Periodic Table
(Here is a webpage that contains several different "forms" of periodic table.  Click here!)
—Organizing the elements according to their properties (into groups or families: vertical columns) by Dmitri Mendeleev.

—Organizing the order of their nuclear charge (the atomic number) by Henry Moseley.

—Each raw or period ends with a noble/inert gas. (Does this suggest anything?)

Three major categories: metals, non-metals, and metalloids (NOTE: Transition elements are metals!!)

OR in three groups: s-block (first two columns), p-block (the right 6 columns except the last column of the noble gases), the d-block (the center 10 columns), and the f-block (the "bottom" two raws) which are determined by e in outermost orbital

Metals: malleable, high electrical (and heat) conductivity
Except for Hg, all the metals are solids at room temperature (25 °C or 298 K)
Other metals with low melting points: Cs, m.p. = 301.55 K; Fr, 300 K; Ga, 302.93 K

Nonmetals: gases (O2, H2, N2, the inert/noble gases He, Ne...), brittle solids (B, C, I, etc.), liquid (Br), poor electrical conductors
    Note that all gas molecules are diatomic, ecept inert gases which are monoatomic.

Metalloids: between metals and non-metals, e.g. semi-conductors (Si and Ge)

Important elements in the human body:
Tables 2.2 and 2.3

Metal ions in biological processes (click here to learn more!)

Electron and Energy

—Electrons absorb energy (light) to "jump" from the ground state to the excited states, and emit energy (light) to relax to the ground state as shown by the brilliant color of fireworks. (A good webpage about the color of fireworks:

—The energy gaps are different among different atoms, thus can be used for characterization of elements (e.g., in stars, cf. Figure 2.10).

—These "electron transitions" are thus considered "quantized", i.e., energy levels are not continuous but well-defined steps.

Shells: Principal energy levels/states identified by the principal quantum number 1 (1st raw, can fill up to 2 electrons), 2 (2nd raw, 8 electrons), 3 (3rd raw, 8 electrons), 4 (4th raw, 18 electrons), …

Subshells: "sub-energy levels" within shells, identified by s (can fill up to 2 electrons), p (6 electrons), d (10 electrons), and f (14 electrons)


Electronic structure (configuration) of atoms

Location of electron:

—It is impossible to "pinpoint" the location of an electron.

—The "location" of an electron is described by orbitals, which is a region in space where there is a good chance of finding the electron.
        (Here is an image of the dz2 orbital of Cu.  Click here to view the image!)

—Shapes of orbitals? (cf. P. 54)

    s     1 orbital

    p     3 orbitals (px, py, and pz)

    d     5 orbitals

     f     7 orbitals

The ways in which electrons are arranged in the orbitals:

1. The Aufbau principle: Electrons enter orbitals of lowest energy first, like from the 1st shell then the 2nd

2. The pauli exclusion principle: An atomic orbital may be occupied by at most two electrons. (This two electrons must have opposite spins.)

3. The Hund's rule: Electrons occupy orbitals of the same energy level (degenerate orbitals, like the 3 p orbitals) with maximum number of unpaired electrons. Thus, the 3 electrons in p subshell are arranged in the way as px1 py1 pz1, instead of px2 py1 pz0 or other ways.

Electron configuration and periodic table
First, locate the element in the "Blocks";

e.g., N, p block; Ni, d block

s block: 1A and 2A groups—Na, [Ne] 3s1; Mg, [Ne] 3s2
p block: 3A through 7A groups—O, [He] 2s2 2p4
d block: transition metals—Fe, [Ar] 4s2 3d6
f block: inner transition metals—Pm, [Xe] 6s2 4f5

Second, locate the element in the row (period);

1st row — 1s
2nd — [1s] 2s 2p
3rd — [1s 2s 2p] 3s 3p
4th — [1s 2s 2p 3s 3p] 4s 3d 4p; etc.

e.g., N, 2nd row: 1s2 2s[?] 2p[?]

Third, filling electrons from low to high energy levels

e.g., Ni,  (number of electrons = 28)
4th row [1s 2s 2p 3s 3p] 4s 3d 4p,
d-block: [1s 2s 2p 3s 3p] 4s 3d 4p
8th position: 1s2 2s2 2p6 3s2 3p6 4s2 3d8

NOTE: An element can have different electron configurations by gaining or losing electrons.
e.g., Na: 1s2 2s2 2p6 3s1; Na+: 1s2 2s2 2p6 (3s0)
        O: 1s2 2s2 2p4; O2–: 1s2 2s2 2p6

The two elements are quite active, whereas the ions are not such as in NaCl and CaCO3. (Why? Which element has a similar electron configuration as the ions?)