Modern periodic law

Introduction:
Of the 109 elements known today, 92 are natural and the remaining are man made. As it was very difficult to study the chemistry of each and every element and its related compounds, in the 19th century  Chemists like Dobereiner, Newlands, Dinitri Mendeleef and Moseley thought of a proper classification of elements and did monumental work in analysing properties of these elements and classifying them .

Modern Periodic Law
" The physical and chemical properties of elements are periodic functions of their atomic numbers " .
It is a fact that atomic number is more related to the nucleus and it gives the number  of protons in the nucleus whereas  a few physical and many chemical properties of elements are related to their electronic configuration but not to the number  of electrons merely. Hence, periodic law states as

" The physical and chemical properties of elements are periodic functions of their electronic configurations "
If elements are arranged in the increasing order of their atomic numbers, the physical and chemical properties are repeated according to their electronic configurations at regular intervals.
For example: hydrogen (H0, lithium (Li), sodium (Na), potassium (K), rubidium  (Rb) are their electronic configurations as shown in table below. The atomic numbers (Z) of these elements are  1,3,11,19,37 respectively.

Alkali Metals and their electronic configurations

H ( Z = 1 )           1s1 Li ( Z = 3 )           1s2 , 2s1 Na ( Z = 11 )        1s2 , 2s2 , 2p6 , 3s1 K ( Z = 19 )          1s2 , 2s2 , 2p6 , 3s2 , 3p6 , 4s1 Rb ( Z = 37 )         1s2 , 2s2 , 2p6 , 3s2 , 3p6 , 3d10 , 4s2 , 4p6 , 5s1

Since these elements have similar electronic configurations in their  outer  shells, ns¹ ( 'n'   is the principle number ) , they show similar physical  and chemical properties. They lose electron to get stability  and show a common  oxidation state of +1.  They are all monovalent. Their oxides, hydroxides, halides, etc.., show similar chemical properties. Thus, the properties of elements are  periodic  functions of their electronic configurations.

Present form of peridic table

Present form of periodic table is also known as Long form of periodic table.
Neils Bohr constructed the long form of the periodic table based on the electronic configurations of the elements. In the table, the vertical columns are  termed as "groups" and the horizontal rows as "periods". There are 18 groups and 7 periods in the periodic table.  All the elements are arranged in the increasing order of their atomic numbers and the atomic number increases by one unit from one element to the immediate next element as we move, from left to right in the periodic table .

Differentiating electron :
The electron by which the electron configuration of the given element differs from that of its preceding element is called differentiating electron is the last entering electron of its atom .
Periods in the long form of the periodic table : In each period, the differentiating electron enters the s-orbital in the element p-orbital in the last element . When s and p orbitals are completely filled with electrons as ns²np⁶  , the element is the noble gas with stable configuration of octet .

Periods and electronic configurations

Period	Energy Levels	Number of elements in the period
1st period 2nd period 3rd period 4th period 5th period 6th period 7th period	1s 2s                                   2p 3s                                   3p 4s                         3d      4p 5s                         4d      5p 6s            4f          5d      6p 4s            5f          6d       -	2 8 8 18 18 32 -

                        Table : Periods and energy levels in periodic table

Period	First element	Electronic configuration	Last element	Electronic configuration
1 2 3 4 5 6 7	H Li Na K Rb Cs Fr	1s1 [He] 2s1 [Ne] 3s1 [Ar] 4s1 [Kr] 5s1 [Xe] 6s1 [Rn] 7s1	He Ne Ar Kr Xe Rn -	1s [He] 2s 2p [Ne] 3s 3p [Ar] 3d 4s 4p [Kr] 4d 5s2 5 [Xe] 4f 5d 6s 6p -

