AP Chem Handouts

AP Chemistry

Last Minute Tips + Reminders

• Answer 10 multiple choice questions per every 15 minutes.

• Spend more time on the problems you know rather than the problems you don’t know

• Do the multiple choice in three passes.

• Ones you know that are quick

• Ones that take time

• The ones you didn’t know

• On the free response questions, read every single word of each part of the FRQ very carefully before you start to write down your answer.

• After you have finished your free response question, read it back to yourself. Did you answer the question that was being asked?

• If they ask you to make a choice (higher, lower, increase, decrease, etc.) make the choice and write this part down FIRST… BEFORE you start justifying your answer

• Don’t write too much!

• If part (a) looks too confusing to you, then skip it and come back to it later. You can answer part (b) before part (a). Label your answers clearly!

• If your answer is ΔH, ΔS, ΔG or Eo, make sure your answer includes the proper SIGN and the proper UNITS.

• Stating a trend is NOT a justification.

• Trends are usually about nuclear charge and distance (Think about Coulomb’s Law)

• Atoms don’t “want” or “need” anything.

• Breaking attractions (bonds, IMFs, nucleus/valence electrons) REQUIRES energy (endothermic).

• If the question says to justify your answer in terms of IMFs, mention ALL of the IMFs that each substance experiences not just the strongest force.

• Do NOT say “like dissolves like” when justifying why two substances dissolve or mix together.

• If the question says, “Draw ALL resonance structures,” chances are there’s more than one Lewis Structure.

• Thermodynamically favorable: ΔGo < 0, Eo > 0, K > 1

• At equilibrium, ΔG = 0, E = 0. Both essentially indicate how far away from equilibrium the system is.

• Horizontal trends: they both have the same energy level (subshell), but one of them has more protons (nuclear charge).

• Vertical trends: one of them has valence electrons in an orbital with a higher value of n; the orbital is high in energy and therefore further from the nucleus.

• If the question asks you to compare 2 substances make sure you refer to TWO substances.

• All equations must be balanced for atoms AND charge.

• For net ionic equations: dissociate aqueous, strong electrolytes, and eliminate spectator ions.

• Read the question. Answer the question.

• Thermodynamics tells you IF it will happen. Kinetics tells you how QUICKLY.

• The other thing that changes the value of K is changing the temperature (or adding a catalyst).

• Raising the temperature does NOT affect ΔH or the activation energy.

• The acidic species in a buffer neutralizes added base.

• Weak acid + strong base titration: pH = pKa at the halfway point

• When pH = pKa, [HA] = [A-]

• Increasing the temperature increases both the forward and reverse reaction rates. It just increases the endothermic direction’s rate more.

• Lab error normally looks for an increase or decrease in experimental values.

• Do NOT use 22.4 L unless you are actually at STP.

• STP = standard temperature and pressure is NOT standard state (standard state = 25℃, 298K)

• Don’t use 22.4 L if you have aqueous solutions.

• Sometimes the data tells you something that is different than what you would have predicted (i.e. CCl4 has stronger IMFs than HCl) GO WITH THE DATA.

• Mass spec is all about isotopes.

• In multiple choice math questions, round/approximate. You NEVER need long division.

• PES is all about Coulomb's law and electron configurations. (Photoemission Spectroscopy)

• Never trust printed structures for bond angles! Count the electron domains (Aka not drawn to scale!)

• Never multiply potentials (cell, reduction, or oxidation potentials aren’t affected by coefficients like thermochemistry formation values are [ Delta G, H, and S are affected by moles])..

• If you titrate 10 mL of 1 M weak acid and titrate 10 mL of a 1 M strong acid, these samples would require the SAME amount of 1 M NaOH to reach the equilibrium point (i.e. volume of titrant is based on moles, not acid strength).

• Discuss Q to explain a shift in equilibrium (example: saying “Lechatelier’s Principle” will get you NOTHING!) If Q increases, reaction goes towards the reactants. The same goes for a voltaic cell voltage change. If Q increases, the voltage decreases.

• If you have a reaction with a HUGE value of K, assume the reaction essentially goes to completion, and treat it like a normal stoichiometry/limiting reactant problem; you do not need to make an ICE table.

• Acid-base reactions proceed from the strong side to the weak side. Look at the magnitude of K to determine which way the reaction is favored (high = product favored, low [negative power] = reactant)

• In a Lewis structure, carbon will make four bonds total and hydrogen will only form one.

• Tetrahedral is not a bond angle. (around 109.5 is the tetrahedral bond angle depending on unshared pairs)

• When in doubt, find moles!

