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Applied Practice in

IMF, Liquids, and Solids

AP* Chemistry Series RESOURCE GUIDE

Volume 6

*AP is a registered trademark of the College Entrance Examination Board, which was not involved in the production of, and does not endorse, this product. Pre-AP is a trademark owned by the College Entrance Examination Board.

Copyright © 2009 by Applied Practice, Ltd All rights reserved. No part except the Student Practices

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Printed in the United States of America.

APPLIED PRACTICE Resource Guide


Teacher Notes and Strategies

A Note for Teachers.............................................................. 5

Teaching Strategies ............................................................... 8

Glossary of Terms............................................................... 14

Student Practices

Multiple-Choice Questions ................................................. 19

Intermolecular Forces.................................................. 20

Changes in State .......................................................... 24

Properties of Liquids ................................................... 29

Properties and Types of Solids.................................... 34

Descriptive Chemistry................................................. 38

Free-Response Questions.................................................... 43

Answer Key and Explanations

Multiple-Choice Answer Key ............................................. 53

Multiple-Choice Answer Explanations............................... 57

Free-Response Answers and Scoring Guides ..................... 69

*AP is a registered trademark of the College Entrance Examination Board, which was not involved in the production of, and does not endorse, this product.

Teacher Notes and Strategies

for


A NOTE FOR TEACHERS

The Applied Practice in AP Chemistry series was designed for use by teachers as an instructional supplement to major units in the AP Chemistry curriculum. This series was also conceived as a resource for teachers in preparing students for the AP Chemistry Exam. As you teach each unit, your students will have the opportunity to practice and to develop those skills required on the exams. Each book in the series includes:

• Teaching notes and strategies • Glossary of terms • 75 multiple-choice questions replicating Section I of the AP Chemistry Exam • Multiple-choice answer keys and answer explanations • 6 free-response questions replicating Section II of the AP Chemistry Exam • Free-response answer keys and scoring guide

We offer a few suggestions and explanations to help you receive the maximum benefit from our materials:

1. Applied Practice booklets do not purport to duplicate exactly an Advanced Placement Examination. However, questions are modeled on those typically encountered on these exams. Thus, students using these materials will become familiar and comfortable with the format, question types, and terminology of Advanced Placement Examinations.

2. Each Applied Practice booklet focuses on one topic within the AP Chemistry

curriculum. These booklets are excellent resources for teachers and their students. Their unique format includes questions designed for use during the initial teaching of the required topics. Other questions are exceptional for the review phase of the course, as students pull the entire year together leading up to the AP Chemistry Exam. The AP exam often will require knowledge in multiple content areas on the same question.

3. You have the option of using the Applied Practice booklets for your own lesson

and test preparation or, if you so choose, students may work through an Applied Practice test booklet on their own as they progress though the course. The students can check their own answers with the answer key and read the answer explanations provided in the teacher edition, conferring with the teacher as needed.

4. The order of topics in the Applied Practice booklets has been organized to follow a logical progression that is similar to the sequence in many of the most widely selected AP chemistry textbooks. You will find that they can easily be adapted to whatever sequence you find most productive at your school.

5. The free response questions in each topic were created to provide practice questions similar to both those given in part A of the AP Chemistry Exam, which allows use of a calculator, and those given in part B, in which no calculator is allowed. In a few cases, the specific content is best assessed with a combination of both types.

6. Due to the emphasis on laboratory experience in the College Board’s AP Chemistry program, the Applied Practice booklets in AP Chemistry frequently include laboratory-based questions appropriate to the subtopic addressed. A required laboratory-based question does appear on the AP Chemistry Exam. While most Applied Practice booklets in the AP Chemistry series do contain laboratory-based free-response questions, some topics do not lend themselves to the College Board-recommended laboratory experiments. However, each Applied Practice booklet does contain multiple-choice questions related to both laboratory and descriptive chemistry. Only one of the six free-response questions included on the AP Chemistry Exam is laboratory based.

7. Each booklet includes a glossary of terms that applies to the vocabulary of that

particular topic. 8. If the teacher wishes to replicate the conditions under which students will take the

actual AP Chemistry Exam, he or she should understand the following about multiple-choice versus free-response questions when using Applied Practice booklets: When answering multiple-choice questions (AP Exam, Section I) students are not allowed the use of a calculator, and the only reference information available to them is a periodic table (with only symbol, mass number, atomic number) and a small table of abbreviations/symbols used in the questions. When answering free-response questions (AP Exam, Section II), much more information is available to the student. In addition to the periodic table, a table of standard reduction potentials in aqueous solutions and a relatively complete list of equations, constants, and abbreviations/symbols are provided.

COPYRIGHT NOTICE

The copyright law of the United States (Title 17, United States Code) governs the making of photocopies or other reproductions of copyrighted material. Reproduction of individual worksheets from this booklet, excluding content intended solely for teacher use, is permissible by an individual teacher for use by his or her students in his or her own classroom. Content intended solely for teacher use may not be reproduced, stored in a retrieval system, or transmitted in any way or by any means (electronic, mechanical, photocopying, or otherwise) without prior written permission from Applied Practice. Reproduction of any portion of this booklet for use by more than one teacher or for an entire grade level, school, or school system, is strictly prohibited. By using this booklet, you hereby agree to be bound by these copyright restrictions and acknowledge that by violating these restrictions, you may be liable for copyright infringement and/or subject to criminal prosecution.

TEACHING STRATEGIES

Following are some suggestions for using Applied Practice materials as you work to help your students develop mastery in answering AP-style questions. GENERAL The AP Chemistry Exam will be a challenge for all students. In this section, we have provided strategies for answering the questions themselves as well as some suggestions for teachers to keep in mind when presenting the material. 1. A certain amount of rote memorization is required. We offer the following

suggestions to help students prepare for this:

• Mnemonic devices are excellent tools for memorizing hard facts. For example, in Kings Play Chess On Funny Green Squares, the first letter of each word represents the divisions of the classification system in binomial nomenclature: kingdom, phylum, class, order, family, genus, and species.

• Students are also encouraged to create flashcards; a tried and true method that can help them be successful.

Items that students are required to know for the AP Chemistry Exam include: • names and formulas of polyatomic ions • names and formulas of strong acids and bases, as well as common acids and bases • the metric prefixes from pico to giga • the Greek prefixes in molecular nomenclature: mono, di, tri, tetra, penta, etc. • the solubility rules for ionic compounds • the identity of the seven diatomic elements • the phases of all elements at room temperature (25oC) • the SI quantities and their units • rules for significant figures • rules for assigning oxidation numbers to atoms in compounds

2. Students will see many graphs in both the multiple-choice and free-response

questions. In a few instances, they may be required to produce the graph from given information. Teachers should incorporate interpretation of scientific data written in graphs during instruction both in the classroom and the laboratory. The more graphs that students are required to produce during the year, the easier the questions containing this type of information will be for them on the exam.

MULTIPLE-CHOICE QUESTIONS (MC): 1. The questions in these booklets are valuable for both the initial teaching of the

concepts of chemistry and for preparing students for the style of questions and the level of difficulty encountered on the AP Chemistry Exam. In particular, some subtopics are more fundamental than others. While these subtopics are required in order for students to be able to comprehend more advanced material later on, they would not necessarily appear “as is” in a question on the AP exam. However, we have included multiple-choice questions over this material to support instruction, to provide an expanded test bank of questions, and to familiarize students with this type of assessment.

2. The question most commonly asked by students preparing for the AP exam is: “How

many multiple-choice questions should I answer in Section I?” To answer this question, we will consider what is required for a student to receive a score of “3” (which is labeled as “qualified” by the College Board and considered “passing”; thus, this score is a reasonable goal for most students.) Judging by released past exams for which full data is available, a student typically must score approximately 52–58 points out of the160 possible on Sections I and II combined in order to receive a “3.” Thus, it is reasonable to assume that to pass the AP Chemistry Exam, students need to earn 26–29 points or more from Section I. For each question answered incorrectly, a 0.25 point penalty will be subtracted from the score. Typically, if students are answering more questions correctly than incorrectly, they should continue to increase the number of questions they answer, for it takes four incorrect answers to cancel one correct answer. Most students reach a point at which answering additional questions merely lowers their score. Successful students over the years typically answer the following numbers of questions corresponding to these eventual scores:

• 35–50 questions: “3” • 45–60 questions: “4” • 55–70 questions: “5”

This formula by no means guarantees a particular score; each student is unique in his or her approach to answering questions. This is simply offered as a guideline; an average number of questions a student may choose to answer out of the 75 total questions in an effort to attain a particular achievement level.

