Block (periodic table)

A block of the periodic table of elements is a set of adjacent groups. The term appears to have been first used by Charles Janet.^[1] The respective highest-energy electrons in each element in a block belong to the same atomic orbital type. Each block is named after its characteristic orbital; thus, the blocks are:

• s-block
• p-block
• d-block
• f-block
• g-block (hypothetical)

The block names (s, p, d, f and g) are derived from the spectroscopic notation for the associated atomic orbitals: sharp, principal, diffuse and fundamental, and then g which follows f in the alphabet.

The following is the order for filling the "subshell" orbitals, according to the Aufbau principle, which also gives the linear order of the "blocks" (as atomic number increases) in the periodic table:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, ...

For discussion of the nature of why the energies of the blocks naturally appear in this order in complex atoms, see atomic orbital and electron configuration.

The "periodic" nature of the filling of orbitals, as well as emergence of the s, p, d and f "blocks" is more obvious, if this order of filling is given in matrix form, with increasing principal quantum numbers starting the new rows ("periods") in the matrix. Then, each subshell (composed of the first two quantum numbers) is repeated as many times as required for each pair of electrons it may contain. The result is a compressed periodic table, with each entry representing two successive elements:

1s 2s 2p 2p 2p 3s 3p 3p 3p 4s 3d 3d 3d 3d 3d 4p 4p 4p 5s 4d 4d 4d 4d 4d 5p 5p 5p 6s 4f 4f 4f 4f 4f 4f 4f 5d 5d 5d 5d 5d 6p 6p 6p 7s 5f 5f 5f 5f 5f 5f 5f 6d 6d 6d 6d 6d 7p 7p 7p

Periodic table

There is an approximate correspondence between this nomenclature of blocks, based on electronic configuration, and groupings of elements based on chemical properties. The s-block and p-block together are usually considered as the main group elements, the d-block corresponds to the transition metals, and the f-block are the lanthanides and the actinides. However, not everyone agrees on the exact membership of each set of elements, so that for example the Group 12 elements Zn, Cd and Hg are considered as main group by some scientists and transition metals by others. Groups (columns) in the f-block (between groups 2 and 3) are not numbered.

In periodic tables organized by blocks, helium is placed next to hydrogen, instead of on top of neon as in tables organized by chemical properties. This is because helium is in the s-block, with its outer (and only) electrons in the 1s atomic orbital. In addition to the blocks listed in this table, there is a hypothetical g-block which is not pictured here. The g-block elements can be seen in the expanded extended periodic table. Also, lutetium and lawrencium are placed under scandium and yttrium to reflect their status as d-block elements^[2] (although it has been argued that lanthanum and actinium should instead hold these positions, as they have no electrons in the 4f and 5f orbitals, respectively, while lutetium and lawrencium do).^[3]

Blocks in the periodic table
Group →	1	2 (+He)		3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18 (−He)
↓ Period
1	1 H	2 He
2	3 Li	4 Be												5 B	6 C	7 N	8 O	9 F	10 Ne
3	11 Na	12 Mg												13 Al	14 Si	15 P	16 S	17 Cl	18 Ar
4	19 K	20 Ca		21 Sc	22 Ti	23 V	24 Cr	25 Mn	26 Fe	27 Co	28 Ni	29 Cu	30 Zn	31 Ga	32 Ge	33 As	34 Se	35 Br	36 Kr
5	37 Rb	38 Sr		39 Y	40 Zr	41 Nb	42 Mo	43 Tc	44 Ru	45 Rh	46 Pd	47 Ag	48 Cd	49 In	50 Sn	51 Sb	52 Te	53 I	54 Xe
6	55 Cs	56 Ba		71 Lu	72 Hf	73 Ta	74 W	75 Re	76 Os	77 Ir	78 Pt	79 Au	80 Hg	81 Tl	82 Pb	83 Bi	84 Po	85 At	86 Rn
7	87 Fr	88 Ra		103 Lr	104 Rf	105 Db	106 Sg	107 Bh	108 Hs	109 Mt	110 Ds	111 Rg	112 Cn	113 Uut	114 Fl	115 Uup	116 Lv	117 Uus	118 Uuo
				57 La	58 Ce	59 Pr	60 Nd	61 Pm	62 Sm	63 Eu	64 Gd	65 Tb	66 Dy	67 Ho	68 Er	69 Tm	70 Yb
				89 Ac	90 Th	91 Pa	92 U	93 Np	94 Pu	95 Am	96 Cm	97 Bk	98 Cf	99 Es	100 Fm	101 Md	102 No
s-block p-block d-block f-block Background color shows the block of the periodic table Primordial From decay Synthetic Border shows natural occurrence of the element

