Explain and Predict—Practice Free Response Questions for AP® Chemistry

Materials

24 Student Booklets (10 AP^® Chemistry practice free-response questions per booklet, with space for answers)

Teacher Tips

The booklet may be incorporated into an AP^® Chemistry course myriad ways. You may choose to assign a weekly question as the AP Chemistry Exam approaches, or you may choose to assign the 10 questions as one assignment. The booklet is meant to prepare students for the AP Chemistry Exam in a manner befitting teachers’ unique students/classes.
The booklet includes three 4-point questions and seven 10-point questions. Students should spend approximately 7–8 minutes answering each 4-point question and approximately 20 minutes answering each 10-point question. The booklet’s questions cover topics from the six AP Chemistry big ideas and individual questions often cover multiple big ideas.
The answers provided herein represent strong answers that would likely receive full credit.
Reproduction permission for only the student booklet is granted only to the science teacher who purchased AP7711 from Flinn Scientific, Inc., No part of the booklet or the accompanying Teacher’s Notes publication may be included on any website.

Further Extensions

Alignment with the Curriculum Framework for AP^® Chemistry

Question 1

Essential Knowledge
2C1: In covalent bonding, electrons are shared between the nuclei of two atoms to form a molecule or polyatomic ion. Electronegativity differences between the two atoms account for the distribution of the shared electrons and the polarity of the bond.
2C2: Ionic bonding results from the net attraction between oppositely charged ions, closely packed together in a crystal lattice.
2C4: The localized electron bonding model describes and predicts molecular geometry using Lewis diagrams and the VSEPR model.

Science Practices
1.4: The student can use representations and models to analyze situations or solve problems qualitatively and quantitatively.

Learning Objectives
2.18: The student is able to rank and justify the ranking of bond polarity on the basis of the locations of the bonded atoms in the periodic table.
2.21: The student is able to use Lewis diagrams and VSEPR to predict the geometry of molecules, identify hybridization, and make predictions about polarity.
2.22: The student is able to describe the relationships between the structural features of polar molecules and the forces of attraction between the particles.

Question 2

Essential Knowledge
1B1: The atom is composed of negatively charged electrons, which can leave the atom, and a positively charged nucleus that is made of protons and neutrons. The attraction of the electrons to the nucleus is the basis of the structure of the atom. Coulomb’s Law is qualitatively useful for understanding the structure of the atom.
1B2: The electronic structure of the atom can be described using an electron configuration that reflects the concept of electrons in quantized energy levels or shells; the energetics of the electrons in the atom can be understood by consideration of Coulomb’s Law.
1C1: Many properties of atoms exhibit periodic trends that are reflective of the periodicity of electronic structure.
1C2: The currently accepted best model of the atom is based on the quantum mechanical model.
1D1: As is the case with all scientific models, any model of the atom is subject to refinement and change in response to new experimental results. In that sense, an atomic model is not regarded as an exact description of the atom, but rather a theoretical construct that fits a set of experimental data.

Science Practices
1.1: The student can create representations and models of natural or man-made phenomena and systems in the domain.
1.2: The student can describe representations and models of natural or man-made phenomena and systems in the domain.
1.4: The student can use representations and models to analyze situations or solve problems qualitatively and quantitatively.

Learning Objectives
1.5: The student is able to explain the distribution of electrons in an atom or ion based upon data.
1.6: The student is able to analyze data relating to electron energies for patterns and relationships.
1.7: The student is able to describe the electronic structure of the atom, using PES data, ionization energy data, and/or Coulomb’s Law to construct explanations of how the energies of electrons within shells in atoms vary.
1.9: The student is able to predict and/or justify trends in atomic properties based on location on the periodic table and/or the shell model.
1.12: The student is able to explain why a given set of data suggests, or does not suggest, the need to refine the atomic model from a classical shell model with the quantum mechanical model.

Question 3

Essential Knowledge
1E1: Physical and chemical processes can be depicted symbolically; when this is done, the illustration must conserve all atoms of all types.
6B1: Systems at equilibrium respond to disturbances by partially countering the effect of the disturbance (Le Chatelier’s principle).
6C1: Chemical equilibrium reasoning can be used to describe the proton-transfer reactions of acid–base chemistry.
6A3: When a system is at equilibrium, all macroscopic variables, such as concentrations, partial pressures, and temperature, do not change over time. Equilibrium results from an equality between the rates of the forward and reverse reactions, at which point Q = K.

Science Practices
1.4: The student can use representations and models to analyze situations or solve problems qualitatively and quantitatively.
6.4: The student can make claims and predictions about natural phenomena based on scientific theories and models.
7.2: The student can connect concepts in and across domain(s) to generalize or extrapolate in and/or across enduring understandings and/or big ideas.

