Adrian Dingle
Chemistry Teacher and Author
http://www.adriandingleschemistrypages.com

Carolina Staff

Updated March 2019

If you're reading this, you're probably an AP® Chemistry teacher who has survived a semester of the curriculum. Congratulations! Perhaps you had a to-do list that looked something like this:

• Incorporate 6 inquiry labs into the chemistry course
• Submit an updated syllabus for the audit
• Learn about PES

So some things may not have gone entirely smoothly, but how can you ease your burden going forward? Well, in addition to knowing the content that is still ahead, you're also aware of the content that isn't in your future. There are 22 exclusion statements that appear in the course and exam description that are designed to tell you what's "out."

Here's a quick guide to the stuff you can omit. These exclusion statements are found in the appropriate Essential Knowledge sections of the AP® Chemistry Course and Exam Description, available for download at AP Central. Note: These comments are my own interpretations and come with the disclaimer that they could significantly change over time, such as when more exam questions and knowledge about the course come to light.

Exclusions quick guide

• Memorization of exceptions to the Aufbau principle. This article, "The trouble with the Aufbau principle," might shed some light on these exceptions.
• Assignment of quantum numbers. Electronic configuration is still "in," so you'll still be teaching the principal quantum number.
• Phase diagrams, colligative properties and calculations of molality, percent by mass, and percent by volume. Molality goes to the grave with colligative properties, but it's difficult to have discussions about an empirical formula (and impossible when dealing with the real-world inquiry H₂O₂ / MnO₄- titration that is Investigation 8) without talking about percent by mass!
• Other cases of weaker hydrogen bonding (i.e., those outside of H connected to N, O, or F). There's no change here (i.e., hydrogen bonding is still "in") and this statement has only been included because of the recent redefinition.
• Knowledge of specific types of crystal structures and the study of the specific varieties of crystal lattices for ionic compounds. Crystal structures have never been on a contemporary exam.
• The use of formal charge to explain why certain molecules do not obey the octet rule. Don't misinterpret this—formal charge is still "in" and could be used to distinguish between 2 structures that do obey the octet rule, but it will not be used to explain why a Lewis structure with an expanded octet might be preferred over an alternative one with just an octet.
• Learning how to defend Lewis models based upon assumptions about the limitations of the models. Since it is linked to EK:2C4f and EK:2C4d, for now I interpret this to mean that formal charge is not going to be used to justify a Lewis structure with odd numbers of electrons.
• An understanding of the derivation and depiction of hybrid orbitals. Who has been teaching "derivation and depiction" of hybrid orbitals anyway? Speaking of hybridization, sp, sp², and sp³ are still "in," but any involvement of d-orbitals in hybridization is "out."
• Other aspects of molecular orbital theory, such as recall or filling of molecular orbital diagrams. I have never taught any MO since it has been "out" for years; however, EK:2C2i suggests that maybe the interpretation of an MO diagram could be "in." This is an interesting development and worth keeping an eye on, although remember, every question must be associated with a specific Learning Objective and it's tricky to find one that would fit.
• Lewis acid-base concepts and language of reducing and oxidizing agents. Say what? Any kids leaving an AP® course that I have taught will know what Lewis acids and bases are, and they will be using the terms "oxidizing" and "reducing agent."
• Labeling electrodes as positive or negative. The assignment of charge to electrodes in electrolysis and cells is confusing, and it doesn't add much to understanding the electrochemistry itself. Nernst equation. Careful, because although quantitative stuff about non-standard conditions in cells has gone, qualitative reasoning is still in play; you'll need to apply Le Châtelier's principle to non-standard cells and maybe even deal with a concentration cell.
• Calculations involving the Arrhenius equation. As an exclusion, this means little since calculations have rarely ever been features of contemporary exams, but expect the useful application of the equation in a y = mx + b plot to determine the activation energy and qualitative aspects to remain "in."
• Collection of data pertaining to the experimental detection of a reaction intermediate. Precisely who has been "collecting data" about intermediates to build evidence in support of 1 mechanism over another as an exam skill?
• Numerical computation of the concentration of each species in the titration curve for polyprotic acids. Still expect calculations for monoprotic acids and qualitative treatments of polyprotics.
• Computing the change in pH resulting from the addition of an acid or base to a buffer. Buffer calculations are still "in" (Henderson-Hasselbalch remains among the Equations and Constants), so I interpret this to mean that a question where addition of acid or base to buffer that is formed outside the context of a titration is "out."
• The production of the Henderson-Hasselbalch equation by the algebraic manipulation of the relevant equilibrium constant expression. The equation was (and still is) given with the exam; the derivation of the equation is "out."
• Memorization of other "solubility rules." A pity since some rote learning is good; competent chemists do not need to google basic facts!
• Computations of solubility as a function of pH and computations of solubility in such solutions. These 2 statements seem to be saying that quantitative, common ion calculations that involve H⁺ and/or OH^– are "out," as are calculations relating to EK:6C3f. I think that one could also interpret them to mean that all quantitative common ion treatments are "out," but like so much of what is written above, only time will tell.

