Chemistry Solutions

The Advanced Placement Program Reading (AP Reading) is an annual event in June, where AP teachers and college faculty gather to score all the AP exams taken by students during the previous month.

I always enjoy my experiences serving at the AP Chemistry Readings because they help me to learn and improve as a teacher. I especially value the opportunities to discuss student responses with my colleagues while considering the seemingly innumerable ways students find to approach answering the same questions.

After each AP Reading, I compile all my notes—both written and mental—to determine how I might use those insights to develop strategies that help students deepen their content understanding and ultimately improve performance on the AP exam.

I have served as a Reader, a Table Leader, and a Question Leader over the past sixteen AP Chemistry Readings. In July 2026, I was the Question Leader for this year’s Question 5. The article that follows is a collaboration with Table Leaders from the six other exam questions. The goal is to bring you advice and perspectives gathered during the 2026 Reading. Each leader gives a brief overview describing topics within the question they scored where students notably excelled or struggled. They each then provide suggestions based on their experiences for teachers looking to improve how their students perform.

This article supplements the in-person and online presentations given by the Chief Reader for AP Chemistry Dr. Aileen Beard. The “AP Chemistry – Results from the 2026 Exam” presentation was given as part of the George R. Hague Memorial AP/IB Chemistry Symposium at the 2026 Biennial Conference on Chemical Education (BCCE) in Madison, WI. A similar presentation will also be shared as an AACT webinar on Wednesday, September 23, 2026.

I recommend viewing the following information while reading this article:

• 2026 AP Chemistry Exam Free Response Questions
• My preliminary draft version of the 2026 scoring guidelines; the official scoring guidelines have not yet been released at the time of this writing
• My Summary of Mistakes and Misconceptions, a point-by-point list to supplement this article

Question 1 – Ryan Johnson, Table Leader

Ryan Johnson is a veteran AP Chem teacher with ten years of experience and has served as an AP Reader and Table Leader for eight years. He is now a college chemistry instructor at the United States Air Force Academy and Pikes Peak State College in Colorado Springs, CO.

• Question 1 centered on the compound potassium chloride, KCl. Major topics assessed were atomic structure, thermochemistry, and solubility equilibria.

Figure 1. Bohr models of potassium atom and ion. Both contain the same number of protons to attract the electrons. Potassium atoms have four occupied energy levels, while potassium ions have three occupied energy levels. The atom has a larger radius than the ion because it has more occupied energy levels, and its electrons are therefore farther from the nucleus.

Atomic Structure (Part A): Writing an electron configuration was an accessible point for most students, with common mistakes relating to correctly interpreting the total number of electrons in the K⁺ ion.

Students correctly identified the potassium atom as having a larger radius than the potassium ion (Figure 1), but many responses lacked sufficient justification. Many students attempted to justify the choice using principles of coulombic attraction or periodicity, but they failed either to include structural features or to provide any comparative reasoning.

Tips for Teachers:

• Make clear to students the distinction between an electron configuration, an orbital diagram, and a Bohr model drawing.
• Direct students to practice writing the complete electron configurations for both atoms and ions.
• Remind students frequently, “Electrons are negative!” to help them remember to pause and consider the whole system of charges that makes up an atom or ion. Removal of (negative) electrons results in a positive species.
• Bohr-style models are rarely necessary to include in a response, but they can be helpful during instruction. Points to emphasize:
  • Energy levels are not physical spaces and, as a model, are available to any atom or ion, whether occupied or not.
  • Replace phrases like “contains more energy levels” with “contains more occupied energy levels.”
• Relative atom and ion sizes can usually be compared by addressing two features: (1) the number of protons in each nucleus and (2) the number of occupied energy levels. For example, students can correctly use coulombic attraction ( ) to compare atoms and/or ions in different scenarios:
  • Separation (r) as the dominant factor: typical when two nuclei contain the same number of protons, while one species contains electrons farther away than the other (as was the case in this question)
  • Magnitude of charge (q) as the dominant factor: typical when two atoms or ions contain electrons in the same number of energy levels and have a different number of protons in their nuclei
  • Electron-electron repulsion as the dominant factor: typical when comparing atoms or ions of the same element, with equal numbers of protons in their nuclei and equal numbers of occupied energy levels but with different numbers of electrons in the outer shells of the two.

Thermochemistry (Parts B-E): Students analyzed calorimetry data from the dissolution of potassium chloride and performed some calculations.

It was very clear that the vast majority of students (just like years and years of my own chemistry students!) had experienced mostly exothermic reactions and not endothermic reactions. The data in the question show a temperature decrease upon dissolving, and therefore dissolution of KCl is an endothermic process. This distinction was important because it could affect how students worded their responses in several parts of the question.

