and bonding.

A chemist performs the same tests on two white crystalline solids, A and B. The results are show in the table below.

Solid A Solid B

Melting Point High, 801≈C Low, decomposesat 186≈C

Solubility in H2O(grams per 100.0 g in H2O at 0≈C)

35.7 3.2

Electrical Conductivity(in aqueous solution)

Good conductor Nonconductor

1. The results of these tests suggest that:

A. both solids contain only ionic bondsB. both solids contain only covalent bondsC. solid A contains only covalent bonds and solid B contains only ionic bondsD. solid A contains only ionic bonds and solid B contains only covalent bonds

Answer: D

2. Which element has the highest electrical conductivity when a large amount of it is present?

A. MgB. HC. HeD. Cl

Answer: A

C4.3i Explain why ionic solids have higher melting points than covalent solids. (For example, NaF has a melting point of 995°C while water has a melting point of 0° C.)

1. Which of the following observations about a certain pure solid would indicate most strongly that the solid is ionic?

A. Its water solution is a good conductor of electricity. B. It is composed of small white crystals. C. It has a density greater than 1.0 gram/cm3. D. It has a high melting point.

Answer: D


C5.4c Explain why both the melting point and boiling points for water are significantly higher than other small molecules of comparable mass (e.g., ammonia and methane).

1. Water has unique properties because of its interparticle forces. Thinking about this property, why does water have significantly higher melting and boiling points than other small molecules of comparable mass such as ammonia and methane?

Answer: Water has strong interparticle forces, (strength of the attractive forces between particles) which are held together tightly. A large amount of energy is required to break these forces, allowing the particles to evaporate. This results in a low vapor pressure and a high normal boiling point.

C5.4d Explain why freezing is an exothermic change of state.

1. Which of the following processes is exothermic?

A. water boilingB. photosynthesisC. ice meltingD. a car battery charging

Answer: C

2. Which phase change is exothermic?

A. H2O(s) H2O(l)

B. H2O(l) H2O(s)

C. H2O(s) H2O(g)

D. H2O(l) H2O(g)

Answer: A

3. Distinguish between exothermic and endothermic reactions, and give an example of each, using the three states of water (H2O).

C5.4e Compare the melting point of covalent compounds based on the strength of IMFs (intermolecular forces)

1. Which of these substances has the lowest melting point?

A. LiBrB. CaOC. COD. CH3OH

Answer: C


QUARTER THREEUnit 5

NOMENCLATURE AND FORMULA STOICHIOMETRY

C4.1a Calculate the percent by weight of each element in a compound based on the compound formula.

1. Calculate the percent composition of Calcium in CaCO3 .

A. 30%B. 40%C. 60% D. 80%

Answer: B

2. Calculate the percent composition of Phosphorus in P2O5.

A. 25.8%B. 30.2%C. 38.7%D. 43.6%

Answer: D

3. Calculate the percent composition of Sodium in NaN3.

A. 28%B. 35%C. 42%D. 58%

Answer: B

4. Calculate the percent composition of Nitrogen is NH3.

A. 45.7%B. 56.2%C. 68.5%D. 82.3%

Answer: D


C4.1b Calculate the empirical formula of a compound based on the percent by weight of each element in the compound

1. Calculate the empirical formula of a compound composed of 38.67 % C, 16.22 % H, and 45.11 %N.

Answer: 3.220 mole C, 16.09 mole H, 3.219 mole N - therefore: CH5N

2. A compound is 43.64 % P and 56.36 % O. What is the empirical formula?

Answer: P2O5

3. Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?

Answer: C4H5N2O1

C4.1c Use the empirical formula and molecular weight of a compound to determine the molecular formula.

1. A compound is known to be composed of 71.65 % Cl, 24.27% C and 4.07% H. Its molar mass is known (from gas density) to be 98.96 g. What is its molecular formula?

Answer: Cl2C2H4

2. The empirical formula for a substance is CH2. If the molecular mass of the substance is 56, the molecular formula is –

A. C2H4

B. C3H6

C. C4H8

D. C5H10

Answer: C

C4.2A Name simple binary compounds using their formulae.

1. What is the name of the compound, Na2S:

A. Sodium sulfateB. Sodium sulfideC. Sodium persulfateD. Sodium trisulfide

Answer: B


2. What is the name of the compound, BaO:

A. Barium HydroxideB. Barium dioxideC. Barium oxideD. Barium oxylate

Answer: C

3. What is the name of the compound, CsCl:

A. Cesium chlorideB. Cesium dichlorideC. Cesium chlorateD. Cesium trichloride

Answer: A

4. What is the name for the compound, Al2O3

A. Aluminum dioxideB. Aluminum oxideC. Aluminum oxylateD. Aluminum dioxide

Answer: B

C4.2B Given the name, write the formula of simple binary compounds.