3. Another question related to multiple-choice questions on Section I of the AP Chemistry Exam is, “Should I guess?” Because of the penalty for wrong answers (-0.25) and the fact that there is no penalty for questions not answered (just no points earned), it is advisable not to answer questions if the student has little or no knowledge of the information addressed in the question. However, if students can eliminate one or more of the answer choices, it may be a good idea to select an answer, as statistically the odds are then in their favor overall for those questions.

4. A general strategy for multiple-choice questions involves time management. Since the exam limits students’ time to 90 minutes for Section I, they will need to learn to be efficient. While some students simply work more quickly than others, virtually all students find that time is a limiting factor on the AP Chemistry Exam. The ability to read and comprehend quickly is a tremendous advantage. Another skill that will help students overcome the time factor is reading and assessing the level of difficulty of each multiple-choice question quickly. Those questions that students believe they can answer but will take quite a bit of time to analyze should be skipped over initially. Students can mark those questions and return to them later, provided there is enough time. The vast majority of questions answered correctly by students are those for which they can determine the answer almost immediately. Those questions that consume the most time are also those that are most likely to be answered incorrectly.

5. As students approach AP-level question sets, teachers should prepare students to be

satisfied with a lower percentage of correct answers than is typical for them up to this point in their educational experience. Most schools around the country consider 90% an “A” grade, and many AP chemistry students are disappointed when earning test scores in the 70% or 80% range. However, the AP Chemistry Exam is scored in such a way that a successful “passing” score is between 34% and 50%, and a “5”—the highest score possible—is achieved with any point total above approximately 63%. Thus, it will take some time for students to adjust psychologically. Students should understand that a typical passing rate nationwide on the AP Chemistry Exam is approximately 56% of those taking the exam. It is definitely an achievement for a high school student to find success with a “3,” “4,” or “5” on one of the most challenging AP exams. One way to help students cope with a lower percentage of correctly-answered questions is to point out that identifying one’s weakest areas is actually very beneficial when taking the exam. Students who can identify these questions quickly and spend little time on them can accumulate precious minutes to use on other questions, thus raising their score.

6. The percentage scores on AP exams noted above refer to data published by the

College Board on their released questions and will vary from year to year, as individual students and classes of students vary in their abilities from year to year. Also keep in mind is that when students are being assessed on tests in class during the school year, they are often assessed on a concentrated, narrow amount of information – helpful for formative assessment and preparation, but not necessarily characteristic of a final AP exam administered after course has been completed. While students may achieve a higher degree of success on AP-level multiple-choice questions when they are tested immediately at the end of a unit; teachers are encouraged to factor in the level of difficulty and drastically reduced raw score that is regarded as highly successful by the College Board when assessing the students’ efforts.

7. AP chemistry students typically do struggle with AP-level questions early in the course. AP-level questions require a broad background in the overall chemistry curriculum, and are also more difficult than those most students have encountered up to this point. It is advisable to provide answers to practice questions for students, especially early in the school year. They will find it helpful to read the question, work out an answer, and then get immediate feedback. Providing answer keys allows students to work backward from the answer and develops their understanding of how to attack multiple-choice questions at this level of difficulty. The answer explanations provided in the Applied Practice booklets are particularly useful in this regard. In addition to allowing some access to answer keys (unless, of course, you will be using these questions for assessment), teachers should plan to include time in class for guided practice of these questions, working them out with the students. Collaboration among students in small learning groups is also a useful technique. Chemistry is much easier to learn when students are able to talk about the content and reason out answers together.

8. A simple but useful tip for students is to underline, or in some way highlight, key

terms in the questions. Many times, a student will choose an incorrect answer because he or she simply misread the question.

FREE-RESPONSE QUESTIONS (FRQs): 1. One of the most common reasons for producing an unsatisfactory response to a free-

response question on an AP Chemistry Exam is that the student did not answer the specific question that was asked. Often, students write information that is factually correct but does not apply to the question, or their response relates to information in the question without actually answering the question. Free-response questions often present a great deal of information when only small bits of that information are actually needed to answer each of many different parts of the question. Students should read the entire question before they begin to write. They should underline, or in some way highlight, exactly what is being asked in each part of the question. After writing their response, students should proofread their work to be sure they answered the question specifically. Often, students will look back over their answers when their papers are returned and realize their answers make little sense, often responding, “But I meant to say ____.” They will likely experience greater success when they take time to read over their answers to verify they actually wrote what they intended to say.

2. Students should be specific in their answers and should understand the difference

between a definition and information that justifies or explains their answer. In most cases, students should use complete sentences. A few questions have parts that merely require a list of items, and others ask for additions to diagrams, but for the most part,

students will need to write coherently. A one-word answer is not sufficient to earn credit.

3. Students should be sure they understand the difference between an answer that

justifies their point and one that merely cites a correlation. A common example on past tests is a question about what determines the relative strength of London dispersion forces between molecules. Larger-sized atoms and molecules have more electrons and are more easily “ionizable”; thus, they are more likely to form dispersion forces. As atoms and molecules get larger, their nuclei get heavier, and thus the atomic and molecular weights increase along with dispersion forces. However, while increasing mass correlates to increasing strength of attraction between molecules, it does not explain or justify the increase.

An analogy illustrating this concept would be a graph of children’s mass with their level of intelligence. As they age, their brains develop at the same time their bodies grow and become heavier. Thus, older children are both smarter and heavier than younger children. But it would be incorrect to say that the children are becoming smarter because they are becoming heavier.

4. When teaching concepts in chemistry, it is tempting to make atoms, ions, and

molecules anthropomorphic. Humans tend to more easily understand concepts that are explained in a way that is relevant to the way we think and behave. Teachers must be careful not to instruct the students about what the atoms “want” or “need.” Often, students will write answers about an atom that is “happy” because it got the electron it “needed.” Even though students may have come to an appropriate level of comprehension on this concept, they will not receive credit when citing the emotions of a piece of matter on the AP Chemistry Exam.

5. Students should write out words fully, unless the abbreviations are universally

accepted. While their teachers may have come to understand a particular student’s shorthand, students are likely to lose out when the AP exams are graded.

6. Students should develop the habit of avoiding pronouns in their written answers, such

as beginning their first sentence with “it.” In many questions, there are dozens of things to which “it” could refer, and students will receive no credit for their answer if they use vague pronouns.

7. Using the free-response questions in the Applied Practice books is a good way to

develop students’ writing skills. Answers, explanations, and scoring guides for each of the six questions in an Applied Practice booklet are provided. The more writing students do, the better they will become at answering the questions in Section II of the AP exam. Although it takes time to grade students’ writing, the time will be well worth it in terms of their improvement and eventual success. Free-response questions are of three general types: large calculation questions that require good problem solving skills, equation writing, and “essay” types that require good writing skills.

The first type of question will be seen on Section A of Part II of the exam, and the last two types will be found on Section B of Part II.

8. When answering free-response questions that require calculations (section II, part A,

questions 1, 2, and 3), students should develop these good habits from the beginning of the year:

Always list any mathematical formula to be used (i.e. PV=nRT) at the beginning of

the question. Never put numerical information in an equation without first writing that equation down first. The vast majority of all mathematical formulas to be used are listed in the reference sheets provided on the AP Chemistry Exam in section II. Some knowledge of mathematical formulas will be required on section I (multiple- choice), so it is good for students to know them and not always rely on the reference sheets.

While students’ work must be shown, typically much less work is required than that

required by most high school math teachers. If the student lists the mathematical formula, substitutes the correct numerical values from the question into the proper location, and clearly marks the answer, the student does not need to write out each of the various algebraic steps. However, there is no disadvantage in scoring if students are more confident showing their work. They simply need to be coached in the amount of time to use on each calculation question. Some students will need to write fewer steps if time becomes a factor overall.

Once the students calculate and write a final answer, they should do the following

each time:

• Double check the number of significant figures in the answer. They must be within one significant figure up or down from the correct number.