S-block

The s-block is on the left side of the periodic table and includes elements from the first two columns, the alkali metals (group 1) and alkaline earth metals (group 2), plus helium.

Most s-block elements are highly reactive metals due to the ease with which their outer s-orbital electrons interact to form compounds. The first period elements in this block, however, are nonmetals. Hydrogen is highly chemically reactive, like the other s-block elements, but helium is a virtually unreactive noble gas.

S-block elements are unified by the fact that their valence electrons (outermost electrons) are in the s orbital. The s orbital is a single spherical cloud which can contain only one pair of electrons; hence, the s-block consists of only two columns in the periodic table. Elements in column 1, with a single s-orbital valence electron, are the most reactive of the block. Elements in the second column have two s-orbital valence electrons, and, except for helium, are only slightly less reactive.

P-block

The p-block is on the right side of the periodic table and includes elements from the six columns beginning with column 13 and ending with column 18. Helium, though being in the top of group 18, is not included in the p-block.

The p-block is home to the biggest variety of elements and is the only block that contains all three types of elements: metals, nonmetals, and metalloids. Generally, the p-block elements are best described in terms of element type or group.

P-block elements are unified by the fact that their valence electrons (outermost electrons) are in the p orbital. The p orbital consists of six lobed shapes coming off a central point at evenly spaced angles. The p orbital can hold a maximum of six electrons, hence why there are six columns in the p-block. Elements in column 13, the first column of the p-block, have one p-orbital electron. Elements in column 14, the second column of the p-block, have two p-orbital electrons. The trend continues this way until we reach column 18, which has six p-orbital electrons.

Metals

P-block metals have classic metal characteristics: they are shiny, they are good conductors of heat and electricity, and they lose electrons easily. Generally, these metals have high melting points and readily react with nonmetals to form ionic compounds. Ionic compounds form when a positive metal ion bonds with a negative nonmetal ion.

Of the p-block metals, several have fascinating properties. Gallium, in the 3rd row of column 13, is a metal that can melt in the palm of a hand. Tin, in the fourth row of column 14, is an abundant, flexible, and extremely useful metal. It is an important component of many metal alloys like bronze, solder, and pewter.

Sitting right beneath tin is lead, a toxic metal. Ancient people used lead for a variety of things, from food sweeteners to pottery glazes to eating utensils. It has been suspected that lead poisoning is related to the fall of Roman civilization,^[4] but further research has shown this to be unlikely.^[5][6] For a long time, lead was used in the manufacturing of paints. It was only within the last century that lead paint use has been restricted due to its toxic nature.

Metalloids

Metalloids have properties of both metals and nonmetals, but the term 'metalloid' lacks a strict definition. All of the elements that are commonly recognized as metalloids are in the p-block: boron, silicon, germanium, arsenic, antimony, and tellurium.

Silicon is perhaps the most famous metalloid. It is the second most abundant element in Earth's crust and one of the main ingredients in glass. It is used to make microchips for computers and other electronic devices. So valuable is silicon to the technology industry that Silicon Valley in California is named after it.