• Reduction at the cathode, oxidation at the anode. (OIL RIG, Cathode is positive so it needs to Gain Electrons)

• MV = moles

• Get a good night’s sleep! (Preferably the entire week before the test!)

• Do the easy free response questions first! (During Timed tests you need to get those points!)

• Ions have a charge; Na and Na+ are not the same. (Na makes water explode[Unstable] , Na+ is safe to eat [stable # of electrons])

• Inert gases and catalysts DO NOT shift an equilibrium. [They just lower activation energy to start it]

• If the data shows that the half-life is constant over time, it’s first-order kinetics for that substance.

• Burets measure liquids more precisely than beakers or graduated cylinders.

• Nonpolar molecules can contain polar bonds.

• Electrons travel in the wire from anode to cathode. Ions flow through the salt bridge: anions toward the anode, cations toward the cathode.

• The greater the difference in electronegativity between two atoms, the more polar the bond is.

• If x = solubility of an ionic solid (in mol/L):

• Ksp for AgCl = (x)(x) = x2

• Ksp for PbCl2 = (x)(2x)2 = 4x3

• Lattice energy and Coulomb's law: MgO beats NaF (similar ionic radii, greater charge wins) NaF beats KCl (same charges, smaller ionic radii wins).

• Larger electron cloud = more LD = more polarizable electron cloud.

• Larger Ka = more ionizable.

• Stronger IMFs, higher BP, lower VP, and greater viscosity.

• If you complete an equilibrium calculation and find “x,” you’re not done! x may not be the answer that the question is asking for!

• HX versus HY… If HX is the stronger acid, the Y- is the stronger base.

• If you start to spend too long on a multiple choice question, remember that 10 of the 60 multiple choice questions don’t even contribute to your score. Efficiency is key!

• H-H is not a hydrogen bond and the covalent bonds inside of a water molecule are not hydrogen bonds.

• When explaining trends, do not talk about their placement on the table; talk about things in the atoms like Zeff, nuclear pull, distance between nucleus and valence electrons, etc.

• You would read the volume on this buret as 5.65 mL.... NOT 6.35 mL.

• H-bonds only occur between an H already bonded to N, O, F and another N, O, F.

• If you are clueless about rounding off your answer, three significant figures is plenty most of the time.

• Don’t panic… watch the time!

• “Because it has a full shell” is not a valid justification for any periodic trends.

• Don’t use filling an orbital as justification for anything--don’t say, “it wants to be full” or “happy”

• Equilibrium constant values (K).

• Reverse the reaction, the new K is the reciprocal of the old K.

• Multiply the reaction by 2, the new K is the old K squared.

• Add two reactions together...multiply (K1) x (K2).

• % error does not mean the same thing as % yield.

• High activation energy = slow reaction rate; low activation energy = fast reaction rate.

• Electroplating shortcut… grams = (MM)(amps)(time) divided by (n)(F)

• zero order: [A] vs. t; first order: ln [A] vs. t; second order: 1/[A] vs. t

• Absolute value of the slope = k... Don’t forget to include the proper UNITS for your k value!

• Don’t confuse rate constant units with rate units.

• Any two gases at the same temperature have the same KE but not the same velocities

• Larger = size; Heavier = mass; More = amount (like moles). Don’t mix them up!

• LDF depend on size, not mass!

• Lone pairs on the central atom doesn’t mean that the molecule is polar (think XeF4 and KrF2)...dipoles not canceling makes the molecule polar.

• Ideal gas conditions = high T low P; deviations also occur when gases are too sticky (polar) or too large (LD).

• Avoid the use of the word “it” refer to substances by name so we know exactly what you meant to say.

• Use electron domains to help with determining shapes (double bonds represent a single domain).

• Bond energy is reactants minus products, heat of formation is products minus reactants.

• LEO goes GER; An ox red cat; OIL RIG

• Always consider formal charge when drawing molecules--obey the octet rule first!

• M1V1 = M2V2 not on the formula sheet but extremely useful in dilutions and neutralizations.

• The more (+) Ered is the reduction and the less (+) the oxidation reaction

• The side of the electron cloud matters for LDF (larger = mole easily polarized = stronger IMF), not for dipole-dipole. Polarity is determined by the difference in electronegativity of the atoms in the bond (the greater the difference in electronegativity = more polar = stronger IMFs.

• Beware the pressure equilibrium expression!

• If it is Kp, don’t use brackets.

• When calculating Kp, the units of pressure must be in atmosphere.

• Adding an inert gas does not affect equilibrium, because the PARTIAL pressures remain unchanged.

• Chromatography: affinity for the mobile phase and affinity for the stationary phase.