1. What is the formula for Calcium fluoride?

A. CaFB. CaF2

C. Ca2F3

D. Ca3F4

Answer: B

2. What is the formula for Potassium phosphide?

A. KP2

B. K2P3

C. K3PD. K3P2

Answer: C


3. What is the formula for Lithium oxide:

A. LiOB. LiO2

C. Li2OD. LiO3

Answer: C

4. What is the formula for Magnesium chloride:

A. MgCl2

B. Mg2ClC. Mg2Cl3

D. MgCl

Answer: A

C4.2c Given a formula, name the compound.

1. The chemical name of Mg3(PO4)2 is:

A. Magnesium phosphateB. Magnesium phosphideC. Magnesium phosphiteD. Magnesium phosphotate

Answer: A

2. The chemical name of Cd(MnO4)2 is:

A. Cadmium manganateB. Cadmium manganiteC. Cadmium phosphateD. Cadmium permanganate

Answer: D

3. The chemical name of Al(NO3)3 is:

A. Aluminum nitrateB. Aluminum nitrideC. Aluminum nitriteD. Aluminum trinitrate

Answer: A


4. Which of the following is the formula for Ammonium hydroxide?

A. Al2O3

B. AmO2

C. NH3OHD. NH4OH

Answer: D

C4.2d Given the name, write the formula of ionic and molecular compounds.

Which of the following is the balanced chemical equation for the reaction shown above?

C4.2e Given the formula for a simple hydrocarbon, draw and name the isomers.

1. The formula for pentane is C5H12, draw two isomers.

Answer: 2-methylbutane and 2,2 dimethylpropane

C4.6a Calculate the number of moles of any compound or element given the mass of the substance.

1. If you weigh out 20 grams of NaOH and place it in a 250 ml volumetric flask. Fill the flask up to the designated mark on the neck. Determine the Molar concentration of this solution.

Answer: 1M NaOH

2. A sample of NaNO3 weighing 8.5 grams is placed in a 500 ml volumetric flask and distilled water was added to the mark on the neck of the flask. Calculate the Molarity of the resulting solution.

Answer: 0.2M NaNO3


3. How would you prepare 100 mL of 0.25 M KNO3 solution?

Answer: Add 2.7 g of KNO3 to 100 ml distilled water

4. A chemist dissolves 98.4 g of FeSO4 in enough water to make 2.000 L of solution. What is the molarity of the solution?

Answer: (98.4g) / (2,000 g H2O)(1 mol) / (151.8 g/mol) = 0.324 molarity

C4.6b Calculate the number of particles of any compound or element given the mass of the substance.

1. How many atoms are contained in 97.6 g of platinum (Pt)?

A. 5.16 × 1030

B. 3.01 × 1023

C. 1.20 × 1024

D. 1.10 × 1028

Answer: B

2. How many atoms are in a chromium sample with a mass of 13 grams? A. 1.5×1023

B. 3.3×1023

C. 1.9×1026

D. 2.4×1024

Answer: A


Unit 6EQUATIONS & STOICHIOMETRY

C3.4A Use the terms endothermic and exothermic correctly to describe chemical reactions in the laboratory.

1. Which of the following are examples of an exothermic chemical reaction? Check all that apply.

A. photosynthesisB. burning a piece of woodC. freezing water into iceD. none of the above

Answer: B

C3.4c Write chemical equations including the heat term as a part of equation or using H notation.

1. Which of the following equations correctly shows the reaction of methane and oxygen?

A. CH4 + O→CH4O ∆H -890.3 KJ/molB. CH4 +O2→CO2 + H2O ∆H 890.3 KJ/molC. CH4 +O2→CO2 + H2O ∆H -890.3 KJ/molD. CH4 + O→CH4O ∆H 890.3 KJ/mol

Answer: C

C5.2A Balance simple chemical equations applying the conservation of matter

1. __CCl4 +__ HF =>__ CCl2F2 +__ HCl

Answer: 1CCl4 + 2HF 1CCl4 + 2HCl

2. __ HCl +__ Na3PO4 =>__ H3PO4 +__ NaCl

Answer: 3HCl + Na3PO4 H3PO4 + 3NaCl

3. __A1 + __C12 =>__A1C13

Answer: 2Al + 3Cl2 3Cl2 + 2AlCl3

4. __HNO3 +__ NO =>__ NO2 +__ H2O

Answer: 2HNO3 + 1NO 3NO2 + 1H2O


5. C3H6 + O2 CO2 + H2O When the equation above is balanced and all coefficients are reduced to their lowest whole-number values, the coefficient for H2O is:

A. 2B. 3C. 4D. 6

Answer: D (2C3H6 + 9O2 6CO2 + 6H2O)

C5.2B Distinguish between chemical and physical changes in terms of the properties of the reactants and products.

