Atomic number 68

Introduction :
The element in the periodic table with atomic number 68 is Erbium.  Erbium is a element which belongs to lanthanide series in the periodic table with atomic number 68 and the symbol is Er.  Naturally Erbium is found along with other elements on earth.  It is a rare earth element which is associated with mineral gadolinite from ytterby.  It has optical fluorescent properties which is useful in laser applications like laser optical amplifier. Glasses or crystals which are doped with the element atomic number 68 (erbium) are mainly used in amplification media.

Silvery white Erbium:                                         Erbium chloride showing pink fluorescense
                                                                        under sunlight
            

Occurrence of element with atomic number 68:
In the earth crust erbium has 2.8mg/kg concentration where as in sea water it has 0.9ng/L.  This element found bound with monazite sand ores.  The sources of erbium are xenotime and euxenite.

Monazite sand:



Isotopes of element with atomic number 68:
Erbium which occurs naturally is composed of 6 stable isotopes, Er-162, Er-164, Er-166, Er-167, Er-168, and Er-170 with Er-166 being the most abundant isotope.

Properties of Element with atomic number 68:

Physical properties of Element with atomic number 68:

1. It appears as silvery white but under sun light it shows pink fluorescence.
2. Erbium has density 9.006g/cm³.
3. It decomposes at 1802K.
4. It boils at 3141K.
5. Heat of fusion is 19.90kJ/mol.
6. Heat of vaporization is 280kJ/mol.
7. Specific heat capacity at 25⁰C is 28.12J/mol/K.
8. Below 19K it is ferromagnetic, between 19-80K it is antiferromagnetic and above 80K it is paramagnetic in nature.  

Chemical Properties of Erbium with atomic number 68:
1. Erbium metal burns to form erbium (III) oxide.
    4 Er + 3 O₂ → 2 Er₂O₃
2. It reacts slowly with cold water and quickly with hot water and forms erbium oxide.
    2 Er (s) + 6 H₂O (l) → 2 Er (OH) ₃ (aq) + 3 H₂ (g)
3. Erbium reacts with all the halogens
    2 Er (s) + 3 F₂ (g) → 2 ErF₃ (s) [pink]
    2 Er (s) + 3 Cl₂ (g) → 2 ErCl₃ (s) [violet]
    2 Er (s) + 3 Br₂ (g) → 2 ErBr₃ (s) [violet]
    2 Er (s) + 3 I₂ (g) → 2 ErI₃ (s) [violet]
4. Erbium dissolves readily in dilute sulfuric acid forms yellow [Er (OH₂)₉]³⁺ hydration complexes
    2 Er (s) + 3 H₂SO₄ (aq) → 2 Er³⁺ (aq) + 3 SO2−4 (aq) + 3 H₂ (g)

 Applications of element with atomic number 68:

1. Oxide erbium which has pink color is used as a colorant for glass, porcelain and cubic zirconia. These glasses are used in sunglasses and some jewelry.
2. These are used in neutron –absorbing  control rods.
3. When erbium is doped  with optical silica glass fibers, they are used widely in optical communications.
4. These are also used to create fiber lasers.
5. These are also used in metal welding and cutting applications.
6. Erbium ions have large variety  of medical applications.

Atomic Number 4

Introduction :
Beryllium is the compound which as atomic number 4 and mass number 9.01218.  Beryllium has electronic configuration 1s², 2s².  Beryllium has oxidation state +2.  Group 2 elements are Be, Mg, Ca, Sr, Ba, Ra are alkaline earth elements.  Atomic number 4 is a exception : it does not react with water or steam, and its halides are co-valent bond with beryllium.  All the alkaline earth metal as two electrons in the outermost shell, so filled electron shell is achieved by the lose of two electrons to form doubly charged positive ions.  Atomic number 4 is a bivalent element.  It is found naturally only combined with other minerals.  Notable gemstone which has Atomic number 4 include beryl (aquamarine, emerald) and chrysoberyl.  Atomic number is a steel – gray, strong,  light weight brittle alkaline earth metal.



 Beryllium (White grey metal).

1)      Atomic Radius of Group 2 elements: Atomic radius increases down the group.  Atomic radius of Ra is less than Ba.

