Report of Brief Discussion of Periodic Table Elements.pdf

Brief Discussion of Periodic Table Elements

1. Introduction

 We know that 118 elements have been discovered so for. A table in which similar elements are placed in the same group is called a periodic table. A periodic table is the most significant achievement in the field of chemistry. It serves the students to know the properties of elements in a systematic way. It is the basic framework to study the periodic variations in the physical and chemical behavior of elements or their compounds. The modern periodic table is used to organize all the known elements. Elements are arranged in the table by increasing the atomic number. In the modern periodic table, each element is represented by its chemical symbol (Levi and Rosenthal, 1984; Kitaev, 2009). Columns of the periodic table are called groups. Elements in the same group have similar properties. For chemistry students and teachers: The tabular chart on the right is arranged by an Atomic number. The first chemical element is Hydrogen and the last is Ununoctium. Please note that the elements do not show their natural relation towards each other as in the Periodic system. To summarize, the periodic table is important because it is organized to provide a great deal of information about elements and how they relate to one another in one easy-to-use reference. The table can be used to predict the properties of elements, even those that have not yet been discovered. The periodic table got its name from the way the elements are arranged in rows which are called periods. The columns of the table are called groups, some of which have specific names, such as the noble gases and the halogens. ... Going down the periodic table, the number of atomic orbitals increases by one for each row(Flanigen et al., 1986). An organized table of chemical elements such as oxygen and carbon is an example of the periodic table. Dmitri Mendeleev published the first periodic table in 1869. ... Based on the work of physicist Henry Moseley, the periodic table was reorganized based on increasing atomic number rather than on atomic weight. The revised table could be used to predict the properties of elements that had yet to be discovered. The periodic table got its name from the way the elements are arranged in rows which are called periods. The columns of the table are called groups, some of which have specific names, such as the noble gases and the halogens.  Going down the periodic table, the number of atomic orbitals increases by one for each row(Dal Corso, 2014).

2. Historical background

 Many chemists have been working for the periodic classification of elements. In this field the names of AL-RAZI, DOBERENIER, NEWLAND, & MENDELEEV, (MENDELEJEFF) are very important. (Brooks, 2015).

2.1 AL-RAZI

 AL-RZI was a great Muslim chemist. He classified the elements on the bases of their physical and chemical properties. His work proved as a basic tool for his followers(Bonifacio, 2012).

2.2 DOBERENIER

  DOBERENIER was a German chemist. In 1829 he arranged the similar elements into the group of three. These groups are called Tirades. 

            e.g          Li—Na---K                             Cl---Br---I

In a tirade the atomic mass of the central elements is the arithmetical mean of the other two.

It is called LAW OF TRIADS.

e.g              In                  Li---Na---K            Na=7+39/2=23amu.

2.3 NEWLAND:

 Newland was an English chemist. In 1864 Newland arranged the elements in the increasing order of their atomic masses. He gave a law called “LAW OF OCTAVES”(Cowling et al., 1996)This law states “If elements are arranged in the increasing order of their atomic masses then properties of every eighth element are similar to the first one”  eg Li is similar to Na, Be is similar to Mg and Boron is similar to Aluminum. The law of octaves is not valid for all elements.

2.4 Work of Henry Moseley

In 1913 Henry Moseley developed the concept of atomic number; Moseley determined the frequencies of X-rays emitted as different elements were bombarded with high-energy electrons. He found that each element produces X rays of a unique frequency; further he found that the frequency generally increased as the atomic number increased. He arranged the X-ray frequency in order by assigning a unique whole atomic number, called an atomic number; to each element. Moseley identified the atomic number with the number of protons in the nucleus of the atom and the number of electrons in the atom(Mendeleev and Moseley, 2014). The concept of atomic number clarified some problems in the early model of the periodic table which was based on atomic masses (Moseley, 1913).

2.5 MENDELEEV

Mendeleev was a Russian chemist. He classified the elements regularly. He presented the first regular periodic table of elements. In Mendeleev’s periodic table there were eight groups and twelve periods. The vertical column of elements is called groups(Kibler, 2007). The horizontal rows of elements are called periods(Bensaude-Vincent, 1986). Mendeleve periodic table is defined as ‘If elements are arranged in ascending (increasing) order of their atomic masses. Their chemical properties repeat periodically. Mendeleev’s table provides a base for modern classification of elements” Although it has some confusion. The work of Moseley proved beyond doubt that the properties of elements are well explained and most of the anomalies and defects of Mendeleev’s periodic table disappear if the basis of the classification is changed from atomic masses to atomic numbers. This formulates the modern periodic law. Which is states, the physical and chemical properties of elements are the periodic function of their atomic numbers. Most people think Mendeleev invented the modern periodic table. Dmitri Mendeleev presented his periodic table of the elements based on increasing atomic weight on March 6, 1869, in a presentation to the Russian Chemical Society. The major difference is that the elements in Mendeleev's periodic table were arranged by atomic mass and the modern periodic table arranges elements by atomic number (Bensaude-Vincent, 1986).

Defects of Mendeleev’s periodic table:

Mendeleev’s table does not explain the structure of atoms. In this table position of K, Ar, and Ni, Co is against the periodic law. In Mendeleev's table many positions were vacant. These positions for unknown elements .eg Germanium (Ge), Gallium (Ga).Zn, Cd, Hg, and alkaline earth metals have very different properties .they are placed in the same group. It is against the periodic law.

In this table the position of Lanthanides and Actinides is against the periodic law(Beaumont et al., 2008; Zvyagintseva and Shalimov, 2014).

Improvement in Mendeleev periodic table:

 In 1913, Mosley discovered the atomic number. After the discovery of the atomic number the periodic table was improved. This improvement is based upon MODERN PERIODIC LAW. This law states(Blanie, 1981; Fernandez et al., 2007).

“If the elements are arranged in ascending order of their atomic numbers, their chemical properties repeat periodically” By modern periodic law confusions (defects) in Mendeleev table are removed.  All the isotopes of an element have the same atomic number. So they have only one position in the periodic table. The position of misfit pairs (K, Ar, and Ni, Co ) has been corrected. An extra group (group 8) for noble gases has been introduced. The position of Zn, Cd, Hg, and alkaline earth metals has been corrected by introducing two types of subgroups A and B. The alkaline earth metals are placed in group IIA and Zn, Cd, Hg are placed in group IIB (Bensaude-Vincent, 1986).