Figure 2. Particle diagrams can help students analyze energy flow in various scenarios. These images show possible ways to represent a process occurring in an imperfectly insulated calorimeter. If the water and surrounding air start at the same temperature, then choice A best represents an endothermic dissolution.

Many students struggled with integrating all the parts of the calorimetry experiment. Students generally knew to use q = mc∆T to calculate the magnitude of thermal energy transfer, but used incorrect variations for the mass. As the process involves a solution (single phase), the mass should include everything in the solution (water + dissolved salt). Students also had an incomplete understanding of the relationship between thermal energy (q) and enthalpy of solution (ΔH_soln).

Conceptually, students clearly needed more practice with identifying heat flows throughout a chemical process and within a system-and-surroundings set-up. See Figure 2.

This may be complicated by the fact that dissolving occurs at the ion-water interface, even though it takes place within a solvent containing many water molecules not involved in the initial thermodynamic process. This combines the system and surroundings into a single phase (the solution)—so it is easy to confuse the meaning of the solution’s temperature decrease.

One part of the question proposed an identical experiment with only a change from potassium chloride to rubidium chloride, both having approximately equal values of ΔH_soln.

With similar heats of solution, the scenario required students to realize that a given mass of RbCl contains fewer moles of substance than the same mass of KCl, due to the larger molar mass of RbCl. A sample with fewer moles would have less heat transfer during dissolution, leading to a smaller change in temperature.

Many responses included some correct ideas but failed to connect them into an explanation. Some stated that because everything else was equal, the change in temperature must also be the same. A few students also tried to answer this question through differences in lattice energy, citing the difference in potassium and rubidium ion sizes. While lattice energy arguments may be relevant in a different scenario, the differing number of moles in this scenario limits the validity of those arguments.

Tips for Teachers:

• Expose students to different kinds of calorimetry problems and experiments, while including both exothermic and endothermic reactions.
• Ask students to draw a picture of their lab calorimeter—or give choices as in Figure 2—and use arrows to show the direction of heat flow between the reaction—or dissolving substance—and the surrounding solvent, as well as between the calorimeter and the surroundings.
• Show students that the desired units can be like a key for what needs to be done with the problem. The “kJ/mol” noted in the prompt shows that the heat (kJ) must be divided by the moles! The rest just becomes a quick calculation.
• Help students practice error scenarios and predicting how lab data would change with different variables to help familiarize students with the Science Practices being assessed. Comparing the same experiment by using calorimeters made of different materials (e.g., a paper cup and a Styrofoam cup) would help students to better understand error analysis as described in this question.
• In a laboratory setting, the solution’s temperature does not stop changing when the process is complete because it will eventually reverse back as thermal equilibrium is re-established with the room. This fact could be emphasized by providing lab analysis or practice questions that show data or graphs extending beyond the initial increase or decrease of temperature.

Figure 3. In Question 1 Part C(i) initial and final temperatures have three significant figures. The final calculation uses ΔT, which contains only two.

• Remind students that a calculated answer may have fewer significant figures than any single data point due to addition or subtraction within a calculation. Adding and subtracting can change the number of significant figures. For example, when solving for q = mcΔT, the values for mass, heat capacity, and temperatures in the data table all have three significant figures, but the ∆T value used in the calculation has only two significant figures. See Figure 3.

Solubility (Part F): The final part of Question 1, which proved challenging for students, tested their understanding of solubility and K_sp. Writing a net ionic equation for RbCl dissolution was a generally accessible point for students, but the expected errors were still common with chemical species being represented incorrectly, such as without the appropriate charges.

Students struggled with the relationship between molar solubility and K_sp. Some thought K_sp was simply an abbreviation for molar solubility. Others over-generalized equilibrium principles and applied them incorrectly. Some students started with a proper mathematical setup but did not follow it through to an answer, as though they had abandoned the idea as a plausible strategy.

In addressing how the solubility of RbCl would differ in a KCl solution from its solubility in pure water, the most common mistake was to invoke Le Châtelier’s Principle, saying that the equilibrium will “shift toward the reactants.”

Although equilibrium principles are at play during the dissolving process, Le Châtelier’s Principle only applies when a system that is already at equilibrium is somehow changed. In this question, no initial equilibrium exists; the new equilibrium established once enough RbCl is added will simply settle at a different position—specific concentrations of all species—than it would if chloride ions were not already present.