Element	Be	Mg	Ca	Sr	Ba	Ra
A.R.(pm)	112	160	197	211	222	215

   Where A.R.  = Atomic Radius in pico metre.

2)      Density of beryllium is 1.848g/cm³.

3)      Ionisation energy (I.E.) of Atomic number 4:
       1^st I.E.  =     899.5 KJ/mol.
       2^nd I.E.  =   1757.7 KJ/mol.
       3^rd I.E.  =  14848.7KJ/mol.
        3^rd I.E. is very large because the electron is present near to the nucleous and strongly binded to the nucleous.  Therefore it requires more energy to remove that electron which is ns² configuration.

4)      Elecrtropositive character of  Group 2 elements: Electropositive character increases down the group.

5)      Metallic properties of Atomic number 4: Atomic number 4 has low density (1.85 times that of water), high melting point= 1287⁰C.  High temperature stability and low co-efficient of thermal expansion, These conditions is suitable for aerospace material,  Atomic number 4 is a important component of planned space telescope because of its relatively high transparency to X-ray.

6)Reaction of Atomic number 4 with air: Beryllium does not burn unless it is in the form of dust or powder.  Be has a thin layer of Beryllium oxide on its surface which prevent any new oxygen getting at the underlying Be to react.
               2Be_(s)      +    O_2(g)     →       2BeO_(s)

7)      Diagonal relationship of Atomic number 4: Beryllium shows diagonal relationship with Aluminium.  It is the similarity between the first element of a group with the second element in the next higher group.

8)      Ores of Beryllium:
a)      Betrandite      (Be₄si₂O₇(OH)₂)
b)      Beryl              (Al₂Be₃Si₆O₁₈)
c)      Chrysoberyl   (Al₂BeO₄)
d)     Phenakite        (Be₂SiO₄)
      Precious form of Beryl are aquamarine, bixbite and emerald.

Nomenclature of elements with atomic number > 100

Introduction :
In the periodic table the elements are arranged in the order of the atomic number.  Atomic number is nothing but it is the number of proton or number of electron in the neutral atom. It is denoted by the letter Z.

The modern periodic law states that “ the physical and chemical properties of the elements are periodic function of their atomic numbers.”
Periods constitutes a series of elements whose atoms have the same number of  electron shell i.e., principal quantum number (n). There are seven periods and each period starts with a different principal quantum number.

Transuranium elements are elements beyond uranium (Z=92) they all synthethic elements.   Element with atomic number 100 is called as Fermium and the elements beyond these element in the periodic table is called as transfermium element.  The element after atomic number 100 is normally named as famous scientist.  But there is a lot of ambiguity in the names.  In each country the name will differ.  Some of the elements will be having more than one names

For example Kurchatovium with atomic number 104 is also called as Rutherfordium
Neilsbohrium also known as Borium has atomic number of 107.

The discovererc of these elements normally decide the name of the element.  Since same element can be made in the laboratory by more htan one scientist more than one name existed

Some other common names or nomenclature of elements with atomic number > 100 are given as

Atomic Number	Name
Z = 105	Dubnium
Z = 106	Seaborgium
Z = 108	Hassnium
Z = 109	Meiternium

IUPAC nomeclature of elements with atomic number > 100

To over come this ambiguity International Union of Pure and Applied Science (IUPAC) has recommended a nomeclature for elements with atomic number greater than 100.
IUPAC has given the name using the Latin words for their numbers.  The root words for the names is given in the table

Numerical	Root Name
0	nil
1	un
2	bi
3	tri
4	quad
5	pent
6	hex
7	sept
8	oct
9	en

Name (Nomenclature) of some of the elements with atomic number > 100

Atomic number	Name of the element	Symbol
Z = 101	Unnilunnium	Unu
Z = 102	Unnilbium	Unb
Z = 103	Unniltrium	Unt
Z = 104	Unnilquadium	Unq
Z = 105	Unnilpentium	Unp
Z = 106	Unnilhexium	Unh
Z = 107	Unnilseptium	Uns
Z = 108	Unniloctium	Uno
Z = 109	Unnilennium	Une
Z = 110	Ununnilium	Uun
Z = 111	Unununium	Uuu
Z = 112	Ununbium	Uub
Z = 113	Ununtrium	Uut
Z = 114	Ununquadium	Uuq
Z = 115	Ununpentium	Uup
Z = 116	Ununhexium	Uuh
Z = 117	Ununseptium	Uus
Z = 118	Ununoctium	Uuo
Z = 119	Ununennium	Uue
Z = 120	Unbinilium	Ubn