• Double check the calculations. • Make sure they have labeled their answer as to which quantity it is (do not leave

answer as “x=0.050”). A label helps ensure they are answering the question correctly.

• Always include units (except for a few that do not use units, such as specific gravity and the equilibrium constant, Keq, etc).

• Box in the final answer. This not only helps the grader, but it provides an easy way for the students to check their work to ensure they have finished each part of each question. Also, answers from previous sections are often used on subsequent parts of the question. It is easier to find this information if it is boxed.

GLOSSARY OF TERMS allotrope—distinct forms in which a chemical element occurs in the same physical state, each of which differs in its physical properties Born-Haber cycle—a thermochemical cycle, based on Hess’s Law, of chemical reactions used for calculating either the energy required to break down a crystalline solid into its constituent ions (lattice energy), assumes 100% ionic character in the substance charge density—how concentrated an electric charge is on an ion or dipole moment,

related to intensity of attraction/repulsion as determined by Coulomb’s Law: E=kQ1Q2

r,

where Q represents the electric charges of the ions and r is the span between them. Smaller ions with higher charges have greater charge density, thus attract and repel more strongly covalent bond—attachment of atoms (non-metals) held together by sharing a pair of valence electrons; overlap of valence orbitals between two atoms critical point—the temperature above which no amount of pressure can liquefy a gas, associated with the pressure that would liquefy it at that temperature dipole-dipole forces—an intermolecular force (IMF) in which the oppositely charged ends of two polar molecules attract dipole moment—the particular area of positive or negative charge in a molecule, measured in debyes electron affinity—the energy change when an electron is added to a neutral atom in the gaseous state to form a negative ion endothermic— a chemical reaction in which heat is absorbed enthalpy of atomization—the energy required to convert an elemental substance into separate atoms equilibrium—a system where opposite reactions (such as phase changes: melting, freezing) occur at equal rates exothermic— a chemical reaction that produces heat hydrogen bonding—an intermolecular force (IMF), not a true “bond”, between a hydrogen atom in one molecule, and the oxygen, nitrogen, or fluorine atom (which must be covalently bonded to a hydrogen) in another molecule; the strongest of the IMFs

ion-dipole interaction—the electrostatic attraction (attraction between opposite charges) between an ion and the oppositely charged dipole moment in a polar molecule; the way that ionic compounds dissolve in polar water ionic bonding—relatively strong attraction between oppositely charged ions ionization energy—the energy required to remove an electron from an atom in the gaseous state; can be successive, the first ionization energy is the energy to remove the first electron from a neutral atom, the second ionization energy is the energy to remove the second electron from a 1+ cation, etc intermolecular forces—relatively weak attractions between molecules, allowing them to form liquids and solids at lower temperatures and/or higher pressures than most other substances kinetic molecular theory—the theory that explains matter as having particles in constant motion and describes ideal gases and their collisions lattice energy—the energy required to break down a crystalline solid into its constituent ions London dispersion forces—(induced dipole attractions) attraction between temporary dipoles created when electron distribution becomes instantaneously unbalanced, allowing for the opposite charges to attract between two molecules; weakest of intermolecular forces (IMFs); increase in strength with increased size of molecule macro-molecular—giant molecules forming crystalline solids held together with covalent bonds; the hardest of all materials melting point—the temperature at which a solid changes from crystalline form into a liquid form metallic bond—the interaction of delocalized electrons and metal nuclei in a metallic substance; electrons are attracted to many nuclei at the same time, being free to move throughout the material normal phase change point—the temperature at which the phase change occurs when the substance is at “normal” atmospheric pressure of 1.0 atm. periodicity—a regular pattern of physical and chemical properties when they are arranged according to atomic number polarity—having areas of opposite charge in the same molecule; containing dipoles polarizable—the ability to be changed from a nonpolar structure to a polar structure, relates to the size of the molecule; the ability to acquire polarity

solute—the substance that gets dissolved into a solution solvent—the substance that does the dissolving when a solution is formed STP—standard temperature and pressure) the values of 1.0 atmospheres of pressure and 273 K (0oC) used as reference conditions due to the variability of gas volumes at different conditions sublimation—a phase change when a solid is turned directly from a solid into a vapor surface tension—the energy needed to increase the surface area of a liquid by a unit amount; the tendency to become as small as possible (perfectly spherical) due to unbalanced attractive forces being exerted more strongly towards the center of a liquid sample triple point—the temperature and pressure conditions for any substance at which the solid, liquid and vapor(gas) phases are in equilibrium unit cell—the smallest unit of atoms forming faces that are parallelograms, which can be stacked in three dimensions to form a crystal volatile liquid—a liquid that has a relatively low boiling point and a relatively high vapor pressure at normal temperatures vapor—the gaseous phase of any substance that is normally a solid or liquid vapor pressure—the partial pressure above a liquid specifically due to the evaporated liquid particles, measured when liquid-vapor equilibrium has been reached in a closed system at a given temperature viscosity—the resistance to flow of liquids and gases; increases with greater intermolecular forces

Student Practices

for


Multiple-Choice Questions

Intermolecular Forces

The following answer choices can be used in questions 1-3. Each answer may be used once, more than once, or not at all. (A) London dispersion forces (B) Hydrogen bonding (C) Dipole-dipole intermolecular forces (D) Ionic bonding (E) Covalent bonding 1. The interaction that occurs between molecules of water but NOT between molecules

of chlorine 2. The interaction that occurs between molecules of hydrogen chloride but NOT

between molecules of chlorine 3. The interaction that accounts for the increasing melting and boiling points of the

halogens on descending group 17 4. When sodium chloride dissolves in water, which of the following statements is true? I. As part of the process, the lattice energy of the sodium chloride must be

overcome. II. The attraction between the sodium and chloride ions and the water molecules can

be described as an ion-dipole interaction. III. The polarity of the water molecules is essential to the process. (A) I only (B) II only (C) III only (D) I and II only (E) I, II, and III

5. Which of the following molecules will have London dispersion forces that form some part of the intermolecular attractions present?

I. Fluorine, F2 II. Ammonia, NH3 III. Hydrogen fluoride, HF (A) I only (B) III only (C) I and II only (D) II and III only (E) I, II, and III 6. Which of the following has the substances listed in order of increasing boiling point? I. Water < nitrogen < ammonia II. Nitrogen < water < chlorine III. Chlorine < bromine < iodine (A) I only (B) II only (C) III only (D) II and III only (E) I, II, and III 7. Which of the following statements is true on descending group 17? I. The boiling points of the halogen molecules decrease. II. The molecules become more polarizable. III. The dipole-dipole intermolecular interactions increase. (A) I only (B) II only (C) III only (D) II and III only (E) I, II, and III

8. All of the following substances exhibit hydrogen bonding between their molecules EXCEPT

(A) water (B) ethanol (C) ammonia (D) hydrogen chloride (E) ethylamine 9. Which of the following is the most soluble in a polar solvent? (A) Ethanol (B) Methane (C) Hexane (D) Carbon tetrachloride (E) Heptane 10. What type of intermolecular bonding would be expected between molecules of

carbon monoxide?

(A) Ionic (B) Covalent (C) Hydrogen bonding (D) Dipole-dipole (E) Ion-dipole 11. Which of the following shows intermolecular forces in order of decreasing strength?

(A) Hydrogen bonds > London dispersion > dipole-dipole (B) Hydrogen bonds > dipole-dipole > London dispersion (C) Dipole-dipole > hydrogen bonds > London dispersion (D) London dispersion > dipole-dipole > hydrogen bonds (E) Dipole-dipole > London dispersion > hydrogen bonds 12. All of the following molecules exhibit dipole-dipole intermolecular bonding

EXCEPT (A) HCl (B) HBr (C) Cl2 (D) CO (E) NO2

13. All of the following are true of London dispersion forces EXCEPT (A) they generally increase with a molecule’s increased polarizability (B) they generally decrease among the noble gases as group 18 is descended (C) they generally increase among the diatomic halogen molecules as group 17 is

descended (D) they are the weakest of the three intermolecular forces that also include hydrogen

bonding and dipole-dipole interactions (E) they generally increase as the number of electrons and surface area of molecules

increase 14. The existence of hydrogen bonding between the molecules of a substance is generally

associated with all of the following EXCEPT (A) increased viscosity (B) increased melting points (C) increased boiling points (D) decreased vapor pressures (E) molecules that contain double bonds 15. Between pairs of which of the following molecules would one expect the strongest

intermolecular forces? (A) Hydrogen (B) Carbon dioxide (C) Carbon tetrachloride (D) Nitrogen (E) Oxygen

Changes in State The following diagram can be used in questions 16-18. Each letter may be used once, more than once, or not at all.