Tips for Teachers:

• Emphasize that net ionic equations are a way of identifying only the species that are changing during the process being represented. Although water plays a role in the dissolving process, its molecules do not change form and are not included in the net ionic equation for a dissolution.
• Challenge students to recognize when Le Châtelier’s Principle is useful (e.g., when addressing what happens when equilibrium is disturbed), and when it is not (e.g., when equilibrium is not the initial condition).
• Address molar solubility as a maximum amount of dissolving before studying K_sp. This can be as simple as dissolving table salt in water—as a demo or a short lab—to determine the maximum moles that can be dissolved. The molarity of the saturated solution—its molar solubility—can be calculated and used to introduce the K_sp equation.
• Use a quick demo or add a short piece onto an existing K_sp lab to show students an example of the common-ion effect on solubility.
• Use lots of particle diagrams when analyzing equilibrium systems! These diagrams can help students have a better mental framework to rely upon when asked about new systems. Check out these AACT equilibrium resources:
  • Animation Activity: Equilibrium – short animation with student worksheet
  • Equilibrium in a Beaker – students use manipulatives to model various systems
  • Le Châtelier's Principle Particulate View – teams work to draw particle models representing various equilibrium shifts

Question 2 – Karen Compton, Table Leader

Karen Compton is entering her twenty-third year of teaching and seventeenth year teaching AP Chemistry. She currently teaches at Plano East Senior High School in Plano, TX, and has been an AP Reader and Table Leader for ten years.

• Question 2, which studies a variety of chromium-oxygen compounds and ions, assesses principles of molecular bonding, oxidation-reduction reactions and electrochemistry, and kinetics

Figure 4. A completed Lewis structure for the chromate ion, CrO42-.

Molecular bonding (Part A): It became clear that students need practice identifying molecular geometry (Topic 2.7). Using model kits, for example, can help deepen their understanding as they build molecules, ions, etc. and analyze the molecular geometry. The expanded valence structure in chromate (CrO₄^2-) may have been unfamiliar to some students, as responses showed much confusion about how electrons might be distributed to meet the description in the question. See Figure 4.

Tips for Teachers:

• Direct students to practice drawing Lewis structures for molecules and ions, taking care to show the proper number of electrons. Teach them that Lewis structures are a type of model and that the model is only valid when all electrons are shown. Lone pairs of electrons have significance when using the two-dimensional Lewis structure models to predict three-dimensional shapes and to consider how different molecules might interact with each other.
• Focus instruction on characteristics of atoms that drive the relative number of electrons each atom may accommodate. Formal charges can be used to consider the likelihood of a given electron arrangement:
  • More probable structures have fewer and lower-magnitude formal charges.
  • Atoms with high electronegativity are more likely to accommodate negative formal charges.
  • Only atoms containing theoretical d-sublevels (period three and beyond) can accommodate more than eight electrons
• Check out AACT’s resource topic pages on Lewis Structures and Molecular Geometry.

Redox and Electrochemistry (Parts B-D): This section was relatively low-scoring. Evidently, students needed more practice with net ionic equations. Some responses showed all reactants and products as monatomic ions; some used correct species but neglected to show the strong acid as dissociated; and several others displayed the correct equation but without coefficients to balance it.

Similarly, many students correctly identified the reaction as non-redox but failed to justify their answer using oxidation states—or claimed the oxidation state remained the same without offering any evidence. When asked to use a calculation of ΔG to justify the thermodynamics, some students cited the correct equation with no math or explanation. Others seemed to not realize the relationship between E and ΔG and therefore tried to use other equations related to ΔG (but not related to this prompt), such as “products – reactants” or ΔH – TΔS.

Additionally, some students mixed up the relationship between the sign of ΔG and favorability. For the electroplating question, many responses correctly used from the reference sheet to solve for q, but the work essentially ended there. Others confused electron moles with plated metal moles.

Tips for Teachers:

• Teach net ionic equations for different reaction types. Students should be able to do the following:
  • Identify whether species should be written as ions or compounds in the given reaction.
  • Write the correct chemical formula for each reactant and each product, including ion charges when applicable.
  • Eliminate any spectator ions.
  • Balance the equation so that both the number of each type of atom and the total charge match on each side.
• Briefly discuss electricity as a flow of charge, helping students distinguish the number of electrons involved (based on current) from the amount of chemical substance reacting as a result.
• While reviewing before the AP Exam, give students a clean copy of the formulas and equations sheet and challenge them to recall and explain scenarios where each type of equation might be used. Doing so can help to address mistakes students made in Question 2 with various ΔG relationships.
• Reinforce with students the importance of paying attention to the information provided in the question. This will help them avoid losing points because they didn’t follow the directions.
  • I credit Jamie Benigna, AP Chemistry Director of Curriculum, Instruction, and Assessment, for giving me the abbreviations: RTQ (read the question), ATQ (answer the question), and ATFQ (answer the full question), which I frequently use on student feedback to remind them of the nuances they can miss when rushing through their assessments.
• Model and stress the importance of using units in all calculations. Encourage students to write all numeric values substituted in for a mathematical equation, rather than only writing the equation as symbols that lead to a final numeric answer. This makes it easier to demonstrate proper chemical understanding in a way that is more likely to earn points.
• Remind students that dimensional analysis can be their friend with many types of calculations. Teach them how to use the units to guide or check their calculations!