Learning Objectives
2.8: The student can draw and/or interpret representations of solutions that show the interactions between the solute and solvent.
5.16: The student can use Le Chatelier’s principle to make qualitative predictions for systems in which coupled reactions that share a common intermediate drive formation of a product.
6.3: The student can connect kinetics to equilibrium by using reasoning about equilibrium, such as Le Chatelier’s principle, to infer the relative rates of the forward and reverse reactions.
6.8: The student is able to use Le Chatelier’s principle to predict the direction of the shift resulting from various possible stresses on a system at chemical equilibrium.
6.11: The student can generate or use a particulate representation of an acid (strong or weak or polyprotic) and a strong base to explain the species that will have large versus small concentrations at equilibrium.
6.20: The student can identify a solution as being a buffer solution and explain the buffer mechanism in terms of the reactions that would occur on addition of acid or base.

Question 4

Essential Knowledge
1E2: Conservation of atoms makes it possible to compute the masses of substances involved in physical and chemical processes. Chemical processes result in the formation of new substances, and the amount of these depends on the number and the types and masses of elements in the reactants, as well as the efficiency of the transformation.
3A2: Quantitative information can be derived from stoichiometric calculations that utilize the mole ratios from the balanced chemical equations. The role of stoichiometry in real-world applications is important to note, so that it does not seem to be simply an exercise done only by chemists.
3B3: In oxidation–reduction (redox) reactions, there is a net transfer of electrons. The species that loses electrons is oxidized, and the species that gains electrons is reduced.

Science Practices
2.1 The student can justify the selection of a mathematical routine to solve problems.
2.2 The student can apply mathematical routines to quantities that describe natural phenomena.
3.1 The student can pose scientific questions.
3.2 The student can refine scientific questions.
3.3 The student can evaluate scientific questions.
4.2 The student can design a plan for collecting data to answer a particular scientific question.
5.3 The student can evaluate the evidence provided by data sets in relation to a particular scientific question.

Learning Objectives
1.20: The student can design, and/or interpret data from, an experiment that uses titration to determine the concentration of an analyte in a solution.
3.3: The student is able to use stoichiometric calculations to predict the results of performing a reaction in the laboratory and/or to analyze deviations from the expected results.
3.8: The student is able to identify redox reactions and justify the identification in terms of electron transfer.
3.9: The student is able to design and/or interpret the results of an experiment involving a redox titration.

Questions 5

Essential Knowledge
3C2: Net changes in energy for a chemical reaction can be endothermic or exothermic.
4A1: The rate of a reaction is influenced by the concentration or pressure of reactants, the phase of the reactants and products, and environmental factors such as temperature and solvent.
4A2: The rate law shows how the rate depends on reactant concentrations.
4B1: Elementary reactions can be unimolecular or involve collisions between two or more molecules.
4B3: A successful collision can be viewed as following a reaction path with an associated energy profile.
4C1: The mechanism of a multistep reaction consists of a series of elementary reactions that add up to the overall reaction.
4C2: In many reactions, the rate is set by the slowest elementary reaction, or rate-limiting step.

Science Practices
2.1 The student can justify the selection of a mathematical routine to solve problems.
4.2 The student can design a plan for collecting data to answer a particular scientific question.
5.1 The student can analyze data to identify patterns or relationships.
5.3 The student can evaluate the evidence provided by data sets in relation to a particular scientific question.
6.1 The student can justify claims with evidence.
6.2 The student can construct explanations of phenomena based on evidence produced through scientific practices.
7.2 The student can connect concepts in and across domain(s) to generalize or extrapolate in and/or across enduring understandings and/or big ideas.

Learning Objectives
3.11: The student is able to interpret observations regarding macroscopic energy changes associated with a reaction or process to generate a relevant symbolic and/or graphical representation of the energy changes.
4.1: The student is able to design and/or interpret the results of an experiment regarding the factors (i.e., temperature, concentration, surface area) that may influence the rate of a reaction.
4.7: The student is able to evaluate alternative explanations, as expressed by reaction mechanisms, to determine which are consistent with data regarding the overall rate of a reaction, and data that can be used to infer the presence of a reaction intermediate.

Question 6

Essential Knowledge
5A2: The process of kinetic energy transfer at the particulate scale is referred to in this course as heat transfer, and the spontaneous direction of the transfer is always from a hot to a cold body.
5B1: Energy is transferred between systems either through heat transfer or through one system doing work on the other system.
5B2: When two systems are in contact with each other and are otherwise isolated, the energy that comes out of one system is equal to the energy that goes into the other system. The combined energy of the two systems remains fixed. Energy transfer can occur through either heat exchange or work.
5B3: Chemical systems undergo three main processes that change their energy: heating/cooling, phase transitions, and chemical reactions.
5B4: Calorimetry is an experimental technique that is used to measure the change in energy of a chemical system.

Science Practices
2.2 The student can apply mathematical routines to quantities that describe natural phenomena.
2.3 The student can estimate numerically quantities that describe natural phenomena.
4.1 The student can justify the selection of the kind of data needed to answer a particular scientific question.
4.2 The student can design a plan for collecting data to answer a particular scientific question.
5.1 The student can analyze data to identify patterns or relationships.
6.1: The student can justify claims with evidence.
6.2: The student can construct explanations of phenomena based on evidence produced through scientific practices.
7.2: The student can connect concepts in and across domain(s) to generalize or extrapolate in and/or across enduring understandings and/or big ideas.