Pres

sure

Temperature

1 atm

C

D

B

A

E

16. The point on the phase diagram where sublimation could be occurring 17. The point on the phase diagram where the substance is only in the gaseous phase 18. The point on the phase diagram where cooling the substance at a constant pressure

will cause the substance to change from an equilibrium mixture of liquid and solid to that of just a solid

19. Which of the following is true for a pair of volatile liquids, A and B, if they have approximately the same molecular mass but A has stronger intermolecular forces than B?

I. Under identical conditions, B will have a higher boiling point than A. II. Under identical conditions, B will have a higher vapor pressure than A. III. If molecules of A experience hydrogen bonding, it is possible that molecules of

B only experience London dispersion forces. (A) I only (B) II only (C) III only (D) II and III only (E) I, II, and III 20. In calculating the total energy needed to heat a lump of ice from 263 K to 288 K at 1

atm, which of the following constants are needed? I. Enthalpy of fusion for H2O II. Specific heat capacity of ice III. Specific heat capacity of water (A) I only (B) III only (C) I and II only (D) II and III only (E) I, II, and III 21. If a substance with a liquid phase less dense than its solid phase is subjected to

increasing pressure at a constant temperature, which of the following can occur? I. Melting II. Sublimation III. Condensation (A) I only (B) II only (C) III only (D) II and III only (E) I, II, and III

22. If a substance is cooled at a constant pressure, which of the following can occur? I. Freezing II. Condensing III. Sublimation (A) I only (B) II only (C) III only (D) I and II only (E) I, II, and III 23. If a solid substance that has a triple point that is above 1 atm, is heated at STP, what

would one observe? (A) The solid would melt. (B) The solid would melt and then turn into a vapor. (C) The solid would turn directly into a vapor. (D) The solid would turn into a vapor and then turn into a liquid. (E) No change would be observed. 24. Which of the following substances would one expect to have the highest vapor

pressure at STP? (A) Water (B) Hexane (C) Decane (D) Carbon tetrachloride (E) Heptane

Questions 25-27 are based upon the phase diagram shown below.

Pres

sure

in a

tm

Temperature in K

1.0

273 373

25. At a pressure of 1 atm and a temperature of 300 K, the substance will be (A) a solid (B) a liquid (C) a gas (D) an equilibrium mixture of liquid and gas (E) an equilibrium mixture of solid, liquid, and gas 26. A temperature of 373 K at a pressure of 1 atm represents (A) the triple point of the substance (B) the critical point of the substance (C) the normal boiling point of the substance (D) the normal freezing point of the substance (E) the normal melting point of the substance

27. At a pressure of 1 atm and a temperature lower that 273 K, which of the following statements is false?

(A) The substance will be in the solid phase. (B) The solid form of the substance would float in the liquid form of the substance. (C) The density of the solid is less than the density of the liquid phase. (D) Heating the substance will cause sublimation. (E) Cooling the substance will cause no change of phase. 28. The point on a phase diagram at which the liquid-gas phase equilibrium line crosses

the 1 atm pressure line is called the (A) triple point (B) critical point (C) normal boiling point (D) normal freezing point (E) normal melting point 29. The point on a phase diagram at which the particles of a gaseous substance cannot be

liquefied by pressure alone is called the (A) triple point (B) critical point (C) normal boiling point (D) normal freezing point (E) normal melting point 30. The point on a phase diagram at which solid, liquid, and gas phases of a substance

exist in equilibrium with one another is called the (A) triple point (B) critical point (C) normal boiling point (D) normal freezing point (E) normal melting point

Properties of Liquids The following answer choices can be used in questions 31-33. Each answer may be used once, more than once, or not at all. (A) Surface tension (B) Viscosity (C) Kinetic molecular theory (D) Hydrogen bonding (E) London dispersion forces 31. Describes the resistance of a liquid to flow 32. Describes the phenomenon that is caused by the tendency of a liquid to create the

smallest surface area 33. Occurs when hydrogen atoms in a molecule are directly connected to N, O, or F

atoms 34. Which of the following properties of water can be directly attributed to the hydrogen

bonds that are present? I. Increased viscosity II. Elevated boiling point III. Polarity (A) I only (B) II only (C) III only (D) I and II only (E) I, II, and III 35. Which of the following is true of H2O? I. It expands on freezing. II. It has a solid phase that is less dense than the liquid phase. III. It has a gas phase that is more dense than the liquid phase. (A) I only (B) III only (C) I and II only (D) II and III only (E) I, II, and III

36. The specific heat capacity of liquid H2O is approximately twice the value of the specific heat capacity of solid H2O. This means

I. that more energy is required to raise the temperature of 1 g of water by 1 K

compared to the energy required to raise the temperature of 1 g of ice by 1 K II. that if the same amount of energy is absorbed by 1 g of water and 1 g of ice, the

temperature of the water will increase by a greater amount that the temperature of the ice

III. that if 1 g of water is cooled by 10 K, compared to 1 g of ice being cooled by 10 K, more heat will release from the water

(A) I only (B) II only (C) III only (D) I and III only (E) I, II, and III 37. Water has intermolecular forces that are hydrogen bonds, whereas H2S has dipole-

dipole intermolecular forces. Which of the following statements is true? I. Measured under identical conditions, water has a lower boiling point than H2S. II. Measured under identical conditions, water has a higher vapor pressure than H2S. III. Water molecules are attracted to one another to a greater extent than H2S

molecules are attracted to one another. (A) I only (B) II only (C) III only (D) I and II only (E) I, II, and III 38. Which of the following would one expect to be the most soluble in water? (A) Methane (B) Propane (C) Hexane (D) Potassium bromide (E) Chlorine

39. When H2O is cooled from a temperature above 373 K to 290 K at 1 atm pressure, which of the following statements is false?

(A) The average kinetic energy of the particles decreases. (B) A change of state occurs. (C) A liquid is formed. (D) Condensation occurs. (E) The amount of hydrogen bonding present dramatically decreases. 40. Water is a relatively poor conductor of electricity because (A) it has a boiling point of 373 K (B) it will decompose to yield hydrogen gas and oxygen gas (C) it has a solid which is less dense than its liquid (D) it undergoes very little ionization (E) it is a polar substance 41. When compared to the other dihydrides of group 16, oxygen’s compound is found to

have an anomalously high boiling point. What factor accounts for this? (A) The large atom mass of O compared to the other elements of group 16 (B) The high ionization energy of O (C) The strength of the covalent bonds between O and H (D) The strength of the O-O double bond in the oxygen molecule (E) The intermolecular hydrogen bonds present in oxygen’s dihydride

Questions 42-44 are based upon the phase diagram for H2O shown below.

Pres

sure

in a

tm

Temperature in K

1.0

273 373

42. The temperature of 273 K can be described as (A) the temperature of the critical point for water (B) the temperature of the normal boiling point of water (C) the temperature of the normal freezing point of water (D) the temperature of the triple point for water (E) the temperature at which water forms an equilibrium mixture of gas and liquid 43. At a pressure of 1 atm, solid H2O, when heated, will melt to form a liquid and then

boil to form a gas. Which feature of the phase diagram confirms this sequence of phase changes?

(A) The triple point is below 1 atm. (B) The critical point is above the triple point. (C) The critical pressure is above 1 atm. (D) The boiling point is 373 K. (E) The freezing point is 273 K.

