THE BIG BAD REVIEW QUESTIONS! YEAH!
1. Where do you find exceptions to the ionization energy trend within periods? Why do they occur?
Generally exceptions are found when new sublevels are begun (as when moving from Mg to Al) or
when moving one electron past a half full sublevel (as when moving from N to O). Electrons in sublevels by
themselves, or that are added to an orbital that already contains an electron are easier to remove than
expected, so ionization energy is lower than expected.
2. Discuss the relationship between the last electrons placed in an atom and the position of each of the
sublevel regions on the periodic table.
The last electron will always correspond to the sublevel block in which the element is located. For
example, nitrogen ends in 2p3 and nitrogen is in the “p” block.
3. Which family would you expect to have an extremely high 3rd ionization energy? Why?
The alkaline earth metals. They have two valence electrons, and these are relatively easy to remove.
The third electron would need to be removed from a lower energy level where there would be less shielding,
so the third electron would be much more difficult to remove. Ionization energies for an individual element
dramatically increase when changing energy levels.
4. Why is there a discrepancy between the ionization energy of phosphorous and sulfur?
Ionization energy for sulfur is lower than phosphorous because the sulfur has 4 electrons in its p
sublevel. This is one more than half full, and the 4th electron is repelled by the electron that is already in the
orbital.
5. Why is there a discrepancy between the atomic size of zinc and gallium?
Gallium is just a bit larger than zinc because the last electron is in the p sublevel all by itself, and is
repelled by all of the electrons that are “below” it.
6. Where do you find exceptions to the electron affinity trends within periods? Why do they occur?
Electron affinity tends to increase as you move across a period, but there are exceptions when
sublevels are full, and half full. Full and half full sublevels tend not to gain electrons.
7. Why is there a discrepancy between the electronegativity of chromium and manganese?
Manganese has a full 4s sublevel and a half full d sublevel. This makes Mn less likely to attract
electrons.
8. Discuss how shielding effect and effective nuclear charge affect each of the trends we discussed.
This is important!
Within a period, increasing nuclear charge (number of protons) causes atomic radius (and ionic
radius) to decrease, which causes electronegativity, electron affinity, and ionization energy to increase.
Within a family, increasing shielding effect diminishes effective nuclear charge even though there are more
protons. This causes the trends above to be reversed.
9. Why does manganese not have an electron affinity? Discuss its electron configuration!
Mn has a full 4s and a half full 3d. This makes it unlikely to gain electrons due to full and half full
sublevels.
10. List the following in order of increasing atomic radius: S-2, Cl-1, K+1, Na+1
Protons
Electrons
S-2
16
18
Cl-1
17
18
K+1
19
18
Na+1
11
10
The sodium cation is the smallest as it has the fewest energy levels. The other three are isoelectronic, so the
ion with the most protons will be the smallest and the ion with the least protons will be the largest. So, the
overall order is Na ion, K ion, Cl ion and S ion.
11. List the following in order of increasing ionization energy: Cl, Ar, Cl-1, K, K+1
Cl
Ar
Cl-1
K
K+1
Protons
17
18
17
19
19
Electrons
17
18
18
19
18
The K atom has the most energy levels, so it will have the lowest ionization energy. The other particles have
their outer shell electrons in the same energy level, so we must look at protons and electrons. K+1 has the
most protons affecting 18 electrons, so it should be the smallest and have the highest ionization energy. This
leaves two chlorines and the argon. The Cl-1 has only 17 protons holding 18 electrons so it would be the
largest of the three and have the lowest IE of the three. Cl and Ar follow the normal periodic trend of size
decrease, so Ar will be the smallest of the two. So the final should be:
K
Cl-1
Cl
Ar
K+1
(note: measurement techniques can vary and complications in radius measurement can occur due to the state of matter in which the measurements are
taken. We will not concern ourselves with these potential problems in this course).
12. Half full sublevels are stable - give an example of this using two elements and a trend!
Ionization energy decreases between N and O.
13. Explain the electron affinity discrepancy between fluorine and chlorine!
The added electron in fluorine is added to an outer shell that is crowded due to its small size. This
means that there is some repulsion from the 7 electrons that are already in the outer shell. In Cl, the outer
shell is larger (further from the nucleus) so it is less crowded, less repulsion from the existing 7 electrons, and
a greater energy release.
14. Ionization energy should increase as we progress through the periodic table, but it actually fluctuates
throughout the transition metals - what phenomenon best explains this?
Repulsion by electrons that are in the complicated d orbital shapes.
15. List the following in order of increasing electron affinity - Mg, Na, I, I-1!
Mg has full sublevel, Na has an open spot in its s sublevel, and I- has a full p sublevel. We know that
Mg has a positive electron affinity and that sodium has a negative affinity, so of these three, Na has the
highest affinity. The iodide ion is large, and the outer shell is full, but it has a proton electron imbalance, so
it should have a high positive affinity. The order of increase is most likely to be Iodide, Mg, and then Na. It is
difficult to determine between iodide and Mg, so don’t worry too much about these two.