Colligative properties and determination of molar mass

Introduction :

The vapour pressure of solution decrease when a non volatile solution is added to a volation solvent. There are many properties of solution which are connected with this decreasing of vapour pressure. These are the relatively of vapour pressure of the solvent, depression of freezing points of the solvent. Elevation of the boiling point of the solvent, osmotic pressure of the solution. Everyone these of the property depend on the numeral of solute particle irrespective of their environment relative.

Relative lowering of vapour pressure:

In colligative properties and determination of molar mass, the vapour pressure of a solvent in solution is less than that of the pure solvent. Raoult recognized that the lower of vapour pressure depends simply on the concentration of the solute particles and it is dependent of their individuality.

`p_(1)=x_(1)p_(1)^(0)`

The reaction of the colligative properties in the vapour pressure of solvent is given as:

`Deltap_(1)=p_(1)^(0)-p_(1)=p_(1)^(0)-p_(1)^(0)x_(1)`

           =`p_(1)^(0)(1-x_(2))`

In a solution colligative properties containing several non volatile solutions, the lowering of the vapour pressure depends sum of the mole fraction of different solutions.

`(Deltap_(1))/(p_(1)^(0))` =`(p_(1^(0))-p_(1))/(p_(1)^(0))` =`x_(2)`

                The expression on the left hand side of the equation as mentioned earlier is called relative lowering of vapour force and is equal to the mole division of the solution of the colligative properties. The above equation can be determinations as:

`(p_(1)^(0)-p_(1))/(p_(1)^(0))` =`(n_(2))/(n_(1)+n_(2))`

                Here n₁ and n₂ are the number of mole of solvent and solute respectively present in the solution.

For dilute solution n₂<<n₁ hence neglection n₂ in the denominator we have

`(p_(1)^(0)-p_(1))/(p_(1)^(0))` =`(w_(2)xxM_(1))/(M_(2)xxW_(1))`

Here w₁ and w₂ are the mass and M₁ and M₂ are the molar mass of the solvent and solute correspondingly.

Elevation of boiling point:

In the colligative determination of molar mass, the vapour pressure of a liquid increase with enhance of temperature.It boils at the warmth at which its vapour pressure is identical to the atmospheric strain. The determination molar boiling point of a solution is always higher than that of the boiling point of the pure solvent.

Depression of freezing points:

The lowering vapour pressure of a solution cause a lowering of the freezing points compared to that of the determination  pure solvent in molar mass. The freezing points of the substance, the solids phase are the dynamic equilibrium with the liquid phases. The freezing points of the substance may be defined as temperature at which the vapour pressure of the substance in its liquid phase is equal to the vapour pressure in the solids state.

Mole chemistry

Introduction:
All substance is made up of smallest particle, called atom. In chemistry, the smallest particle, used for calculation, is said to be a mole. Any chemical equation or a chemical expression is given with moles of a substance under consideration.

Mole Definition:
Mole is the smallest unit used in all calculations in chemistry. Mole is a unit of measurement, which gives the same number of chemical entities (atoms, molecules, ions, electrons), as in number of atoms in 12 grams of carbon.

Mole concept is used in calculation of concentrations of solutions, in the calculation of molecular mass, etc.

Molar mass of a substance is defined as “mass per mole” of a substance.
Mole of a substance can also be defined as: “one mole of a substance contains Avogadro number of molecules or atoms”.
The value of Avogadro’s number is – 6.023 x 10²³
This is given by:

Mole =
Mole Calculation

Many calculations, regarding a chemical compound can be obtained from the mole concept.
Mole of a substance is used to calculate-
a.     Grams of a substance, if molar mass is known.
b.     Molar mass of a substance, if grams are known.
c.      Molarity, molality, mole fraction.
d.     Number of atoms present, with the help of Avogadro’s number.

Thus, moles play a very important role in chemical calculations.