 

Figure 2 Different parts of the periodic table

3. THE MODERN PERIODIC TABLE

 The modern periodic table is based upon the modern periodic law. This law states, “ If elements are arranged in ascending order of their atomic numbers, their chemical properties repeat periodically” in the modern periodic table all elements are arranged in ascending order of their atomic numbers .some essential features of the modern periodic table are given below(Seaborg, 1996; Scerri, 1997).

3.1 Groups

 The vertical columns of elements in the periodic table are called groups. There are eight groups. They are represented by Roman numbers I to VIII. Each group is divided into two subgroups A and B. The subgroups A contain typical or normal elements. The subgroup ‘(B)’ Contains transition elements. The elements of a group have similar properties (Hume-Rothery, 1930).

North American labels the group with Roman numerical which have A and B designations. Europeans use a similar convention that numbers the columns from 1A to 8A and then from IB through 8B, thereby given the label 7B (or VIIB) instead of 7A to the group headed by fluorine F, and thus some columns have A and B interchanged, To eliminate this confusion, the international union of pure and applied chemistry (IUPAC) suggested a new convention that numbers the group from 1 through North American and the IUPAC conventions. When we refer to an element by its periodic group, we will use the traditional North American convention. Elements in a group have similar chemical properties. The properties of elements in one group are differing from the properties of elements in other groups. Elements within periods have that change progressively across the table. A period contains elements to increasing atomic numbers. A new period is begun to place elements of similar properties in vertical columns. An element can be located in the periodic table by listing and group numbers(Ellwood and Partin, 2016).

3.2 Periods

 The horizontal rows of elements in the periodic table are called periods. There are seven periods. They are represented by Arabic numbers 1 to 7(Jensen, 1982).

3.2.1 FIRST PERIOD

 The first period contains two elements. They are Hydrogen and Helium.

3.2.2 SECOND AND THIRD PERIOD

 The second and third period contains eight elements each. They are called short periods. The elements of these periods are normal (representative) elements and belong to A-sub groups.

 In these periods the properties of every eighth element are similar to the first element. For example, Lithium of 2nd period is similar to the sodium of the third period is similar to chlorine of the third period(Suhail and Ali, 2017).

3.2.3 4TH & 5TH PERIODS

 The 4th and 5th periods contain eight elements each. There are called elements, and ten transition elements in these 18 elements.In these periodic properties repeat after ‘(18)’ elements. For example K19 is similar to Rb 37.(Berry, 2006).

3.2.4 6TH PERIOD

 The 6th period contains 32 elements. It is the first very long period. It contains 8 normal elements10 transition elements and 14 inner transition elements. A new set of fourteen elements that start after La57 are called Lanthanides. All lanthanides have similar properties. They are placed at the bottom of the periodic table. The lanthanides are also called inner transition elements.

3.2.5 7TH PERIOD: 

THE 7TH period is the last in the periodic table. It is incomplete so for. It is the second very large period in the table. It contains two normal elements (F87 and Ra88). 10 Transition and,( 14 inner transition elements). A set of fourteen elements that follow AC89 are called Actinides (Berry, 2006).

4. Blocks in the periodic table

 All the elements in the periodic table are divided into four blocks (S BLOCK, P BLOCK, D BLOCK, AND F BLOCK) (Palmacci and Seeberger, 2004).

4.1 S-block: In the atoms of these elements, the differentiating (last) electron enters the ns orbitals of the outermost shell which is being progressively filled. Consequently, the valence shell electronic configuration of these elements varies from ns1 to ns2. The elements of IA and IIA are called S-block elements because their valence electrons are present in S- orbital. The elements of group IA, IIA H and He belong to this block since the valence shell configuration of group IA and H is ns1 and group IIA and He is ns2. The s-block elements lie on the extreme left of the periodic table and consist of active metals (except H and He). The properties of s-block elements depend on the number of electrons present in ns orbital (Scerri et al., 2018).

4.2 P-block: In the storms of these elements, the last electron enters in p orbital of the outermost shell; hence they are called p-block elements. In the atoms of these elements the ns orbital is filled and hence the valence shell configuration of these elements varies from ns2 p1 to ns2 p6. The elements of group IIIA to VIIA are called p-block elements because their valence electrons are present in p-orbitals. These elements lie at the extreme right of the periodic table and consist of some metals. Non-metals, metalloids, and inert gases (Ochs and Yin, 2001).The s- and p- block elements (i.e. group elements) in the periodic table are called the representative or main group elements. The elements generally show distinct and fairly regular variations in their properties with the atomic number  (Chen and Rowell, 2007).

4.3 D-block: In these elements, either in their atomic state or in any of their oxidation state the last electron enters the (n-1) d orbital of the penultimate shell. In other words in these elements (n-1) d orbital is being progressively filled. Hence these elements are called b-block elements. The electronic configuration for the two outer shells of these elements can be represented by (n-1) d1-(10 ns1-2). With the expectation of Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, Pt, and in the atoms of these elements the ns-orbitals are filled (ns2 configuration). These elements are placed in the middle of the periodic table between s-and p-block elements (Imyanitov, 2014).  The transition elements are called d-block elements because their valence electrons are present in d-orbital. The transition elements are classified into four series corresponding to the filling of 3d, 4d, 5d, and 6d, orbitals of the (n-1)shell of these atoms as follows (Khazan, 2011).

 1st Transition series (3rd series).  21Se through 29Cu. In the atoms of these elements the last electron enters the 3rd orbitals.

 2nd Transition series (3rd series).39Y through 47 Ag.

 3rd Transition series (5d series). 57La and 72Hf through 79Au.

 4th Transition series (6d series). (Incomplete)  89 Ac and elements 104 through 112. These are the elements of the 7th period (incomplete period). Strictly speaking the group IIB elements, Zn, Cd, and mercury are not d-transition metals since their “last” electron goes into s-block: but they are usually discussed with the d-transition metals because their chemical properties are similar.   

4.4 F-block: In these elements, either in their atomic state or in any of their common oxidation state, the last electron enters into the (n-2) f orbitals of the (n-2) shell. In other words, in these elements (n-2) f orbitals are being progressively filled. Hence these elements (n-2) f-orbitals may be either 4f orbitals or 5f orbitals. The inner transition elements are located between-group IIIB and IVB in the periodic table and are usually show in two separate rows below the main part of the body. The f-subshell has a maximum population of 14 electrons, so there are 14 Lanthanides and 14 Actinides,  The Lanthanides, and Actinides (inner transition elements) are called f-block elements because their valence electrons are present in f-block. The f-block elements are of two types. (Palmacci and Seeberger, 2004; Scerri et al., 2018)4f-series or inner Transition series (Lanthanides): This series has 14 elements 58Ce through 71Lu.5f-series or 2nd inner Transition series (Actinides): 90 the through 103Lr.All of the actinides are radioactive and form neptunium (elements 93) on, are synthetic. They have been produced in nuclear reactors or by using particle accelerators.