• Bonds breaking endothermic, bonds forming exothermic, delta H is the net sum of this.

• When adding two half reactions together, the electrons must cancel out.

• If you reverse the reaction, change the sign of Eo.

• If you double it, do NOT double the voltage.

• Don’t use shielding as a justification for size increasing down a group

• Don’t forget your units for the rate law constant

• Or anything…

• A reaction with a large K is said to be product favored, which is not the same as saying “the reaction shifts toward the right.” If there’s no stress imposed on the system, there’s no Le Chatelier.

• Redefining the meaning of a term is NOT a justification.

• Example #1: “Mg has a higher 1st ionization energy that Na because it requires more energy to remove the valence electron.”

• Example #2: “HF has a higher boiling point than HCl because it requires more energy to convert HF from a liquid into a gas.”

• Don’t confuse the atom and ion or the molecule and formula unit.

• Diluting a salt solution increases the percent ionization.

• Q versus K!

• Write less, but say more.

• Use Q vs. K, not Le Chatelier, as justifications

• When in doubt, use the Kelvin temperature!

From Adrian Dingle #2: (the items that are lined out are no longer required knowledge)

1. Organic amines like methylamine, CH3NH2, are weak bases since the lone pair on the N atom can accept H+.

2. Nickel (II) salts are green.

3. Positive Ecell values go with negative ∆G values and very large K values.

4. When [H+] in solution is < [OH–] the solution is basic (and vice-versa).

5. When [H+] = [OH–] the solution is neutral.

6. Positrons are essentially, ‘positive electrons’.

6. A large Rf value suggests that the component has a high affinity for the mobile phase.

7. Beers’ law can only be applied to colored salts.

8. In dynamic equilibrium, the forward and backward reactions do not stop, they just occur at the same rate.

9. When considering macro changes in entropy, look at how the number of gas moles change.

10. Delta H = SUM Enthalpy products – SUM Enthalpy of reactants, ONLY works if you are dealing with FORMATION enthalpies.

11. Two phases exist on the horizontal part of a cooling or heating curve.

12. Oxygen re-lights a glowing splint.

13. Equilibrium constants are constant at constant temperature.

14. Ionic solids have strong ionic bonds that are electrostatic (Coulombs law) and as a result have high melting and boiling points.

15. Changing phase in molecular substance involves breaking IMF’s, NOT covalent bonds.

16. Weak acids only partially dissolve, so their colligative properties are less dramatic than fully soluble salts.

16. Capillary action and surface tension can be explained in terms of intermolecular forces.

17. Si and SiO2 have giant structures similar to diamond.

18. Si and As are used in semi-conductors.

19. When calculating molality, it often helps to assume a volume of 1L and calculate the separate mass of solute and solvent from there.

19. The dependance of the rate constant on activation energy and temperature is explained by the Arrhenius equation.

20. Hydrogen fluoride is a weak acid.

21. Hydrogen fluoride etches glass.

22. CFC’s (chlorofluorocarbons) are implicated in climate change.

23. Osmotic pressure is a colligative property.

23. An exothermic reaction does work on the surroundings.

24. Ozone is O3.

25. STP for gases is 273 K and 1 atm.

26. ‘Standard’ conditions in thermochemistry usually means 298 K.

27. Kinetic energy of gases depends on their Kelvin temperature.

28. Lead metal can be used to stop some radioactivity.

28. The mass spectrum of a monotomic element contains a peak for each isotope.

29. Sulfur can exist as S8 molecules and phosphorus can exist as P4molecules.

30. Pure solids, liquids and gases are never ionized in NIE’s, but soluble salts and strong acids and bases IN SOLUTION, are.

31. 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)

32. B and O, and Al and S, have slightly lower first ionization energies than we expect, but for different reasons.

33. Conjugate acid and base pairs differ in their formula only by H+.

34. A very strong acid (e.g., HCl) will have a very weak conjugate base (Cl–).

35. Add concentrated acids and bases to large volumes of water, NOT the other way around.

36. Carbon dioxide makes limewater milky.

37. Carboxylic acid + alcohol gives ester + water.

37. Heat is a transfer of energy from a high energy system, to a low energy system, in order to ultimately achieve thermal equilibrium.

38. The Delta H for the formation of an element is zero (nothing changes).

39. The Delta S for the formation of an element is zero (nothing changes).

40. …BUT elements have ABSOLUTE entropies that are NOT zero.