44. As pressure increases (A) the boiling point of H2O increases and the freezing point decreases (B) the boiling point of H2O decreases and the freezing point decreases (C) the boiling point of H2O increases and the freezing point increases (D) the boiling point of H2O decreases and the freezing point increases (E) the boiling point and the freezing point of H2O remain the same 45. What property of water contributes directly to its ability to dissolve table salt? (A) Its color (B) Its covalent nature (C) Its melting point (D) Its boiling point (E) Its polarity

Properties and Types of Solids The following answer choices can be used in questions 46-48. Each answer may be used once, more than once, or not at all. (A) LiF (B) LiBr (C) NaCl (D) NaBr (E) KBr 46. The ionic substance with the largest magnitude for lattice energy 47. The ionic substance with ions that have the highest charge densities 48. The ionic substance with the highest melting point 49. Which of the following includes only molecular substances? I. CO2, I2, H2O II. CO2, C(diamond), C(graphite) III. NaCl, MgCl2, CO2 (A) I only (B) II only (C) III only (D) I and II only (E) I, II, and III 50. Which of the following lists the ionic compounds in order of increasing lattice

energy? I. KBr < KCl < MgO II. KCl < CaO < MgO III. NaI < NaBr < CaO (A) I only (B) III only (C) I and II only (D) II and III only (E) I, II, and III

51. If the lattice energy of magnesium oxide is found to be -3781 kJ mol-1, which of the following statements is true?

I. The lattice energy of NaCl will be of higher magnitude. II. The energy associated with the process, MgO(s) Mg2+

(g) + O2-(g) will be equal

to +3781 kJ mol-1. III. The lattice energy of MgS will be of higher magnitude. (A) I only (B) II only (C) III only (D) II and III only (E) I, II, and III 52. Large magnitudes for lattice energies are associated with I. increasing the charge on the ions II. decreasing the size of the ions III. higher melting points for ionic substances (A) I only (B) II only (C) III only (D) I and II only (E) I, II, and III 53. Which of the following has, on average, one-eighth of a particle at each corner of the

unit cell and no other particles present in the unit cell? (A) Face-centered cubic (B) Simple cubic (C) Body-centered cubic (D) Rhombic (E) ABC packing 54. All of the following form ionic solids EXCEPT (A) Sodium chloride (B) Magnesium oxide (C) Magnesium fluoride (D) Polyethylene (E) Lithium bromide

55. The diagram below shows

(A) a rhombic crystal (B) a face-centered crystal (C) a body-centered crystal (D) a crystal with 6:6 co-ordination (E) a triclinic crystal 56. Which of the following equations is associated with the term “lattice energy”? (A) Na(s) Na(g) (B) Na(g) Na+

(g) + e- (C) Na+

(g) + Cl-(g) NaCl(s)

(D) Na(s) + ½Cl2(g) NaCl(s) (E) Na(g) + ½Cl2(g) NaCl(s) 57. Which of the following shows lattice energies in the correct order? (A) LiF > NaF > KF (B) KF > NaF > LiF (C) KF > LiF > NaF (D) CaO > MgO > BeO (E) MgO > BeO > CaO 58. On average, how many atoms does a face-centered cubic unit cell have? (A) 4 (B) 6 (C) 8 (D) 12 (E) 14

59. Which of the following compounds is likely to have the closet match of theoretical lattice energy (as predicted by a Born-Haber cycle calculation) and an experimentally determined value?

(A) NaCl (B) NaBr (C) NaI (D) AgBr (E) AgI 60. Which of the following statements in relation to lattice energy is true? (A) On descending group 1, the lattice energies of the fluorides decrease in

magnitude. (B) On descending group 1, the lattice energies of the fluorides remain

approximately constant in magnitude. (C) On descending group 17, the lattice energies of the halides of sodium increase in

magnitude. (D) On descending group 2, the lattice energies of the oxides increase in magnitude. (E) Lattice energies cannot be predicted as a function of periodicity.

Descriptive Chemistry The following answer choices can be used in questions 61-63. Each answer may be used once, more than once, or not at all. (A) High melting points (B) Good conductors of heat (C) Excellent conductors of electricity (D) High vapor pressure (E) Good conductivity in water solution 61. A property associated with an ionic solid in the solid state 62. A property associated with both the molten and aqueous state of an ionic compound

but not associated with the solid state of an ionic compound 63. A property that is associated with ionic solids but NOT molecular solids 64. The giant network of covalent bonding in diamond causes it to I. have a very high melting point II. have an allotropic form III. act as a lubricant (A) I only (B) II only (C) III only (D) I and II only (E) I, II, and III 65. Iodine forms a molecular solid where the molecules are attracted to one another via

London dispersion forces. Which of the following is true as a direct result of its structure?

I. It will sublime at room temperature. II. It is a soft solid that can easily be crushed. III. It is a dark purple/black–colored solid. (A) I only (B) III only (C) I and II only (D) II and III only (E) I, II, and III

66. Solids with ionic bonds are I. good conductors of electricity in all phases II. soluble in polar solvents such as water III. insoluble in non-polar solvents (A) I only (B) II only (C) III only (D) II and III only (E) I, II, and III 67. A solid with very weak intermolecular forces will I. have a low melting point II. have a high viscosity III. be very hard (A) I only (B) II only (C) III only (D) I and II only (E) I, II, and III 68. Which of the following accounts for the fact that aluminum oxide is used in the lining

of furnaces? (A) It is a covalent compound with a high melting point. (B) It is an ionic compound with a high melting point. (C) It is an excellent conductor of electricity when solid. (D) It is an excellent conductor of heat. (E) It is an ionic compound with a low melting point. 69. Which of the following is associated with graphite? (A) It can conduct electricity in one plane only. (B) It is an ionic compound. (C) It is a covalently bonded compound. (D) It forms an ionic lattice. (E) It is the hardest substance known.

70. Ice is an example of a solid with all of the following attributes EXCEPT (A) it is a hydrogen-bonded solid (B) it is a solid that is less dense than the corresponding liquid form (C) it is a solid with an anomalously high melting point (D) it is a solid with an ionic lattice structure (E) it is a molecular solid 71. A solid that has its particles held together with which of the following interactions

would be expected to have the lowest melting point? (A) Metallic bonds (B) Hydrogen bonds (C) Covalent bonds in a giant network (D) London dispersion forces (E) Ionic bonds 72. All of the following properties are associated with ionic solids EXCEPT (A) high conductivity when solid (B) high solubility in polar solvents (C) high conductivity when aqueous (D) high conductivity when molten (E) high melting points 73. Molecular solids tend to have low melting points because (A) their covalently bonded molecules have low molar masses (B) the covalent bonds holding the atoms together are strong (C) the intermolecular forces are weak (D) the molecules are held together in giant networks (E) the covalent bonds holding the atoms to one another are weak 74. Which of the following is not a property generally associated with metallic bonding? (A) Excellent conductors of heat (B) Excellent conductors of electricity (C) High melting points (D) High boiling points (E) High vapor pressures

75. The extreme hardness of diamond can be attributed to the fact that (A) it is ionic (B) it has a high melting point (C) it forms a covalently bonded, giant network of atoms (D) it forms an allotrope in graphite (E) it s a poor conductor of electricity

Free-Response

Questions

1. (a) By constructing Born-Haber cycles (or another method), use the data in the table

below to calculate the enthalpy of formation of lithium fluoride and the enthalpy of formation of beryllium fluoride.

All values in

kJ mol-1 Li Be F LiF BeF2

Enthalpy of atomization +161 +326 +112 - -

1st Ionization Energy +520 +901 - - -

2nd Ionization Energy +7298 +1757 - - -

1st Electron Affinity - - -342 - -

Lattice Energy - - - -1031 -3461

(b) Considering only the ionization energies of lithium, explain why Li does not form

LiF2. (c) Explain why the lattice energy of beryllium fluoride is approximately 3.5 times

the value given for lithium fluoride. (d) Explain why the value for the lattice energy of beryllium fluoride predicted by the

Born-Haber cycle is significantly different from the experimentally found value.

2. Solids can be classified in a number of ways. Four such ways are listed below. In each case give an example of such a solid and identify the specific type(s) of particle present, the type of attraction between those particles, and one property that the solid as a whole possesses.

(a) Ionic (b) Molecular (c) Metallic (d) Giant atomic network

3. A 5.00 g sample of H2O is removed from a freezer and heated from an initial temperature of -5.00 oC through to a final temperature of 102.0 oC.

(a) Given the following data, calculate the amount of energy that has to be absorbed

by the H2O during the whole process.

