—Organizing the elements.in 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: http://cc.oulu.fi/~kempmp/colours.html)

—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 (3^rd raw, 8 electrons), 4 (4^th 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 d_z2 orbital of Cu.  Click here to view the image!)

—Shapes of orbitals? (cf. P. 54)

    s     1 orbital

    p     3 orbitals (p_x, p_y, and p_z)

    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 1^st shell then the 2^nd…

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 p_x¹ p_y¹ p_z¹, instead of p_x² p_y¹ p_z⁰ 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; O^2–: 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?)