Chemistry Mole Problems-
Example – 1:
A sample of magnesium hydroxide contains 12 grams of the substance. Calculate the number of moles of Mg(OH)₂ present.
Answer:
Molar mass of Magnesium hydroxide is: 58.32

Moles of Magnesium hydroxide = Mass in grams / Molar mass
                                                         = 12 grams / 58.32 grams/mole
                                                        = 0.206 moles.

Example – 2
13.65 moles of methane gas was obtained in a reaction. Find the mass in grams of methane.

Answer:
Molar mass of methane is 16.04 grams/mole.
Moles of methane = Grams of methane / Molar mass
Grams of methane = Moles of methane x Molar mass
                                    = 13.65 moles x 16.04 grams / mole
                                    =   218.94 grams of methane
Mole Problems Chemistry
Finding number of atoms, with Avogadro number
Example – 3
Calculate the number of atoms present in 4.2 moles of Sodium.
Answer:

1 Mole of a substance contains Avogadro number of atoms/molecules/ions.
Therefore, Avogadro’s constant =

We have – 4.2 moles of Sodium.

Number of atoms of Sodium present =


    4.2 moles of Na x   = 25.29 x 10²³ atoms of Na

Example – 4
 There are 3.01 x 10³² molecules of carbon dioxide present. Calculate:

i)                   Number of moles

ii)                Number of grams of CO₂.

Answer:

i)  To calculate the number of moles-

Moles of Carbon dioxide

= 3.01 x 10³² atoms of CO₂ x

= 4.99 x 10⁹ moles of Carbon dioxide

ii)                 To calculate the mass in grams of carbon dioxide –

Mass in grams = Moles of CO2 x Molar mass

                          = 4.99 x 10⁹moles x 44 grams / mole

                           = 219.56 x 10⁹ grams of Carbon dioxide

Molar Mass

Two very important entities of elements that recognise its characteristics are its atomic number and atomic mass. Similarly, mol mass of a compound is very important and is used in almost all stoichiometric calculations of a compound.

Molar Mass Definition

Molecular mass of a compound is the sum of atomic masses of all elements present in it.

Molecular mass is a physical property of a compound. It is denoted by M.

Molar Mass-

Molecular mass can also be calculated using the mass of a substance and the amount of substance present in moles.

Molecular mass = Grams of a substance / Moles

Mol mass is expressed in grams per mole.

Mol mass is used to find the moles of a substance, when its mass is given.

How to Calculate Molar Mass –

Molar mass of a substance can be found from the atomic masses of the elements present in it. Atomic masses of the individual elements can be obtained from the periodic table.

To calculate mol mass, we need to follow the following steps:

1. Write the Formula for Molar Mass of the compound whose mol mass is to be calculated.

Example – Magnesium chloride

MgCl₂

2. Find the subscripts/number of each element present in the compound.

1 x Mg + 2 x Cl

3. Multiply the number of each element present with the atomic mass of that element.

1 x 24.305 (Mg) 2 x 35.453 (Cl)

= 24.305 = 70.906

4. Finally, put all the atomic masses and their multiplied values together and sum it up.

Mg + 2 x Cl

= 24.305 + 70.906

= 95.211 grams/mole

Molar Mass of Water –

Formula of water is H₂O.

To calculate the mol mass of water, we need to have the atomic masses of hydrogen and oxygen.

Atomic mass of hydrogen = 1.008 grams/mole

Atomic mass of Oxygen = 15.994 grams /mole

There are two moles of Hydrogen. Thus, 2 x 1.008 = 2.016

Mol mass of water = 2.016 + 15.994 = 18.01 grams/mole.

Example of molar mass calculations:

To calculate the mol mass of Sodium hydroxide:

Formula of sodium hydroxide is NaOH

There are 1 mole of sodium, Na, 1 mole of hydrogen, and 1 mole of Oxygen.

To find the mol mass, we can add the atomic masses of all the elements.

Atomic mass of Sodium = 22.989 grams/mole

Atomic mass of oxygen = 15.994 grams/mole

Atomic mass of Hydrogen = 1.008 grams/mole

Mol mass of Sodium hydroxide = Na + O + H

= 22.989 + 15.994 + 1.008 = 39.991 grams/mole.