When math is not your strong suit

Math is not my strong point. Frankly, that puts me in good company with a number of AP® Chemistry students. For those kids, a lack of math acumen can undermine what could otherwise be some very good progress in chemistry. What can we do to prevent that shortcoming from damaging the scores of those students?

When I was an undergraduate, I took a course entitled Math for Chemists. It was for those of us who were strictly "chemists" but in need of some targeted pointers that might help us to overcome some of the mathematical challenges associated with physical chemistry. Math for Chemists helped me to navigate some scary moments amidst wave functions, the Schrödinger equation, nasty derivatives, and Hamiltonian operators. Ever since that experience, I've always felt that some similar, stealthy tips, albeit at a much lower level, have the potential to be really quite useful for the chemistry student who, like me, will perhaps never learn to love mathematics.

Don't be alarmed

I'm not about to enter into a discussion about the mathematical niceties of quantum mechanics or integrals here, but what I will do is offer math pointers that even I can give to students if they are struggling with a particular quantitative aspect. The list is not an effort to teach math by any stretch of the imagination, nor will it necessarily lead to any kind of enhanced mathematical understanding, but it does represent a list of mental shortcuts that just might unlock some points for a few students on the exam.

Nothing here is groundbreaking or perhaps things that you haven't read in the appendix of a chemistry textbook, but a gentle reminder never hurt; and you may be surprised by how reinforcing a few of these simple relationships can help students save chemistry points on the exam. Beyond that, and in terms of multiple-choice questions, some of these tips remain invaluable since the students are bereft of a calculator, and estimation and mental arithmetic remain crucial skills on that part of the exam.

This list is for students—I know that you know this stuff. Some of it falls under the heading of general math tips, and some is more specifically related to chemistry, but all should prove useful in the areas suggested, and perhaps beyond.

Top 10 math pointers

1. Logs. What's a log? It's a button on the calculator, a "function" if you will. No more, no less, it converts 1 number to a different number.
   
   –log (1 × 10^–4) = 4, –log (1 × 10^–3) = 3, etc.
   
   This means that the –log of a number, such as 5 × 10^–4, that's somewhere between those 2 values (bigger than the first but smaller than the second), is between 3 and 4. (Acids, bases, and buffers)
   
   "p" simply means "–log." (Acids, bases, and buffers)
   
   The log of a number less than 1 is negative, and that of a number greater than 1 is positive. (Yes, I know Nernst has gone, but this could still be useful in a Henderson-Hasselbalch calculation.) (Acids, bases, and buffers)
   
