Name _______________________ Period ______ date due?_________ Pts. ________
Study Guide Chapter 10
(you may use your own paper to answer these questions if you wish)
1. Use the periodic table to predict oxidation numbers.
What is the most common oxidation number for the following elements within an ionic
compound?
(a) O -2
(b) K +1
(c) F -1
(d) Ba +2
(e) P -3
2. (a) Describe the process by which positive ions form.
electron(s) are lost and the atom/ion shrinks.
(b) Describe the process by which negative ions form.
electron(s) are gained and the atom/ion enlarges.
3. Why is it possible for transition elements to have multiple positive oxidation numbers?
Transition metals can lose electrons from different energy levels/sublevels in that they
give off electrons from their outer s sublevel, but most can lose a multiple number of d
sublevel electrons as well. Because of this many of the d-block/transition metals have
multiple oxidation numbers.
4. How is the periodic table used to predict the relative (on your own paper)
(a.) density: Iridium (Z = 77) has the highest density.
(b) atomic radii: Francium (Z = 87) is the largest atom.
(c) ionic radii: Ions increase in size as we move toward the bottom left of the periodic table
(toward francium). However negative ions are larger than positive ions.
(d) first ionization energy: Francium (Z = 87) has the lowest first ionization energy.
5. Define: (on your own paper)
(a) shielding effect: more electrons between the nucleus and the outer level electrons which lessens
(shields), the attraction of the nucleus on the outer electrons.
(b) ionization energy: The energy needed to remove an electron from an atom or ion.
(c) ion: The charge on a single atom.
(d) cation: A positive ion.
(e) anion: A negative ion.
(f) effective nuclear charge: How well an atom is able to hold on to its most loosely held electron.
6. Explain the differences in multiple ionization energies.
Each time an electron is removed from an atom the remaining electron(s) is/are held
more tightly to the nucleus. Therefore each time an electron is removed the next
electron removed requires more energy to remove it.
7. How does effective nuclear charge determine
(a) atomic radii: The greater the effective nuclear charge the smaller the atom.
(b) ionization energy: The greater the effective nuclear charge the higher the ionization
energy.
8. How is the periodic table used to predict the relative reactivity of metals? nonmetals?


Metals increase in reactivity as we move towards francium
Nonmetals increase in reactivity as we move towards fluorine. (Remember that noble
gases are stable.)
9. Relate the reactivity of (a) metals and (b) nonmetals related to their effective nuclear charge,
atomic size and the ability to gain or lose electrons.


Large metal atoms with low effective nuclear charges are more reactive because it is
easier to remove electrons from them.
Small nonmetal atoms with high effective nuclear charges are more reactive because it is
easier for them to capture electrons from other atoms.
10. Know how some elements exhibit more than one possible oxidation number.
 Use the electron configuration of Vanadium (Z = 23) to predict two common oxidation
numbers for this element.
The electron configuration of Vanadium is [Ar]4s23d3. We would therefore
predict Vanadium to exhibit oxidation numbers ranging from +2 ([Ar]3d3)
to +5 ([Ar]).
11. Determine the oxidation number for each atom/ion within the following substances.
(a) SrI2
Sr2+
I-
(b) H2SO4
H+
S6+
O2-
(c) Cu(NO3)2
Cu2+ N5+
O2-
(d) MnS2
Mn4+ S2-
(e) (NH4)2CO3
N3-
H+
C4+
O2-