Learning Objectives
5.4: The student is able to use conservation of energy to relate the magnitudes of the energy changes occurring in two or more interacting systems, including identification of the systems, the type (heat versus work), or the direction of energy flow.
5.5: The student is able to use conservation of energy to relate the magnitudes of the energy changes when two nonreacting substances are mixed or brought into contact with one another.
5.6: The student is able to use calculations or estimations to relate energy changes associated with heating/cooling a substance to the heat capacity, relate energy changes associated with a phase transition to the enthalpy of fusion/vaporization, relate energy changes associated with a chemical reaction to the enthalpy of the reaction, and relate energy changes to PΔV work.
5.7: The student is able to design and/or interpret the results of an experiment in which calorimetry is used to determine the change in enthalpy of a chemical process (heating/cooling, phase transition, or chemical reaction) at constant pressure.

Question 7

Essential Knowledge
1E1: Physical and chemical processes can be depicted symbolically; when this is done, the illustration must conserve all atoms of all types.
2A1: The different properties of solids and liquids can be explained by differences in their structures, both at the particulate level and in their supramolecular structures.
2A3: Solutions are homogenous mixtures in which the physical properties are dependent on the concentration of the solute and the strengths of all interactions among the particles of the solutes and solvent.
3B2: In a neutralization reaction, protons are transferred from an acid to a base.
5E4: External sources of energy can be used to drive change in cases where the Gibbs free energy change is positive.
6B1: Systems at equilibrium respond to disturbances by partially countering the effect of the disturbance (Le Chatelier’s principle).
6C1: Chemical equilibrium reasoning can be used to describe the proton-transfer reactions of acid–base chemistry.
6C3: The solubility of a substance can be understood in terms of chemical equilibrium.
6D1: When the difference in Gibbs free energy between reactants and products (ΔG°) is much larger than the thermal energy (RT), the equilibrium constant is either very small (for ΔG° > 0) or very large (for ΔG° < 0). When ΔG° is comparable to the thermal energy (RT), the equilibrium constant is near 1.

Science Practices
1.4: The student can use representations and models to analyze situations or solve problems qualitatively and quantitatively.
6.1: The student can justify claims with evidence.
7.2: The student can connect concepts in and across domain(s) to generalize or extrapolate and/or across enduring understandings and/or big ideas.

Learning Objective
2.8: The student can draw and/or interpret representations of solutions that show the interactions between the solute and solvent.
2.15: The student is able to explain observations regarding the solubility of ionic solids and molecules in water and other solvents on the basis of particle views that include intermolecular interactions and entropic effects.
5.12: The student is able to use representations and models to predict the sign and relative magnitude of the entropy change associated with chemical or physical processes.
5.13: The student is able to predict whether or not a physical or chemical process is thermodynamically favored by determination of (either quantitatively or qualitatively) the signs of both ΔH° and ΔS°, and calculation or estimation of ΔG° when needed.
5.15: The student is able to explain how the application of external energy sources or the coupling of favorable with unfavorable reactions can be used to cause processes that are not thermodynamically favorable to become favorable.
5.16: The student can use Le Chatelier’s principle to make qualitative predictions for systems in which coupled reactions that share a common intermediate drive formation of a product.
6.24: The student can analyze the enthalpic and entropic changes associated with the dissolution of a salt, using particulate level interactions and representations.

Question 8

Essential Knowledge
2A2: The gaseous state can be effectively modeled with a mathematical equation relating various macroscopic properties. A gas has neither a definite volume nor a definite shape; because the effects of attractive forces are minimal, we usually assume that the particles move independently.
6A3: When a system is at equilibrium, all macroscopic variables, such as concentrations, partial pressures, and temperature, do not change over time. Equilibrium results from an equality between the rates of the forward and reverse reactions, at which point Q = K.
6A4: The magnitude of the equilibrium constant, K, can be used to determine whether the equilibrium lies toward the reactant side or product side.
6B1: Systems at equilibrium respond to disturbances by partially countering the effect of the disturbance (Le Chatelier’s principle).
6B2: A disturbance to a system at equilibrium causes Q to differ from K, thereby taking the system out of the original equilibrium state. The system responds by bringing Q back into agreement with K, thereby establishing a new equilibrium state.

Science Practices
1.4: The student can use representations and models to analyze situations or solve problems qualitatively and quantitatively.
2.2: The student can apply mathematical routines to quantities that describe natural phenomena.
2.3: The student can estimate numerically quantities that describe natural phenomena.
6.4: The student can make claims and predictions about natural phenomena based on scientific theories and models.

Learning Objectives
2.6: The student can apply mathematical relationships or estimation to determine macroscopic variables for ideal gases.
6.4: The student can, given a set of initial conditions (concentrations or partial pressures) and the equilibrium constant, K, use the tendency of Q to approach K to predict and justify the prediction as to whether the reaction will proceed toward products or reactants as equilibrium is approached.
6.5: The student can, given data (tabular, graphical, etc.) from which the state of a system at equilibrium can be obtained, calculate the equilibrium constant, K.
6.6: The student can, given a set of initial conditions (concentrations or partial pressures) and the equilibrium constant, K, use stoichiometric relationships and the law of mass action (Q equals K at equilibrium) to determine qualitatively and/ or quantitatively the conditions at equilibrium for a system involving a single reversible reaction.
6.8: The student is able to use Le Chatelier’s principle to predict the direction of the shift resulting from various possible stresses on a system at chemical equilibrium.

