Chemistry — AP Chemistry Review

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This deck includes 59 flashcards covering pes spectrum, atomic radius, radius change, and related concepts. Use it to review key Chemistry ideas, focus on weak cards, and prepare for your exam with StudyLess.

1. 01

   What does each peak represent in a photoelectron spectroscopy (PES) spectrum?

   A subshell

2. 02

   How do you determine the relative number of electrons from a PES spectrum?

   By the relative peak height (e.g., the first peak is $1s^2$)

3. 03

   How does atomic radius change across a period from left to right?

   It decreases due to an increasing effective nuclear charge drawing valence electrons closer to the nucleus.

4. 04

   How does atomic radius change down a group on the periodic table?

   It increases due to added sublevels and the shielding effect.

5. 05

   Why is a cation smaller than its neutral atom?

   It has lost valence electrons while the nuclear charge remains the same, resulting in a stronger inward pull on fewer electrons.

6. 06

   Why is an anion larger than its neutral atom?

   Additional electrons increase electron-electron repulsion, and the constant nuclear charge cannot pull them in as effectively.

7. 07

   What is ionization energy?

   The energy needed to remove an electron.

8. 08

   How does ionization energy trend across the periodic table?

   It increases from left to right due to greater effective nuclear charge and decreases down a group due to increased shielding.

9. 09

   What is electronegativity?

   The ability of an atom to attract electrons toward itself.

10. 10

    Which element has the highest electronegativity value on the AP exam scale?

    Fluorine (4.0)

11. 11

    What is electron affinity?

    The energy change when an atom becomes negatively charged (always negative overall).

12. 12

    What does the Aufbau Principle state?

    Electrons fill the lowest energy orbitals first.

13. 13

    What does Hund's rule state about orbital filling?

    Each sublevel must have one electron before any are doubled up.

14. 14

    What does the Pauli Exclusion Principle state?

    No two electrons in the same atom can have identical values for all four of their quantum numbers.

15. 15

    According to Coulomb's law, what makes ionic attractive force stronger?

    Larger ionic charge and smaller ionic radius (shorter distance between ions).

16. 16

    How does bond order affect bond energy?

    Smaller bonds require more energy to break; double bonds have more bond energy than single bonds.

17. 17

    What is lattice energy?

    Energy released when ions form a crystalline solid; smaller ions with higher charge have more lattice energy.

18. 18

    What is the difference between substitutional and interstitial alloys?

    Substitutional alloys have metal atoms of roughly the same size, whereas interstitial alloys have smaller atoms filling the spaces between larger metal atoms.

19. 19

    What is the formula to calculate reaction enthalpy from bond enthalpies?

    $$\Delta H^{\circ}_{reaction} = \sum \Delta H_{bonds\ broken (reactants)} - \sum \Delta H_{bonds\ formed (products)}$$

20. 20

    What is the Duet Rule in Lewis structures?

    It occurs between Hydrogen and Helium where they share 2 electrons.

21. 21

    What are the common incomplete octet exceptions for Lewis structures?

    Boron can only have 6 electrons, and Beryllium can only have 4 electrons.

22. 22

    Which elements can form an expanded octet?

    Atoms in row 3 or above (such as Sulfur, Phosphorus, and Arsenic).

23. 23

    How is formal charge calculated?

    Actual valence electrons minus (number of bonds + unshared valence electrons).

24. 24

    What is a resonance structure?

    When there is more than one valid Lewis structure for a compound, where the actual bond length is an average of the structures.

25. 25

    What does VSEPR stand for?

    Valence Shell Electron Pair Repulsion

26. 26

    How many sigma and pi bonds are in a single bond, a double bond, and a triple bond?

    Single bond = 1 sigma; double bond = 1 sigma and 1 pi; triple bond = 1 sigma and 2 pi.

27. 27

    Rank the general intermolecular force (IMF) strength from weakest to strongest.

    London dispersion forces < dipole-dipole < hydrogen bonding < ion-dipole < ionic bonding.

28. 28

    What causes London dispersion forces?

    An instantaneous dipole in one atom induces an uneven electron distribution in a neighboring atom.

29. 29

    What defines a hydrogen bond?