1. When performing an experiment which of the following would ensure that chemical changes have occurred?

A. the product is a different shape than the reactantB. the product and the reactant are different physical statesC. the product has no odor the reactant has a strong odorD. the product has a greater volume than the reactant

Answer: B

2. Which of the following is a way to recognize a chemical change?

A. the substance has changed shapeB. the substance has changed from liquid to gasC. a new substance has formedD. the change is reversible

Answer: C

3. Which of the following indicates a chemical change has occurred?

A. a gas forms when two liquids are mixedB. a powder dissolves in a liquidC. a solid has melted into a liquidD. a piece of clay has a new shape

Answer: A


C5.2d Calculate the mass of a particular compound formed from the masses of starting materials.

1. For the following balanced chemical reaction: H2O + C => H2 + COHow many grams of H2 will be produced by the reaction of 82.48 grams of H2O?

Answer: 9.16g H2

2. For the following balanced chemical reaction: 2Na + 2H2O => 2NaOH + H2

How many grams of NaOH will be produced, if 0.39 grams of H2 are produced?

Answer: 15.6g NaOH

3. For the following balanced chemical reaction: CaCN2 + 3H2O => CaCO3 + 2NH3

How many grams of NH3 will be produced, if 12.23 grams of CaCO3 are produced?

Answer: 4.16 g NH3

4. For the following balanced chemical reaction: 4KO2 + 2CO2 => 2K2CO3 + 3O2

How many grams of O2 will be produced by the reaction of 72.78 grams of KO2?

Answer: 24.6 g O2

C5.2e Identify the limiting reagent when given the masses of more than one reactant.

1. For the following balanced chemical reaction, if 84.32 grams of Na2O2 are reacted with 18.22 grams of H2O, what is the limiting reagent?

2Na2O2 + 2H2O => 4NaOH + O2

A. Na2O2

B. H2O

Answer: B

2. For the following balanced chemical reaction, if 24.46 grams of CaCN2 are reacted with 19.32 grams of H2O, what is the limiting reagent?

CaCN2 + 3H2O => CaCO3 + 2NH3

A. CaCN2

B. H2O

Answer: B


3. For the following balanced chemical reaction, if 68.13 grams of C8H18 are reacted with 296.28 grams of O2, what is the limiting reagent?

2C8H18 + 25O2 => 16CO2 + 18H2O

A. C8H18

B. O2

Answer: A

4. For the following balanced chemical reaction, if 23.04 grams of C4H10 are reacted with 62.65 grams of O2, what is the limiting reagent?

2C4H10 + 13O2 => 8CO2 + 10H2O

A. C4H0

B. O2

Answer: B

C5.2f Predict volumes of product gases using initial volumes of gases at the same temperature and pressure.

1. Blaise decides to build a barometer. He can’t find any mercury in his workshop and decides to use water instead. Assume that the density of water is 1.00×103 kg/m3. If the atmospheric pressure is 7.60×102 torr, how tall must his barometer be in order to obtain an accurate reading?

Answer: First, convert the atmospheric pressure to Pascals. 7.60×102 torr = 101, 325 Pa. Now that all the variables are in SI units, rearrange P = ghρ to P/(gρ) = h and plug the variables into the equation.

=

= 10.3 m


Thus Blaise’s barometer must be taller than 10.3 meters. A comparable mercury barometer would be 0.76 m tall.

2. In a fit of inspiration and hot air, you have blown the world’s largest balloon. Your two cousins, Bongo the 300 lb. gorilla and Jeeves the 70 lb. weakling, both want to climb to the top of your balloon. When Bongo goes, he goes in style. He wants to lay down on his 1 m by 5 m air bed at the summit. Jeeves proposes to bounce on the top of the balloon on his pogo stick, whose head has an area of 0.001 m2. Assume that the masses of the bed and pogo stick are negligible, and that their occupants’ weights are evenly distributed upon them.

You know that your balloon can sustain 200 more kPa of pressure on its surface before it pops. Assuming both can make it to the top without damaging the balloon (or themselves), which cousin(s) should you allow to climb?

Answer: P = F/A, so the first thing we need to do is convert everything to the appropriate units. Let’s use SI units. 1 lb = 0.454 kg, and F = (mass)×(9.8 m/s2), so Bongo and Jeeves exert forces of 1330 and 311 Newtons, respectively. P = F/A, so Bongo has a pressure of PBongo =

= 270 Pa. Jeeves exerts a pressure of PJeeves = (311 N)×(0.001 m2) = 310 kPa. You should allow Bongo on, but not Jeeves.

3. A sample of an element has a volume of 78.0 mL and a density of 1.85 g/mL. What is the mass of the sample (measured in grams)?

Answer: .025 g

4. The graph shows the pressure of an ideal gas as a function of its volume. According to the graph, increasing the volume from 100 mL to 150 mL –

A. decreases the pressure by 80 kPaB. decreases the pressure by 160 kPaC. increases the pressure by 80 kPaD. increases the pressure by 180 kPa

Answer: A


C5.6b Predict single replacement reactions.