Each electron in the atom is distinguished by four quantum numbers . Among the four quantum numbers, the principle quantum number ( n 0 defines the main energy levels of the shell. In the periodic table, as the shells are filled with electrons in the order n=1, n=2 , ...... , the periods are formed serially .
In the 1st period as the 1st energy level is filled with two  electrons, there are only two elements. They are hydrogen and helium. In the second period, the differentiating electron enters the 2s orbital in lithium and the second energy level is completely filled 2s²2p⁶ with electrons in neon. Thus the second energy level is now completely filled with eight electrons  and hence the second period has eight elements . In the third period , the differentiating electron enters 3s orbital in sodium. After 3s, then 3p gets filled with differentiating electron upto argon. Then the differentiating electrons does not enter 3d orbitals but enters 4s. Hence the 3rd period also has only eight elements .
In the 4th period , the energy level 4s is 1st filled i.e. from potassium. After 4s is filled the differentiating electron enters 3d orbitals from scandium to zinc. and these ten elements are called transition elements. then the differentiating electron enters into the 4p orbitals upto krypton. Krypton attains stability due to completely filled orbitals. Thus, the 4th period has 18 elements .
In the fifth period, the energy  levels are 5s, 4d, 5p are filed in a similar order with differentiating electrons as in the 4th period. The period starts with rubidium and ends wiyth xenon. The fifth period has 18 elements which include second transition series.
In the sixth period, the energy levels 6s, 4f, 5d, 6p get filled with differentiating electrons in the same order. After 6s is filled, the differentiating  electrons enter the 4f orbitals to form the 1s inner transition series with an exception of La that gets the differentiating e^-1 into 5d. 14 elements from cerium to lutetium get their differentiating electrons into 4f orbitals. These elements are therefore called 'f' block elements. Then 5d get filled, to give the third transition series and next 6p orbitals get filled upto radon. The 6th period is the longest period with a total of 32 elements. The 7th period is incomplete and has about 20 elements. The period includes the second inner transition series from thorium  to Lawrencium.

Periodic table cesium

Introduction:

• Lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs) and francium (Fr) constitutes Group 1 of the periodic table.
• Robert Wilhelm Bunsen and Gustav Robert Kirchhoff found the metal cesium in the year 1860.
• Cesium has chemical symbol ‘Cs’.
• Periodic table cesium has atomic number 55 and mass number 132.9054
• Periodic table cesium belongs to Group 1 and Period 6
• Group 1 elements are called alkali metals because their oxides and hydroxides dissolve in water to produce strong alkali’s.
• Group1 elements are highly reactive and highly electropositive elements.
• Francium is a radioactive element.

• Electron configuration cesium:

     1S², 2S², 2P⁶, 3S², 3P⁶, 3d¹⁰, 4S², 4P⁶, 4d¹⁰, 5S²,5P⁶, 6S¹

• Electrons per energy level: 2, 8, 18, 18, 8, 1
• Shell model:

Properties of periodic table cesium:

The properties of alkali metals are similar because they belong to the same group in the periodic table and possess similar electronic configuration.  A few of these properties are discussed below.
1)      Size: Atomic radius increases as we move down the alkali metal group because an extra shell of electrons is added down the group.  Hence, lithium is the smallest atom and francium is the largest atom among alkali metals.  Francium is an unstable element.

Element	Li	Na	K	Rb	Cs
Atomic radius (pm)	134	154	196	211	235

2)      Electronic configuration: All alkali metals have a similar electronic configuration with only one electron in the outermost s-orbital.  The electronic configuration of alkali metals are given by:

• Li:  [He] 2s¹

• Na: [Ne] 3s¹

• K:  [Ar] 4s¹

• Rb: [Kr] 5s¹

• Cs: [Xe] 6s¹

• Fr: [Rn] 7s¹

General electronic configuration group 1 elements can be written as [Rare gas] ns¹.
3)      Density: Alkali metals are the light metals.  Their density is low because of larger atomic volumes.  However, the density goes on increasing down the group because an increase in mass predominates over increase in volume down the group.  Li, Na, K are lighter than water.  Density of sodium and potassium are 0.972 g/cm³ and 0.867 g/cm³ respectively.  The density of potassium is less than that of sodium because of  unusual increase in atomic size of potassium.  The density of cesium is 1.67 g/cm³.
 4)      Ionization energy: Alkali metals have low ionization energies because the last electron is present in the outermost s-orbital and the removal of electron is easy.  As we move down the group, there is decrease in the ionization energy.

Elements	Li	Na	K	Rb	Cs
Atomic radius (pm)	520	495	418	403	376

5)      Electropositive character: Due to low ionization energy, alkali metals are highly electropositive.  They have a tendency to lose electrons readily and change to M⁺ ion.
M → M⁺ + e^-
As ionization energy decreases down the group, electropositive character increases from lithium to cesium.  Metals like Cs and Rb lose electrons even they expose to light.  This property is known as photoelectric effect.
6)      Oxidation states: All alkali metals exhibit an oxidation state of +1 because they have only one electron in the outermost shell which is lost during the reactions.  +2 state is not observed because the second ionization of alkali metals is very high than the first ionization energy.
7)      Metallic properties: The metallic character of alkali metals increases from Li to Cs due to low ionization energy.  They show strong metallic character.  Hence , They show strong metallic character.  Hence they are highly electropositive and less electronegative.  As they have only one unpaired electron, metallic bond is weak.  Hence they are soft metals and can be easily deformed.
8)      Flame test: When cesium metal is heated strongly in a non-luminous flame, the salts are ionized.  Cesium ions are later excited and the electrons jumps to high energy levels.  When these electrons return back, they emit radiations which fall on the visible region.  Periodic test cesium emits pale violet color to the flame.
9)      Reducing property: Alkali metals are powerful reducing agents because they have very low reduction potentials.  This indicates that electrons can be released from them very easily.

