Periodic Table

The periodic law refers to the periodic or systematic variation of physical and chemical properties when elements are arranged in the increasing order of atomic number. The organization of elements into different sets of properties gives rise to what is known as the periodic table. The horizontal row refers to a period of elements with distinctive properties whereas the vertical column refers to a group of elements with similar properties.

There are specific names and notations for different groupings of elements. The three major classes of elements are (a) the representative elements (b) the transition elements (c) the rare earth elements.

Metals and Nonmetals

A bold zig-zag (i.e. "red staircase") line runs from top to bottom of the Periodic Table beginning to the left of boron (B) and ending between polonium (Po) and astatine (At). This line acts as the boundary between metals, to the left, and nonmetals, to the right. Elements straddling the boundary such as B, Si, Ge, and As have properties intermediate between metals and nonmetals and are often termed metalloids.

Atomic Number and Atomic Mass

The atomic number (Z) or the number of protons in the nucleus (i.e. the nuclear charge) and the atomic mass of each element are available from the periodic table. More detailed periodic tables may provide information such as electron arrangement, relative sizes of atoms, and most probable ion charges.

MAIN GROUP OR REPRESENTATIVE ELEMENTS

The representative elements are denoted with the last letter of "A" in the group names and there are eight major groups. Some of the more notable groupings are listed as follow:

Group IA (alkali metals) consists Li, Na, K, Rb, Cs, Fr

Group IIA (alkali earth metals) consists of Be, Mg, Ca, Sr, Ba, Ra.

Grous IIIA, IVA, and VA elements are located close to the "staircase" that divides the metals on the left and the non-metals on the right. Most elements located along the "staircase" are classified as metalloids, e.g. B, Si, Ge, As, Sb, and Po.

Group VIA consists of O, S, Se, Te, and Po.

Group VIIA (halogens) consists of F, Cl, Br, I, and At.

Group VIIIA (rare gases) consists of He, Ne, Ar, Kr, Xe, and Rn.

TRANSITION AND RARE EARTH METALS

The transition elements occupy 4 rows or periods in the periodic table which includes

1) the first row of transition elements from atomic number of 21 (Sc) to 30 (Zn);

2) the second row transition elements from atomic number of 39 (Y) to 48 (Cd);

3) the third row transition elements from atomic number of 72 (Hf) to 80 (Hg) and 57 (La);

4) the fourth row transition elements from atomic number of 104 (Rf) to 109 (Mt) and 89 (Ac).

The rare earth elements consist of lanthanides with the atomic number of 58 (Ce) to 71 (Lu) and actinides with the atomic number of 90 (Th) and 103 (Lr).

Electronic Configuration

Electrons surrounding the nucleus of any atom can be classified as valence and inner-core electrons.

Valence electrons are located in the outermost energy levels which determine the reactivity of elements by losing, gaining, or sharing electrons.

Inner-core electrons generally do not participate in chemical reactions because they are held more tight by the protons in the nucleus by attractive forces.

The number of valence electrons in an atom is the same as the group number of the representative elements.

The energy level (n) of the valence electrons corresponds to the period number.

Valence Electrons

Outermost electrons in an atom are valence electrons. For representative elements, the number of valence electrons in an atom corresponds to the number of the family in which the atom is found. Metals tend to have fewer valence electrons, and nonmetals tend to have more valence electrons. The energy levels are symbolized by n, with the lowest energy level assigned a value of n = 1 and the highest energy level having a value of n = 7.

1. Each energy level may contain up to a fixed maximum number of electrons. Two general rules of electron configuration are based upon the periodic law:

2. The number of valence electrons in a neutral atom equals the group number for all representative (A group) elements.

3. The energy level (n =1,2, etc.) in which the valence electrons are located corresponds to the period in which the element may be found.

Helium is an exception to rule 1, above. It cannot have eight valence electrons, since all of its electrons are in the n=1 level which has a maximum capacity of only two electrons.

Energy Levels and Sublevels

The principal energy levels are designated n =1,2,3 and so forth. The number of possible sublevels in a principal energy level is also equal to n. When n =1 there can be only one sublevel; n =2 allows two sublevels, and so forth. The total electron capacity of a principal level is 2(n)².

The sublevels or subshells have unique orbital shapes and increase in energy in the order of s<p<d<f

Both the principal energy level and type of sublevel are specified when describing the location of an electron (e.g. 1s, 2s, 2p etc.) The first principal energy level (n=1) has one possible subshell, 1s. The second principal energy level (n=2) has two possible subshells; 2s and 2p. The third principal energy (n=3) has three possible subshells: 3s, 3p, 3d. The fourth principal energy level (n=4) has four possible subshells: 4s,4p, 4d, and 4f.