1. Given the reactants: Fe + CuSO4, predict the products in a single replacement reaction:

A. Fe2SO4 + CuB. FeSO4 + 2CuC. FeSO4 + CuD. Fe3SO4 + Cu

Answer: C

2. Given the reactants Cd + 2HCl predict the products in a single replacement reaction:

A. CdCl + H2

B. CdCl2 + H2

C. CdCl3 + 2H2

D. CdCl + 2H2

Answer: B

3. The figure below represents a reaction.

What type of reaction is shown?

A. synthesisB. decompositionC. single displacementD. double displacement

Answer: A

4. The equation, Fe + CuSO4 → FeSO4 + Cu is an example of:

A. decomposition reactionB. oxidationC. single replacementD. double replacement

Answer: C


Unit 9THERMOCHEMISTRY SOLUTIONS

C2.1c Compare qualitatively the energy changes associated with melting various types of solids in terms of the types of forces between the particles in the solid.

1. The melting point for salt is 800.8°C and the melting point for sugar is 146°C. What can be concluded from this information?

A. The intermolecular forces in sugar release more energy than those of saltB. Sugar molecules require more energy than salt to break apartC. The intermolecular forces of salt release more energy than those of sugarD. Salt molecules require more energy than sugar to break apart

Answer: D

C2.2d Explain convection and the difference in transfer of thermal energy for solids, liquids, and gases using evidence that molecules are in constant motion.

1. The illustration below shows a hot-air balloon. The pilot can change the altitude of the hot-air balloon by changing the temperature of the gas inside the balloon. When the gas is heated, the balloon rises.

Which of the following best explains this phenomenon?

A. Heating the gas reduces its pressure. B. Heating the gas decreases its density. C. Heating the gas decreases its molecular

motion. D. Heating the gas reduces the frequency of the

gas molecules’ collisions.

Answer: B


C3.1c Calculate the ΔH for a chemical reaction using simple coffee cup calorimetry.

1. When 12.29 g of finely divided brass (60% Cu, 40% Zn) at 95.0oC is quickly stirred into 40.00 g of water at 22.0oC in a calorimeter, the water temperature rises to 24.0oC. Find the specific heat of brass.

A. 334.7 JB. -334.7 JC. 0.38 J g-1 oC-1.D. 0.98 J g-1 oC-1.

Answer: B

C3.1d Calculate the amount of heat produced for a given mass of reactant from a balanced chemical equation.

Figure A Figure B

1. Which of the figures is depicting an endothermic reaction?

A. Figure AB. Figure BC. Both Figures A and BD. Neither Figure A nor B

Answer: B

Use this balanced equation to answer the following question: Fe2O3 + 2Al → Al2O3 + 2Fe

2. 18.2 g of Aluminum (Al) reacts with Fe2O3 according to the reaction: ∆Hrxn = -853.2 kJ/mol Fe2O3. How much heat is produced?

Answer: 18.2g(1m Al/27g Al)(-853.2kJ/1mAl) = -575.12kJ heat


C3.4g Explain why gases are less soluble in warm water than cold water.

1. When stirred in 30°C water, 5 g of powdered potassium bromide, KBr, dissolves faster than 5 g of large crystals of potassium bromide. Which of the following best explains why the powdered KBr dissolves faster?

A. Powdered potassium bromide exposes more surface area to water molecules than large crystals of potassium bromide.

B. Potassium ions and bromide ions in the powder are smaller than potassium ions and bromide ions in the large crystals.

C. Fewer potassium ions and bromide ions have been separated from each other in the powder than in the crystals.

D. Powdered potassium bromide is less dense than large crystals of potassium bromide.

Answer: A

C4.7a Investigate the difference in the boiling point or freezing point of pure water and a salt solution.

1. At which temperature do KBr and KNO3 have the same solubility?

A. 27°CB. 48°CC. 65°CD. 80°C

Answer: B


The freezing point of a liquid is the temperature at which a liquid changes to a solid. The table below lists the freezing point of four liquids.

Table adapted from Leonard Bernstein, Martin Schachter, Alan Winkler, and Stanley Wolfe, Concepts and Challenges in Physical Science, 3rd edition. ©1991 by Globe Book Company.

2. Which of the liquids in the table would be LEAST likely to freeze outside during the winter in Michigan?

A. waterB. glycerinC. seawaterD. ethyl alcohol

Answer: D

3. An ice-skating rink has tubes under its floor to freeze the water. Salt water is cooled well below the freezing point of water and pumped through the tubes to freeze the water in the rink. Why can the salt water be cooled so low without freezing?

A. Salt has a very low freezing point.B. Adding salt to water lowers its freezing point.C. Movement of the salt water through the tubes keeps it in the liquid state. D. The salt water is constantly absorbing energy from its surroundings.

Answer: B

C5.4A Compare the energy required to raise the temperature of one gram of aluminum and one gram of water the same number of degrees.

1. In order to raise the temperature of 200g of water and 200g of aluminum 5 degrees each:

A. The same amount energy is required for eachB. More energy is needed to raise the temperature of waterC. More energy is needed to raise the temperature of aluminumD. It is impossible to determine which needs more energy

Answer: C


C5.4B Measure, plot, and interpret the graph of the temperature versus time of an ice-water mixture, under slow heating, through melting and boiling.