5. Classification of Elements based on their complete and incomplete electron shell (BOHR’s classification)

 There are four types of elements based on the presence of the complete and incomplete electron shell in their atoms. (Cartledge, 1928)

5.1 Noble Gases

 In the atoms of these elements the outermost shell is filled. The atoms of these elements have ns2p6 configuration except He which has 1s2 configuration.ns2p6 configuration is stable and Hence these elements generally do not take part into ordinary chemical reactions (Ogilvie and Wang, 1992). These elements are present in group VIIA of the periodic table (Nicklass et al., 1995).

5.2 Representative or Normal elements

In most of these elements the outermost shell is partially filled while the inner shells are filled. The electronic configuration of the outermost shell of the atoms of these elements varies from 1s1 to ns2p6. The elements of group IA, IIA, IIIA, IVA, VA, and VIIA belong to this type. The configuration of the outermost shell of these elements is bellowed(Spayd et al., 1978).

Group IA (Li-Fr) ---ns1                                 ;    Group VA (N-Bi)—ns2p3

Group IIA (Be-Ra)—ns2                               ;    Group VIA (O-Po)—ns2p4

Group IIIA (B-TI) ---ns2p1                           ;    Group VIIA (F-At) ---ns2p5

Group IVA(C-Pb) ---ns2p2

The configuration given above of these elements shows that s and p block elements except noble gases constitute representative or normal elements. The group containing these elements is also known as the main group of the periodic table. These elements tend to acquire ns2 np6 configuration in the outermost shell by the loss or gain of electrons  (Xia et al., 2003).

5.3 Transition Elements

 In the atoms if these elements two outermost shells nth and (n-1) Th are partly filled while the remaining inner shells are filled. The last two shells of these elements namely outermost and penultimate shells are incomplete.

The last shell contains one or two electrons and the penultimate shell may contain more than eight up to eighteen electrons. Their outermost electronic configuration is similar to d-block elements i.e.  (n-1) d1-10 ns1-2.According to the latest definition of transition elements those elements which have partly filled d-orbitals in a neutral state or any stable oxidation state are called transition elements.    According to this definition Zn, Cd, and Hg (IIB group) are d-block elements but not transition elements because these elements have d10 configuration in neutral as well as instable +2 oxidation state  (Dolg et al., 1987; Kresse and Hafner, 1994).

5.4 Inner transition elements

 (a)  In these elements last three shells i.e. last, penultimate, and pre-penultimate shells are incomplete.

(b) These are related to IIIB i.e. group 3.

(c)  The last shell contains two electrons. Penultimate shell may contain eight or nine electrons and pre-penultimate shell contains more than 18 upto32 electrons.

(d)  Their outermost electronic configuration is similar to ¦-block elements i.e. (n-2)1-14  (n-1)s2 (n-1)p6 (n-1)d0-2ns2

(e) Elements of the seventh period after atomic number 92 (i.e. actinides) are synthetic elements and are called transuranic elements. (Fidelis and Mioduski, 1981; Krishnamurthy, 1996) . According to their shared physical and chemical properties, the elements can be classified into the major categories of metals, metalloids, and nonmetal.  Metals are generally shiny, highly conducting solids that form alloys with one another and salt-like ionic compounds with nonmetals (other than noble gases). A majority of nonmetals are colored or colorless insulating gases; nonmetals that form compounds with other nonmetals feature covalent bonding. In between metals and nonmetals are metalloids, which have intermediate or mixed properties (Ćirić-Marjanović, 2013). 

6. Metals nonmetals and metalloids

 Metal and nonmetals can be further classified into subcategories that show graduation from metallic to non-metallic properties when going left to right in the rows. The metals may be subdivided into the highly reactive alkali metals, through the less reactive alkaline earth metals, lanthanides, and actinides, via the archetypal transition metals, and ending in the physically and chemically weak post-transition metals. Nonmetals may be simply subdivided into the polyatomic nonmetals, being nearer to the metalloids and show some incipient metallic character; the essentially nonmetallic diatomic nonmetals, nonmetallic and the almost completely inert, monatomic noble gases. Specialized groupings such as refractory metals and noble metals are examples of subsets of transition metals, also known and occasionally denoted. (Appenroth, 2010).Placing elements into categories and subcategories based just on shared properties is imperfect. There is a large disparity of properties within each category with notable overlaps at the boundaries, as is the case with most classification schemes. Beryllium, for example, is classified as an alkaline earth metal although its amphoteric chemistry and tendency to mostly form covalent compounds are both attributes of chemically weak or post-transition metal. Radon is classified as a nonmetallic noble gas yet has some cationic chemistry that is characteristic of metals. Other classification schemes are possible such as the division of the elements into mineralogical occurrence categories, or crystalline structures. Categorizing the elements in this fashion dates back to at least 1869 when Hinrichs wrote that simple boundary lines could be placed on the periodic table to show elements having shared properties, such as metals, nonmetals, or gaseous elements (Appenroth, 2010).

6.1 Metals

 The elements which tend to lose electrons, and from positive, ions are called metals. They are good conductors of heat and electricity. They form basic oxides which give bases when dissolved in water. All the elements on the left-hand side in the center and at the bottom of the periodic table (De Boer et al., 1988; Silver, 1998).

6.2 Nov Metals

 The elements which tend to gain electrons and form negative ions are call Nonmetals is a poor conductor of heat and electricity. They form acidic oxides which give acid when dissolved in water. The elements on the top right corner of the periodic table are non-metals. All the gases in the periodic table are non-metals (Tansel et al., 2000).

6.3 Metalloids

 The elements which have properties of metals, as well as non-metals, are called Metalloids or semi-Metals. They form amphoteric oxides. The oxides which have both acidic and basic properties are called amphoteric oxides.in the periodic table lower members of group IIIA to VA are metalloids (Si, As, Te) (Tansel et al., 2000).