Freezing point of H2O = 0.00oC; Boiling point of H2O = 100.oC ΔHfusion of H2O = 6.01 kJ mol-1 ΔHevaporation of H2O = 40.7 kJ mol-1

Specific heat capacity of ice = 2.05 J g-1 K-1; Specific heat capacity of water = 4.18 J g-1 K-1 Specific heat capacity of steam = 2.08 J g-1 K-1

(b) Calculate the percentage of the total energy required for the overall process in (a)

that is used just in heating the liquid water. (c) Calculate the mass of ice at 0.00oC, that can be melted in 30 minutes by a heater

that produces 8000 kJ of energy per hour but in such a way that only 70% of the heat produced is absorbed by the ice.

4. Using principles of chemical bonding theory, explain each of the following observations. In each case discuss both substances.

(a) Argon has a higher boiling point than helium.

(b) The lattice energy of magnesium oxide is significantly greater than the lattice energy of sodium fluoride.

(c) Solid sodium is an excellent conductor of electricity but solid sodium chloride is

an insulator. (d) Bromine has a higher boiling point than hydrogen chloride.

5. Consider the phase diagram for a particular substance shown below.

Temperature

Pres

sure

1 atm

S

T

(a) Carefully describe the changes that take place if a solid sample of the substance is

heated at a pressure of 1 atmosphere. (b) What does point S represent? What characteristics does the substance have at

point S? (c) Carefully describe the nature of the substance at all points along the line ST. (d) When starting at a point on the line ST, describe the change that occurs when the

substance is subjected to an increase in pressure at a constant temperature. (e) If point S was below the 1 atm line rather than above it, and a gaseous sample of

the substance was cooled at constant atmospheric pressure, what changes would occur?

6. Using principles of chemical bonding theory, explain each of the following observations. In each case discuss both substances.

(a) Silicon dioxide has a melting point that is approximately 2000 degrees higher

than that of oxygen. (b) Water and ammonia have very similar molecular masses, but water is a liquid at

room temperature while ammonia is a gas. (c) Iodine is a solid that will not readily dissolve in water whereas potassium iodide

is highly soluble in the same solvent. (d) At STP, chlorine is a gas but bromine is a liquid.

Answer Key and Explanations

for


Multiple-Choice

Answer Key

Spiral-bound teacher pages may not be shared or reproduced.

Do n

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opyMULTIPLE-CHOICE ANSWER KEY

1. B

2. C

3. A

4. E

5. E

6. C

7. B

8. D

9. A

10. D

11. B

12. C

13. B

14. E

15. C

16. B

17. E

18. C

19. D

20. E

21. C

22. D

23. C

24. B

25. B

26. C

27. D

28. C

29. B

30. A

31. B

32. A

33. D

34. D

35. C

36. D

37. C

38. D

39. E

40. D

41. E

42. C

43. A

44. A

45. E

46. A

47. A

48. A

49. A

50. E

51. B

52. E

53. B

54. D

55. C

56. C

57. A

58. A

59. A

60. A

61. A

62. C

63. E

64. A

65. C

66. D

67. A

68. B

69. A

70. D

71. D

72. A

73. C

74. E

75. C


Multiple-Choice

Answer Explanations


Do n

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INTERMOLECULAR FORCES 1. (B) Hydrogen bonding. Water and chlorine both have London dispersion forces between their molecules, but hydrogen bonding only occurs when hydrogen atoms are directly, covalently bonded to nitrogen, oxygen, or fluorine atoms. 2. (C) Dipole-dipole intermolecular forces. Water and chlorine both have London dispersion forces between their molecules, but dipole-dipole intermolecular forces only occur when a molecule has atoms with different electronegativities which create dipoles that do not cancel through symmetry. 3. (A) London dispersion forces. Halogen molecules have no dipoles (therefore no dipole-dipole interactions) and no hydrogen bonding. LDFs are the only intermolecular force. 4. (E) I, II, and III. Polar water molecules surround Na and Cl ions, and the hydration enthalpy exceeds the lattice energy, causing the sodium chloride to dissolve. 5. (E) I, II, and III. LDFs are present between ALL covalent molecules. 6. (C) III only. Water has the strongest intermolecular force (hydrogen bonding). Ammonia is also hydrogen bonded, but its hydrogen bonding is weaker than that of water (less electronegative atom attached to the H). The other molecules have LDFs as intermolecular forces which increase with size due to greater polarizability. 7. (B) II only. London dispersion intermolecular forces increase with size due to greater polarizability. Increased intermolecular forces increase the boiling point. 8. (D) hydrogen chloride. Hydrogen bonding only occurs when hydrogen atoms are directly, covalently bonded to nitrogen, oxygen, or fluorine atoms. 9. (A) Ethanol. Water is a polar solvent that will dissolve polar solutes. Ethanol is significantly polar and exhibits hydrogen bonding that significantly increases its solubility in water since it can hydrogen bond with the solvent (water). 10. (D) Dipole-dipole. Dipole-dipole intermolecular forces occur when a molecule has atoms with different electronegativities which create dipoles that do not cancel through symmetry. 11. (B) Hydrogen bonds > dipole-dipole > London dispersion. This is statement of fact. 12. (C) Cl2. Dipole-dipole intermolecular forces occur when a molecule has atoms with different electronegativities which create dipoles that do not cancel through symmetry. The two atoms in a chlorine molecule have identical electronegativities.


Do n

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opy13. (B) they generally decrease among the noble gases as group 18 is descended.

London dispersion forces increase with size due to greater polarizability. 14. (E) molecules that contain double bonds. Hydrogen bonding is the strongest intermolecular force, and it increases the attractive forces between molecules. This increases the viscosity, melting point, and boiling point of a substance and decreases vapor pressure. 15. (C) Carbon tetrachloride. London dispersion forces increase with size due to greater polarizability.


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CHANGES IN STATE 16. (B) See diagram. Point B lies on the line on the phase diagram that represents an equilibrium between gas and solid. 17. (E) See diagram. Point E lies in the area of the phase diagram that represents the gaseous phase. 18. (C) See diagram. Point C lies on the line on the phase diagram that represents an equilibrium between liquid and solid. Increasing the pressure at a constant temperature will cause the liquid to solidify. 19. (D) II and III only. The same molecular mass suggests that their London dispersion forces will be approximately the same. Stronger intermolecular forces will lead to higher boiling points and lower vapor pressures. London dispersion forces are weaker than hydrogen bonds. 20. (E) I, II, and III. At 1 atm and 263 K, H2O is a solid. Calculating the energy required to heat the ice to its melting point (273 K) requires knowledge of the specific heat capacity of ice. Calculation of the energy required to melt the ice requires knowledge of the enthalpy of fusion. Calculating the energy required to heat the water to 288 K requires knowledge of the specific heat capacity of water. 21. (C) III only. Increasing pressure at a constant temperature will cause particles to come closer together and in turn cause a change of state where the new phase has less energy than the former. 22. (D) I and II only. Decreasing temperature at a constant pressure will cause particles to come closer together and in turn cause a change of state where the new phase has less energy than the former. 23. (C) The solid would turn directly into a vapor. See phase diagram associated with Q16-18. 24. (B) Hexane. High vapor pressures are associated with weak intermolecular forces. Water is hydrogen bonded whereas the other molecules have only weaker London dispersion forces. London dispersion forces increase with size due to greater polarizability. Hexane is the smallest of the remaining molecules. 25. (B) a liquid. A pressure of 1 atm and a temperature of 300 K on the phase diagram lies in the portion where only the liquid phase exists. 26. (C) the normal boiling point of the substance. This point lies on the line that represents an equilibrium between liquid and gas and, as such, defines the normal (boiling temperature at normal pressure) boiling point of the substance.


Do n

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opy27. (D) Heating the substance will cause sublimation. Since the triple point is below 1

atm, heating the solid will cause it to melt before forming a gas. 28. (C) normal boiling point. Statement of fact. 29. (B) critical point. Statement of fact. 30. (A) triple point. Statement of fact.