Elements	Li	Na	K	Rb	Cs
R.P. (volts)	-3.05	-2.71	-2.93	-2.99	-2.99

Where R.P.  = Reducing potential.
Lower the R.P.  Greater the reducing potential.  Hence alkali metals react readily with water to liberate hydrogen.
2Cs + 2H₂O → 2CsOH + H₂↑
10)  Reaction of periodic table cesium with air: When alkali metals are exposed to air, they tarnish rapidly due to formation of oxides on the surface.  Hence they are kept under kerosene to protect from the action of air.
When Cesium metal are heated in air they burn vigorously to form cesium superoxide.
Cs + O₂  →  CsO₂
11)  Diagonal relationship between periodic table Cesium and radium: The diagonal element relationship is the similarity between the first elements of a group with the second element in the next higher group.

Uses of periodic table Cesium:

• Cs is used as a catalyst in the hydrogenation of few organic compounds.
• Cs metal can be used in ion propulsion systems.  Although not usable in the earth’s atmosphere.
• Periodic table cesium is used in atomic clocks.
• Cs has high electron affinity.  Therefore it is used in electron tubes.
•  Cs is used in photoelectric cells and vaccum tubes.
• Periodic table cesium is used in Infra red lamps. 

W on the periodic table

Introduction :
Let us discuss about the symbol ‘W’ on periodic table. Though Peter Woulfe examined the mineral wolframite and concluded the presence of a new substance, it was Juan Jose and Fausto d'Elhuyar from Spain who purified tungsten in 1783. In Swedish tungsten refers to “heavy stone”. Let us explore more on ‘W’ on periodic table.

Fig:- tungsten

Some Properties of 'W' on periodic table:

Tungsten, which is also called Wolfram and denoted by W on periodic table, has an atomic weight of 183.85 with an atomic number of 74 making it a transition metal on the periodic table. This gives it an electronic configuration of [Xe] 6s² 4f¹⁴ 5d⁴ .
Cutting of a pure tungsten can be done with a saw, spun, drawn, forged, and extruded though impure tungsten is very brittle. Tungsten is a lustrous metal grayish-white in color. Tungsten is known to have the highest melting point compared to all other metals and the lowest vapor pressure of the metals. Because of this property, traditional processes such as smelting cannot refine tungsten. It has very high tensile strength at temperatures exceeding 1650°C. Oxidation of tungsten takes place at increased temperatures. It has very low effect on itself when reacted with acids, as it is highly corrosion- resistant. Very few metallic acids have the strength to attack tungsten.
Physical properties

• Density(g/cc) – 19.3
• Melting Point (K): 3680
• Boiling Point (K): 5930
• Appearance: tough gray to white metal
• Atomic Radius (pm): 141
• Atomic Volume (cc/mol): 9.53
• Tensile strength 50000 – 75000 @1000 degree Celsius, psi
• Reflectivity is 62%

Uses of Tungsten:

Filaments of electrical lamps and picture tubes of television are made from tungsten. It is used in multiple high temperature applications like metal evaporation materials, target for X-rays etc. Some compounds of tungsten are also used in fluorescent lighting such as magnesium tungstenates. Some of the tungsten compounds are also used in mixture of paints. Some lubricants used at high temperatures are also made of tungsten compounds.

Conclusion for tungsten that is 'W' on periodic table:

Therefore, from the above discussion, we can conclude that tungsten is one of strategic and indispensible metals and it is represented by the symbol ‘W’ on periodic table.

Breaking atomic bonds

Solids are characterized by incompressibility, rigidity and mechanical strength. This represent the molecules, atoms or ions that make up a solids which are closely packed. They are join together by strong cohesive forces and cannot move at random. Thus, in solids we have well ordered molecular atomic or ionic arrangements. Thus, it is extremely hard to break atomic bonds between these molecules.

