Periodic Table-Classification of Elements, Chemistry tutorial


There are basically 103 identified elements, a few of which are familiar and others are rare. Remember that the atoms are built of particles of three types: protons, neutrons and electrons. The nucleus of each and every atom is made up of protons and neutrons. The number of protons (that is, the atomic number) finds out the electric charge of the nucleus and the total number of neutrons and protons (that is, the mass number) finds out its mass. In a neutral, atom the number of electrons surrounding the nucleus is equivalent to the atomic number.

The physical and chemical properties of an element are administered by the number and arrangement of the electrons. Some of the attempts have been made since the year 1817 to group elements altogether based on recurring properties like atomic weight. The most significant step in the growth of the periodic table was published in the year 1869 by Dmitri Mendeleev, who made a detailed study of the relation between the atomic weights of the elements and their chemical and physical properties. The term periodic implies recur at regular period.

The preliminary arrangement has now been largely substituted following new knowledge regarding electronic structure of atoms. The present periodic table is mainly based on the recurrence of characteristic properties whenever elements are arranged in order of increasing atomic number.

The Periodic Law and Periodic table:

The periodic law defines that the properties of chemical elements are not random, however based on the structure of the atom and differ with the atomic number in a systematic manner. In another words, the properties of the elements are periodic function of their atomic number. Whenever elements are methodically arranged in order of increasing atomic number, some of the characteristics reappear at regular intervals.

The periodic reappearance or recurrence of properties of the elements by means of increasing atomic number can be clearly illustrated by arranging the elements in a table, termed as the periodic table of the elements. The periodic table illustrates the arrangement of elements in seven horizontal rows and eight vertical columns as illustrated in table shown below:

2194_Periodic Table.jpg

Table: Periodic table

Description of the Periodic Table:

The periodic table comprises of rows of elements termed as periods and columns of elements are termed as groups. The horizontal rows of the periodic table comprises of a very short period (having hydrogen and helium, atomic number 1 and 2), two short periods of 8 elements each, two long periods of 18 elements each, a very long period of 32 elements, and an incomplete period. The elements in the period encompass the similar number of shells and the number of valence electrons increases gradually by one across the period from left to right. For illustration, members of period 2 are Li, Be, B, C, N, O, F and Ne having 3, 4, 5, 6, 7, 8, 9 and 10 electrons correspondingly.

For all the members of the period, the extra electron is added to the second shell therefore the name period 2. In general, each and every period begins with an element having one electron in its outermost shell (example: Li, Na, K) and ends by an element whose outermost shell is fully filled (example: He, Ne and Ar - the inert or noble elements). The properties of elements change in a systematic manner via a period. For illustration the first members of each and every period are all light metals which are reactive chemically, and this metallic character reduce across the periods that ends by the un-reactive inert gases.

The elements which appear in a vertical column belong to the similar group or family. They have the similar number of outer electrons or valence electrons and have closely associated physical and chemical properties. The groups IA and HA elements are generally positioned at the left side of all the periods and IRA, IVA, VA, VIA VIIA and 0 (at times termed as VHIA) elements are at the right side. The central elements of the long periods, called the representative elements have properties differing from those of the elements of the short periods. The transition metals are placed in groups IIIB, IVB, VB, VIB, VIIB, VIIIB, IB and IIB.

The very long periods (6 and 7) are compressed to the table by eliminating 14 elements each, termed as the inner transition or rare-earth metals (Z = 58 to Z = 71 → lanthanides; and Z = 90 to Z = 103 → actinides and representing them separately beneath the table. The lanthanides that are in period 6, beginning from lanthanum (La) and ending with Lutetium (Lu) are rare metals which illustrate a great resemblance to one other. The actinides that are in period 7, beginning with actinium (Ac) are termed as artificial elements as they don't occur naturally however are made up all through nuclear reactions. They are unstable and short-lived.

The periods:

The periods are numbered from 1 to 7 and the elements in the similar period encompass the similar number of electron shells. Period 1 elements have one electron shell (K); period 2 elements encompass two electron shells (K, L); period 3 elements encompass three electron shells (K,L,M); and so on. The number of valence electrons in the atoms of the elements in the similar period increase gradually by one from left to right.

Across a specific period, there is a gradual change in chemical properties. For illustration, metallic properties reduce across the period whereas non-metallic features increases. The first three members of any period (Groups 1 to 3), apart from period 1 are metals whereas those of Group 4 to 7 and 0 are non-metallic in behavior. By using period 3 as an illustration, sodium, magnesium and aluminium are metallic and form mostly ionic compounds and basic oxides. To the right of the period, phosphorus, sulphur and chlorine are non-metallic and form mostly covalent compounds and acidic oxides.

The main groups:

Elements in the groups IA to VIIA and group 0 are usually termed to as the main Group elements. Hydrogen is positioned in group IA for ease only due to the single electron however does not encompass alike features with other members of the group.

1) Group IA: the alkali metals: Lithium (Li), Sodium (Na), Potassium (K) and so on are light metals that are extremely reactive chemically that is, reacts vigorously with cold water to discharge hydrogen gas The atoms of each and every element encompass only one electron, which they readily contribute and are strong reducing agents. They react via losing this valence electron to form ionic or electrovalent bonds. The alkali metals are exceptional conductors of electricity as the valence electrons are mobile.