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PROPERTIES OF LIQUIDS 31. (B) Viscosity. Statement of fact. 32. (A) Surface tension. Statement of fact. 33. (D) Hydrogen bonding. Statement of fact. 34. (D) I and II only. Polarity is caused by a difference in electronegativity between O and H atoms and not as a consequence of hydrogen bonding. Hydrogen bonding is the strongest intermolecular force and increases the attractions between the molecules. The increased attraction causes a greater viscosity and boiling point. 35. (C) I and II only. I and II are statements of fact. Statement III is false. 36. (D) I and III only. Application of q = m c ΔT where the c values are different for each substance. 37. (C) III only. Hydrogen bonding is a stronger intermolecular force than dipole-dipole and, as such, boiling points will be higher and vapor pressures lower for hydrogen bonded substances. 38. (D) Potassium bromide. Water is a polar solvent which will tend to dissolve polar solutes. Since it is ionic, KBr is the only polar solute. 39. (E) The amount of hydrogen bonding present is dramatically decreased. The boiling point of H2O is 373 K and the freezing point is 273 K. This means that gaseous H2O will become a liquid under the circumstance described. 40. (D) it undergoes very little ionization. For a liquid to be an electrolyte it must have ions present; water forms very few ions. 41. (E) The intermolecular hydrogen bonds present in oxygen’s dihydride. Oxygen’s dihydride is water, the only group 16 dihydride that undergoes hydrogen bonding. Increased strength of intermolecular forces accounts for increases in boiling point. 42. (C) the temperature of the normal freezing point of water. Statement of fact; at 1 atm, the liquid solid equilibrium line is crossed. 43. (A) The triple point is below 1 atm. If the triple point were above 1 atm, the solid substance would sublime on heating.


Do n

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opy44. (A) the boiling point of H2O increases and the freezing point decreases. Moving

vertically along the solid liquid equilibrium line on the phase diagram results in a right to left movement on the temperature axis (temperature goes down). Moving vertically along the liquid gas equilibrium line on the phase diagram results in a left to right movement on the temperature axis (temperature goes up).

45. (E) Its polarity. Table salt (NaCl) is ionic and requires a polar solvent. Water is a polar solvent which will tend to dissolve polar solutes.


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PROPERTIES AND TYPES OF SOLIDS 46. (A) LiF. The lithium and fluoride ions are the smallest of the ions present in the compounds listed, and since all of these ions have charges of either +1 or -1, this will mean they have the highest charge densities. Their small size means they can get closer together than any other pair, and thus LiF will have the lattice energy with the largest magnitude. 47. (A) LiF. The lithium and fluoride ions are the smallest of the ions present in the compounds listed, and the fact that all of these ions have charges of either +1 or -1 means they have the highest charge densities. 48. (A) LiF. The lithium and fluoride ions are the smallest of the ions present in the compounds listed, and since all of these ions have charges of either +1 or -1, they have the highest charge densities. Their small size means they can get closer together than any other pair; therefore, LiF will have the lattice energy with the largest magnitude. The largest attraction (lattice energy) will cause the highest melting point. 49. (A) I only. Carbon (in the form of graphite or diamond) is a giant atomic network solid. NaCl and MgCl2 are ionic. 50. (E) I, II, and III. The smallest of the ions with the smallest charges will have the highest charge densities. Their small size means they can get closer together than any other pair and thus will have the lattice energy with the largest magnitude. 51. (B) II only. NaCl has ions with smaller charges, so the attraction will not be as great. Statement II is the opposite process of lattice energy and thus will have an equal and opposite enthalpy change. Sulfur ions are less charge dense than oxygen ions, so the attraction between S and Mg ions will be lesser and have a small magnitude for lattice energy.

52. (E) I, II, and III. Lattice energy = 1 2q qr

. As the charges of the ions increase and

their sizes decrease, their attraction increases. Melting points increase with greater attraction (greater lattice energies). 53. (B) Simple cubic. Cubic crystals have particles at each corner of a cube (i.e. 8). Each particle is shared among seven other unit cells, so one eighth of it is devoted to each unit cell. 54. (D) Polyethylene. Polyethylene is a polymer. All of the other compounds consist of a metal with a non-metal and thus will be ionic. 55. (C) A body-centered crystal. Statement of fact.


Do n

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opy56. (C) Na+

(g) + Cl-(g) NaCl(s). Statement of fact.

57. (A) LiF > NaF > KF. The smallest of the ions with the smallest charges will have the highest charge densities. Their small size means they can get closer together than any other pair and therefore will have the lattice energy with the largest magnitude. 58. (A) 4. Face-centered crystals have particles at each corner of a cube (i.e. 8) and one particle in the center of each face of a cube. Each particle at the corners is shared among

seven other unit cells, so one-eighth of it is devoted to each unit cell; 18 = 18

⎛ ⎞⎜ ⎟⎝ ⎠

. Each

particle on the faces is shared between one other unit cell, so one-half of it is devoted to

each unit cell; 16 = 32

⎛ ⎞⎜ ⎟⎝ ⎠

. 1 (from corners) + 3 (from faces) = 4 (total).

59. (A) NaCl. Born-Haber cycle calculations assume 100% ionic character of the compound. Sodium chloride has the least degree of covalent character (is the most ionic); as a result, the calculated values are closest to the experimentally found values. 60. (A) On descending group 1, the lattice energies of the fluorides decrease in magnitude. Descending any group causes the ions to get bigger and charge densities to decrease. The attraction between ions becomes less, and the lattice energy decreases.


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DESCRIPTIVE CHEMISTRY 61. (A) High melting points. Statement of fact; ionic bonding is very strong. 62. (C) Excellent conductors of electricity. Ionic substances are only good conductors when their ions are free to move. In the solid state this is not possible, but when molten or aqueous, the ions are free to move. 63. (E) Good conductivity in water solution. Polar molecules are soluble in water but do not ionize to any great degree. Soluble ionic compounds ionize significantly, allowing the individual ions to be mobile and thus conduct electricity. 64. (A) I only. Strong covalent bonds in a giant network lead to very high melting points. Carbon does have an allotropic form (graphite) which is used as lubricant, but this is not a consequence of the network bonding in diamond. 65. (C) I and II only. Weak intermolecular forces lead to very low melting points and soft solids. Iodine is a dark purple/black color, but this is not a consequence of it being a molecular solid with weak intermolecular forces. 66. (D) II and III only. Ionic substances are only good conductors when their ions are free to move. In the solid state this is not possible, but when molten or aqueous, the ions are free to move. Polar solvents can penetrate the ionic lattice and surround the ions (charged particles) by forming ion-dipole attractions. This is not possible with non-polar solvents. 67. (A) I only. Statements II and III are associated with strong interactions. 68. (B) It is an ionic compound with a high melting point. A furnace lining must be able to withstand extremely high temperatures. The strength of the ionic bonding means aluminum oxide has a very high melting point and can be used in this application. 69. (A) It can conduct electricity in one plane only. Statement of fact; graphite has a layered structure where the electrons can move freely along the plane but cannot jump between layers. 70. (D) it is a solid with an ionic lattice structure. No ions are present in solid H2O. 71. (D) London dispersion forces. London dispersion forces are the weakest intermolecular forces, causing substances that have them as the only intermolecular force to have low melting and boiling points. 72. (A) high conductivity when solid. Ionic substances are only good conductors when their ions are free to move. In the solid state this is not possible, but when molten or aqueous, the ions are free to move.


Do n

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opy73. (C) the intermolecular forces are weak. London dispersion forces are the weakest

intermolecular forces, causing substances that have them as the only intermolecular force to have low melting and boiling points. 74. (E) High vapor pressures. Metallic bonding is strong and metals tend to be solids at room temperature. As a result, they are not volatile and do not easily vaporize, leading to low vapor pressures. 75. (C) it forms a covalently bonded, giant network of atoms. Strong covalent bonds in a giant network lead to very high melting points. Diamond has such a structure.


