AP CHEMISTRY CHAPTER 10 HIGHLIGHTS, NOTES AND OBSERVATIONS
Many representative metal atoms lose electrons to achieve a noble gas
configuration, that is, a filled "p" sublevel. Many also gain electrons to achieve this
stable configuration. Examples are the alkali metals, alkaline earth metals and group
three elements (losing 1, 2, and 3 electrons respectively) and the oxygen family and the
halogens (gaining 2 and 1 electrons respectively). Different ions that achieve the same
electron configuration are said to be isoelectronic. Different ions that achieve the
same charge but differ in the total number of electrons (the principle quantum number is
different) are said to be isovalent ( K+ and Na+ for example).
The previous ideas work best (or only) for the main group or representative
elements. Transition metals do not (with some exceptions) form ions that achieve a
noble gas configuration. We cannot tell whether transition metals lose "s" or "d"
electrons when forming positive ions.
We do know that metal ion spectral
measurements indicate that there are no "s" orbital electrons in the ground state
configuration. This corresponds nicely with the observation that the most common
positive ion charge is 2+ for transition metals! (Examine your Periodic Table)
Remember that the "s" and "d" orbital energies lie very close to each other. The
electron configuration of any transition metal ion should NOT have s electrons. This is
a common tested fact. Many transition metal ions also have the ability to form more
than one type of ion due to the stability of half filled d and f orbitals. For example, one
can rationalize Fe3+. Take 3 electrons away (two from the s orbital and one from the d
orbital) and you end up with 1s22s22p63s23p64s23d5 with each of the 5 d electrons
occupying its own d orbital (Hund's rule). Of course, how in the world can you
rationalize the schizophrenic behavior of manganese?
You should be familiar with the definitions of atomic radius: covalent radius,
ionic radius and metallic radius. Be able to discuss the variation of atomic radii
within a group, within a period and within a transition series. The concept of effective
nuclear charge, Zeff, is worth a one-time careful read through. A very nice analogy for
shielding is to examine the section entitled, Visualizing Concepts, and problem 7.1 on
the bottom of page 291.
Be able to identify the radii change when atoms accept or lose electrons and be
able to explain this phenomenon.
Ionization Energy relates only to gaseous atoms. (This is a frequently tested
fact.) Be familiar with trends in ionization energies and why these trends are fairly easy
to rationalize in terms of electron configurations, shielding and size. Ionization energy
may be thought of in terms of ∆H, (kJ/mol). Note Figure 7.10. You should be able to
explain the "glitches" in the trends as related to the filling of s, p or d orbitals.
Explain the data in Table 7.2 in terms of valence and core electrons. There are
neat testing problems related to this chart.
Electron Affinity is a measure of the change in energy when a gaseous atom
gains an electron (NEVER LOSS OF AN ELECTRON). If energy is released, the event
is exothermic and E.A. is a negative quantity. If E.A. is largely negative, one could say
that the addition of an additional electron imparts increased stability to the species. If
the value is not listed or positive then the gain of an electron is not desirable in terms of
energy. Your text highlights that the addition of a second electron is always an
endothermic event due to fact that the 2 nd electron approaches a now negative ion. Use
the data in Figure 7.12 to help. In the next chapter we will relate the concept of
electronegativity trends to both ionization energy and electron affinity.
You should be able to write, or at least recognize the simple equations denoting an
ionization energy and electron affinity event.
Note the terms diamagnetic and paramagnetic and be able to rationalize these
properties in light of electron configurations. We shall see later that the idea of
unpaired electrons in molecules can impart magnetic properties (the famed but
complicated Molecular Orbital Theory).
You should be aware of general trends in physical and chemical properties as
you move down and across the periodic table. Most important are being able to explain
why the melting point of the Alkali metals and Alkaline Earth metals decreases as you
move down the Periodic Table AND why the boiling points of Halogens and the Noble
gases increases as you go down the Periodic Table. Of high importance is the
reducing abilities of Group 1A and 2A metals (especially on water) and the oxidizing
abilities of the Halogens, especially F2 and Cl2.
Know the terms acidic oxides, basic oxides and amphoteric and the typical
chemical reactions these substances experience. Yes, the number of categories of
equations that you should recognize just formally went up! You should know the
reactions that Alkali and Alkaline Earth metals undergo with water. You should know
the reactions of both metal oxides and nonmetal oxides when placed in water. You
should know the typical reactions of metal oxides with acids and nonmetal oxides with
bases. Your book highlights several of these reaction types.