Atomic structure and chemical bonding , Basics of material science, General studies , UPSC ESE

Nucleons

• mass of a proton = mass of  neutron = 
• 
  mass of an electron = 

Atomic mass number  

• A = number of protons + number of neutrons 
• sodium --- 11
• charge of an electron =
• 
  charge of a proton = 

Atomic number 

• Z is  number of electrons or number of protons in a neutral atom. 

Atomic weight 

• It is defined as average relative weight of atom as compared to the weight of one atom of carbon 12 and expressed in terms of atomic mass unit. 

Atomic mass unit 

• It is defined as one twelfth of the mass of an atom of carbon 12. 
  

Mole (SI unit) 

• 1 mole is the amount of substance that contains as many particles or entities as their atoms in 12 gram of carbon 12 isotope. 
• The mass of one mole of substance in gram is called molar mass. 
  

Avogadro number 

• Avogadro proposed that equal volumes of gases at the same temperature and pressure should contain equal number of molecules. 
• The number of entities or atoms in a mole is known as Avogadro number. 

Elements 

• There are three types of elements 

1. Isotopes :same atomic number but different atomic mass. 
2. Isobars : same atomic mass but different atomic number. 
3. Isotones : same number of neutrons but different atomic number. 

• Energy is inversely proportional to stability. 
• As nucleons size increase ,stability decrease ,energy increase ,emit radiations (three types of radiations) called radioactive materials. 

Comparison between alpha ,beta and gamma rays 

Alpha rays 

• It consists of two protons and neutrons. 
• Its mass is equal to 4 AMU. 
• emitters of alpha particle are heavier elements. 
• It is positively charged particle. 
• It has highest ionisation potential. 
• It travels with very slow speed. 
• It has least penetration power. 
• It cannot penetrate even the outer layer of human skin. 

Beta rays 

• High energy in nucleus ,break neutron. 
• These are extremely energetic electrons emitted from the nucleus. 
• It is negatively charged particles. 
• Ionization potential is less than alpha particles. 
• These particles travel with almost speed of light. 
• These are having higher penetration power than alpha particle. 
• These can easily penetrate human skin but cannot penetrate thin metal sheet. 

Gamma rays 

• These are high frequency electromagnetic waves. 
• These are massless and chargeless. 
• These are having negligible ionization potential. 
• These travel with a speed of light. 
• These are having highest penetration power .
• They can penetrate thick layer of concrete. 

Atomic theories 

Dalton's atomic theory

• Matter is made up of very thin tiny particles called atoms which are indivisible structure. 
• Atoms can neither be created nor destroyed. 
• All atoms of a particular element are identical in all respects including mass and physical properties etc whereas atoms of different elements differ in their properties. 
• Compounds are formed when atoms of different elements combined in a fixed ratio law of chemical composition. 

Thomson's atomic model 

• JJ Thomson proposed that an atom is a sphere filled with positively charged metal distributed uniformly with sufficient electrons are embedded in it so the atom is electrically neutral. 

Rutherford's atomic model 

• 
  On the basis of his famous alpha particle scattering experiment Rutherford proposed the nuclear model of atom ,according to this model the positive charge and most of the mass of the atom potentially concentrated in extremely small region known as nucleus. 
• The electrons are revolving around the nucleus with identical path as the planets are moving around the sun. 

Bohr's atomic model 

• Bohr made the following assumption in his theory 

1. An atom consists of nucleus which contain protons and neutrons and hence most of the mass of the atom is concentrated in to nucleus. Electrons revolve around the nucleus in same circular orbit. 
2. The electrons revolve only in certain non radiating orbits called as stationary orbit. There is no loss of energy in these orbit when electron revolve ,Bohr found that electron can only occupy an orbit for which the angular momentum of moving electron is an integral multiple of h/2pie where h is Planck's constant. 
   
3. When an electron jumps from an orbit of higher energy to another orbit of lower energy than energy is released in the form of radiations and vice versa. The amount of energy released or absorbed is the difference of energies into orbit. 
   

Electron energy through Bohr model 

1. 
   The greater the distance of an electron from the nucleus higher is its total energy. 
2. An electron orbiting very close to the nucleus in the first shell is tightly bound to the nucleus and possesses very small amount of energy so it would be difficult to knock out this electron from the nucleus. On the other hand an electron orbiting  far from the nucleus in the outermost shell is loosely bound to the nucleus and possesses greater amount of energy so this electron can be easily knocked out of its orbit. This is the reason why valence electron participate in chemical reactions and chemical bonding etc. 