    A strong dipole-dipole interaction between a hydrogen atom bonded to F, O, or N and another F, O, or N atom.

30. 30

    What is the Ideal Gas Law equation?

    $PV = nRT$

31. 31

    What are the core assumptions of the Kinetic Molecular Theory (KMT) for ideal gases?

    Gas molecules have no volume, exert no attractive or repulsive forces, and undergo perfectly elastic collisions.

32. 32

    Under what conditions do real gases deviate most from ideal behavior?

    High pressure, low temperature, and low volume.

33. 33

    What is the triple point on a phase diagram?

    The condition of temperature and pressure where all three phases coexist in equilibrium.

34. 34

    What is the critical point on a phase diagram?

    The point beyond which a gas can no longer be liquefied regardless of pressure.

35. 35

    What rule governs solubility interactions between substances?

    "Like dissolves like" (substances with similar intermolecular interactions tend to be soluble in each other).

36. 36

    What is Beer's law equation for colored solutions?

    $A = abc$ (Absorbance = molar absorptivity × path length × concentration)

37. 37

    What are the steps to write a net ionic equation?

    Start with the balanced molecular equation, write the complete ionic equation, and cross off spectator ions.

38. 38

    What is the equivalence point in a titration?

    The point where exactly enough titrant has been added to react completely with all of the analyte.

39. 39

    How do Bronsted-Lowry acids and bases function?

    An acid donates a proton ($H^+$) and a base accepts the proton.

40. 40

    What does LEO says GER stand for in redox reactions?

    Loss of Electrons is Oxidation; Gain of Electrons is Reduction.

41. 41

    What is the oxidation number of a single uncombined element?

    0

42. 42

    What is a precipitation reaction?

    A reaction where two aqueous ionic compounds combine to form an insoluble solid.

43. 43

    What factors affect reaction rate?

    Concentration of reactants, presence of catalysts, temperature, pressure, and surface area.

44. 44

    What determines the overall order of a rate law?

    The sum of the exponents ($m + n + \dots$) in the rate law expression.

45. 45

    In a reaction mechanism, which step determines the overall rate law?

    The slowest step (rate-determining step).

46. 46

    What is required for a successful collision according to the collision model?

    The correct orientation and sufficient energy to break bonds.

47. 47

    How do catalysts increase reaction rates?

    By providing an alternative reaction pathway with a lower activation energy.

48. 48

    What does the First Law of Thermodynamics state?

    Energy cannot be created or destroyed; the total energy of an isolated system is constant.

49. 49

    What are four different ways to find enthalpy change ($\Delta H$)?

    $q = mc\Delta T$, bond enthalpies, standard enthalpies of formation ($\Delta H_f^{\circ}$), and Hess's Law.

50. 50

    What does Hess's Law state about reversing a reaction?

    Flipping a reaction reverses the sign of its $\Delta H$ and inverts its equilibrium constant ($1/K$).

51. 51

    What is entropy ($\Delta S$)?

    A measure of disorder or randomness in a system.

52. 52

    Rank the standard entropy states of matter from lowest to highest.

    $S_{solid} < S_{liquid} < S_{gas}$

53. 53

    What does the Third Law of Thermodynamics state?

    The entropy of a perfect crystal at $0\text{ K}$ is zero.

54. 54

    What is the Gibbs Free Energy equation relating enthalpy, temperature, and entropy?

    $$\Delta G = \Delta H - T\Delta S$$

55. 55

    What condition of $\Delta G^{\circ}$ indicates a thermodynamically favored (spontaneous) reaction?

    $\Delta G^{\circ} < 0$

56. 56

    What is the mnemonic device for remembering where oxidation occurs in an electrochemical cell?

    AN OX (Anode = Oxidation)

57. 57

    What is the mnemonic device for remembering what happens at the cathode's mass?

    FAT CAT (Cathode's mass increases because cats gain weight)

58. 58

    How do you identify the cathode from a table of standard reduction potentials?

    The half-reaction with the largest $E^{\circ}$ value is the cathode (gets reduced).

59. 59

    What is the relationship equation between Gibbs free energy and cell potential?

    $$\Delta G^{\circ} = -nFE^{\circ}$$