Mol mass of sodium hydroxide is taken approximately for calculations as 40 grams/mole.

Metals Nonmetals and Metalloids

Periodic table consists of an array of elements, with different metallic properties. Some are completely metallic, making them soft or hard metals. Other major type includes the non-metals, which have completely different physical and chemical properties from metals and are easily distinguishable.

Some elements have properties in-between that of a metal and a non-metal. These types of elements are termed as ‘metalloids’.

Characteristics of Metals Nonmetals Metalloids-

Common characteristics through which a metal, non-metal and a metalloid can be differentiated are:

Metals:

An element is called as metal, when, in the process of forming an ionic bond, it donates electrons, to form a positive ion. Thus, the main characteristics of a metal is that, it should have very less first ionization energy, or energy due to removal of outermost electron.

Some common properties of metal are:

1. Metals are mostly solids, hard or soft. Metals of first two groups of the periodic table are soft solids, while transition metals are hard. Mercury is the only metal, which is a liquid.
2. They are malleable and ductile.
3. They transfer heat and electricity due to the presence of ions in their structure. There is a special type of bond called as metallic bond, which gives metals all these distinctive properties.

Non-metals:

An element is said to be a non-metal, when it shows electronegative property than electropositive property. They accept electrons to form an ionic bond. Their first ionization energy is very high.

1. Non-metals are mostly liquids, gases or in some case, amorphous solids.
2. If they are solids, they are brittle solids, and are not malleable and ductile.

Metalloids:

These are elements which have properties in-between that of metals and non-metals. Metalloids are called as semi-metals. They have lustre like metals, but do not conduct electricity.

Metalloids find use as semiconductors. Metalloids are placed with the non-metals in 14^th, 15^th and 16^th Group of the periodic table.

List of Metals Nonmetals and Metalloids

Some of the metals, metalloids and non-metals are listed below:

Metals	Metalloids	Non-metals
Copper- Cu	Silicon -Si	Oxygen-O
Iron- Fe	Germanium - Ge	Chlorine –Cl
Mercury Hg	Antimony -Sb	Nitrogen – N
Cadmium Cd	Arsenic - Sb	Carbon – C
Sodium Na	Tellurium -Te	Sulfur – S
Calcium Ca		Phosphorus -P
Chromium Cr		Bromine -Br

Is Gold a Metal Nonmetal or Metalloid–

Gold, Au, is one of the transition elements. It has an atomic number of 79 and is placed in group 11 of the table. It is a transition metal.

Is Sodium a Metal Nonmetal or Metalloid–

Sodium is placed in the first group of the periodic table. It is a soft metal. Sodium is one of the alkali metals.

Is Calcium a Metal Nonmetal or Metalloid-

Calcium, a white amorphous solid, is a metal, because of its electron donating property. Calcium is an alkaline earth metal.

Ionic Compound

Elements combine together to form compounds. Chemical compounds are of many types, depending upon the bond present in them. Corresponding to the two ways by which any two atoms rearrange to form a compound, two types of bonds are formed.

1. Ionic bond
2. Covalent bond.

Ionic bond or electrovalent bond is established by the transfer of one or more valence electrons from one atom to the other.

Thus, ionic bond is a chemical bond formed between two atoms by the transfer of one or more valence electrons from one atom to the other.

This bond is also called as a polar bond.

Formation of an ionic bond:

Formation of an ionic bond can be explained using the following example:

Consider an atom A , which has two electrons in its outermost shell. Another element B, has 6 electrons in its outermost shell.

The atom A has two electrons in excess, to make it to the Noble gas electronic configuration, while atom B has two electrons less to make it to that level.

Now, atom A gives two of the excess electron to atom B, and by this, atom A attains a completely filled shell, while atom B, having attained the required amount of electrons, also reaches the noble gas configuration.

They therefore form an ionic bond between themselves.