1. Between points 2 and 3, energy is being used to:

A. heat waterB. evaporate waterC. heat water vaporD. melt ice

Answer: D

C5.5e Relate the melting point, hardness, and electrical and thermal conductivity of a substance to its structure.

1. Rank the following compounds by increasing melting point:

(I)C2H6, (II)C2H5OH, (III)C2H5F

A. I, II, IIIB. II, III, IC. III, I, IID. II, I, IIIE. None of the above

Answer: A


QUARTER FOURUnit 10

ACIDS-BASES

C5.7A Recognize formulas for common inorganic acids, carboxylic acids, and bases formed from families I and II.

1. Which of the following is an example of a base

A. HClB. HC2H3O2

C. NaOHD. H2O

Answer: C

2. Which of the following is and example of an acid

A. HClB. NaOHC. Mg(OH)2

D. H2O

Answer: A

C5.7B Predict products of an acid-based neutralization

1. If it takes 1.0 gram of solid NaOH (a strong base) to change the pH of a solution containing a strong acid from 1 to 2, how much solid NaOH (in g) would be needed to change the pH of this solution from 2 to 3? You may assume that there is no change in the volume of the solution.

Answer: 1.0 x 10-1

2. Calculate the pH of a solution prepared by dissolving 2.13 g of NaOH (a strong base) in enough water to produce 100 mL of solution.

Answer: 13.7:

3. Calculate the concentration (in M) of HCl that would be needed to make an aqueous solution of HCl with a pH of 4.3.

Answer: 5.0 x 10-5


4. Calculate the pH of a solution prepared by dissolving 5.9 x 10-5 moles of HCl in 1.0 L of pure water.

A. 3.77 B. 4.77 C. 4.23 D. 10.23

Answer: D

Analyzing a Titration Experiment Given: Becky is conducting an experiment in which she measures 30 ml of an acid of unknown concentration. She adds an indicator solution to it, then slowly adds a basic solution of a known concentration from a buret until the acid solution is neutralized. She measure and records the pH of the solution periodically throughout the experiment. 1. What laboratory technique is Becky using?

A. electrolysisB. distillationC. titrationD. electrophoresis

Answer: C 2. Which reaction would fit this experiment?

A. CuSO4 + 2NaCl --> Na2SO4 + CuCl2

B. 2H2O --> 2H2 + O2

C. KOH + HCl --> H2O + KClD. 2H2O + BaSO4 --> H2SO4 + Ba(OH) 2

E. All of the reactions listed

Answer: C 3. All bonds holding atoms together within both the reactants and products in Becky’s

experiment are of what type?

A. MetallicB. IonicC. CovalentD. Ionic and CovalentE. Hydrogen

Answer: D


C5.7C Describe tests that can be used to distinguish an acid from a base.

1. Which is the best indicator for giving a color change at the equivalence point?

A. Bromocresol greenB. Indigo carmineC. Neutral redD. Phenolphthalein _

Answer: D

C5.7D Classify various solutions as acidic or basic, given their pH.

1. The table below contains data for water samples from four sources.

Nancy analyzed water samples from several sources: rainfall, a nearby creek, a swimming pool, and her kitchen faucet. She recorded her data in the table.


Which sample was most acidic?

A. rainB. creekC. poolD. faucet.

Answer: A

2. If we gradually increase the concentration of a strong acid in water without changing the volume significantly, then the:

A. concentrations of hydronium and hydroxide ion will both increase. B. concentrations of hydronium and hydroxide ion will both decrease. C. hydronium ion concentration will decrease and the hydroxide ion concentration

will increase. D. hydronium ion concentration will increase and the hydroxide ion concentration

will decrease. E. concentrations of hydronium and hydroxide ion will be unaffected.

Answer: D

5.7E Explain why lakes with limestone or calcium carbonate experience less adverse effects from acid rain than lakes with granite beds.

1. Why do limestone lakes experience less adverse effects from acid rain than lakes with granite beds?

A. limestone is more acidic and more likely to reactB. limestone is more basic and neutralizes the rainC. granite is more basic and more likely to reactD. granite is more acidic and neutralizes the rain.

Answer: B

C5.7f Write balanced chemical equations for reactions between acids and bases and perform calculations with balanced equations.

1. How many liters of 2.5 M HCl are required to exactly neutralize 1.5 liters of 5.0 M NaOH?

A. 1.0 LB. 2.0 LC. 3.0 LD. 4.0 L

Answer: C


2. How many milliliters of 0.600 M H2SO4 are required to exactly neutralize 100 milliliters of 0.300 M Ba(OH)2?

A. 25.0 mLB. 50.0 mLC. 100.0 mLD. 200.0 mL

Answer: B 3. How many milliliters of a 0.20 M KOH are needed to completely neutralize 90.0 milliliters

of 0.10 M HCl?