Some solids like sodium chloride NaCl₂, sulphur S, and sugar (carbohydrates),  besides being incompressible and rigid, have also characteristic geometrical forms. Such substances are known be a crystalline solids. The X-ray crystallography studies reveal that their ultimate particles such as molecules, atoms or ions are arranged in unusual pattern throughout the entire three-dimensional (3D) network of crystal. This definite and ordered arrangement of molecules, atoms or ions lengthens over a large distance making it extra difficult in breaking atomic bonds.

The natural history of the Inter-atomic Force resulting in the breaking atomic bonds

Basically, an atom consists of a tiny positively charged body, located at its center called as nucleus. The nucleus, though small have all the protons and neutrons. Since the mass of an atom entirely owing to the presence of protons and neutrons, it is evident that almost the entire mass of an atom resides in the nucleus.

Between the atoms or ions or molecules the inter-atomic bonds is present. This type of breaking atomic bonds is set up by equilibrium between attractive and repulsive forces with the remaining force being zero (0). When the breaking atomic bonds is at stabile. It is evident that the atoms are far apart from the attractive forces between these molecules so it will govern and when they are very close packed together the repulsive force will becomes higher; both these help in ruining away as the separation increases. This report proves that the breaking atomic bonds result from inter-atomic force, as a function of atom separation. Fig 1: Representation of bonds between two molecules

Breaking atomic bonds is also characterized by bond energies

Enthalpy formation of the bond.

Bond energy for any particular type of bond in a compound may be defined as the average amount of energy required to dissociate (to break) one mole, viz., Avogadro’s number of bonds of that type present in the compound. Bond energy is also called the enthalpy of formation of the bond.

Sub-atomic particles

Introduction

Subatomic particles when considered in physics and chemistry refers to  the smaller particles which makes up the  nucleons and atoms.  Two types of subatomic particles are present in nature. The first one is the elementary particles, which are not composed of other particles, and composite particles. Particle physics and nuclear physics make a study on these particles their interactions. This term describes the behavior of matter and energy at the molecular scales of quantum mechanics. According to uncertainty principle it has been concluded that analyzing of particles at different scales will require a statistical approach.

Types of subatomic particles

Elementary particles

Elementary particles present in the Standard Model include

• Six "flavors" of quarks: up, down, bottom, top, strange, and charm;
• Six types of leptons are present namely electron, electron neutrino,  tau, tau   neutrino, muon, muon neutrino,;
• Twelve gauge bosons (force carriers) are present namely the three W and Z bosons of the weak force, the photon of electromagnetism, and the eight gluons of the strong force.

Composite particles

The bounded states of two or more elementary particles form a composite particle. For instance, two up quarks and one down quark constitute a proton, while two neutrons and two protons make up the atomic nucleus of helium-4. The composite particles consist of all hadrons, a group composed of baryons (e.g., protons and neutrons) and mesons (e.g., pions and kaons). Hundreds of subatomic particles are known till date. The cosmic rays interacting with matter produces the majority of sub atomic particles or they are produced in particle accelerators by scattering processes.

Energy of subatomic particles

According to Einstein’s hypothesis, matter and energy are analogous. Matter can be expressed in terms of energy and vice-versa is also possible. Energy can be transferred by only two types of mechanisms which are known as waves and particles. Light can be expressed both as particles and waves. This type of paradox is termed as Wave-particle Duality Paradox. It has been established that all particles also have an associated wave nature. This is true for both elementary and compound particles. Few laws have been derived which explain how particles collide and interact.

Atomic radii

Atomic  radii  may be defined as  the distance between the nucleus and the outermost  electronic level of the atom. Since electrons are considered as the negatively charged electronic cloud there is no well defined boundary  for an atom.The diffused  nature of the electron cloud  makes it difficult  to give exact definition of  atomic size or atomic radii.

Introduction: Atomic radii

Thus  the atomic radii is an arbitrary  concept and is influenced by the nature of neighbouring atoms.

Types of atomic radii

As  there  is   no exact definition for the atomic radius, a number of radii have been defined for an atom. They are  Covalent radius, Crystal radius (otherwise called as metallic radius)  Vander Waal radius (otherwise called Collision radius). Let us learn one by one.

Covalent radius

 Covalent  Radius:
       Covalent radius  is used to measure the  atomic radii of  non- metals. The atomic  radius of  a non- metal is calculated from the  covalent bond length. In case of  homonuclear diatomic molecules ( type AA) , like F₂, Cl₂,Br₂ ....etc half of the covalent bond length is taken as atomic radius. For example the value of  Cl - Cl  bond idstance is 1.98 A^o  half of the distance 0.99 A^o is taken as the  atomic radius of  chlorine
       Another example: measuring the atomic radius of  carbon in diamond. The value of  C- C bond distance in the diamond is 1.54 A^o half of the  distance 0.77 A^o  is considered as the  atomic  radius of carbon atom.