Question 9

Essential Knowledge
3B3: In oxidation–reduction (redox) reactions, there is a net transfer of electrons. The species that loses electrons is oxidized, and the species that gains electrons is reduced.
3C3: Electrochemistry shows the interconversion between chemical and electrical energy in galvanic and electrolytic cells.
5E3: If a chemical or physical process is not driven by both entropy and enthalpy changes, then the Gibbs free energy change can be used to determine whether the process is thermodynamically favored.
5E4: External sources of energy can be used to drive change in cases where the Gibbs free energy change is positive.

Science Practices
2.2 The student can apply mathematical routines to quantities that describe natural phenomena.
2.3 The student can estimate numerically quantities that describe natural phenomena.
6.1 The student can justify claims with evidence.
6.2 The student can construct explanations of phenomena based on evidence produced through scientific practices.
6.4 The student can make claims and predictions about natural phenomena based on scientific theories and models.

Learning Objectives
3.8: The student is able to identify redox reactions and justify the identification in terms of electron transfer.
3.12: The student can make qualitative or quantitative predictions about galvanic or electrolytic reactions based on half-cell reactions and potentials and/or Faraday’s laws.
5.14: The student is able to determine whether a chemical or physical process is thermodynamically favorable by calculating the change in standard Gibbs free energy.
5.15: The student is able to explain how the application of external energy sources or the coupling of favorable with unfavorable reactions can be used to cause processes that are not thermodynamically favorable to become favorable.

Question 10

Essential Knowledge
2D1: Ionic solids have high melting points, are brittle, and conduct electricity only when molten or in solution.
3A1: A chemical change may be represented by a molecular, ionic, or net ionic equation.
3B2: In a neutralization reaction, protons are transferred from an acid to a base.
4A1: The rate of a reaction is influenced by the concentration or pressure of reactants, the phase of the reactants and products, and environmental factors such as temperature and solvent.
4A2: The rate law shows how the rate depends on reactant concentrations.
4B2: Not all collisions are successful. To get over the activation energy barrier, the colliding species need sufficient energy. Also, the orientations of the reactant molecules during the collision must allow for the rearrangement of reactant bonds to form product bonds.
5A1: Temperature is a measure of the average kinetic energy of atoms and molecules.

Science Practices
3.3: The student can evaluate scientific questions.
5.3: The student can evaluate the evidence provided by data sets in relation to a particular scientific question.
6.1: The student can justify claims with evidence.
6.2: The student can construct explanations of phenomena based on evidence produced through scientific practices.
7.2: The student can connect concepts in and across domain(s) to generalize or extrapolate in and/or across enduring understandings and/or big ideas.

Learning Objectives
2.3: The student is able to use aspects of particulate models (i.e., particle spacing, motion, and forces of attraction) to reason about observed differences between solid and liquid phases and among solid and liquid materials.
2.8: The student can draw and/or interpret representations of solutions that show the interactions between the solute and solvent.
3.6: The student is able to use data from synthesis or decomposition of a compound to confirm the conservation of matter and the law of definite proportions.
4.1: The student is able to design and/or interpret the results of an experiment regarding the factors (i.e., temperature, concentration, surface area) that may influence the rate of a reaction.

Answers to Questions

Question 1
The WO₂F₄^2– molecular anion is an octahedron with a W atom bound to two O atoms and four F atoms. The W=O bonds are double bonds and the W–F bonds are single bonds. The W atom has six electrons available for bonding (4 points).

Draw a polar isomer and nonpolar isomer to represent the WO₂F₄^2– anion using dashes and wedges to represent bonds behind and in front of the plane (2 points).
{13781_Answers_Figure_1}
Provide a chemical explanation as to why the W atom forms double bonds to the oxygen atoms and not the F atoms. That is, why does W form shorter/stronger bonds to the O atoms (1 point)?
Owing to the inherent electronegativity difference between O and F, the W forms shorter, more covalent bonds with the O atoms than with the F atoms. The electronegativity difference between W and O is smaller than the electronegativity difference between W and F. As a result, the W=O bond is more covalent in character.
Unlike the WO₂F₄^2– anion, the NbOF₅^2– anion does not exist in a nonpolar form. Explain (1 point).
The NbOF₅^2– anion contains a single oxygen atom. As a result, there can be no symmetrical arrangement of terminal atoms about the central Nb atom.

Question 2
The following are experimentally-determined values, in kJ/mol, of successive ionization energies for a third period element: IE1 578, IE2 1820, IE3 2750, IE4 11600. That is, 578 kJ of energy are required to remove the outermost electron from 1 mole of atoms of the element, 1820 kJ of energy are required to remove the next (of the remaining electrons) valence electron, and so on (4 points).