A. 25 mLB. 45 mLC. 90 mLD. 180 mL

Answer: B

C5.7g Calculate the pH from the hydronium ion or hydroxide ion concentration.

1. What is the pH of a solution that contains 2.5 x 10-5 M hydronium ion?

Answer: -log(2.5 x 10 – 5) = pH 4.602

2. What is the molar concentration of hydroxide ion, [OH-], in a water solution that contains 2.5 x 10-4 hydronium ion (H3O+)?

Answer: 4 x 10 – 11 (OH-)

3. What is the molar hydroxide ion concentration, [OH-], in a solution with pH = 7.63?

Answer: inverse log (7.63) = 2.3 x 10 –8

4. For each hydrogen/hydronium/hydroxide ion concentration given, determine the pH and indicate whether the sample is acidic, basic, or neutral.

A. [H+] = 1 X 10-10 MB. [H+] = 1 X 10-2 MC. [OH-] = 1 X 10-8 MD. [H3O+] = 1 X 10-7 M

Answer: a. 10 = basicb. 2 = acidicc. 8 = basicd. 7 = neutral


C5.7h Explain why sulfur oxides and nitrogen oxides contribute to acid rain. Angie had heard that acid rain was due to coal burning power plants. these power plants heat water to make steam that is then used to generate electricity. During the combustion (burning) process, sulfur in the coal combines with oxygen to make sulfur dioxide (SO2). Upon leaning the power plant smokestack, the sulfur dioxide combines with water vapor in the atmosphere to make sulfurous acid (H2SO3). Angie wanted to investigate the acidity of rain samples downwind from a coal burning power plant. She collected 100 milliliter (mL) rainwater samples from the following locations:

Angie added 1 drop of phenolthalein solution to each rainwater sample to use as an endpoint indicator. She then titrated each beaker with 0.1 M (molar) sodium hydroxide (NaOH) solution until the phenolthalein turned pink in color. She gathered the following results from her experiment:

Sample mL of NaOH needed toreach endpoint

A 1.8 2.4C 5.6

1. All of the following are reasons why acid rain concerns people. What would be the major

reason the results of Angie’s experiment would be of interest to an agricultural community that was situated 75 kilometers downwind from the power plant?

A. Acid rain could increase the yield (pounds/acre) of crops by releasing nutrients locked up in the soil.

B. Acid rain could have a long-term effect on materials and people’s health.C. Acid rain could affect the population’s health for a few days if a bad storm

occurred.D. Acid rain could cause weaknesses in concrete poured when it rained, creating

structural problems in new buildings.E. Acid rain could slowly decay the limestone found in old buildings.

Answer: B


2. Why did Angie use sodium hydroxide (NaOH) to titrate the rain water samples?

A. Because phenolthalein is a specific indicator for NaOH and not other bases.B. Because NaOH would neutralize the acid present, allowing the indicator to

change color.C. Because the large amount of acid present required a strong base.D. Because NaOH, being a strong base, is not affected by the dust particles found in

rain water.

Answer: B 3. Based on the results, how much acid did sample "C" contain compared to normal rainwater?

A. About 5.6 times as much as normal rain.B. About 3.8 times as much as normal rain.C. About 3.2 times as much as normal rain.D. Cannot be compared to normal rainwater because of inadequate sampling.

Answer: E


Unit 11REDOX-EQUILIBRIUM

C5.3a Describe equilibrium shifts in a chemical system caused by changing conditions (Le Chatelier’s Principle).

1. The reaction for the synthesis of ammonia: N2(g) + 3H2(g) ↔ 2NH3(g) is exothermic. Increasing the temperature applied to the system:

I. increases the amount of NH3.II. decreases the amount of NH3.

III. changes the value of Keq.IV. does not change the value of Keq.

A. I and III only B. II and III only C. I and IV only D. II and IV only

Answer: B

2. Consider the system below at equilibrium. Which of the following changes will shift the equilibrium to the right?

N2(g) + 3H2(g) → 2NH3(g) + 92.94 kJ

I. Increasing the temperature II. Decreasing the temperatureIII. Increasing the pressure on the system

A. I only B. II only C. III only D. I and III E. II and III

Answer: E


3. The value of the equilibrium constant, K, is dependent on:

I. The temperature of the systemII. The concentration of the reactants

III. The concentration of the productsIV. The nature of the reactants and products

A. I, IIB. II, IIIC. III, IVD. I and IVE. I, II, and IV

Answer: D

C5.3b Predict shifts in a chemical system caused by changing conditions (Le Chatelier’s Principle).

1. Which action will drive the reaction to the right?

A. heating the equilibrium mixtureB. adding water to the systemC. decreasing the oxygen concentrationD. increasing the system’s pressure

Answer: D

2. The reaction shown below occurs inside a closed flask. What action will shift the reaction to the left?

A. pumping CO gas into the closed flaskB. raising the total pressure inside the flaskC. increasing the NO concentration in the flaskD. venting some CO2 gas from the flask

Answer: B


C5.3c Predict the extent reactants are converted to products using the value of the equilibrium constant.