List of Ionic Compound Formula-

Lists of ionic-compounds with their formula are:

S.No	Ionic-compound	Formula
	Sodium chloride	NaCl
2.	Potassium iodide	KI
3.	Lithium iodide	LiI
4.	Aluminium oxide	Al2O3

Ionic Compound Example-

Some examples of ionic compound are:

Magnesium oxide - MgO – Mg²⁺, O^2-

Calcium Fluoride – CaF₂ – Ca²⁺, 2F^-

Aluminium Fluoride – AlF₃ – Al³⁺, F^3-

Properties of an Ionic Compound-

1. Ionic-compounds are three dimensional solids, with well-defined geometrical pattern.
2. Ionic solids conduct electricity when they are in water solution or in the fused state (molten state).
3. They are quite hard, have low volatility and high melting and boiling point.
4. Ionic solids are soluble in polar solvents, due to dissociation of their ions.
5. Ionic solids are very stable and have very high density.

Is NaCl an Ionic Compound-

Sodium chloride, NaCl is ionic in nature.

The formation of sodium chloride is as follows:

Na has an electronic configuration of 2, 8, 1. The last electron, in the third shell, has to be removed, for it to attain noble gas configuration.

Chlorine has a configuration of 2, 8, 7. It needs one extra electron, which it gains from Sodium, thereby getting the required magic number of ‘8’.

Is Salt an Ionic Compound–

Sodium chloride is also called as common salt. Other than this, most of the compounds, commonly known as salts are formed from the neutralization reaction of an acid and a base. They are all ionic-compounds, because they dissociate into ions in their solution.

Bonding and molecular structure

Syllabus : 

Valency electrons, the octet rule. Electrovalent and covalent bonds with examples. Properties of electrovalent and covalent compounds. Limitation of octet rule (examples), coordinate covalent bonds (examples).

Directionality of covalent bonds, shapes of polyatomic molecules (examples), concept of hybridization of atomic orbitals (qualitative pictorial approach): sp³, sp² and sp hybridizations with typical examples. Tetrahedral space model of-carbon atom, single-bond, double- bond and triple - bond involving carbon atom with examples a and 7t bonds.

Valence shell Electron Pair Repulsion (VSEPR) concept (elementary idea) - shapes of H₂0, H₂S,
CH₄, NH₃, C0₂, N0₂ and S0₂ molecules. Concept of resonance (elementary idea), resonance structures (examples). Elementary idea about electronegativity, bond polarity and dipole moment. Hydrogen bonding (inter - & intra molecular structures) and its effects on physical properties (mp, gp, and solubility).

Double salts and complex salts, and coordination compounds (examples only), coordination number (examples with C.N 4 and 6 only).

Valency Electrons and the Octet rule

When details of the electronic configurations of the elements came to be known, it was found tha the arrangement of energy levels in different orbits round the nucleus was different for different elements. Chemical union between atoms to form molecules are energetically favoured only if the chemical combination leads to lowering of energy of the system i.e, if the energy of the combined atoms i.e., the product molecule is less than the sum of the energies of the reactant molecules.

Chemical combination of atoms involves electrons of the two atoms and nuclei take no active part in chemical combination. In each individual atom, the outermost electrons have the highest energy amongst all its electrons. So the lowering of energy must be through the interactions of the outermost electrons of the two interacting atoms. Thus chemical combination in all probabilities, should involve only the outermost electrons of the participating atoms and the interactions should be such that it causes lowering of energy w.r.t. to the initial condition when the atoms lay separated from each other.

A close study of chemical properties of the different classes of elements led us to the fact that the noble gases were the most chemically inactive species amongest all the elements. They had no tendency to combine with themselves or with other elements. They were so inert that they even did not like to form molecules by the combination of two atoms. Inert gas molecules are monatomic — or in other words their atoms do not form molecules. They are devoid of any chemical affinity. mSo it may be seen in the light of our discussion above that their outer electronic configurations are most stable amongest all the elements and no further stabilization is possible by further interaction among themselves or with electrons in other elements.

The extra nuclear electrons present in n successive inert gases are 2,10,18, 36,54 and 86. These numbers were termed magic numbers as the presence of electrons in any one of these numbers in an atom gives special stability to the atom. This stability is lost if this number is changed even by 1 unit on either side.