 

  

(Nonmetals, on the right side of the periodic table, are very different from metals. Their surface is dull and they don't conduct heat and electricity. As compared to metals, they have low density and will melt at low temperatures. Elements that have properties of both metals and nonmetals are called metalloids. The periodic table on the left separates elements into three groups: the metals (green in the table), nonmetals (orange), and metalloids (blue). Most elements are metals. They are usually shiny, very dense, and only melt at high temperatures. Their shape can be easily changed into thin wires or sheets without breaking).

7. Periodic trends and patterns

7.1 Electron configuration

The electron configuration of an atom is the representation of the arrangement of electrons distributed among the orbital shells and subshells. Commonly, the electron configuration is used to describe the orbitals of an atom in its ground state, but it can also be used to represent an atom that has ionized into a cation or anion by compensating with the loss of or gain of electrons in their subsequent orbitals. The electron configuration or organization of electrons orbiting neutral atoms shows a recurring pattern or periodicity. 

The electrons occupy a series of electron shells (numbered 1, 2, and so on). Each shell consists of one or more subshells (named s, p, d, f, and g) (Luder, 1943) As atomic number increases, electrons progressively fill these shells and subshells more or less according to the Made lung rule or energy ordering rule, as shown in the diagram. (Ma et al., 2002). The electron configuration for neon, for example, is 1s2 2s2 2p6. With an atomic number of ten, neon has two electrons in the first shell, and eight electrons in the second shell; there are two electrons in the s subshell and six in the p subshell.

 In periodic table terms, the first time an electron occupies a new shell corresponds to the start of each new period, these positions being occupied by hydrogen and the alkali metals. Since the properties of an element are mostly determined by its electron configuration, the properties of the elements likewise show recurring patterns or periodic behavior, some examples of which are shown in the diagrams below for atomic radii, ionization energy, and electron affinity. It is this periodicity of properties, manifestations of which were noticed well before the underlying theory were developed, that led to the establishment of the periodic law (the properties of the elements recur at varying intervals) and the formulation of the first periodic tables. (Elfeky et al., 2004; Ching et al., 2006).(Babor, 1944; Scerri, 1997).

Before assigning the electrons of an atom into orbitals, one must become familiar with the basic concepts of electron configurations. Every element on the Periodic Table consists of atoms, which are composed of protons, neutrons, and electrons. Electrons exhibit a negative charge and are found around the nucleus of the atom in electron orbitals, defined as the volume of space in which the electron can be found within a 95% probability. The four different types of orbitals (s,p,d, and f) have different shapes, and one orbital can hold a maximum of two electrons. The p, d, and f orbitals have different sublevels, thus can hold more electrons.

As stated, the electron configuration of each element is unique to its position on the periodic table. The energy level is determined by the period and the number of electrons is given by the atomic number of the element. Orbitals on different energy levels are similar to each other, but they occupy different areas in space. The 1s orbital and 2s orbital both have the characteristics of an s orbital (radial nodes, spherical volume probabilities, can only hold two electrons, etc.) but, as they are found in different energy levels, they occupy different spaces around the nucleus. Each orbital can be represented by specific blocks on the periodic table. The s-block in the region of the alkali metals including helium (Groups 1 & 2), the d-block are the transition metals (Groups 3 to 12), the p-block are the main group elements from Groups 13 to 18, and the f-block are the lanthanides and actinides series.                                                                                    

       

                    

 

7.2 Atomic radii

 Atomic radii vary in a predictable and explainable manner across the periodic table. For instance, the radii generally decrease along each period of the table, from the alkali metals to the noble gases; and increase down each group. The radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period. These trends of the atomic radii (and of various other chemical and physical properties of the elements) can be explained by the electron shell theory of the atom; they provided important evidence for the development and confirmation of the quantum theory. (Suresh and Koga, 2001; Cordero et al., 2008).The electrons in the 4f-subshell, which is progressively filled across the lanthanide expected and that are almost identical to the atomic radii of the elements immediately above them. ( series, are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out. The elements immediately following the lanthanides have atomic radii that are smaller than would be Putz et al., 2003)Hence hafnium has virtually the same atomic radius (and chemistry) as zirconium, and tantalum has an atomic radius similar to niobium, and so forth. This is known as the lanthanide contraction. The effect of the lanthanide contraction is noticeable up to platinum (element 78), after which it is masked by a relativistic effect known as the inert pair effect. The d-block contraction, which is a similar effect between the d-block and p-block, is less pronounced than the lanthanide contraction but arises from a similar cause. (Putz et al., 2003). 

 

7.3 Ionization energy

 The first ionization energy is the energy it takes to remove one electron from an atom, the second ionization energy is the energy it takes to remove a second electron from the atom, and so on. For a given atom, successive ionization energies increase with the degree of ionization. For magnesium as an example, the first ionization energy is 738 kJ/moll and the second is 1450 kJ/mol. Electrons in the closer orbitals experience greater forces of electrostatic attraction; thus, their removal requires increasingly more energy. Ionization energy becomes greater up and to the right of the periodic table. (Politzer et al., 2002; Politzer et al., 2010)Large jumps in the successive molar ionization energies occur when removing an electron from a noble gas (complete electron shell) configuration. For magnesium again, the first two molar ionization energies of magnesium given above correspond to removing the two 3s electrons, and the third ionization energy is a much larger 7730 kJ/moll, for the removal of a 2p electron from the very stable neon-like configuration of Mg2+. Similar jumps occur in the ionization energies of other third-row atoms. (Politzer et al., 1991; Chakraborty et al., 2010).

 

7.4 Electronegativity

  Electronegativity is the tendency of an atom to attract a shared pair of electrons. An atom's electronegativity is affected by both its atomic number and the distance between the valence electrons and the nucleus. The higher its electronegativity, the more an element attracts electrons. It was first proposed by Linus Pauling in 1932. In general, electronegativity increases on passing from left to right along a period, and decreases on descending a group. Hence, fluorine is the most electronegative of the elements, while cesium is the least, at least of those elements for which substantial data is available(Little Jr and Jones, 1960; Iczkowski and Margrave, 1961). There are some exceptions to this general rule. Gallium and germanium have higher electronegativity than aluminum and silicon respectively because of the d-block contraction. Elements of the fourth period immediately after the first row of the transition metals have unusually small atomic radii because the 3d-electrons are not effective at shielding the increased nuclear charge, and smaller atomic size correlates with higher electronegativity. The anomalously high electronegativity of lead, particularly when compared to thallium and bismuth is an artifact of electronegativity varying with oxidation state: its electronegativity conforms better to trends if it is quoted for the +2 state instead of the +4 state (Skinner and Pritchard, 1953; Scerri et al., 2018).