   
2. Add exponents when multiplying, and subtract when dividing. (Equilibrium)
3. Reversing a chemical equation at equilibrium creates a new K value that's the reciprocal of the original. (Equilibrium)
4. Multiplying the stoichiometric coefficients of a chemical equation at equilibrium creates a new K value that's the original raised to the power of the multiplier. (Equilibrium)
5. K_w = [ H⁺ ][ OH^– ] = 1 × 10^–14 is the "same" equation as pK_w = pH + pOH = 14 because of #1 and #2 above and when taking logs of things multiplied together, they become summed. The same is true of the relationship between the acid dissociation constant K_a and the Henderson-Hasselbalch equation. (Acids, bases, and buffers)
6. Units matter! You had better realize that 4.95 × 10^–7 m = 495 nm, etc. For example, where a wavelength has been calculated in m, but multiple-choice answers have been reported in nm (and vice-versa). This is a simple but important thing to remember. (Atomic structure/electrons and of course all quantitative aspects of the course)
7. Units matter! Delta H values are usually recorded in units of kJ mol ^–1 and delta S values are usually presented in units of
   J mol ^–1 K ^–1. This matters when using ΔG = ΔH – ΤΔS. Make sure that you have converted 1 of the values to the units of the other before calculating ΔG. (Entropy, enthalpy and Gibbs free energy, and of course all quantitative aspects of the course)
8. Units matter! They can help you to determine which value of R to use (there are 3 on the Equations and Constants sheet) in different situations. Use 8.314 J mol ^–1 K^–1 when dealing with "energy situations," and 0.08206 L atm mol ^–1 K^–1 or 62.36 L torr mol ^–1 K^–1 (depending on the units of pressure) when gases are involved. (Thermochemistry, electrochemistry, gases, and of course all quantitative aspects of the course)
9. Dimensional analysis can help a great deal when keeping track of units. For example, in #8 above, when using P V = n R T to calculate the temperature of a certain number of moles of an ideal gas, with a pressure given in atm and a volume given in L, the only R that makes sense in terms of units is 0.08206 L atm mol^–1 K^–1. Why? Well, because in terms of units the following is true:
   
   
   
   
10. Estimation remains an important skill. On the multiple-choice section of the exam you do not have access to a calculator but can still be asked questions that involve calculations. For example, using estimation to realize that (99.8)(1.01) / (0.08206)(350) is approximately the same as (100)/(30), which in turn is equal to a number between 3 and 4, can help you to choose from a list of potential answers without having to do any calculations. That's a lot easier than calculating the actual value of (99.8)(1.01) / (0.08206)(350) in your head.
    
    Of course, I could go on about significant figures and rounding and checking answers, too, but then I'd be getting a long way away from chemistry.

Math still matters

Some may say that, because AP® Chemistry has arguably moved away from the more quantitative aspects of the past, these tips are perhaps less important than they once were. I see it differently.

First, not all of the math has gone away. Logs, exponents, and the ability to estimate are still very relevant. Second, by continuing to use some of the old, quantitative relationships that have actually been removed from the Equations and Constants sheet, one can actually aid the understanding of concepts that have shifted entirely to a qualitative treatment.

Two such examples are root mean square speed and Graham's law of effusion and diffusion. These equations are no longer given with the exam and this means that a quantitative treatment of them is not something that we should expect to see in future exam questions, but qualitative aspects of them are definitely still in play. What does that mean for these 2 examples? Well, not much more than knowing that with a greater molar mass, u_rms decreases; with increasing temperature, u_rms increases; and that heavier particles tend to effuse and diffuse more slowly than lighter ones.

The argument that lab work helps to illustrate theory extends and evolves into math work can help to illustrate theory. It's all well and good saying that heavier particles move more slowly, but (even with what some might criticize as being no more than mindless plugging and chugging) we can aid that understanding and cement it with some calculations. Plug the molar masses of 2 different gases into either of those mathematical formulas, pick up a calculator, and you'll see that the resultant numbers bear out what the theory says. For that reason alone, I'll continue to use some of the depreciated quantitative equations in my classes.

So, you've made it through teaching the AP Chemistry curriculum.