1. If at a given temperature the equilibrium constant for the reaction

is Kp, the equilibrium constant for the reactioncan be represented as:

A. B.

C. D. E.

Answer: C

2. The value of the equilibrium constant, K, is dependent on:

V. The temperature of the systemVI. The concentration of the reactants

VII. The concentration of the productsVIII. The nature of the reactants and products

A. I, IIB. II, IIIC. III, IVD. I and IVE. I, II, and IV

Answer: D

C5.6a Balance half-reactions and describe them as oxidations or reductions.

1. __KClO4 +__H3ASO3 => __KCl +__H3AsO4

Answer: 1, 4, 1, 4, reduction

2. __H2S + __MnO4- + __H+ => __S + __Mn2++ __H2O

Answer: 1, 1, 2, 1, 1, 4, reduction


C5.6c Explain oxidation occurring when two different metals are in contact.

1. Identify the oxidizing agent in the following reaction.

Cu2+ (aq) + Mg (s) → Cu (s) + Mg2+ (aq)

Answer: Cu loses e- (oxidizes)

2. Write the equation for the redox reaction that occurs when a piece of iron (Fe) metal is dipped in a solution of copper (II) sulfate (CuSO4).

Answer: Fe2+ = CuSO4 → CuII2+ + SO4

3. Identify which reactant is reduced and which is oxidized in each of the following reactions.

a. 2Al (s) + 3Cu2+ (aq) → 2Al3+ (aq) + 3Cu (s)b. 2Cr3+ (aq) + 3Zn(s) → 2Cr(s) + 3Zn2+ (aq)

Answers: a. Al oxidizes and Cu reducesb. Cr oxidizes and Zn reduces

4. Silver (Ag) is a commonly used metal that easily tarnishes. The silver reacts with hydrogen sulfide (H2S) in the air. This reaction produces silver sulfide (Ag2S), a dull brownish compound, and hydrogen gas (H2).

a. Write a balanced equation for this reaction and identify the reaction type. b. Explain why silver tarnishes faster in a heated room than in an unheated room. c. Describe how you could slow down this reaction or prevent it from occurring.

Answers: a. 2Ag + H2S Ag2S + H2

b. Silver tarnishes faster in a heated room than an unheated room because particles have more kinetic-molecular energy and as a result have more collisions.

c. The reaction could be slowed or prevented by lowering the temperature in the room.


C5.6d Calculate the voltage for spontaneous redox reactions from the standard reduction potentials.

1. Using the following equation calculate the voltage for the redox reactionZn Zn2+ + 2e- 0.76vCu Cu2+ + 2e- -0.34v

A. 1.95vB. -1.95vC. .42vD. -.42v

Answer: C

C5.6e Identify the reactions occurring at the anode and cathode in an electrochemical cell.

1. Which of the following half-cell reactions describes what is happening at the anode in the diagram?

A. Zn Zn2+ + 2e -

B. H2 2H+ + 2e -

C. 2Cl - Cl2 + 2e -

D. SO4- S + 2O2 + 6e -

E. 2H+ + 2e- H2

Answer: A


Unit 12THERMODYNAMICS

C2.2e Compare the entropy of solids, liquids, and gases.

1. Consider the combustion reaction between octane and oxygen in the cylinder of an automobile. The reaction generates CO2, H2O and heat, and at the same time the expanding gases push up the piston thereby powering the car. If the system is defined as the contents of the cylinder (not including the piston or the cylinder) which of the following statements is true for the system in this process?

A. The change in internal energy (ΔE) is negativeB. The change in internal energy (ΔE) is positiveC. The change in internal energy (ΔE) is zeroD. Not enough information is given to unambiguously determine the sign of ΔE

Answer: A

2. Which of the following processes is likely to have the most positive entropy change per mole under standard conditions?

A. Freezing of water to ice B. A sample of helium gas compressed to half its original volume at constant

temperature C. The vaporization of liquid trichloromethane D. The cooling of graphite from 77 K to 4.2 K E. The transformation of graphite to diamond at room temperature

Answer: C

C2.3a Explain how the rate of a given chemical reaction is dependent on the temperature and the activation energy.

1. Each of two solutions are mixed separately, and both solutions are found to be the same temperature. The two solutions are mixed, and a thermometer shows that the mixture’s temperature has decreased. Which of the following statements is true?

A. The chemical reaction is exothermic. B. The chemical reaction is absorbing energy. C. The chemical reaction is releasing energy. D. The energy released could be found by multiplying the temperatures together. E. The energy absorbed by the solution is equal to the difference in temperature of

the solutions.

Answer: B


2. Which of the following events is least likely to occur with an increase in temperature for the reaction given? (With ΔH = -45.9 kJ/mol.)