 

7.5 Electron affinity

The electron affinity of an atom is the amount of energy released when an electron is added to a neutral atom to form a negative ion. Although electron affinity varies greatly, some patterns emerge. Generally, nonmetals have more positive electron affinity values than metals.  The electron affinities of the noble gases have not been measured conclusively, so they may or may not have slightly negative values. Electron affinity generally increases across a period. This is caused by the filling of the valence shell of the atom; a group 17 atom releases more energy than a group 1 atom on gaining an electron because it obtains a filled valence shell and is, therefore, more stable. (Zollweg, 1969).A trend of decreasing electron affinity going down groups would be expected. The additional electron will be entering an orbital farther away from the nucleus. As such this electron would be less attracted to the nucleus and would release less energy when added. In going down a group, around one-third of elements are anomalous, with heavier elements having higher electron affinities than their next lighter congeners. Largely, this is due to the poor shielding by d and f electrons. A uniform decrease in electron affinity only applies to group 1 atoms (Gutsev and Boldyrev, 1984)(Hotop et al., 1973).

 

7.6 Metallic character:

The lower the values of ionization energy, electronegativity, and electron affinity, the more metallic character the element has. Conversely, the nonmetallic character increases with higher values of these properties. Given the periodic trends of these three properties, the metallic character tends to decrease going across a period (or row) and, with some irregularities (mostly) due to poor screening of the nucleus by d and f electrons, and relativistic effects tend to increase going down a group (or column or family). (Edwards and Sienko, 1983) Thus, the most metallic elements (such as cesium and francium) are found at the bottom left of traditional periodic tables and the most nonmetallic elements (oxygen, fluorine, chlorine) at the top right. (Westlake and Ockers, 1970) The combination of horizontal and vertical trends in metallic character explains the stair-shaped dividing line between metals and nonmetals found on some periodic tables, and the practice of sometimes categorizing several elements adjacent to that line, or elements adjacent to those elements, as metalloids (Wan et al., 2005).

 

8. Linking or bridging groups

From left to right across the four blocks of the long- or 32-column form of the periodic table are a series of linking or bridging groups of elements, located approximately between each block. (Dai et al., 2006) These groups, like the metalloids, show properties in between, or that are a mixture of, groups to either side. (Beaumont et al., 2008) Chemically, the group 3 elements, scandium, yttrium, lanthanum, and actinium behave largely like the alkaline earth metals or, more generally, s block metals but have some of the physical properties of d block transition metals (de Zea Bermudez et al., 2001). Lutetium and lawrencium, at the end of the f block, may constitute another linking or bridging group. Lutetium behaves chemically as a lanthanide but shows a mix of lanthanide and transition metal physical properties. Lawrencium, as an analog of lutetium, would presumably display like characteristics. (Bolt and Tebbe, 1988; Olkhovyk and Jaroniec, 2005) The coinage metals in group 11 (copper, silver, and gold) are chemically capable of acting as either transition metals or main group metals. The volatile group 12 metals, zinc, cadmium, and mercury are sometimes regarded as linking the d block to the p block. Notionally they are d block elements but they have few transition metal properties and are more like their p block neighbors in group 13. (de Zea Bermudez et al., 2001; Gao et al., 2009) The relatively inert noble gases, in group 18, bridge the most reactive groups of elements in the periodic table—the halogens in group 17 and the alkali metals in group 1.(Castricum et al., 2008).

9. Placement of hydrogen and helium

Simply following electron configurations, hydrogen (electronic configuration 1s1) and helium (1s2) should be placed in groups 1 and 2, above lithium (1s22s1) and beryllium (1s22s2). While such a placement is common for hydrogen, it is rarely used for helium outside of the context of electron configurations: When the noble gases (then called "inert gases") were first discovered around 1900, they were known as "group 0", reflecting no chemical reactivity of these elements known at that point, and helium was placed on the top of that group, as it did share the extreme chemical inertness seen throughout the group. As the group changed its formal number, many authors continued to assign helium directly above neon, in group 18; one of the examples of such placing is the current IUPAC table. The position of hydrogen in group 1 is reasonably well settled. Its usual oxidation state is +1 as is the case for its heavier alkali metal congeners. Like lithium, it has significant covalent chemistry. It can stand in for alkali metals in typical alkali metal structures. It is capable of forming alloy-like hydrides, featuring metallic bonding, with some transition metals (Labarca and Srivaths, 2017) (Scerri, 2004).

Nevertheless, it is sometimes placed elsewhere. A common alternative is at the top of group 17 given hydrogen's strictly univalent and largely non-metallic chemistry, and the strictly univalent and non-metallic chemistry of fluorine (the element otherwise at the top of group 17). Sometimes, to show hydrogen has properties corresponding to both those of the alkali metals and the halogens, it is shown at the top of the two columns simultaneously. Another suggestion is above carbon in group 14: placed that way, it fits well into the trends of increasing ionization potential values and electron affinity values, and is not too far from the electronegativity trend, even though hydrogen cannot show the tetra valence characteristic of the heavier group 14 elements. Finally, hydrogen is sometimes placed separately from any group; this is based on its general properties being regarded as sufficiently different from those of the elements in any other group  (Scerri, 2010).

The other period 1 element, helium, is occasionally placed separately from any group as well. The property that distinguishes helium from the rest of the noble gases (even though the extraordinary inertness of helium is extremely close to that of neon and argon) is that in its closed electron shell, helium has only two electrons in the outermost electron orbital, while the rest of the noble gases have eight (Scerri, 2010).

10. Group 3 and its elements in periods 6 and 7

Although scandium and yttrium are always the first two elements in group 3, the identity of the next two elements is not completely settled. They are commonly lanthanum and actinium, and less often lutetium and lawrencium. The two variants originate from historical difficulties in placing the lanthanides in the periodic table, and arguments as to where the f block elements start and end. It has been claimed that such arguments are proof that, "it is a mistake to break the [periodic] system into sharply delimited blocks". A third variant shows the two positions below the yttrium as being occupied by the lanthanides and the actinides. A fourth variant shows group 3 bifurcating after Sc-Y, into a La-Ac branch, and a Lu-Lr branch. Chemical and physical arguments have been made in support of lutetium and lawrencium but the majority of authors seem unconvinced. Most working chemists are not aware there is any controversy. In December 2015 an IUPAC project was established to make a recommendation on the matter (Jensen, 1982; Guenther, 1987).