Kinetics (Parts E-G): It was difficult to determine whether the low overall performance in these sections resulted from poor understanding of kinetics concepts in general or poor understanding of the types of mathematics that explain the kinetics.

A particular disconnect appeared with translating between graphical representations of kinetics data and the mathematical equations that model that same data. When prompted to use specific “data from Figure 3” (Part E of question two) to support a conclusion, many responses left me wondering which data the students meant—no actual numbers appeared in the explanation.

Other responses made correct claims (e.g., relating first order reactions to a linear plot of ln[Cr₂O₇^2-] vs time) but that was not the plot shown in Figure 3. Students made no connection between what was in the figure to that type of plot!

When using information from a graph to calculate a value, an alarming number of students took the natural log of the natural log of concentration. Yes, you read that correctly: they applied ln a second time to a value that was already the ln of a concentration from the data. With those values now negative, one can imagine the confusion that caused—since the natural log of a negative number is undefined.

Further, most students were unable to interpret the effect of simply changing the initial concentration on the reaction kinetics, possibly because the response had to be drawn on the graph. See Figure 5. Most students would likely be able to correctly state that halving the initial concentration of a first order reaction would proportionally halve the initial rate. The graph forces students to think beyond initial proportionality to consider the pattern in the overall reaction.

Tips for Teachers:

Figure 5. Linear plot of ln(concentration) vs. time, as similar to Figure 3 of exam Question 2. Original data are plotted in blue. The correct starting point for the predicted graph was given (red arrow). Changing the initial concentration does not change the mathematical pattern of decrease. So the new slope should match that of the original, as shown by the red line running parallel to the original plot.

• Give students practice with writing justifications from specific data sets. On the AP exam, a prompt directing students to refer to the given data or graph is specifically testing their ability to make meaning of that data within the topic's scope. Students can perform poorly, even with a strong grasp of a given topic, when they have not had enough experience using the Science Practices within a topic’s content.
• When students are studying kinetics (Unit 5), teach them to articulate the differences in zero order, first order, and second order rate processes in terms of: 
  • describing how the rate changes over time and with varying concentration
    
  • characteristic features of the graphical plot

• Ensure students are proficient with calculations involving the integrated rate laws. This work should include calculating the following values and expressing the answers with the correct units:
• the value of the rate constant k
• the half-life for a first order process
• the concentration of the reactant at a certain point in time during the reaction
• the time required for the initial concentration of the reactant to decrease to a certain value

Question 3 – Matt Ford, Table Leader

Matt Ford has taught AP Chemistry for over twenty years at St. John's Preparatory School in Danvers, MA. He became an AP Reader a decade ago and has been a table leader for the last two years. He recently served a term on the AP Chemistry Development Committee.

• Question 3 explored the properties of nitrous acid (HNO₂₎. The first half of the question focused on acid-base equilibrium and the second half was built around a titration of nitrous acid with sodium hydroxide, NaOH.

Student responses were most successful when the task verb was “identify” or “draw,” as in Parts A, D, and F. When the task verb was “calculate,” the responses demonstrated understanding when the calculation was short and straightforward, as in Parts B(i) and C(i). They were less successful when the calculation became trickier or had more steps, as in Parts B(ii), Part E, and Part G.

Responses to questions requiring students to make and justify a claim, as in Part C(ii) and Part H, varied with some solid answers and some that were incomplete. In general, the responses followed a predictable pattern: student performance decreased when the cognitive load of the task increased.

Acids and Bases (Parts A,D-H): While students could generally identify a Brønsted-Lowry acid-base conjugate pair, many did not earn the point because they omitted charges, thus “identifying” the incorrect species. As the form and charge of a substance are fundamental to this topic, that distinction is important for demonstrating understanding. 

Brønsted-Lowry acid-base chemistry is all about exchanging H⁺ ions (aka “protons”). When the correct charges are not labeled on both species, the response does not clearly show an understanding of these proton transfer reactions.

It was also clear that students struggled with the math of exponents and logarithms. The equation sheet shows how to calculate pH when the concentration of hydronium ions is known (e.g., pH = -log[H₃O⁺]). When instead using a pH value to calculate [H₃O⁺], this rearranges to [H₃O⁺] = 10^-pH. Some students incorrectly used “e” or “E” rather than the base of “10” for this calculation. 