An orbital is a specific region of a subshell containing a maximum of two electrons. The sublevel contains only one orbital, the p sublevel contains three orbitals, the d sublevel contains five orbitals, and the f sublevel contains seven orbitals. Each orbital may be empty, contain one electron, or be filled, containing two electrons. The s, p, d, and f sublevels have maximum capacities of 2,6, 10, and 14 electrons.

Each type of orbital has its own characteristic shape. The s orbital is spherical and the orbital size increases as the period number increases (i.e. 1s < 2s< 3s.......<7s). There exist three kinds of p orbitals, each identical in shape (often modeled as a dumbbell). They differ only in their orientation in space, along the hypothetical x-axis (p_x), y-axis (p_y), and z-axis (p_z). The d and f orbitals are more complex and are important in accommodating the valence electrons of d- and f-orbitals, respectively.

Each atomic orbital has a maximum capacity of two electrons. The electrons are perceived to spin on an imaginary axis. Usually the electron with a clockwise spin is designated with the quantum number of +1/2 and the one with anticlockwise spin the -1/2. The two electrons in the same orbital must have opposite spins are referred to as paired electrons.

Aufbau Principle

Aufbau Principle refers to the order of filling up the energy levels with electrons from the lowest to the highest energy.

The principal energy levels (n=1,2,….7) can further be subdivided into sublevels (s, p, d, f). Examples: K-shell (n=1) has only one 1s sublevel; L-shell (n=2) has 2s and 2p sublevels, M-shell (n=3) has 3s, 3p, and 3d sublevels, N-shell (n=4) has 4s, 4p, 4d, and 4-f orbitals.

For each sublevel, there are different number of orbitals (e.g. s, p, d, f each has 1, 3, 5, 7 orbitals, respectively) with each orbital being able to accommodate 2 electrons of opposite spins.

The maximum number of electrons (2n²) that can be present in K, L, M, and N-shells are 2, 8, 18, and 32, respectively.

The increasing order of orbital energy-levels is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d.

Octet Rule and Electronic Configuration

The octet rule states that, in chemical reactions, elements tend to gain, lose, or share their valence electrons to achieve the stable electron configuration of the nearest inert gas (i.e. 8 electrons for Ne, Ar, Kr, Xe, and Rn).

Elements forming ionic compounds follow the octet rule by losing and gaining electrons for the positive cations and the negative anions, respectively (e.g. NaCl, MgO, CaF₂, FeCl₃, KBr, ZnSe, Li₂O).

Elements forming covalent compounds follow the octet rule by sharing electrons (e.g. CO, and CO₂)

Trends of Properties in Periodic Table

Atomic size decreases from left to right and increases from top to bottom in the periodic table.

Cations are smaller than their parent atoms but anions are larger than their parent atoms.

The ionic radius of a given element decreases as the positive charge increases (Al⁺>Al²⁺>Al³⁺) but increases as the negative charge increases (O^2->O^-).

The ionization energy (Na à Na⁺ + e) decreases going down a group but increases going from left to right; the electron affinity decreases going down the group and increases from left to right in a period.

Ionization Energy

The energy required to remove and electron from an isolated atom in the gas phase is the ionization energy. The magnitude of the ionization energy correlates with the strength of the attractive force between the nucleus and the outermost electron.

As we go down a group, the ionization energy decreases, since the atom's size is increasing. The outermost electron is progressively farther from the nuclear charge and hence easier to remove.
As we go across a period, atomic size decreases, as the outermost electrons are closer to the nucleus, more tightly held, and more difficult to remove. Therefore, the ionization energy must increase.

A correlation does indeed exist between trends in atomic size and ionization energy. Atomic size decreases form bottom to top of a group and left to right in a period. Ionization energies increase in the same periodic way. Note also that ionization energies are highest for the noble gases; this accounts for the extreme stability and non-reactivity of the noble gases.

Electron Affinity

The energy released when a single electron is added to a neutral atom in the gaseous state is known as the electron affinity. Electron affinity is a measure of the ease of forming negative ions. A large value of electron affinity (energy released) indicates that the atom becomes more stable as it becomes a negative ion (through the process of gaining an electron).

Periodic trends for electron affinity are as follows:

1. Electron affinities generally decrease as we go down a group.

2. Electron affinities generally increase as we go across a period.

Be aware that exceptions to these general trends do exist.

Chemical Bonding

The combination or bonding of different elements to form chemical compounds can generally be classified into ionic or covalent bonding

In ionic bonding, there is transfer of electrons from metals with low ionization energy (i.e. Group 1A, 2A) to non-metals with high electron affinity (i.e. Group 6A and 7A).

In covalent bonding, the electrons are shared between 2 or more elements which do not have strong tendency to lose or gain electrons.

Both ionic and covalent compounds can be indicated by Lewis electron structures.