11. Lanthanum and actinium

Lanthanum and actinium are commonly depicted as the remaining group of 3 members. It has been suggested that this layout originated in the 1940s, with the appearance of periodic tables relying on the electron configurations of the elements and the notion of the differentiating electron. The configurations of caesium, barium and lanthanum are [Xe] 6s1, [Xe] 6s2 and [Xe] 5d16s2. Lanthanum thus has a 5d differentiating electron and this establishes it "in group 3 as the first member of the d-block for period 6". A consistent set of electron configurations is then seen in group 3: scandium [Ar] 3d14s2, yttrium [Kr] 4d15s2 and lanthanum [Xe] 5d16s2. Still in period 6, ytterbium was assigned an electron configuration of [Xe] 4f135d16s2 and lutetium [Xe] 4f145d16s2, "resulting in a 4f differentiating electron for lutetium and firmly establishing it as the last member of the f-block for period 6". Later spectroscopic work found that the electron configuration of ytterbium was in fact [Xe] 4f146s2. This meant that ytterbium and lutetium—the latter with [Xe]4f145d16s2—both had 14 f-electrons, "resulting in a d- rather than an f- differentiating electron" for lutetium and making it an "equally valid candidate" with [Xe]5d16s2 lanthanum, for the group 3 periodic table position below yttrium. Lanthanum has the advantage of incumbency since the 5d1 electron appears for the first time in its structure whereas it appears for the third time in lutetium, having also made a brief second appearance in gadolinium (Jensen, 1982; Scerri et al., 2018).

In terms of chemical behavior, and trends going down group 3 for properties such as melting point, electronegativity, and ionic radius, scandium, yttrium, lanthanum, and actinium are similar to their group 1–2 counterparts. (Lavelle, 2008) In this variant, the number of f electrons in the most common (trivalent) ions of the f-block elements consistently matches their position in the f-block. For example, the f-electron counts for the trivalent ions of the first three f-block elements are Ce 1, Pr 2, and Nd 3 (Scerri et al., 2018).

12. Lutetium and lawrencium

In other tables, lutetium and lawrencium are the remaining groups of 3 members. Early techniques for chemically separating scandium, yttrium, and lutetium relied on the fact that these elements occurred together in the so-called "yttrium group" whereas La and Ac occurred together in the "cerium group".(Jensen, 1982) Accordingly, lutetium rather than lanthanum was assigned to group 3 by some chemists in the 1920s and 30s. Several physicists in the 1950s and '60s favored lutetium, in light of a comparison of several of its physical properties with those of lanthanum. (Jensen, 2015),  This arrangement, in which lanthanum is the first member of the f-block, is disputed by some authors since lanthanum lacks any f-electrons. It has been argued that this is not a valid concern given other periodic table anomalies—thorium, for example, has no f-electrons yet is part of the f-block. As for lawrencium, its gas phase atomic electron configuration was confirmed in 2015 as [Rn] 5f147s27p1. Such a configuration represents another periodic table anomaly, regardless of whether lawrencium is located in the f-block or the d-block, as the only potentially applicable p-block position has been reserved for nihonium with its predicted configuration of [Rn] 5f146d107s27p1 (Jensen, 1982; Tansel et al., 2000).

Chemically, scandium, yttrium, and lutetium (and presumably lawrencium) behave like trivalent versions of the group 1–2 metals. On the other hand, trends going down the group for properties such as melting point, electronegativity, and ionic radius, are similar to those found among their group 4–8 counterparts (Jensen, 2015). In this variant, the number of f electrons in the gaseous forms of the f-block atoms usually matches their position in the f-block. For example, the f-electron counts for the first five f-block elements are La 0, Ce 1, Pr 3, Nd 4 and Pm 5.(Jensen, 1982)

13. Lanthanides and actinides

A few authors position all thirty lanthanides and actinides in the two positions below the yttrium (usually via footnote markers). This variant, which is stated in the 2005 Red Book to be the IUPAC-agreed version as of 2005 (several later versions exist, and the last update is from 1st Dec. 2018), emphasizes similarities in the chemistry of the 15 lanthanide elements (La–Lu), possibly at the expense of ambiguity as to which elements occupy the two group 3 positions below yttrium, and a 15-column wide f block (there can only be 14 elements in any row of the f block) (Spayd et al., 1978; Levi and Rosenthal, 1984) (Sanderson, 1964).

La-Ac and Lu-Lr

In this variant, group 3 bifurcates after Sc-Y into a La-Ac branch and a Lu-Lr branch. This arrangement is consistent with the hypothesis that arguments in favor of either Sc-Y-La-Ac or Sc-Y-Lu-Lr based on chemical and physical data are inconclusive. As noted, trends going down Sc-Y-La-Ac match trends in groups 1−2 whereas trends going down Sc-Y-Lu-Lr better match trends in groups 4−10. (Chakraborty et al., 2010).(Sanderson, 1964) .(Fenton et al., 1984)The definition of a transition metal, as given by IUPAC, is an element whose atom has an incomplete d sub-shell, or which can give rise to actions with an incomplete d sub-shell. By this definition all of the elements in groups 3–11 are transition metals. The IUPAC definition, therefore, excludes group 12, comprising zinc, cadmium, and mercury, from the transition metals category.

Some chemists treat the categories "d-block elements" and "transition metals" interchangeably, thereby including groups 3–12 among the transition metals. In this instance the group 12 elements are treated as a special case of transition metal in which the d electrons are not ordinarily involved in chemical bonding. The 2007 report of mercury (IV) fluoride (HgF4), a compound in which mercury would use its d electrons for bonding, has prompted some commentators to suggest that mercury can be regarded as a transition metal. Other commentators, such as Jensen, have argued that the formation of a compound like HgF4 can occur only under highly abnormal conditions; indeed, its existence is currently disputed. As such, mercury could not be regarded as a transition metal by any reasonable interpretation of the ordinary meaning of the term. (Zhao and Truhlar, 2006)Still other chemists further exclude the group 3 elements from the definition of a transition metal. They do so on the basis that the group 3 elements do not form any ions having a partially occupied d shell and do not, therefore, exhibit any properties characteristic of transition metal chemistry. In this case, only groups 4–11 are regarded as transition metals. (Zamir, 1965) Though the group 3 elements show few of the characteristic chemical properties of the transition metals, they do show some of their characteristic physical properties (on account of the presence in each atom of a single d electron).