You've overcome the trials and tribulations of the Big Ideas, the Enduring Understandings, the Essential Knowledge, and the Learning Objectives. You pressed on through inquiry labs, particulate diagrams, and Coulombic forces. Depleted by an unusually large number of snow days, you're perhaps scrambling to tie up some loose ends since it's likely been quite a tumultuous year, even for the most experienced. So how can we take stock, remain calm, and set the kids up for exam success?

75 tips for success

For many years, I have offered my kids a last-minute list of "tiny tips" that I think could serve them well going into the exam. A few of the tips are the type of random factoids that occasionally appeared on multiple-choice exams of the past, where perhaps a color or an environmental application was tested for a single point. How relevant that kind of thing will be going forward remains to be seen, but others on the list are reminders of fundamentally important concepts.

1. The speed of a chemical reaction is not related to the equilibrium position.
2. Hydrogen bonding is an INTERmolecular force, not an INTRAmolecular bond.
3. Electrolysis is only necessary when a reaction is non-spontaneous with a positive Delta G.
4. Le Châtelier's principle is not an explanation in itself. A shift in position to reduce an external stress is, as is an understanding that Q must equal K at equilibrium.
5. Periodic trends are not explanations.
6. Potassium manganate(VII) and sodium dichromate(VI), when in acid, are common oxidizing agents.
7. Wash a buret with the solution that it will be dispensing in the titration, and fill the tip.
8. Gases behave ideally when at relatively low pressures and relatively high temperatures.
9. Catalysts increase the rate of the forward and the backward reaction.
10. Common ions make slightly soluble salts even less soluble.
11. Kp expressions include ONLY gas partial pressures.
12. "Optimal" buffers (that can absorb acid and base equally well), have pH = pKa.
13. Clean up an acid spill with a carbonate, not an equally corrosive, strong base.
14. Writing the full electronic configuration of an atom can help to explain differences in ionization energies and the patterns in PES data.
15. Transition metal ions are often colored in solution.
16. Reduction always takes place at the cathode.
17. Fluorine always has an oxidation number of –1.
18. The bigger the pK_a, the weaker the acid; the bigger the K_a, the stronger the acid.
19. When using R = 0.0821 in P V = n R T, pressure must be in atm, temperature in K, and volume in L.
20. On the exam, use the FULL atomic masses printed on the periodic table.
21. When predicting shape, a double bond counts as only 1 area of electron density.
22. It is unlikely that any numerical answer on the AP exam will ever require 10 significant figures!
23. Since C and H have a similar electronegativity, hydrocarbons are largely non-polar.
24. Polarity in organic molecules helps them to be soluble in water.
25. Only the first bond of a double or triple bond is counted in hybridization. The others are pi bonds formed by the overlap of UNhybridized p orbitals.
26. Breaking bonds within reactants is an ENDOTHERMIC process.
27. Alcohols are soluble because they can H-bond with water, NOT because they have a hydroxide group—they DON'T!
28. Ions travel through the salt bridge, not electrons.
29. Net ionic equations must balance charge as well as atoms.
30. A graph of [ X ]^–1 versus time gives a straight line for a second order reaction.
31. Bromine and mercury are liquids at room temperature and iodine is a solid.
32. Transition metals lose their s electrons first.
33. Always use temperature in Kelvin in gas calculations.
34. The units of Delta H are usually expressed in kJ, and Delta S usually in J; therefore, they must be converted to a single, common unit, if using them to calculate Delta G.
35. Orders of reaction can be fractions.
36. Organic amines like methylamine, CH₃NH₂, are weak bases since the lone pair on the N atom can accept H⁺.
37. Positive E_cell values go with negative Delta G values and very large K values.
38. When [ H⁺ ] in solution is < [ OH^– ] the solution is basic (and vice versa).
39. The condition for neutrality is [ H⁺ ] = [ OH^– ], not pH = 7.