Sommerfeld model 

• 
  Mass of electron is changing.
•  According to sommerfield ,  the path of the electron revolving is an elliptical orbit. The velocity of the electron is an elliptical orbit varies at different parts of the orbit. This causes the proportional variation in the mass of moving electron. To deal with these variables first two quantum numbers are introduced. 

Quantum numbers 

First quantum number (n) 

• Also known as principal quantum number. 
• n = 1,2,3,.... 
• It gives energy of electrons. 
• It represent shells or orbits(K,L,M,....). 
• It gives distance of an electron from nucleus. 
• It is only quantum number related with bohr model. 

Second quantum number (l) 

• Also known as angular or azimuthal quantum number. 
• It gives angular momentum of electrons. 
• It represent subshells (s,p,d,f,...). 
• L = 0,1,....,(n-1)  

Third quantum number (ml)

• It gives energy states in subshells. 
• There are (2l + 1) values of ml ranging from -l to + l. 

Fourth quantum number 

• Also known as a spin quantum number. 
• Electrons can have either clockwise or anticlockwise spin ,hence spin quantum number can be 
  

Modern concept of atomic model or electron cloud model or wave mechanical model

• Modern concept is based on quantum mechanical theory or wave mechanical model. 
• It takes into account the Heisenberg uncertainty principle. 
• Heisenberg uncertainty principle:  
• 
  Consequently definite discrete orbits of electrons are not possible so with this model an electron is no longer treated as a particle moving in a discrete orbit rather position  is considered to be the probability of an electrons being at various locations around the nucleus. In other words position is described by a probability distribution or electron cloud. 
• The electrons in an atom are in a continuous motion around the nucleus so we can express the motion of electron in the force field of nucleus by schrodinger wave equation. 
• Schrodinger wave equation
  
• The solution of the equation 1 gives spatial distribution of electron density around the nucleus. Equation 1 can have many solutions but the physically meaningful and acceptable solutions are obtained only when the electrons are permitted to have specified value of energies. These allowed values of energies or the allowed electronic state in an atom is specified by the solution in terms of four quantum numbers .
  
• In three substate of  2p subshell electrons are distributed in three separate directions which are perpendicular to each other. Clearly electrons in these States has directional characteristics. However if the 2p state is full then distribution become spherical. 

Chemical bond 

• The binding forces between atoms or molecules are known as chemical bond. 
  

Ionic bond 

• Ionic bonds are non directional bonds.
• Cations are smaller in size than anions. 
• Crystalline in nature. 
• High-strength. 
• High hardness. 
• Atom movement within the crystal is not possible without breaking bond which makes this material characteristically brittle. 
• High melting point. 
• Non conducting of electricity due to absence of free electrons. 
  

Note: in these materials a small amount of conductivity can be observed due to movement of Ions which is known as ionic conductivity. It has very small value and can be considered equal to zero for all practical purposes. 

Covalent bonds 

• Covalent bonds are directional bonds. 
• Covalent solids do not form closed packed structure. 
  
• Covalent bonds are of two types :
• Polar covalent bond : it is formed between the similar elements,Example H2O ,HCL ,etc. ,And has partial ionic Character also. 
• Non polar covalent bond: bond is formed between identical elements. Example H2 ,cl2 , etc. Bond is purely covalent. 
  

Molecular orbital theory 

• Case 1 consider the case of hydrogen molecule 
• 
  Placement of electrons in molecular orbitals takes place one at a time starting from lowest energy molecular orbital . As long as the number of electrons in bonding molecular orbitals are greater than the number of electrons in antibonding molecular orbitals the molecules remain stable. 
• Case 2 consider the case of O2 molecule. 
  
• Let us assume that bond is formed along x axis . 
  
• 
  
• 
  Demagnetization results when there are all paired electrons in molecular orbitals. 
• Paramagnetism result when there are unpaired electrons in molecular orbitals.

Metallic bond:

• It is the weakest bond while ionic bond is strongest bond among all primary bonds.
• High density.
• Good conductors of electricity.
• Metals have lusture or shine due to presence of electron clouds.
• Plastically deformable.
• High melting point.
  

Secondary bonds

• Molecular bonds
• Vander waal's forces
• Weak and unstable

Dispersion bond

• example : inert gases , non polar molecules ,etc
  

Dipole bond

• They are stronger than dispersion bond but they are weaker than primary bonds.

Hydrogen bond

• It is a special case of dipole bond.
• It is the strongest among all secondary bonds but weaker than primary bonds.