The arrangement of the electrons in the inert gases can be described as follows :
He — 2
Ne — 2-8
Ar — 2-8-8
Kr —2-8-18-8
Ne —2-8-18-18-8

Molecular mass of polymers

Introduction :  
Several simple organic molecules of one or two types combine with each other by chemical bonds forming macro molecules the product is called a polymer. Simple organic molecules which can form polymers by chemical bonding are called monomers while this process of chemical combination is called polymerisation. In any sample or monomers there are very large number of molecules of lower masses same molecular weights and similar physical and chemical properties.

Polymers:

Ehereas in polymer sample having same molecular weights the number of molecules are very small. The polymers have comparatively very high molecular weight but all molecules do not have comparatively very high molecular weight but all molecules do not have identical molecular weights. Polymers prepared from the same monomer in different conditions do not have all the properties identical. Depending on reaction conditions polymer products have different proportions of molecules of lower and higher masses.

Molecular mass of Polymers:
Overall molecules of ethene combine with each other by addition reaction and give polyethene. This reaction at the first stage two molecules of ethene monomer combine together giving a dimer. A third molecule of ethene combines with this dimer giving a trimer and a forth molecule combine further gives a tetramer. This way monomer molecules go on joining and chain becomes longer. As result very large chain is former which is called macro molecule or polymer.
CH₂ = CH₂ `stackrel(CH_2=CH_2)(->)` CH₃-CH₂-CH = CH₂ `stackrel(CH_2=CH_2)(->)`
CH₃-CH₂-CH₂-CH₂-CH=CH₂ `stackrel(CH_2=CH_2)(->)`
CH₃-CH₂-CH₂-CH₂-CH₂-CH₂-CH=CH₂
`stackrel(nCH_2=CH_2)(->)`  [-CH₂-CH₂-]_n

The polymer chain can be lengthen upto certain limit at laboratory cantons. The tendency of this long chain then decrease to combine with further monomers. Thus in any condition polymers resulting from monomers do not increase in weight more than a certain limit. Generally any polymer asmple contains varying chain-lengths, its molecular mass is always an average molecular mass. The molecular mass of a polymer is expressed as number average molecular mass `barM_n` or weight average molecular mass `barM_w`.
`barM_n = (sumN_tM_t)/(sum tN_t)`
`barM_w = (sumN_tM^2_t)/(sumtN_tM_t)`
Where N_t = number of molecules
M_t = molecular mass.

Molecular Orbital Theory

Introduction :
Molecular orbital theory is a polycentric region in space, defined by its size and shape, associated with two or more atoms in a molecule and each has a capacity of two electrons with opposite spins. Thus, in a molecular orbital, electrons are revolving in the field of more than one nucleus. The molecular orbital to explain formation of chemical bond, relative bond strengths, paramagnetic or diamagnetic nature.

Feature of Molecular Orbital Theory:

• The main features of the Molecular Orbital Theory are:
• The atomic orbital of the combine atoms partly cover to form new orbital, called molecular orbital.
• As a result of this, the atomic orbitals lose their being identity.
• Thus in a Molecular Orbital Theory, electrons revolves in the field of more than one nucleus.
• The number of Molecular Orbital is produced is equal to the number of overlap atomic orbitals.
• Maximum capacity of a Molecular Orbital is two electrons with opposite spins.
• Only those atomic orbitals can come together to form Molecular Orbital Theory which has analogous energies as well as proper orientations.
• Molecular orbital theory obtained addition of wave functions of atoms involved,`Psi`(MO)=`Psi`_A + `Psi` _B   is called bonding molecular orbital.
• Molecular orbital theory obtained by subtraction of wave functions of atoms involved  `Psi`  ^*(MO)=`Psi`_A - `Psi`_B is called antibonding molecular orbital.
• Probability of bonding molecular orbital formation greater than that of antibonding molecular orbital formation.
• Molecular theory gives the electron probability distribution around a group of nuclei just gives the electron probability distribution around nucleus.
• The shape of the molecular orbital theory produced depends on the type of the combining atomic orbitals.
• Inner molecular orbital theories which do not take part in bond formation are called non-bonding molecular orbital theory.

 Conditions for the formation of molecular orbitals:

• Any two atomic on combination do not form molecular orbitals.

• In fact, there are certain limitations to the combination of atomic orbitals.

• The energies of combining atomic orbitals should of similar magnitude.

• Thus, a homonuclear diatomic molecule will not be formed. if 1s orbital of one atom overlaps with 2s-orbital of another atom.