40. Beer's law can only be applied to colored salts.
41. In dynamic equilibrium, the forward and backward reactions do not stop, they just occur at the same rate.
42. When considering macro changes in entropy, look at how the number of gas moles changes.
43. Oxygen relights a glowing splint.
44. Equilibrium constants are constant at constant temperature.
45. Ionic solids have strong ionic bonds that are electrostatic (Coulomb's law), and as a result, have high melting and boiling points.
46. Changing phase in molecular substance involves breaking intermolecular forces, NOT covalent bonds.
47. Si and SiO₂ have giant structures similar to diamond.
48. Si and As are used in semiconductors.
49. The p-type semi-conductors are those that are doped with an element with 1 less valence electron (e.g., a group 13 atom added to a group 14 semi-conductor).
50. The n-type semi-conductors are those that are doped with an element with 1 more valence electron (e.g., a group 15 atom added to a group 14 semi-conductor).
51. Hydrogen fluoride is a weak acid, and it etches glass.
52. CFCs (chlorofluorocarbons) are implicated in climate change.
53. STP for gases is 273 K and 1.00 atm.
54. Kinetic energy of gases depends on their Kelvin temperature.
55. Sulfur can exist as S₈ molecules.
56. Phosphorus can exist as P₄ molecules.
57. Pure solids, liquids, and gases are never ionized in NIEs, but soluble salts and strong acids and bases IN SOLUTION are.
58. When considering valence electrons of p block elements, remember to include the outer s electrons as well (e.g., Al has 3 valence electrons, s² and p¹).
59. B and O, and Al and S have slightly lower first ionization energies than we expect, but for different reasons.
60. Conjugate acid and base pairs differ in their formula only by H⁺.
61. A very strong acid (e.g., HCl) will have a very weak conjugate base (Cl^– ).
62. Add concentrated acids and bases to large volumes of water, NOT the other way around.
63. Carbon dioxide makes limewater milky.
64. The Delta H for the formation of an element is 0 (nothing changes).
65. The Delta S for the formation of an element is 0 (nothing changes).
66. . . . BUT elements have ABSOLUTE entropies that are NOT 0.
67. Catalysts provide alternative pathways for reactions that have lower activation energies.
68. Before weighing on electronic balances, allow heated items to cool.
69. A beaker/Erlenmeyer flask is NOT a measuring instrument.
70. Group 1 oxide + water gives the corresponding hydroxide that is soluble.
71. Group 1 metal + water gives the corresponding hydroxide that is soluble AND hydrogen gas.
72. Group 1 hydride + water gives the corresponding hydroxide that is soluble AND hydrogen gas.
73. Equilibrium systems that undergo changes in pressure should only have their gas molecules considered.
74. Lead(II) iodide is a yellow precipitate.
75. The TOTAL area under a Maxwell-Boltzmann distribution curve is the same for a reaction at a high and a low temperature.

Support from Carolina

As you soldier on through, remember that Carolina has a full line of kits, Carolina Investigations® for AP® Chemistry, specifically designed for the curriculum. The 16 kits in this series address the 6 Big Ideas of chemistry and meet the requirements of the lab curriculum. Each kit focuses on a single Big Idea and offers the option to do either a guided activity or an inquiry activity with your students. To help students prepare for the exam, all kits include Big Idea assessment questions that follow the exam's free-response format.

Flexibility to teach AP Chemistry YOUR way

The College Board's AP® Chemistry lab manual contains 16 inquiry labs that teachers can choose from to fulfill the requirement of 6 inquiry labs in the AP® Chemistry course. The materials and quantities needed for these example labs are listed in the document "Materials for AP Chemistry Guided-Inquiry Experiments: Applying the Science Practices.” The document also includes suggestions for the corresponding Carolina Investigations® for AP® Chemistry kits that meet the same Learning Objectives as the College Board example labs.

*AP is a registered trademark of the College Board, which was not involved in the production of, and does not endorse, these products.