Free-Response Answers and

Scoring Guides


Do n

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IMF, LIQUIDS, AND SOLIDS 1. (a) Li(s) Li(g) (161) Li(g) Li+

(g) + e- (520) ½F2(g) F(g) (112) F(g) + e- F-

(g) (-342) Li+

(g) + F-(g) LiF(s) (-1031)

___________________________ Li(s) + ½F2(g) LiF(s) -580 ___________________________ Be(s) Be(g) (326) Be(g) Be2+

(g) + 2e- (901) + (1757) F2(g) 2F(g) (112) (2) 2F(g) + 2e- 2F-

(g) (-342) (2) Be2+

(g) + 2F-(g) BeF2(s) (-3461)

___________________________ Be(s) + F2(g) BeF2(s) -937 ___________________________ All units are kJ mol-1. (4 points: 2 point for each calculation) (b) The large endothermic energy required for the second ionization energy of

Li (i.e., that required to form the Li2+(g) ion) is not compensated for by the

subsequent exothermic processes. (2 points) (c) Lattice energy is dependent upon the charges of the ions present. The

greater the charge, the greater the attraction and the higher the lattice energy. Since Be forms an ion with a charge of +2, and Na forms an ion of +1, the attraction between the Be and F ions will be much greater than that of Li and F ions. (2 points)

(d) Born-Haber cycle calculations assume 100% ionic character of the

compound. Beryllium fluoride has a significant degree of covalent character and, as a result, the calculated values differ from the experimentally found values. (2 points)


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IMF, LIQUIDS, AND SOLIDS 2. There are obviously a number of possible answers here. Typical responses

that would score full credit are given below. Any one property given will be sufficient. One point for any two of the four parts answered correctly in each solid type. Two points for all four.

(a) Sodium chloride (any ionic compound) Lattice of cations and anions Ionic Bonds (strong electrostatic interactions between ions)

High melting and boiling points, insulators when solid, conductors when molten or in solution, soluble in polar solvents (2 points)

(b) Carbon dioxide (any molecular substance) Molecules

London dispersion forces (weak electrostatic interactions between induced, temporary dipoles) (an appropriate intermolecular force) Low melting and boiling points, insulators (2 points)

(c) Aluminum

Closest packed lattice of atoms/ions with a free moving “sea” of electrons Metallic bonding (electrostatic attraction between ions and electrons)

High melting and boiling points, excellent conductors of heat and electricity, malleable and ductile (2 points)

(d) Diamond Giant network of covalently bonded atoms Covalent bonds (shared pairs of electrons creating very strong bonds) Extremely high melting and boiling points, insulator, very hard (2 points)


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IMF, LIQUIDS, AND SOLIDS

3. (a) Heating ice: ⎛ ⎞⎛ ⎞ ⎛ ⎞⎛⎜ ⎟⎜ ⎟ ⎜ ⎟⎜

⎝ ⎠ ⎝ ⎠⎝⎝ ⎠

5.00 ⎞⎟⎠

g 2.05 J 5 K 1 kJg K 1000 J

= 0.05125 kJ

Melting ice: ⎛ ⎞⎛ ⎞ ⎛⎜ ⎟⎜ ⎟ ⎜

⎝ ⎠ ⎝⎝ ⎠

5.00 ⎞⎟⎠

g 1 mol 6.01 kJ18.01 g 1 mol

= 1.67 kJ

Heating water: ⎛ ⎞⎛ ⎞ ⎛ ⎞⎛⎜ ⎟⎜ ⎟ ⎜ ⎟⎜

⎝ ⎠ ⎝ ⎠⎝⎝ ⎠

5.00 ⎞⎟⎠

g 4.18 J 100 K 1 kJg K 1000 J

= 2.09 kJ

Boiling water: ⎛ ⎞⎛ ⎞ ⎛⎜ ⎟⎜ ⎟ ⎜

⎝ ⎠ ⎝⎝ ⎠

5.00 ⎞⎟⎠

g 1 mol 40.7 kJ18.01 g 1 mol

= 11.3 kJ

Heating steam: ⎛ ⎞⎛ ⎞ ⎛ ⎞⎛⎜ ⎟⎜ ⎟ ⎜ ⎟⎜

⎝ ⎠ ⎝ ⎠⎝⎝ ⎠

5.00 ⎞⎟⎠

g 2.08 J 2 K 1 kJg K 1000 J

= 0.0208 kJ

Total Energy = 0.05125 + 1.67 + 2.09 + 11.3 + 0.0208 = 15.1 kJ (4 points)

(b) ⎛ ⎞⎜ ⎟⎝ ⎠

2.09 kJ 10015.1 kJ

= 13.8 % (1 point)

(c) Energy absorbed by ice = 8000 kJ 1 hour 30 mins 701 hour 60 mins 100

⎛ ⎞⎛ ⎞⎛ ⎞⎛⎜ ⎟⎜ ⎟⎜ ⎟⎜⎝ ⎠⎝ ⎠⎝ ⎠⎝

⎞⎟⎠

= 2800 kJ

2800 kJ = ⎛ ⎞⎛ ⎞ ⎛⎜ ⎟⎜ ⎟ ⎜

⎝ ⎠ ⎝⎝ ⎠

x ⎞⎟⎠

g 1 mol 6.01 kJ18.01 g 1 mol

x = 8391 g (3 points)


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IMF, LIQUIDS, AND SOLIDS 4. (a) Argon and helium are monatomic gases that have London dispersion

intermolecular forces. London dispersion forces increase with size as the atoms become larger and more polarizable. Since argon is larger than helium, it will have greater LDFs and therefore a higher boiling point. (2 points)

(b) Lattice energy is dependent upon the charge density of the ions present. The

greater the charge densities, the greater the attraction and the higher the lattice energy will be. Since Mg and O have ions that have a charge of 2, and Na and Cl have ions with charges of 1, the attraction between the Mg and O ions will be much greater than that of similar size but lesser charged Na and Cl. (2 points)

(c) Sodium metal exhibits metallic bonding where a free moving “sea” of

electrons allows the metal to be an excellent conductor. In solid sodium chloride, ions are present. These ions cannot conduct electricity unless they are free to move. In the solid state, the ionic bonded ions cannot move and therefore cannot conduct electricity. (2 points)

(d) Bromine molecules and HCl molecules have London dispersion forces

between their molecules. In addition, HCl molecules have dipole-dipole intermolecular forces. Dipole-dipole forces are stronger than London dispersion forces; HOWEVER, in this case the much more massive bromine molecules have sufficiently large London dispersion forces to overcome the total intermolecular London dispersion and dipole-dipole forces of the HCl molecules. This gives bromine a higher boiling point than HCl. (2 points)


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IMF, LIQUIDS, AND SOLIDS 5. (a) The solid will sublime - change from a solid directly to a gas/vapor. (1 point) (b) The triple point. At this temperature and pressure, all three phases of the

substance can exist in equilibrium with one another. (2 points) (c) Solid and liquid phases are in equilibrium with one another. (1 point) (d) The equilibrium mixture of solid and liquid phases will be converted to a

solid. (1 point) (e) The gas would first form a liquid and then form a solid. (1 point)


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IMF, LIQUIDS, AND SOLIDS 6. (a) Silicon dioxide has a macro-molecular, giant network structure with atoms

held together with very strong covalent bonds. This gives it a very high melting and boiling point. Oxygen molecules are discrete, individual molecules that have weak, London dispersion intermolecular forces between their molecules. These attractions are weak and are easily broken; thus, oxygen has a low melting and boiling point. (2 points)

(b) The similarity in their molar masses suggests that their London dispersion

forces are similar. In addition, both molecules exhibit hydrogen bonding between their molecules. However, the hydrogen bonding in water is stronger than that of ammonia for two reasons: First, the electronegativity of O is greater than that of N creating larger dipoles. Second, since the O atom has two lone pairs as opposed to one on the N atom, it allows a greater number of electrostatic interactions between molecules of water. (2 points)

(c) Water is a polar solvent. Polar solvent will dissolve polar solutes. Iodine is a

non-polar solid (a covalent, linear molecule with atoms of the same electronegativity at either end), but potassium iodide is extremely polar; in fact, it is ionic. (2 points)

(d) Chlorine and bromine are diatomic gases that have London dispersion

intermolecular forces. London dispersion forces increase with size as the molecules become larger and more polarizable. Since bromine is larger than chlorine, it will have greater LDFs and therefore a higher boiling point. (2 points)


Contributors Series Editor David Emmerson is a veteran teacher with more than twenty years’ experience teaching AP Chemistry. He has previously worked on developing an online AP Chemistry course. He holds a B.S. in Biology from Cornell University and an M.A. in Science Education from the State University of New York Writer Adrian Dingle is a chemistry educator with 18 years’ experience and is creator of the award-winning chemistry Web site www.adriandingleschemistrypages.com. He holds a B.Sc. (Hons.) Chemistry and a Postgraduate Certificate in Education from the University of Exeter, England.

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