16. Sulfur has a negative first electron affinity, but a positive second electron affinity, even though it
wants two electrons total.
Explain why this might be!
After the first electron has been added, the second electron is repelled by the electrons that are now
in the valence shell. Kind of like the irregularity between chlorine and fluorine.
17. Why does the ionization energy decrease between Be and B, if B has more protons? Explain!
The last electron in B is in a new energy level.
18. An atom releases energy during a process. Is it now more or less stable?
More energetically stable. Its potential energy has been released.
19. Why is the second ionization energy for magnesium higher than the first, if it wants to give away two
electrons total?
After one electron has been removed, there is a greater effective nuclear charge attracting the
remaining electrons.
20. Ionization energy is ALWAYS positive. What does this tell us about the energy states of atoms vs.
cations?
Since energy is required to remove electrons, cations will be in higher energy states than anions.
21. How many valence electrons does V have? How about Co? How about Ag? What do these elements
have in common with each other?
block.
All should have two valence electrons, although they vary due to their location in the fickle d sublevel
22. What is so unique about the transition elements and their valence states?
Valence states vary greatly in the d sublevel block. They can change within a given element
depending on the reaction conditions.
23. If Br-1, Kr, and Rb+1 all have the same number of electrons, then why is Rb+1 the smallest of the three?
Rb has the most protons.
24. Why is I-1 bigger than Kr if they both have the same number of electrons?
I-1 has fewer protons
25. Why are Bi, Po, Pb, and Sb metals if they are in the p block of elements?
They are the largest in their families. This makes them likely to lose electrons, which is characteristic
of metals.
26. Sn is in group IV, but it is a metal - why is this?
It is large, and tends to lose electrons.
27. Does K or As have a greater effective nuclear charge? What effects in trends do we witness because of
this?
Ar has a greater Zeff (effective nuclear charge. Although K has more protons, it also has an another
energy level. This increases the shielding due to the inner shell electrons and reduces the effective nuclear
charge.
28. Sr has more electrons and protons than Ca, and electrons attract protons, but it is a larger atom explain!
See above explanation. Zeff for Sr is less due to more shielding.
29. Carbon, Nitrogen, and Oxygen - which has the highest effective nuclear charge?
Zeff is greatest for oxygen. The outer shell is the same energy level for all, but oxygen has more
protons.
30. Cr and Cu have different electron configurations listed on the periodic table than what we would
expect - why?
Due to the increased stability (lower energy state) of full and half full sublevels, an electron gets
promoted from 4s to 3d in both elements.
31. Why do transition metals have multiple oxidation states?
Sometimes these elements can lose their d electrons, depending on reaction conditions.
32. Why does Titanium form a +3 or +4 ion? Why does Iron form a +2, +3, or +6 ion?
Ti may lose all of its 4s electrons and 3d electrons, or sometimes 4s and one 3d electron. It also
sometimes loses only the 4s, which creates a +2 ion.
Fe may lose 2 4s electrons, 2 4s and 1 3d electrons, or all of its 3d electrons.
33. As we progress down the periodic table, what should happen to electron affinity? Why, then, between
periods 2 and 3, such as between B and Al, and C and Si, does it increase?
Electron affinity should decrease due to increased shielding effect. The increase between
periods two and three occur because of the “crowdedness” explained previously for fluorine.
34. Why does electron affinity actually decrease between groups I and II? Don’t more protons mean more
attractions to electrons?
Group two has a full s sublevel.
35. If Al has more protons than F, then why does it have a lower electron affinity? (Two reasons)
Al has more shielding effect which means that it has a lower Zeff. Al is further to the left on the
periodic table which also would reduce its Zeff.
36. What does the oxidation state of an element tell us?
The number of electrons that are lost (+ oxidation state), gained (negative oxidation state), or shared
(positive or negative oxidation states depending on what elements are being combined. The more
electronegative element will have the negative oxidation state.
37. An atom loses an electron. Does this process always require energy, always give off energy,
sometimes require energy, sometimes give off energy, or do we not know unless we know what atom we
are talking about?
Losing electrons always requires energy. Ionization energies are always positive.
38. What does it mean to have an exothermic process, or a negative energy, during an atomic process?
A negative energy change means that energy is released, and that the resulting particle is more
energetically stable.
39. Why are there no electronegativities listed for most of the noble gases? Why does Cr have a higher
electronegativity than Mn, if it has less protons pulling on electrons?
a. Noble gases have full outer shells (an octet in the outer shell) and tend not to form chemical bonds.
See explanation for question 7.