14. Groups included in the transition metals 

 The definition of a transition metal, as given by IUPAC, is an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell. By this definition all of the elements in groups 3–11 are transition metals. The IUPAC definition, therefore, excludes group 12, comprising zinc, cadmium, and mercury, from the transition metals category.

Some chemists treat the categories "d-block elements" and "transition metals" interchangeably, thereby including groups 3–12 among the transition metals. In this instance the group 12 elements are treated as a special case of transition metal in which the d electrons are not ordinarily involved in chemical bonding. The 2007 report of mercury (IV) fluoride (HgF4), a compound in which mercury would use its d electrons for bonding, has prompted some commentators to suggest that mercury can be regarded as a transition metal. Other commentators, such as Jensen, have argued that the formation of a compound like HgF4 can occur only under highly abnormal conditions; indeed, its existence is currently disputed. As such, mercury could not be regarded as a transition metal by any reasonable interpretation of the ordinary meaning of the term.

Still other chemists further exclude the group 3 elements from the definition of a transition metal. They do so on the basis that the group 3 elements do not form any ions having a partially occupied d shell and do not, therefore, exhibit any properties characteristic of transition metal chemistry. In this case, only groups 4–11 are regarded as transition metals. Though the group 3 elements show few of the characteristic chemical properties of the transition metals, they do show some of their characteristic physical properties (on account of the presence in each atom of a single d electron). all elements up to oganesson have been discovered, of the elements above hassium (element 108), only copernicium (element 112), nihonium (element 113), and flerovium (element 114) have known chemical properties, and only for copernicium is there enough evidence for a conclusive categorization at present. The other elements may behave differently from what would be predicted by extrapolation, due to relativistic effects; for example, flerovium has been predicted to possibly exhibit some noble-gas-like properties, even though it is currently placed in the carbon group. The current experimental evidence still leaves open the question of whether flerovium behaves more like a metal or a noble gas. (Soddy, 1913) (Frieden, 1972)

15. Elements with unknown chemical properties

 Although all elements up to oganesson have been discovered, of the elements above hassium (element 108), only copernicium (element 112), nihonium (element 113), and flerovium (element 114) have known chemical properties, and only for copernicium is there enough evidence for a conclusive categorization at present. The other elements may behave differently from what would be predicted by extrapolation, due to relativistic effects; for example, flerovium has been predicted to possibly exhibit some noble-gas-like properties, even though it is currently placed in the carbon group. The current experimental evidence still leaves open the question of whether flerovium behaves more like a metal or a noble gas. (Soddy, 1913) (Frieden, 1972).

16. Further periodic table extensions

 It is unclear whether new elements will continue the pattern of the current periodic table as period 8, or require further adaptations or adjustments. Seaborg expected the eighth period to follow the previously established pattern exactly, so that it would include a two-element s-block for elements 119 and 120, a new g-block for the next 18 elements, and 30 additional elements continuing the current f-, d-, and p-blocks, culminating in element 168, the next noble gas. More recently, physicists such as Pekka Pyykkö have theorized that these additional elements do not follow the Made lung rule, which predicts how electron shells are filled and thus affects the appearance of the present periodic table. (Leroy et al., 1993)There are currently several competing theoretical models for the placement of the elements of an atomic number less than or equal to 172. In all of these it is element 172, rather than element 168, that emerges as the next noble gas after oganesson, although these must be regarded as speculative as no complete calculations have been done beyond element 123. (Rinker, 1988) 

17. Element with the highest possible atomic number

The number of possible elements is not known. A very early suggestion made by Elliot Adams in 1911, and based on the arrangement of elements in each horizontal periodic table row, was that elements of atomic weight greater than circa 256 (which would equate to between elements 99 and 100 in modern-day terms) did not exist. A higher—more recent—estimate is that the periodic table may end soon after the island of stability. (Aston, 1942) This is expected to center on element 126, as the extension of the periodic and nuclides tables is restricted by proton and neutron drip lines. Other predictions of an end to the periodic table include at element 128 by John Emsley.at element 137 by Richard Feynman. At element 146 by Yogendra Gambier, and element 155 by Albert Khazan (Lloyd, 1987).

17.1 Bohr model

The Bohr model exhibits difficulty for atoms with atomic numbers greater than 137, as any element with an atomic number greater than 137 would require 1s electrons to be traveling faster than c, the speed of light.  Hence the non-relativistic Bohr model is inaccurate when applied to such an element   (Lloyd, 1987) (Kibler, 2007).

17.2 Relativistic Dirac equation

The relativistic Dirac equation has problems for elements with more than 137 protons. For such elements, the wave function of the Dirac ground state is oscillatory rather than bound, and there is no gap between the positive and negative energy spectra, as in the Klein paradox. (Seaborg, 1996) More accurate calculations taking into account the effects of the finite size of the nucleus indicate that the binding energy first exceeds the limit for elements with more than 173 protons. For heavier elements, if the innermost orbital (1s) is not filled, the electric field of the nucleus will pull an electron out of the vacuum, resulting in the spontaneous emission of a positron.  This does not happen if the innermost orbital is filled, so that element 173 is not necessarily the end of the periodic table.

18. Optimal form

 The many different forms of the periodic table have prompted the question of whether there is an optimal or definitive form of the periodic table. The answer to this question is thought to depend on whether the chemical periodicity seen to occur among the elements has an underlying truth, effectively hard-wired into the universe, or if any such periodicity is instead the product of subjective human interpretation, contingent upon the circumstances, beliefs, and predilections of human observers (Scerri, 2009). An objective basis for chemical periodicity would settle the questions about the location of hydrogen and helium and the composition of group 3. Such an underlying truth, if it exists, is thought to have not yet been discovered. In its absence, the many different forms of the periodic table can be regarded as variations on the theme of chemical periodicity, each of which explores and emphasizes different aspects, properties, perspectives, and relationships of and among the elements (Glawe et al., 2016).