Heteronuclear diatomic  molecule:
      In the calse of heteronuclear diatomic  molecule of  AB type (example CCl_{4 ,} SiC ..etc) bond length  distance d_(A-B) is given by
                   d _{(A -B)}    =  r_(A)  + r_(B)
      r_(A)   and  r_(B) are the  covalent radii of  A and  B  respectively.
     Example:    The experimental value of   d_(C-Cl)  in CCl₄  molecule is  1.76 A^o
                  d _(C-Cl)  =  r _(C) + r_(Cl)
                         r_(C) =  d_(C-Cl) - r_(Cl)
                         r_(C)       =  1.76 A^o -  r_(cl)
                         if  r_(cl)  is  given, then the covalent radius of carbon atom  can be calculated by subtracting the  covalent radius of  Cl from the  d_(C-Cl) bond length.The covalent radius of Cl atom can also be obtained, provided that covalent radius of C atom is known
Crystal Radius:
      It is otherwise called as  Atomic or Metallic radius, and defined as  one half of  the distance between the nuclei of two adjacent metal atoms in the metallic close-packed crystal lattice. For example  the internuclear distance between  two adjacent Na atoms in a crystal of sodium  metal is  3.80 A^oand hence the atomic radius of a Na metal is  half of the  distance, that is  3.80 A^o/ 2                  =  1.90 A^o
      since there  are weak  bonding forces between the metal atoms, the metallic radii are higher than the  single bond covalent  radii and at the same time  the metallic radii are smaller than  the vander Waal radii since the  bonding forces in the metallic crystal  lattice  are much staonger than the vaner waals forces
Vander Waal Radius:
      The name is  derived from theVander  Waal forces which is  found in noble gases.This  type of atomic radii is other wise  called Collision Radius. Tthe distance  between the two non-bonded  isolated  atoms  or the distacnce between  two non-bonded  atoms belonging to two adjacent molecules of an element  in the solid state is called Vander Waals distance  while half of  this  is called vader Waals Radius.
 Example :  The vander Waals distance of Cl₂ molecule =   3.6 A^o  half of  this value is  1.8 A^o  and  1.8 A ^o  is the Vander Waal radius of chlorine  atom.
                   It is to be noted  that the vander Waal radius  of an element  is higher than its covalent radius. Example the measured Vander Waal radius of chlorine is 1.8 A^o  and the  covalent radius  is  0.99 A^o
                  The variation in the atomic radii can be explained as follows.
                  When two chlorine  atoms are  just in contact with each other and  there is no bond between them,  now the distance between nuclei of those two chlorine atoms is called  the vander Waals  distance (3.6 A^o) and  half of it ( 1.8 A^o) is called  vander Waals radius.
                 where as when the electron clouds of the two chlorine atoms merge with each other to form chlorine molecule by forming covalent bond between them, the distance (covalent bond length) between them further decreases and  the distance become 1.8 A^o and half of it  0.99 A^o is the covalent radius.
                 Thus while describing  the atomic radii of various atoms, any of the radii described above can be used.

Dalton's Atomic Theory

When scientists started exploring matter, they realised that matter can be divided into smaller and still smaller particles. What was the ultimate particle like? They discovered that the smallest particle of an element that maintains its chemical identity through all chemical and physical changes is called and 'atom'.

John Dalton (1766 - 1844) can rightly be called the father of the Modern Theory on Atoms. He proposed his Atomic Theory in 1808, i.e., almost 200 years back. He did not have the help of sophisticated instruments that are available today to the scientists. Hence, many of his proposals, have been modified and updated. Over the years, substantial changes have taken place regarding the atomic theory, yet some of the assumptions that Dalton made are still held valid.

Dalton's Atomic Theory

John Daltons Atomic Theory provided a simple theory of matter to provide theoretical justification to the laws of chemical combinations in 1805. The basic postulates of the theory are:

• All substances are made up of tiny, indivisible particles called atoms.
• Atoms of the same element are identical in shape, size, mass and other properties.
• Each element is composed of its own kind of atoms. Atoms of different elements are different in all respects.
• Atom is the smallest unit that takes part in chemical combinations.
• Atoms combine with each other in simple whole number ratios to form compound atoms called molecules.
• Atoms cannot be created, divided or destroyed during any chemical or physical change.