To which element do these successive ionization energies correspond? Justify your answer (2 points).
Ionization energy is the amount of energy required to remove an electron (typically a valence electron) from an atom. Successive ionization energies become increasingly large because as electrons are removed from an atom, the nuclear pull exerted on the remaining electrons increases. The very large difference in magnitude between IE₃ and IE₄ represents the transition from an incomplete valence shell (one without 8 electrons) to a complete valence shell. IE₄ represents the removal of a core electron from an ion with a noble gas configuration. The number of ionizations it requires to reach the complete valence shell, in this case three, reveals that the element is Al.
The electron-shell model of the atom may be invoked to sufficiently explain why successive ionization energies increase in magnitude for an atom. In contrast, the quantum-mechanical model is better suited for explaining how C and O combine to form CO₂. Explain (2 points).
The quantum-mechanical version of the atom better describes how C and O combine to form CO₂. The quantum-mechanical model of the atom tells us that electrons reside in orbitals whose spatial orientations must allow for overlap if a chemical bond is to form between atoms.

Question 3
A buffer solution composed of a weak acid/conjugate base pair, HX/X^–, is blue when the HX:X^– ratio is greater than 1 and green when the HX:X^– ratio is less than 1. When the HX:X^– ratio is 1:1, the solution is colorless. The K_a for HX is equal to 1.2 x 10^–7 (10 points).

Write an equilibrium reaction to describe the HX/X^– buffer (1 point).
{13781_Answers_Equation_1}
What color is the solution when the reactant and product concentrations are given by the K_a? Justify your answer (2 points).
Blue. At equilibrium, the concentration of HX(aq) far exceeds the concentrations of H⁺(aq) and X^–(aq) because K_a of HX(aq) is very small, and the acid is very weak.
Use a particulate-level illustration to represent a 0.1 M solution of HX(aq). Use no more than 10 HX molecules and assume 10% ionization (2 points).
{13781_Answers_Figure_2}
NaOH is added to the HX/X^– buffer until a color change occurs. What color is observed? Justify your answer. (2 points).
Colorless then green. Addition of OH^–(aq) ions will convert a stoichiometric amount of the weak acid, HX(aq), to its conjugate base, X^–(aq). As a result, the HX:X ratio will decrease.
An acid’s dissociation constant, K_a, represents the ratio of the product of the product concentrations to the product of the reactant concentrations at equilibrium. How can a system be at equilibrium when this ratio does not equal 1 (1 point)?
Equilibrium denotes a system in which the rates of the forward and reverse reactions are equal. This does not require the ratio of the product of product concentrations to the product of reactant concentration be 1, only that it remain constant over time.
Assume that the dissociation of HX(aq) into its ions is an exothermic process. What color will a colorless buffer turn when heated? Justify your answer (2 points).
Blue, because the equilibrium will shift to the left according to Le Chatelier’s principle.

Question 4
A student is tasked with carrying out an oxidation-reduction titration to determine the amount of hydrogen peroxide, H₂O₂, in a standard, drug store bottle. The student uses a standardized solution of KMnO₄ as the titrant (10 points).

Write a balanced half-reaction to describe the reduction of MnO₄^–(aq) to Mn²⁺(aq) in acidic solution (1 point).
8H⁺(aq) + MnO4^–(aq) + 5e^– → Mn²⁺(aq) + 4H₂O(l)
Write a balanced half-reaction to describe the oxidation of hydrogen peroxide to gaseous oxygen and aqueous H⁺ ions in acidic solution (1 point).
H₂O₂(aq) → O₂(g) + 2H⁺(aq) + 2e^–
One of the products of the oxidation-reduction reaction of MnO₄^–(aq) in basic solution is MnO₂, an insoluble solid. Why, then, should a basic medium be avoided for this titration (2 points)?
The formation of a solid may hinder the ability to correctly determine when the endpoint is reached because there may be no clear color indication.
The student prepares a 3% H₂O₂ solution for titration by adding 10.0 mL of deionized water, 10.0 mL of 3 M H₂SO₄ and 1.0 mL of the H₂O₂ solution to a 250-mL flask.
1. For which of the three species (deionized water, H₂SO₄, H₂O₂) mentioned above is knowledge of the exact volume added to the 250-mL flask necessary? Justify your answer (1 point).
  It is necessary to know the exact volume of hydrogen peroxide solution because that is the species being studied in the reaction. If the precise volume of this solution is not known, it will limit the precision of the results to fewer significant figures.
2. If the student forgets to swirl the flask and rinse the walls of the flask before the endpoint is reached, how might the calculation of the percent H₂O₂ in the bottle be impacted? Justify your answer (1 point).
  Any hydrogen peroxide or potassium permanganate left on the reaction walls will go unreacted and cause experimental error. Hydrogen peroxide on the side of the flask will be unreacted and will cause the percent H₂O₂ to be lower. Potassium permanganate on the side of the flask will cause the percent H₂O₂ to be higher.
3. Calculate the volume of 0.02 M MnO₄^–(aq) required to completely titrate a 1.00 mL sample of 3% H₂O₂, the density of which is equal to 1.00 g/mL (2 points).
  1.00 mL solution x 1.00 g/mL H₂O₂ = 1.00 g solution
  1.00 g solution x 0.03 g H₂O₂/g solution = 0.0300 g H₂O₂
  0.0300 g H₂O₂ x 1 mol/34.0 g = 8.82 x 10–4 mol H₂O₂
  8.82 x 10^–4 mol H₂O₂ x 2 mol MnO₄^–/5 mol H₂O₂ = 3.53 x 10^–4 mol MnO₄^–
  M = mol/L
  L = moles MnO₄^–/Molarity MnO₄^–
  L = 3.53 x 10^–4 mol/0.02 M = 0.0176 L = 17.6 mL MnO₄^– solution
The permanganate ion is easily precipitated from solution with alkylammonium cations, polyatomic ions with nitrogen centers. One example, tetrabutylammonium (TBA), forms a stable permanganate salt at room temperature. In contrast, tetramethylammonium (TMA) permanganate is extremely sensitive to shock and highly explosive. The act of scraping it from a glass beaker with a metal spatula can initiate an explosion. The significant difference between the two cations is size. The four methyl groups in TMA are significantly smaller than the four butyl groups in TBA. Use this information to explain why tetramethylammonium permanganate is shock-sensitive and tetrabutylammonium permanganate is stable under normal conditions (2 points).
The larger butyl groups block the nitrogen center and prevent a strongly exothermic redox reaction from occurring between the permanganate ion and alkylammonium cation’s center.