Na → Na+ + e-

The alkali metals are made up through electrolysis of the molten hydroxides or chlorides. Due to their reactivity particularly with water, the metals should be kept in an inert atmosphere or under oil.

"The Sodium metal catches fire whenever in contact with water, thus avoid dropping it in the sink in the laboratory".

The metals are very helpful chemical reagents in the laboratory, and they discover industrial utilization in the manufacture of organic chemicals, dyestuffs and tetraethyl lead (that is, the anti-knock agent in gasoline). Sodium is employed in sodium - vapor lamps, and due to its high heat conductivity, in the stems of valves of airplane engines, to conduct heat away from the valve head.

2) Group IIA: the alkaline - earth metals: The metals of group IIA are beryllium (Be) magnesium (Mg), calcium (Ca) and so on and are much harder and less reactive than the alkali metals. They encompass two electrons in their outermost shell and react basically by making divalent ionic bonds. The compounds of all the alkaline - earth metals are identical in composition; they form oxides MO, hydroxides M(OH)2, trioxocarbonates MCO3 and so forth.

Ca → Ca2+ + 2e-

Ca + 2H2O → Ca(OH)2 + H2

3) Group IIIA: Boron (B) is a metalloid-intermediate property between the metals and non-metals; while aluminium (Al) and other members of the group are metals. The members of group are trivalent as each of its atoms consists of three valence electrons and forms electrovalent compounds.

Al → Al3+ + 3e-

Aluminium is the only known element of the group that reacts with steam at 600°C to give hydrogen gas. The oxide and hydroxide of aluminium are amphoteric - they include both basic and acidic properties.

Al(OH)3 + 3H2SO4 → Al2(SO4)3 + 3H2O

Al(OH)3 + NaOH → NaAl(OH)4

4) Group IVA: Carbon (C), Silicon(Si), germanium (Ge), Tin (Sn) and Lead (Pb) are the members of this group. Their atoms each consists of four valence electrons and tend to make covalent compounds. Carbon is a non-metal; silicon and germanium are metalloids whereas tin and lead are metals exhibiting a gradation from non-metallic to metallic character on going down the group.         

Carbon and silicon make more stable +4 state compounds (CO2 and SiO2- Sand) whereas tin and lead form more stable +2 state compounds (SnO and PbO). The compounds of carbon and hydrogen termed as hydrocarbons make a large class of organic compounds employed as fuels example: butane (C4H10) and hexane (C6 H14).

5) Group VA: Nitrogen (N) and phosphorus (p) are the well-known members of this group. Their atoms each comprise of five valence electrons and illustrate two common valences of 3 and 5. Both of them are non-metals. They are electron acceptors in their reactions and make some oxides example: N2O3, N2O5, P4O6 and P4O10

The oxides are acidic and they react via water to form acids.

N2O5 + H2O → 2HNO3 (Trioxonitrate (v) acid)

Nitrogen and phosphorus react by hydrogen to make ammonia (NH3) and phosphine (PH3).

6) Group VIA: Oxygen (O) and sulphur (S) are the familiar members of this group and are both non-metals. They are electron acceptors and are oxidizing agents example:

2Mg + O2 → 2MgO

2Na + S → Na2S

The elements don't attack water however join with hydrogen to give water and hydrogen sulphide

2H2 + O2 → 2H2O

H2 + S → H2S

7) Group VIIA: Fluorine (F), Chlorine (Cl), bromine (Br) and Iodine (I) are the respective members of the group. They are generally known as halogens. They are all non-metals and highly reactive. Their atoms each contain seven valence electrons and they are electron acceptors. The halogens exhibit great similarity in their properties example:

a) The entire are non-metals gases.

b) They exist as diatomic molecules example: I2 and Cl2.

c) They make ionic compounds whenever they react with metals however form covalent compounds by hydrogen.

2Na + Cl2 → 2 NaCl       Ionic

H2 + Br2 → 2 Hbr           Covalent

d) Their hydrides are soluble in water to form acids

HBr (g) + H2O → H3O+ (acidic) + Br

e) Their reactivity reduces down the group example: As fluorine reacts explosively by hydrogen even in the dark, chlorine reacts slowly in the diffuse light, bromine reacts slowly in bright light and iodine reacts partly even in bright light.

8) Group 0: Helium (He), Neon (Ne), Argon (Ar) is the well-known members of this group that are generally termed to as rare gases or noble gases. They have no bonding electrons due to the outermost shell is completely filled therefore the group name zero. The rare gases are thus unreactive and exist freely as monoatomic molecules in the atmosphere (around 1% of the composition of air).

Members of the group show identical properties that are different from those of the halogens which come before them and alkali metals that come after them. This is a confirmation which the end of a period has been reached.

The transition elements:

This is a collection of elements having very similar behavior and is generally positioned between the Groups II and III of the periodic table. All the transition elements have the following characteristics.

1) They are metals having high tensile strength and high melting point.

2) They represent variable valances.

3) They make colored ions.

4) They have the capability to make complex ions.

5) Some of them exhibit catalytic ability. Example: manganese and nickel.

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