In a later question, some claimed that 5.6 x 10^-4 was a larger value than 1.04 x 10^-3, possibly because 5 is greater than 1 or because 4 is greater than 3. This led to either incorrectly claiming the reaction was exothermic or giving the correct claim with flawed justification.

Figure 6. Calculation setup for Part G, where many students had the correct values and were unable to calculate the correct final answer.

Later, when calculating an equilibrium value from other related values, many students arrived at the correct setup but then could not properly execute the calculation. See Figure 6. It was not clear if this was an algebraic error or a calculator error. Taken collectively, these mistakes and misconceptions show that many students lack both the computational skills and conceptual understanding necessary to work with exponents and logarithms

Tips for Teachers:

• Emphasize that charges are very important in chemistry. Whether a species is a neutral atom or a charged ion matters greatly. Unfortunately, many students do not pay attention to these details which can lead to gaps in their understanding of several different concepts.
• Give students practical exposure to explaining and justifying phenomena using acid-base titrations in the laboratory.
• Michael Farabaugh’s video “Particle Diagrams for Acid-Base Titrations” shows one method for this. Links to the worksheets are below the video.
• Students can also draw a model titration curve to use while titrating in the lab, pausing at various points of the titration to consider equilibrium principles, to identify all major species present in the reaction mixture at each point.
• Students can use acid-base indicators for various titrations while also tracking the pH, to observe how the color changes correlate with the pH curve. This can lead to discussions on how to choose the best indicator when pH is not tracked.
• Try out the AACT “Acid-Base Titrations Simulation” that allows for repeated calculation practice in interpreting titration curves.

Equilibrium (Parts B-C): Initial, Change, Equilibrium (ICE) charts are a powerful tool for solving equilibrium problems and organizing data, but no part of Question 3 required a chart. Nonetheless, students attempted to use these charts in multiple ways.

Figure 7. Example of an unnecessary ICE chart setup to solve a simple problem.

Figure 7 shows how an ICE chart could be used to solve Part B(i), but students should recognize that the answer comes from a simple subtraction and that setting up the ICE chart took time that they could have used elsewhere. While some students were able to reach correct answers, they turned one-step solutions into elaborate multi-step solutions, thus using up valuable time.

Most students correctly concluded that the reaction was endothermic by comparing the equilibrium values at two different temperatures. Mistakes on this question were largely related to interpreting exponents, as described above.

Tips for Teachers:

• Include single-step questions on both practice problems and assessments. Many teachers draw their assessments from old AP questions, but using only these questions obscures the ability of our students.
• Assign more practice with equilibrium questions that do not require ICE charts, K expressions, and potentially quadratic formulas. When students see an ICE chart on the board every day, they may start to assume that every time they see a K value and numbers, they must do an ICE chart and then solve an algebraic expression. This type of algorithmic thinking is comforting for students, but it can also lead them astray.
• Provide adequate opportunities for students to:
• perform calculations with the equation pH = –log[H₃O⁺]
• write the equilibrium constant expression, K, for a reaction
• substitute the equilibrium concentration values into the K expression
• analyze relationships between K values for related reactions, as described in the essential knowledge statements for Topic 7.6

Additional Teacher Tips for All Units:

• When possible, assign explanations to accompany mathematical answers. You can also try beginning each topic with conceptual understanding, and bring in the math only after students show proficiency in the concepts.
• Look for ways to increase students’ practice working through various chains of logic. One simple way to do this is to add a question or two to lab activities that require students to use specific data from that lab—combined with contextual understanding—to justify a claim beyond what they directly observed. Claims might be related to:
  • error analysis or considering ways to minimize error
  • how results would change if one substance or component of the experiment is changed
  • how a related experiment might be designed, as compared to the one from the lab
  • how a chemical product of the lab might be characterized or analyzed for purity
• As a general suggestion, I highly recommend teaching students the Claim, Evidence, Reasoning (CER) framework. Although not required on the exam, this can help students to avoid lost points due to incomplete justification. Many resources go into depth about how to teach CER, including this AACT webinar. Table 1 shows examples of using the CER framework for the “justify” prompts from question three.

Table 1. Example answers for Question 3, parts C(ii) and H, using the CER framework. 

Question 4 – Jamie Flint, Table Leader

Jamie Flint has been teaching AP chemistry at Spring Woods High School in Houston, TX for twenty-two years. She has been an AP Reader since 2015 and a Table Leader since 2025.

• Question 4 briefly assessed the bonding, equilibrium reactions, and thermodynamics of white phosphorus, (P₄).

Bonding (Part A): Most students could apply appropriate bonding principles, but many used statements that were either too general or didn’t connect their statements to the comparison. For example, stating that, “a triple bond is the shortest bond,” doesn’t address the other molecule and, thus, does not sufficiently answer the question.