Question 5
A student is tasked with designing an experiment to determine the rate law (i.e., the orders of reaction) for the iodination of acetone. Fortunately, the reaction’s progress can be monitored easily via color change owing to a distinct transition from yellow to colorless. The rate law for the reaction can be written rate = k[acetone]^m[H⁺]ⁿ[I₂]^p. The reaction is zero order with respect to iodine and the rate of reaction may be reported as [I₂°]/Δt. The following materials are available (10 points).

{13781_Answers_Table_1}

How might the student confirm, experimentally, that the reaction order with respect to iodine is zero (2 points)?
The student can vary the concentration of iodine while holding the concentrations of the other reactants constant. The student should note no change in the reaction rate. That is, the time required for the solution to transition from yellow to colorless should not change.
If each experimental observation of a yellow to colorless transition (i.e., the iodination of acetone proceeding to completion), is considered a “run,” what is the fewest number of total runs necessary to determine the two remaining reaction orders, assuming the initial iodine concentration is held constant? Justify your answer (2 points).
Three. For example:

{13781_Answers_Table_2}
The student wishes to change the reaction rate of each run in part b without altering reactant concentrations. Given the available supplies, what can the student do to accomplish this goal? Justify your answer (2 points).
Cool the reactions in an ice water bath because temperature affects reaction rate.
An Internet source claims the iodination of acetone proceeds via a single, elementary step in which two molecules of I₂(aq) combine with one molecule of H⁺(aq) and one molecule of aqueous acetone. Is the claim reasonable? Justify your answer (2 points).
This claim is not reasonable because the experimentally-determined reaction order of iodine is zero. Any mechanism or elementary step that has two iodine molecules, or is bimolecular with respect to iodine, is not feasible.
Assume that the enthalpy of reaction for the iodination of acetone is negative. With this in mind, draw a rudimentary diagram (with energy marked on the y-axis and reaction process marked on the x-axis) to describe the iodination of acetone (2 points).
Drawing will vary. However, the reaction is exothermic and so the products will be lower in energy than the reactants.

{13781_Answers_Figure_3}

Question 6
A student is given the task of determining the specific heat capacity of a chunk of metal of unknown composition. The following materials are available (10 points).