19. Second version and further development

 In 1871, Mendeleev published his periodic table in a new form, with groups of similar elements arranged in columns rather than in rows, and those columns numbered I to VIII corresponding with the element's oxidation state. He also gave detailed predictions for the properties of elements he had earlier noted were missing, but should exist. (Lang and Bradley, 2009) These gaps were subsequently filled as chemists discovered additional naturally occurring elements. It is often stated that the last naturally occurring element to be discovered was francium (referred to by Mendeleev as Eka-caesium) in 1939. Plutonium, produced synthetically in 1940, was identified in trace quantities as a naturally occurring element in 1971. The popular periodic table layout, also known as the common or standard form (as shown at various other points in this article), is attributable to Horace Groves Deming. In 1923, Deming, an American chemist, published short (Mendeleev style) and medium (18-column) form periodic tables. Merck and Company prepared a handout form of Deming's 18-column medium table, in 1928, which was widely circulated in American schools. (Lengler and Eppler, 2007) By the 1930s Deming's table was appearing in handbooks and encyclopedias of chemistry. It was also distributed for many years by the Sargent-Welch Scientific Company. With the development of modern quantum mechanical theories of electron configurations within atoms, it became apparent that each period (row) in the table corresponded to the filling of a quantum shell of electrons. Larger atoms have more electron sub-shells, so later tables have required progressively longer periods. Glenn T. Seaborg, in 1945, suggested a new periodic table showing the actinides as belonging to a second f-block series. In 1945, Glenn Seaborg, an American scientist, suggested that the actinide elements, like the lanthanides, were filling an f sub-level. Before this time the actinides were thought to be forming a fourth d-block row. Seaborg's colleagues advised him not to publish such a radical suggestion as it would most likely ruin his career. As Seaborg considered he did not then have a career to bring into disrepute, he published anyway. Seaborg's suggestion was found to be correct and he subsequently went on to win the 1951 Nobel Prize in chemistry for his work in synthesizing actinide elements. Although minute quantities of some transuranic elements occur naturally, they were all first discovered in laboratories. Their production has expanded the periodic table significantly, the first of these being neptunium, synthesized in 1939. Because many of the transuranic elements are highly unstable and decay quickly, they are challenging to detect and characterize when produced. There have been controversies concerning the acceptance of competing discovery claims for some elements, requiring independent review to determine which party has priority and hence naming rights. In 2010, a joint Russia–US collaboration at Dubna, Moscow Oblast, Russia, claimed to have synthesized six atoms of Tennessee (element 117), making it the most recently claimed discovery. It, along with nihonium (element 113), moscovium (element 115), and oganesson (element 118), are the four most recently named elements, whose names all became official on 28 November 2016 (Demircioğlu et al., 2009).

20. Different periodic tables

20.1 The long- or 32-column table

 The modern periodic table is sometimes expanded into its long or 32-column form by reinstating the footnoted f-block elements into their natural position between the s- and d-blocks, as proposed by Alfred Werner. Unlike the 18-column form this arrangement results in "no interruptions in the sequence of increasing atomic numbers”. The relationship of the f-block to the other blocks of the periodic table also becomes easier to see. Jensen advocates a form of a table with 32 columns because the lanthanides and actinides are otherwise relegated in the minds of students as dull, unimportant elements that can be quarantined and ignored. Despite these advantages the 32-column form is generally avoided by editors on account of its undue rectangular ratio compared to a book page ratio, and the familiarity of chemists with the modern form, as introduced by Seaborg  (Laing, 2008)  (Jensen, 2003).

 

20.2 Tables with different structures

 Within 100 years of the appearance of Mendeleev's table in 1869, Edward G. Mazurs had collected an estimated 700 different published versions of the periodic table. As well as numerous rectangular variations, other periodic table formats have been shaped, for example, like a circle, cube, cylinder, building, spiral, lemniscus octagonal prism, pyramid, sphere, or triangle. Such alternatives are often developed to highlight or emphasize chemical or physical properties of the elements that are not as apparent in traditional periodic tables. (Te Velde and Baerends, 1991) Theodor Benefits spiral periodic table A popular alternative structure is that of Otto Theodor Benfey (1960). The elements are arranged in a continuous spiral, with hydrogen at the center and the transition metals, lanthanides, and actinides occupying peninsulas. Most periodic tables are two-dimensional; three-dimensional tables are known to as far back as at least 1862 (pre-dating Mendeleev's two-dimensional table of 1869). More recent examples include Courtines' Periodic Classification (1925), Wrigley's Lamina System (1949), Giguère's Periodic helix (1965), and Dufour's Periodic Tree (1996). Going one further, Stowe's Physicist's Periodic Table (1989) has been described as being four-dimensional (having three spatial dimensions and one color dimension)  (Roduner, 2006).

 

The various forms of periodic tables can be thought of as lying on a chemistry–physics continuum. Towards the chemistry end of the continuum can be found, as an example,.-Canham's "unruly" Inorganic Chemist's Periodic Table (2002), which emphasizes trends and patterns, and unusual chemical relationships and properties. Near the physics end of the continuum is Janet's Left-Step Periodic Table (1928). This has a structure that shows a closer connection to the order of electron-shell filling and, by association, quantum mechanics. A somewhat similar approach has been taken by Alper, albeit criticized by Eric Sherri as disregarding the need to display chemical and physical periodicity. Somewhere in the middle of the continuum is the ubiquitous common or standard form of the periodic table. This is regarded as better expressing empirical trends in physical state, electrical and thermal conductivity, and oxidation numbers, and other properties easily inferred from traditional techniques of the chemical laboratory. Its popularity is thought to be a result of this layout having a good balance of features in terms of ease of construction and size, and its depiction of atomic order and periodic trends  (Moutevelis and Woolfson, 2009).

Conclusion: 

A periodic table is one of the most important and basic topics of chemistry. It is too much important for a chemist. We can say that without the help of a periodic table brief discussion of elements are very difficult in chemistry. And we also know that without a brief discussion of periodic table elements we cannot realize and simplify the phenomenon and chemical reaction in chemistry. so that’s why a brief discussion of the periodic table and the periodic table is very important in chemistry. We know that 118 elements have been discovered so for. A table in which similar elements are placed in the same group is called a periodic table. A periodic table is the most significant achievement in the field of chemistry. It serves the students to know the properties of elements in a systematic way. It is the basic framework to study the periodic variations in the physical and chemical behavior of elements or their compounds. The modern periodic table is used to organize all the known elements. The periodic table has different properties according to the elements present in it. Every element has different properties as well as different compositions.

The conclusion of this report is to collect all the data about the elements of the periodic table. By collecting this data we can study easily all the aspects of periodic table elements. As we know that periodic table elements have different chemical and different physical properties from each other . so it is difficult to study one by one. We can not briefly study the properties of all the elements without the help of a periodic table. The above report is a complete overview of all the aspects of the periodic table of elements. 

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