Equilibrium (Part B): Students performed very low when they had to use stoichiometry to determine an equilibrium pressure, given both the initial and equilibrium pressures of the other species. Most students that set up an ICE table did so correctly and earned the point. However, a common error was overlooking stoichiometry, which led to a negative P₄ pressure value—an impossibility. Most students were able to properly substitute equilibrium values into the expression to determine the value of the constant. Those that did not succeed either did not show work, substituted values incorrectly, or forgot to square the P₂ pressure (even —though the K_p expression was given in the question.

Tips for Teachers:

• Give students practice solving equilibrium problems that include each of the following:
  • determining the concentrations of each substance at equilibrium when the initial concentrations and the K value are known
  • calculating the value of K when the equilibrium concentrations of each substance are known
• Require students to show all work on assessments. When solving equilibrium problems, students can underline the words initial and equilibrium, reminding them that only equilibrium values can go in an equilibrium expression.
• Remind students to always examine the diagrams that accompany a question so that they understand the question being asked.

Thermodynamics (Part C): Students struggled to interpret how thermodynamic values are combined to determine favorability of a reaction process. See Table 2. Some made the correct choice but failed to explain the relevant values. Others confused the signs associated with favorability. Most students did not earn the point because they failed to make a complete argument or failed to follow the directions within the prompt.

Table 2. Summary of thermodynamic favorability. 

Tips for Teachers:

• Use this AACT activity, “Enthalpy and Entropy as Driving Forces,” to help build students’ understanding of how ΔS and ΔH contribute to the sign of ΔG and what that means in terms of thermodynamic favorability.
• Have students practice making comparisons. Focus instruction and feedback on discussing both species when comparing and checking to ensure they have answered the full question.
  • Instruct them to go back and look at what the question asked and underline those items. For example: Did you agree or disagree? Did you discuss ΔS? Did you discuss ΔG?
• Assign free response questions in which students must explain:
  • why the change in entropy (ΔS) for a certain process is positive or negative
  • how the sign of ΔG° indicates whether a reaction is thermodynamically favorable or unfavorable
  • how the magnitude and signs of ΔH° and ΔS° contribute to the overall sign of ΔG°
  • how the thermodynamic favorability of a reaction may be affected by changes in temperature
• Give students multiple opportunities to choose whether they agree or disagree with a student’s claim and to provide a justification based on experimental evidence.

Question 5 – Michael Farabaugh, Question Leader

Michael Farabaugh has been teaching AP Chemistry for twenty-seven years in Charlottesville, VA. He has served as an AP Reader, Table Leader, and Question Leader since 2011 and has also served on the AP Chemistry Development Committee.

• Question 5 assessed bonding and intermolecular forces (IMFs).

Molecules and Bonding (Parts A-B): Students were asked to identify which bond—among three choices given—is the most polar and to justify their answer.

The most popular incorrect answer was the C-Br bond. Responses from students who chose C-Br seemed to imply that they thought the elements C and Br would have the greatest difference in electronegativity because those two elements are located the farthest apart on the periodic table.

Students who correctly chose the C-F bond as the most polar bond struggled to properly justify their choice. Fluorine, chlorine, and bromine each have higher electronegativity values than carbon. Since fluorine has the highest value of the three halogens being considered, its electronegativity difference from carbon is the greatest of the three, making the C-F bond more polar than either the C-Cl or C-Br bonds.

When justifying a choice, students must remember to specifically address features or characteristics of each species being compared, as many gave only broad statements, rather than complete comparisons.

When asked to explain a difference in bond angles within the same molecule, “using principles of atomic structure and VSEPR theory,” many students incorrectly attempted to use electronegativity or bond polarity as justification. It is possible that students had incorrectly assumed that the previous question would be relevant in some way, leading to these answers.

A surprising number of students did not attempt or quickly abandoned their attempts to answer this question. Students may have seen this question as one that might take too much time to consider and skipped it with the intention of returning to it if time allowed.

Tips for Teachers:

Figure 8. Different models of CBrClF2 from different viewpoints. Top views show a space-filling—or van der Waals radius—view. Bottom views show a wireframe view. Students could be shown these two models and asked why the Br-C-Cl angle is greater than the F-C-F angle. Images made in MolView, https://molview.org.