{13781_Answers_Table_3}

According to the laws of thermodynamics, the heat lost by one substance is exactly equal to the heat gained by the other substance in a closed, two-substance system.
1. Explain how, on a microscopic level, thermal energy (heat) is transferred from a hot chunk of metal to liquid water (at a lower temperature than the metal) in a confined space (1 point).
  The high-temperature metal atoms are on average higher in kinetic energy than the water molecules. The metal atoms’ energy results in translational motion. They move on an atomic scale and collide with neighboring water molecules. As a result, energy transfer occurs.
2. From an experimental design standpoint, explain how the student can heat the metal chunk to an observable/measurable (with a thermometer) temperature? Explain (1 point).
  The student can heat the metal by placing it in a boiling water bath. The student does not have the appropriate equipment to measure the metal’s temperature directly and so must heat indirectly in water and take the heating agent’s temperature as the temperature of the metal. The metal should be left in the boiling water for a time sufficient to ensure heat transfer from the water.
The student measures the temperature of 35.00 g of distilled water: 24.2 °C. The student then places the 2.33 g chunk of metal, heated to 100.2 °C, into the 35.00 g of distilled water.
1. What should the student use to contain the metal/water mixture? Justify your answer (1 point).
  An insulating material such as a styrofoam cup covered with Al foil to minimize heat loss to the surroundings.
2. At some point, following immersion of the hot metal into the room temperature water, a thermal equilibrium between the metal and water is established. How does the student determine the point at which thermal equilibrium is established (2 points)?
  The metal/water mixture reaches a maximum temperature, which then drops as the metal/water mixture loses energy to the surroundings (i.e., the cup and air directly outside the cup). The student must watch for a maximum temperature following addition of the metal to water.
3. The student stirred the mixture as it approached its maximum temperature. Why is this a good idea, from an experimental design perspective (2 points)?
  Heat is transferred between objects by means of contact. If the high kinetic energy atoms (on average) comprising the metal piece contact a greater number of molecules in the water, the heat transfer will be more thorough.
Calculate the experimental value for the specific heat capacity of the unknown metal assuming the maximum temperature of the metal–water mixture was found to be 44.5 °C. C_water = 4.18 J/g x °C (1 point).
q_metal = – q H₂O
m_metal x C_metal x ΔT_metal = – (m_water x C_water x ΔT_water)
m_metal x C_metal x (T_f – T_i) = – [m_water x C_water x (T_f – T_i)]
2.33 g x C_metal x (44.5 °C – 100.2 °C) = –[35.00 g x 4.18 J/g x °C x (44.5° C – 24.2 °C)]
C_metal = –[35.00 g x 4.18 J/g x °C x (44.5 °C – 24.2° C)]/[ 2.33 g x (44.5 °C – 100.2 °C)]
C_metal = 22.8 J/g x °C
The student’s teacher informs him/her that the actual specific heat capacity of the metal is 24.3 J/g x °C. Provide one reason to account for the difference between the student’s experimental value and the actual value. Justify your answer (2 points).
The system is not closed because the cup is not a perfect insulator. Significant heat will be lost to the surroundings and the specific heat capacity of the metal will be underestimated.

Question 7
Dissolution and precipitation are two processes that typically occur with heating and cooling, respectively. That is, to increase an ionic compound’s solubility in water, one must typically increase the solution temperature. In contrast, cooling a solution is often used to facilitate precipitation (10 points).

Consider the dissolution of NaCl in water. Draw a particle-level diagram to illustrate the interactions in an aqueous solution of NaCl. Identify ions with symbols and charges and take care to arrange solution particles in the proper orientation. Use no more than five NaCl formula units and no more than five water molecules (2 points).
{13781_Answers_Figure_4}
Consider the dissolution of a sparingly soluble salt such as Al(OH)₃ in water. The K_sp of Al(OH)₃ in aqueous solution is equal to 1.3 x 10^–33. Qualitatively, what would the addition of 10 g of solid Al(OH)₃ to 100 g of room-temperature water look like (2 points)?
The solid would not dissolve to any appreciable degree and would settle to the bottom of the beaker/container.
Predict whether the enthalpy of dissociation of Al(OH)₃ into its component ions is larger or smaller in value than the enthalpy of dissociation of NaCl. Assume both are positive values. Justify your answer (2 points).
The enthalpy of dissociation, or ΔH_rxn, for the dissociation of Al(OH)₃ into its component ions, is likely a very large, positive number attributable to the fact that a great deal of energy is required to break the intramolecular forces (chemical bonds) in Al(OH)₃ to dissolve it in aqueous solution. In contrast, NaCl would require little energy to dissolve it into solution and so would have a smaller value.
Consider the addition of HCl to an aqueous solution of Al(OH)₃ in which a small amount of the solid has been dissolved. Will addition of HCl to the solution lead to an increase or decrease in the concentration of Al³⁺(aq) ions? Justify your answer (2 points).
Added HCl, or H⁺(aq) ions, will neutralize free OH^–(aq) ions, thereby lowering the concentration of OH^–(aq) in solution. As a result, the equilibrium will shift to the right to produce more OH^–(aq) in response and the Al³⁺(aq) concentration will increase according to Le Chatelier’s principle.
Explain, with reference to the term ΔG, why high temperatures might help to dissolve Al(OH)₃ to an appreciable degree in aqueous solution. Assume that ΔH is positive (2 points).
ΔG = ΔH – TΔS: Reactions occur spontaneously when ΔG < 0 and require intervention to occur when ΔG > 0. As temperature increases, the (–TΔS) term increases in magnitude if ΔS is positive. At high enough temperatures, the negative value of the term exceeds the positive value of the ΔH term and the reaction proceeds.