• Combination of atomic orbitals takes place only, if overlapping takes place to a considerable extent, since greater the overlapping of atomic orbitals, the greater is the build-up of the charge between the nuclei.

• The combining atomic orbitals should have power over the same symmetry about the molecular axis.

Chemical equilibrium animations

Introduction :

The experimental observations of chemical equilibrium tell us that most of the chemical reactions when carried out in closed vessels do not go to completion. Under these a conditions, a reaction starts by itself or by initiation, continues for some time at diminishing rates and ultimately appears to stop. The reactants may still be present but they do not appear to change into products any more. What happens in such case is that the products of the reaction start reacting at the same rate as the reactants. In other words, the rate of the back reaction becomes equal to the rate of the forward reaction.

Characteristic features of chemical equilibrium

Thus, in a given time as much of the products are formed as react back to give the reactants. The composition of the reaction mixture at a given temperature is the same irrespective of the initial state of the system, i.e., irrespective of the fact whether we start with the reactants or the products. The reaction in such conditions is said to be in a state of equilibrium.

The attainment of equilibrium can be recognized by noting constancy of observable properties such as pressure, concentration, density or color whichever may be suitable in a given case.

The relationship between the quantities of the reacting substances and the products formed can be worked out readily with the help of the law of mass action.

The Laws of Mass Action of Chemical equilibrium

The laws of mass action states that the driving force of a chemical reaction is proportional to the active masses of the reacting substances. Assuming that the driving force determines the reaction rate, the law may be stated as follows: The rate at which a substance reacts is proportional to its active mass and the rate of the chemical reaction is directly proportional to the product of the active masses of the reacting substance.

Consider a general reversible chemical reaction

aA + bB ↔ mM + nN

According to the law of mass action, assuming that active masses are equivalent to molar concentrations,

The rate of the forward reaction, r_f α [A]^a [B]^b = K_f[A]^a [B]^b

The rate of the reverse reaction, r_r α [M]^m[N]ⁿ = K_r[M]^m [N]ⁿ

Where Kf and Kr are proportionally constants and square brackets represent the molar concentrations of the entities enclosed. At equilibrium, the rate of the frontward reaction is equal to the rate of the reverse reaction, that is, Kf[A]a [B]b = Kr[M]m [N]n

K_f /K_r = K_eq = [M]^m [N]ⁿ / [A]^a [B]^b

Historical development of the periodic table

We now know more than 100 elements, the elements were classified as metals, metalloids and non metals. Metals are good conductors of heat and electricity, they are shining and they are malleable and ductile. It would be difficult to study individually the chemistry of all the elements and their numerous compounds. The periodic table provides a systematic and extremely useful framework for organizing a lot of information available on the chemical behavior of the elements into a few simple and logical patterns. This gave rise to the necessity of classifying the elements into various groups or families having similar properties.

Introduction to the historical development of the periodic table 

There were many chemists who has contributed their theories for the historical development of periodic table.  Some of them are:
Dobereiner’s Triads
Lother-Meyer’s Arrangement
Newlands Law of Octaves

Dobereiner’s Triads contribution for the historical development of periodic table

In 1829, John Dobereiner (German Chemist) classified elements having similar properties into groups of three. These groups were called triads. According to this law when elements are arranged in the order of increasing atomic mass in a triad, the atomic mass of the middle element was found to be approximately equal to the arithmetic mean of the other two elements. For example lithium, sodium and potassium constituted one triad. However, only a limited number of elements could be grouped into traids.

Newlands Law of Octaves contribution for the historical development of periodic table
In 1865, John Newlands (English Chemist) observed that if the elements were arranged in order of their increasing atomic weights, the eighth element starting from a given one, possessed properties similar to the first, like the eighth note in an octave of music. He called it the law of octaves. It worked well for the lighter elements but failed when applied to heavier elements.

Lother-Meyer’s Arrangement contribution for the historical development of periodic table

In 1869, J. Lother-Meyer in Germany gave a more detailed and accurate relationship among the elements. Lother-Meyer plotted atomic volumes versus atomic weights of elements and obtained a curve. He pointed out that elements occupying similar positions in the curve possessed similar properties.

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.