• When studying molecules, remind students of the periodicity principles they already know and then show three-dimensional models, as in Figure 8, so they can visualize how periodic trends can help them answer bonding questions. Students often isolate topics from each other when learning. So reminders that topics often intersect can help them to make better connections.
• When considering why one bond angle might be larger than another, ask them what they know about the atoms within the bonds. Are they similar or different in size? (Larger atoms experience greater repulsion.) Will they repel each other more than another pair of atoms might? (Lone pairs and electronegativities affect how two atoms will interact across space when near each other.)
  • Although it was not relevant to question 5, a factor that can affect bond angles is the presence of nonbonding (lone) pairs of electrons on the central atom. (For example, the H–N–H bond angle in NH₃ is approximately 107°, which is slightly less than the predicted bond angle of 109.5° associated with sp³ hybridized atoms.)
• As with any comparisons, encourage students to think about what is similar and what is different, as that is usually what will lead them to appropriate reasoning.
• Find an appropriate time to wean students away from using electronegativity charts to determine bond polarity. They should be able to apply their knowledge of periodic trends to determine the relative polarities of bonds. Greater differences in electronegativity generally lead to greater bond dipoles.

Intermolecular Forces (Part C): Given Lewis diagrams for two compounds, students were asked to use the relative strengths of all the intermolecular forces to explain the difference in boiling points. Many students attributed the higher boiling point of CBr₄(l) to its greater mass. Although a correlation can exist between molecular mass and boiling point, it is the larger size of the electron cloud—and thus its greater polarizability—that leads to increased attractive forces and a higher boiling point for CBr₄(l).

Some responses contained generalized or vague language, which hinted at possible confusion between covalent bonds and intermolecular attractive forces. Others ignored—or maybe didn’t notice—the boiling points in the data and stated that CBrClF₂(l) must have the higher boiling point because its molecules are capable of dipole-dipole attraction while those of CBr₄ are not!

• Most of us are guilty of overgeneralizing in the moment—so it is important to show students that the relative strengths of different types of intermolecular forces are only that way when comparing similar molecules. Once students can identify intermolecular forces in different molecules and use them to justify differing properties across substances, challenge them with examples where the expected relationships break down.
• Instruct students to acknowledge that data are correct and that if the data do not agree with their expectations, then they need to seek a different explanation for the data than what they expected. For example:
  • In Question 5, the CBr₄ has the higher boiling point—so it does have the stronger intermolecular forces. Because its structure limits its molecules to only experiencing London dispersion forces, we must conclude that those forces are somehow stronger than the forces between molecules in the other compound which experience both LDF and dipole-dipole forces. The strength of attraction from LDF increases with polarizability—so CBr₄ must have a large enough electron cloud that its induced dipoles create a stronger attractive force than the combined forces between molecules of the other compound.

Question 6 – Diana Gano, Table Leader

Diana Gano just retired after teaching high school chemistry for twenty-four years, with eighteen of those in AP Chem. She has been an AP Reader for seven years, serving as an Early Table Leader the past two years. Diana is actively looking for her next chemistry adventure!

• Question 6 used data from a Beer’s Law experiment to analyze solution and dilution concentrations. Although the vanadium(II) solution used in the experiment would be light purple, the actual color was not relevant for the associated questions.

Solutions and Dilutions (Parts A-C): Students were very successful in demonstrating their particle-level understanding of concentration and dilution. They were less successful though in following the mathematics of a dilution through the multistep process. Most students knew to use the Beer’s Law calibration curve to identify the concentration of a solution from the given absorbance value. However, a high number of simple graph reading errors occurred, where students either missed or added a decimal place when locating the correct value.

The greatest struggle in this question came from the second part of the multistep calculation to determine the molarity of the original, undiluted solution. Although this was a straightforward dilution problem (M₁V₁
= M₂V₂), students were clearly confused by the process of the dilution. Figure 9 represents the process described in Part B of Question 6.

Many used addition or subtraction rather than proportional reasoning with the two volumes. Some made incorrect value substitutions or made other algebraic errors. Overall, a dilution question typically scores very well, but this one did not. This is possibly due to students struggling to make sense of the overall process.

The last part of this question required students to justify whether a particular mistake could account for the calculated concentration being too low. Students struggled to communicate how and why the mistake would lead to the specified concentration error.

The variety of incorrect approaches was remarkable! Confusion occurred among the terms “solute,” “solvent,” and “solution”; dilution was often misinterpreted as an increase in molarity; and many students seemed unfamiliar with volumetric flasks, asking what the “mark” meant.

Some students started with valid arguments but subsequently discounted them due to either incorrect chemistry or lack of connection to the claim. As with many previous AP exam questions, a simpler, shorter answer had a better chance of earning the point here, since longer responses risked contradicting themselves.

Tips for Teachers:

Figure 10. Screenshot of an AACT simulation that can help expose students more to working with a volumetric flask.