41. Catalysts provide different mechanisms/pathways that have lower activation energies.

42. Before weighing on electronic balances, allow heated items to cool.

43. Beakers and Erlenmeyer flasks are NOT measuring instruments.

44. Complex ions that include NH3 ligands, are broken down by acids.

44. Transition metals often have salts which are colored.

45. Group 1 oxide + water gives corresponding hydroxide which is soluble.

46. Group 1 metal + water gives corresponding hydroxide which is soluble AND hydrogen gas.

47. Group 1 hydride + water gives corresponding hydroxide which is soluble AND hydrogen gas.

48. Equilibrium systems that undergo changes in pressure should only have their gas molecules considered.

49. Lead(II) iodide is a yellow precipitate.

49. Remember to filter, wash and dry a precipitate in a gravimetric analysis.

50. The TOTAL area under a Boltzmann distribution curve is the same for a reaction at a high and a low temperature.

From Adrian Dingle #1: (the items that are lined out are no longer required knowledge)

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. Rutherford’s ‘Gold Foil Experiment’ produced evidence of a dense, positively charged nucleus.

5. Le Chatelier’s principle is not an explanation it itself. A shift in position to reduce an external stress (Q versus K), is.

6. Periodic trends are not explanations.

7. Silver, lead and mercury chlorides are the commonly quoted INsoluble chlorides.

7. The solubility rules that you need to know are that sodium, potassium, ammonium and nitrate salts are all soluble in water.

8. Potassium manganate(VII) and sodium dichromate(VI), when in acid, are common oxidizing agents.

9. Orders of reaction can only be found experimentally.

10. Wash a buret with the solution that it will be dispensing in the titration, and fill the tip.

11. Phenolphthalein is pink in basic solution.

12. Gases behave ideally when at relatively low pressures and relatively high temperatures.

13. Ideal solutions are made when the two components have very similar IMF’s.

13. The first ionization energy of an atom corresponds to the lowest energy peak on a PES spectrum. No other ionization energies match any PES peaks.

14. Catalysts increase the rate of the forward and the backward reaction.

15. Common ions make slightly soluble salts even less soluble.

16. Kp expressions include ONLY gas partial pressures.

17. In complex ions and (co-ordination compounds) the ligand is a Lewis base and the transition metal ion the Lewis acid.

17. Very large K values suggest that reactions go to completion and massive ones suggest a practical lack of an equilibrium at all.

18. Kw = (Ka) (Kb).

19. Optimal buffers have pH = pKa.

20. In % error calculations, the actual, accepted value is in the denominator.

21. R-O-R is the general formula for an ether.

21. The existence of a C=C double bond (sigma + pi), prevents rotation and can cause cis/trans isomerization.

22. Clean up an acid spill with a carbonate, not an equally corrosive, strong base.

23. Writing the full electronic configuration of an atom can help to explain differences in ionization energies.

24. Transition metal ions are often colored in solution.

25. Reduction always takes place at the cathode.

26. Lowering of vapor pressure and elevation of boiling point are essentially the same thing.

26. Decreasing the [ ] of a reactant in a REDOX equilibrium/galvanic cell reaction, will force the reaction backward and lower the voltage – and vice-versa.

27. Fluorine always has an oxidation number of  -1.

28. The bigger the pKa, the weaker the acid.

29. The bigger the Ka, the stronger the acid.

30. A carboxylic acid can be represented by R-COOH and RCO2H.

31. When using R = 0.0821 in P V = n R T, pressure must be in atm, temperature in K, and volume in L.

32. Neat handwriting and presentation of your work, CAN make your (and the graders) life easier.

33. On the exam, use the FULL atomic masses printed on the periodic table.

34. When predicting shape, a double bond counts as one area of electron density.

35. It is unlikely that any numerical answer on the AP exam will ever require 10 significant figures!

36. Since C and H have a similar electronegativity, hydrocarbons are largely non-polar.

37. Polarity in organic molecules helps them to be soluble in water, otherwise non-polar organic molecules will dissolve non-polar (covalent) solids.

38. 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.

39. Breaking bonds within reactants is ENDOTHERMIC (+ve).

40. Alcohols are soluble because they can H-bond with water, NOT because they have a hydroxide group – they DON’T!

41. Ions travel through the salt bridge, not electrons.

42. Net ionic equations must balance charge as well as atoms.

43. A graph of 1/[X] versus time gives a straight line for a second order reaction, and it has a POSITIVE slope!

44. Bromine and mercury are liquids at room temperature.

45. Transition metals lose their s electrons first.

46. Always use temperature in K in gas calculations.

47. The units of Delta H and Delta S are often different, and if so, must be converted in a Delta G calculation.

48. The cathode and anode have DIFFERENT charges in a galvanic cell and an electrolytic cell.

49. Orders of reaction can be fractions.

50. Iodine is a solid at room temperature.