Germanium has properties very similar to silicon, yet this element is much more rare. Arsenic is a toxic metalloid that has been used throughout history as an additive to metal alloys, paints, and even makeup.

Noble Gases

Previously called inert gases, their name was changed as there are a few other gases that are inert but not noble gases, such as Nitrogen. The noble gases are located in the far right column of the periodic table, also known as Group Zero or Group Eighteen.

All of the noble gases have full outer shells with eight electrons. However, at the top of the noble gases is helium, with a shell that is full with only two electrons. The fact that their outer shells are full means they rarely react with other elements, which led to their original title of "inert."

Because of their chemical properties, these gases are also used in the laboratory to help stabilize reactions that would usually proceed too quickly. As the atomic numbers increase, the elements become rarer. They are not just rare in nature, but rare as useful elements, too.

Halogens

The second column from the right side of the periodic table, group 17, is the halogen family of elements. These elements are all just one electron shy of having full shells. Because they are so close to being full, they have the trait of combining with many different elements and are very reactive. They are often found bonding with metals and elements from Group One, as these elements in each have one electron.

Not all halogens react with the same intensity. Fluorine is the most reactive and combines with most elements from around the periodic table. As with other columns, reactivity decreases as the atomic number increases.

When a halogen combines with another element, the resulting compound is called a halide. One of the best examples of a halide is Sodium chloride (NaCl).

D-block

The d-block is on the middle of the periodic table and includes elements from columns 3 through 12. These elements are also known as the transition metals.

The d-block elements are all metals which exhibit two or more ways of forming chemical bond. Because there is a relatively small difference in the energy of the different d-orbital electrons, the number of electrons participating in chemical bonding can vary. This results in the same element exhibiting two or more oxidation states, which determines the type and number of its nearest neighbors in chemical compounds.

D-block elements are unified by having in their outermost electrons one or more d-orbital electrons but no p-orbital electrons. The d-orbitals can contain up to five pairs of electrons; hence, the block includes ten columns in the periodic table.

F-block

The f-block is in the center-left of a 32-column periodic table but in the footnoted appendage of 18-column tables. These elements are not generally considered as part of any group. They are often called inner transition metals because they provide a transition between the s-block and d-block in the 6th and 7th row (period), in the same way that the d-block transition metals provide a transitional bridge between the s-block and p-block in the 4th and 5th rows.

The known f-block elements come in two series, the lanthanides of period 6 and the radioactive actinides of period 7. All are metals. Because the f-orbital electrons are less active in determining the chemistry of these elements, their chemical properties are mostly determined by outer s-orbital electrons. Consequently, there is much less chemical variability within the f-block than within the s-, p-, or d-blocks.

F-block elements are unified by having one or more of their outermost electrons from in the f-orbital but none in the d-orbital or p-orbital. The f-orbitals can contain up to seven pairs of electrons; hence, the block includes fourteen columns in the periodic table.

G-block

The g-block is a hypothetical block of elements in the extended periodic table whose outermost electrons are posited to have one or more g-orbital electrons but no f-, d- or p-orbital electrons.

References

1. Charles Janet, La classification hélicoïdale des éléments chimiques, Beauvais, 1928
2. Scerri, Eric. "Mendeleev's table finally completed and what to do about group 3".
3. Lavelle, Laurence. "Lanthanum (La) and Actinium (Ac) Should Remain in the d-Block" (PDF). lavelle.chem.ucla.edu. Retrieved 9 November 2014.
4. Wilford, John Noble (17 March 1983). "ROMAN EMPIRE'S FALL IS LINKED WITH GOUT AND LEAD POISONING". The New York Times. Retrieved 19 January 2016.
5. Killgrove, Kristina (20 January 2012). "Lead Poisoning in Rome - The Skeletal Evidence". Powered by Osteons. Retrieved 19 January 2016.
6. Sumner, Thomas (21 April 2014). "Did Lead Poisoning Bring Down Ancient Rome?". Science Magazine. Retrieved 19 January 2016.