Question 8
A sample of N₂O₅ gas is introduced into a cylinder with a piston. N₂O₅ decomposes and reaches equilibrium with its decomposition products according to the following equation: (10 points)

{13781_Answers_Equation_2}

Assume that N₂O₅ is brown and N₂O₃ is clear/colorless. How, via macroscopic observation, is it possible to determine the time necessary for the system to reach equilibrium? Justify your answer (2 points).
When equilibrium is reached, the color will not change because the reactant and product concentrations are constant.
Write the K_p expression for the equilibrium (1 point).
K_p = (P_O2 x P_N2O3)/P_N2O5
The K_p for the equilibrium is 2.3 x 10^–5. Determine the partial pressures of all gaseous species present if the initial N₂O₅ pressure is 0.12 atm (2 points).
{13781_Answers_Table_4}

K_p = (x²)/(0.12 – x)
2.3 x 10^–5 = x²/0.12
X = (2.76 x 10^–6)^½ = 0.00166
P_N2O5 = 0.12 atm, P_N2O3 = 0.0017 atm, P_O2 = 0.0017 atm
The piston is withdrawn slowly and then held in place. Will the value of K_p increase, decrease, or stay the same before returning to its original value? Justify your answer (2 points).
Increase. Withdrawing the piston will result in a pressure decrease. In response, the system will shift to the side (right) with more relative moles of gas so as to counteract the pressure decrease according to Le Chatelier’s principle.
In another experiment, the flask is charged with N₂O₅, N₂O₃ and O₂ at the following partial, non-equilibrium pressures: 1.1 atm, 0.8 atm and 2.4 atm, respectively. Predict whether the amount of gaseous oxygen in the container will increase, decrease or stay the same. Justify your prediction (2 points).
Q = (P_O2 x P_N2O3)/P_N2O5 = (2.4 atm x 0.8 atm)/1.1 atm = 1.75 = 1.8
Q > K_p, as a result the reaction will shift to the left to reestablish equilibrium and the amount of gaseous oxygen will decrease.
Explain, with reference to microscale events, why the pressure of the system increases when the equilibrium shifts to the right (1 point).
When the equilibrium shifts to the right, additional gas molecules are produced in the closed system. As a result, an increased number of collisions with the container walls (i.e., an increase in pressure) occurs.

Question 9
Metal surfaces are often plated with a layer of chromium metal to improve their appearance. An experiment is set up to determine the viability of plating chromium metal onto copper. Two strips of copper metal are placed into a beaker containing a 1.0 M solution of chromium(III) chloride. The reaction to plate chromium(III) ions onto a copper surface is shown: (4 points)

2Cr³⁺(aq) + 3Cu(s) → 2Cr(s) + 3Cu²⁺(aq)

Determine the standard cell potential, ε˚, and the value of the free energy, ΔG˚, for the above reaction (1 point).
2 x [Cr³⁺(aq) + 3e^– → Cr(s)] E° = – 0.73 V ΔG° = –nFE°
2 x [Cu(s) → Cu²⁺(aq) + 2e^–] E° = – 0.34 V ΔG° = –(6 moles electrons) x (96486C•mol^–1) x (–1.07 V)
E°_cell = –1.07 V ΔG° = 619440•12 J = 619 kJ
Will the chromium ions plate out on the copper strip? Explain why or why not in terms of free energy (1 point).
The chromium ions will not plate onto the copper strips because the value for the Gibbs Free energy of the redox reaction is positive. A spontaneous reaction will only occur if the value of ΔG° is less than zero. The plating of chromium ions onto copper is a nonspontaneous reaction.
The copper strips are now attached to a power supply. Is it necessary to apply a voltage to the copper strips in order to reduce the chromium ions in solution for plating onto the Cu metal? If so, what voltage should be applied? Justify your answer (2 points).
Yes, a voltage of at least 1.07 V is required to plate the chromium ions onto the copper. In order for the chromium ions to plate (be reduced), ΔG° must be less than zero. Therefore, a voltage greater than 1.07 V must be applied to overcome the standard cell potential of –1.07 V.

Question 10
Solid CaCO₃ reacts with HCl(aq) according to the equation given below. A student is tasked with determining the reaction’s rate, which may be monitored via collection of CO₂ gas in a large syringe (10 points).

CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

In a typical experiment, the amount of CO₂ increases rapidly at first and in a linear fashion but then begins to level off after some time. Justify this observation (2 points).
Since the reaction rate is proportional to the concentration of reactant(s), and the concentration of reactant(s) decreases as the reaction proceeds, the rate levels off as more and more of the reactant(s) are consumed.
The reaction of CaCO₃ with HCl takes place in an Erlenmeyer flask to which a syringe is affixed with a rubber adapter. Upon addition of the HCl to the solid CaCO₃ care must be taken to very quickly affix the rubber adapter/syringe to the flask. Why is this the case (1 point)?
Any CO₂ that escapes prior to capping of the Erlenmeyer flask will not be measured. The faster the flask is capped, the less CO₂ will escape and the reaction’s rate better assessed.
Write a general rate law (i.e., one with reaction orders represented by letters) to describe this heterogeneous reaction (1 point).
rate = k[HCl]^x
Reaction rates are typically temperature dependent. Explain why temperature increases typically lead to reaction rate increases (2 points).
As temperature rises, the average kinetic energy of molecules increases and the number of collisions leading to chemical reactions increases.
The CaCO₃ is ground to a fine powder. Will the rate of the reaction increase, decrease, or stay the same? Justify your answer (2 points).
Increase because the smaller, average particle size will produce greater surface area allowing for more collisions between reactants.
Describe an alternative method for monitoring the rate of CaCO₃ decomposition (2 points).
Monitor mass loss over time by carrying out the reaction on a scale.

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