A. The gas particles will move more quickly. B. The reaction will produce more ammonia in a shorter time. C. The reaction will reverse and ammonia will decompose. D. The entropy of the system will increase. E. The equilibrium constant will become smaller.

Answer: B

3. Chemists determine the activation energy for a reaction by

A. measuring product amountsB. measuring ratesC. calculating from bond dissociation energiesD. calculating from ΔH values

Answer: B

C2.3b Draw and analyze a diagram to show the activation energy for an exothermic reaction that is very slow at room temperature.

1. Which letter corresponds to the activation energy of the reaction?

A. A B. B C. C D. Y E. X

Answer: A


2. Which letter corresponds to the change in energy for the overall reaction?

A. A B. B C. C D. Y E. X

Answer: C

3. Which graph represents the reaction shown above?

A. B.

C. D.

Answer: B

C3.1a Calculate the ΔH for a given reaction using Hess’s Law.

1. The specific heat of copper is 0.4 joules/gramoC. How much heat is needed to change the temperature of a 30-gram sample of copper from 20.0oC to 60oC?

A. 1000 JB. 720 JC. 480 JD. 240 J

Answer: C


2. A chemical reaction taking place in a calorimeter causes the temperature to rise 7.5° C. The addition of 50 kJ of energy to the same calorimeter by an electrical heater increases its temperature 2.5° C. What is ΔH of the chemical reaction?

A. 16.7 kJB. -16.7 kJC. -150 kJD. 150 kJ

Answer: C

3. Given the following thermodynamic data at 25° C: 2 HCl(g) + F2(g) --> 2 HF(l) + Cl2(g); ΔH° = -988 kJ/molH2(g) + F2 (g) --> 2 HF(l); ΔH° = -1200 kJ/mol2 H2(g) + O2(g) --> 2 H2O(l); ΔH° = -572 kJ/mol

Calculate the ΔH° of the following reaction: 4HCl(g) + O2(g) --> 2 H2O(l) + 2 Cl2(g)

A. -148 kJ/molB. 996 kJ/molC. -3748 kJ/molD. -2760 kJ/mol

Answer: A

C3.1b Draw enthalpy diagrams for exothermic and endothermic reactions.

1. The diagram above is a potential energy curve for a reaction. Which number represents the effect of a catalyst on the reaction?

A. 1B. 2C. 3D. 4

Answer: B


C3.2a Describe the energy changes in photosynthesis and in the combustion of sugar in terms of bond breaking and bond making.

1. What occurs during this change?

Given the equation representing a reaction:carbon dioxide + water→ glucose + oxygen

A. Energy is absorbed and a bond is broken.B. Energy is absorbed and a bond is formed.C. Energy is released and a bond is broken.D. Energy is released and a bond is formed.

Answer: B

C3.4B Explain why chemical reactions will either release or absorb energy.

2. Write the equilibrium expression for the following reaction:

A. [A]2[B][D]

B.

C.

D.

E.

Answer: D


C3.4d Draw enthalpy diagrams for reactants and products in endothermic and exothermic reactions.

1. The illustrations show a conservation-of-mass experiment. The solution in the beaker lost mass because:

H2SO4 + 2NaOH Na2SO4 + 2H2O

A. materials have less mass at high temperaturesB. the mass of the reactants and products was less than 100 gC. sodium sulfate (Na2SO4) is lighter than airD. some of the water molecules turned into gas

Answer: D


Intermediat

C3.4e Predict if a chemical reaction is spontaneous given the enthalpy (ΔH) and entropy (ΔS) changes for the reaction using Gibb’s Free Energy, ΔG = ΔH - TΔS (Note: mathematical computation of ΔG is not required.)

1. Consider the following reaction at 50° C, given ΔH° = -57.2 kJ/mol and ΔS° = -175.8 J/mol-k:

2 NO2(g) --> N2O4(g)

A. This reaction is nonspontaneous at 50° C, but could be made spontaneous at sufficiently high T.

B. This reaction is nonspontaneous at 50° C, but could be made spontaneous at sufficiently low T.

C. This reaction is spontaneous at these conditions, but could be made nonspontaneous by increasing the T.

D. This reaction is spontaneous at these conditions, but could be made nonspontaneous by decreasing the T.

E. This reaction is spontaneous under all temperature conditions.

Answer: B

2. At constant pressure, the following reaction 2NO2 (g) → N2O4(g) is exothermic. The reaction (as written) is:

A. always spontaneous.B. spontaneous at low temperatures, but not high temperatures.C. spontaneous at high temperatures, but not low temperatures.D. never spontaneous.

Answer: B

C3.4f Explain why some endothermic reactions are spontaneous at room temperature.

The diagram shows an experiment in which crushed coal is heated in a test tube.

1. The flowchart represents what happens in this experiment.

The reaction shown above is –

A. an endothermic reactionB. an exothermic reactionC. a decomposition reactionD. a double-replacement reaction

Answer: B


quarter 1 - svsu · web viewcalculate the concentration (in m) of hcl that would be needed to make...

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