• Remind students, early and often, to answer the question that is asked! If a prompt says, “Do you agree or disagree?” the first part of the response should be “I agree” or “I disagree.” The rest should support that answer. A good format is the CER (Claim, Evidence, Reasoning) method described in Table 1.
  • Tell students that it’s okay to change their minds, but they must then change their original answer to match the justification given.
• Use particle-level diagrams to represent moles and molarity. Give students practice in both analyzing and drawing their own particle diagrams. These diagrams are especially helpful in building quantitative reasoning skills.
• Find time to incorporate simple lab exercises that are not part of a larger experiment—like using volumetric flasks to make solutions and dilutions. This can help students build better lab skills, better understand common lab procedures, and even save time in future labs where the practiced skills are used. The AACT "Preparing Solutions" Simulation (Figure 10) could be used to help reinforce these lab procedures.
• For complicated or multistep processes, consider requiring students to sketch out what is happening so they can learn to more easily track the steps as they work through the math. Sketching visuals does not need to look neat and can be done within a lab notebook or worksheet. Figure 9 shows what this might look like for the dilution problem in this question.
• Make sure students are familiar with using Beer’s Law in the lab. Spectrophotometry is an accessible, highly visual lab technique which has been the prompt for many AP questions, both MCQ and FRQ.
• Incorporate error analysis into your labs and assessments to emphasize the importance of understanding how results can come out incorrectly—and how different measurements or procedures may affect results.

Question 7 – Marc Stephenson, Table Leader

Marc Stephenson has been teaching AP Chemistry for over twenty years at Saugus High School in Saugus, CA. He has served as an AP Reader and Table Leader since 2015.

• Question 7 used the synthesis reaction of sodium oxide to assess thermochemistry and periodicity.

Thermochemistry (Parts A-B): Overall, students did well interpreting enthalpy of formation from the data. This was the most commonly earned point within Question 7. When connecting enthalpy of reaction to stoichiometry and limiting reagent principles, students were less successful.

Some struggled to connect the reaction enthalpy with its stoichiometry. Others failed to determine the appropriate limiting reagent from the data. Many misused the mole ratio or skipped it altogether. Although students only had to determine the proportional amount of heat released based on reactant masses, several made invalid attempts to use q = mcΔT, even though no calorimetry data existed to support it.

Another relatively common mistake was seen when students would appropriately use proportional reasoning with the reaction enthalpy but incorrectly relate it to the sum of all reactant moles, rather than to moles of only the limiting reagent. It seemed as though students were unsure what mol_rxn meant.

Tips for teachers:

• Make sure students show all their work. Too often it appeared as though students were doing work either in their heads or on the calculators. I tell my students I don’t want to see “magic numbers”—numbers that appear from nowhere—in their answers!
• Ensure that students practice with thermochemical and stoichiometric calculations, including how to determine the identity of a limiting reactant.
• Model the use of mol_rxn in class so that students are comfortable with scenarios involving this unit.

Periodicity (Part C): When directed to use Coulomb’s Law in an explanation, it was clear that students generally understood the key points of the law. However, the most frequent mistake was discussing the law in terms of distance from the nucleus to valence electrons, rather than distance between two ions. Students mentioned just about every atomic property commonly encountered with periodic trends, so it was clear that they were either unfamiliar with lattice enthalpy or thinking that the patterns in one property could act to justify patterns in another property.

Tips for teachers:

Figure 11. In the Na2O crystal, the r, or distance, would be from the center of a sodium ion to the center of an oxide ion, as represented by the X marks in the crystal. Benjah-bmm27, “Space-filling model of the sodium oxide (Na2O) unit cell,” Public domain, via Wikimedia Commons.

• Provide opportunities for students to demonstrate their understanding of Coulomb’s Law in situations other than the radius of an atom or ion. It can be applied to other scenarios where two charges are close enough to interact. Using physical models or observing space-filling models online can help students visualize how Coulomb’s Law might apply. See Figure 11.
  • The formula for Coulomb’s law is on the formula sheet with r being defined as “separation” because the law applies to any two charged species.

• Require students to explain the conceptual basis for the mathematical relationships shown by common equations and laws, such as Coulomb’s law.
• Assign practice in evaluating written explanations, such as through peer review of a classmate’s justification or explanation. Teachers can also provide students with flawed justifications or explanations so that students can work in small groups to improve them.

Conclusion

I hope this article has been informative and that you have found practical ways to refine your instruction throughout the school year.

Remember, you don't need to “teach to the test” to help your students succeed. Simply being aware of common mistakes seen during AP exam scoring can help you communicate expectations more clearly—and help your students perform better as a result.

I strive to help my students improve their understanding and become more confident as I support them on their learning journeys. If you have comments or questions about preparing students for success on the AP Chemistry Exam, please feel free to contact me.