Making Sense of Periodic Trends and Correctly and Fully answering Part II AP Questions
On the AP exam, you will most likely be given a “compare and contrast” question for either elements or
compounds. For elements, you will want to pull out your periodic table and use the following
information to help you.
Periodic Trends:
I.
Atomic Radii (The size of an atom)
Normally: Period: As you go across a period, size decreases due to:
1. Same # of p.e.l
(Situation 1)
2. increasing nuclear charge or number of protons (pulling
electrons in).
Group: As you go down a group, size increases due to:
1.
(Situation 2)
Increasing number of principle energy levels (more
and more electrons added onto the outside of the
atom).
Ions:
Comparing the same element as an ion or atom:
(Situation 3)
Metals: lose electrons to become smaller with a positive charge
More + to -, pulling in more.
Nonmetals: gain electrons to become larger with negative charge.
More – to +, less pull from nucleus.
Comparing isoelectronic series:
(Situation 4)
More protons but same # of electrons means smaller size.
II.
Ionization Energy & Electronegativity:
First Ionization Energy: energy to remove outermost electron
Second Ionization Energy: energy to remove 2nd outermost electron
Normally: Period: As you go across a period, ionization energy increases due to: (Sit. 5)
1. Same # of p.e.l
2. Increase in nuclear charge (number of protons)(pulling electrons
more therefore harder to remove)
3. Octet Rule: Group 1 wants to lose electrons, group 17 wants to
gain electrons
Exceptions:
(Sit. 6)
As you go across a period, it is easier to remove an electron from a
partially filled p orbital (higher energy) than a full s orbital (lower
energy level)
(Sit. 7)
As you go across a period, it is easier to remove an electron in period
16 than period 15 due to electron electron repulsion as an electron is
added/paired up with an unpaired electron in the p block.
(sit. 7.5)
Electron affinity: Hund’s rule: 3 unpaired electrons in the p block is a
lower energy state….adding one more electron would cause electron
repulsion (therefore not wanting an electron as much).
Group: As you go down a group, ionization energy decreases due to:
(Sit. 8)
1. More principle energy levels and therefore more shielding and
easier to remove electrons
Second and higher Ionization Energies:
More complicated than 1st ionization energy.
You need to look at the fact that you are removing an electron from a positive ion, the
electron configurations (paired vs. unpaired electrons), and how close it is to an octet.
Comparing 1st vs. 2nd vs. 3rd ionization energies for the same atom
(Situation 9)
Look at how many valence electrons the atom has
-if a metal close to an octet- it is easier to remove
-if it has reached an octet it is harder to remove
-also can be explained by removing an electron from a lower energy
level so harder to remove since closer to the nucleus.
-2nd ionization energy can be higher than 1st slightly due to trying to remove an
electron from a positive ion vs. a neutral atom
Other General Information:
Heisenberg uncertainty principle: One cannot know both the position and the momentum of
an electron.
Quantum Numbers: If you are asked about a specific element’s electrons, use the periodic
table.
N= the period (unless talking about d or f blocks)
L = the blocks (s = 0, p = 1, d = 2, f =3)
Ml = the orbitals (lines in the drawing) ranging from – L to + L
Ms = spin (either + or – 1/2)
Removing electrons from an atom to make an ion:
Take electrons out of the highest (n) energy level first….do not follow Aufbau with the
overlapping 4s and 3d orbitals when removing electrons.
Find the electron and see where it fits with the quantum numbers and remember that
no two electrons can have the exact same #s (Pauli exclusion principle).