1. Title: Metallic Bond Dr
Title: Metallic Bond Dr. Sujata Kundan Assistant Professor Department of Chemistry and Chemical Sciences

2. Metallic Bond:
A metallic bond can be defined as the strong force that results from delocalized electrons (conduction electrons) grouped in a freely moving 'electron sea' that surrounds more than one positively charged metal ion (cation). Because of this sharing of electrons between multiple cations, a better term would be metallic bonding, because there is no such thing as a single metallic bond.
This type of bonding is what makes metals so strong, malleable, conductible, and often shiny (metallic luster). This strong bond is also why there is such a high boiling point and melting point for most metals.
In metallic bonds, the valence electrons from the s and p orbitals of the interacting metal atoms delocalize. That is to say, instead of orbiting their respective metal atoms, they form a "sea" of electrons that surrounds the positively charged atomic nuclei of the interacting metal ions. The electrons then move freely throughout the space between the atomic nuclei.

3. Discoverer/Developer:
William Hume-Rothery was a British metallurgist who graduated with high honors in chemistry from Magdalen College, Oxford. He attained his Ph.D. at Royal School of Mines. After World War II, he returned to Oxford to research intermellatic chemistry. During this time, he tried to find out more about the kind of bonding in metals. This is where he formed his hypothesis of metallic bonding.
Theories of the Metallic Bond:
Electron Sea Model for Metallic Bonding
Band Theory

4. Electron Sea Model for Metallic Bonding:
To account for the bonding in metals, The first idea about this bond was given by Drude (1900) who proposed metals as containing a number of free electrons moving in spaces among atoms like the molecules of the ideal kinetic theory of gases. Lorentz (1923) believed that solid metals consist of lattices of liquid cation spheres with free electrons moving in the interstices. Thus this theory is also known as Drude Lorentz theory.
This model / theory is based on the following characteristic properties of metals:
Low ionization energies: Metals generally have low ionization energies. This implies that the valence electrons of metal atoms are not strongly held by the nucleus. Valence electrons can move freely out of the influence of their kernels (atomic orbit/structure minus valence electrons). Thus, metals have free mobile electrons.
Large number of empty orbitals: It has been observed that in metals a number of valence orbitals remain empty as the number of valence electrons in metals is generally less than the number of valence orbitals.
For example, lithium {(Li, Z = 3) 1s22s1} has 2p-orbitals vacant;
Sodium {(Na, Z = 11) 1s22s22p6 3s1} has 3p-and 5d-orbitals vacant;
Magnesium {(Mg, Z = 12) 1s22s22p6 3s2} has 3p-and 3d-orbitals vacant.

5. The important features of electron sea model are:
The positively charged kernels of metal atoms are arranged in a regular fashion in a metallic lattice.
Loosely held valence electrons, surround each kernel in metallic lattice. Being loosely held to its kernel, the valence electrons enjoy complete freedom in the metallic lattice and are regarded as mobile electrons.
In short, the metal may be regarded as 'a sea of electrons (common pool of electrons) in which there is a three dimensional ordered arrangement of positively charged kernels, surrounded throughout by mobile valence electrons'. This explanation is also responsible for its name electron sea model.
Thus, the simultaneous force of attraction between the mobile electrons and the positive kernels that binds the metal atoms together, is known as metallic bond.

6. This theory explains different metallic properties as:
1. Metallic luster:
The bright luster of metals is due to presence of delocalised mobile electrons.
When light falls on the surface of the metal, the loosely held electrons absorb photons of lights. 
They get promoted to higher energy levels (excited state), oscillating at a frequency equal to that of the incident light. 
These oscillating electrons readily return from the higher to the lower levels of energy by releasing energy, thus becoming a source of light radiations. 
Light appears to be reflected from metal surface and the surface acquires a shining appearance, which is known as metallic luster.

7. 2. Opaqueness:
The light that falls on metals is either reflected or completely absorbed by the delocalised electrons. 
Because of this, no light is able to pass through metals and they are termed as opaque.
3. Melting and boiling points:
Metals have metallic bond strengths, which is intermediate to that of covalent and ionic bonds.
Therefore in general, metals have boiling and melting points in between to that of covalent and ionic compounds.

8. 4. Electrical conductivity:
The presence of mobile electrons causes electrical conductivity of a metal. When a potential difference is applied across the metal sheet, the free mobile electrons in the metallic crystal start moving towards the positive electrode. 
The electrons coming from the negative electrode simultaneously replace these electrons. 
Thus, the metallic sheet maintains the flow of electrons from negative electrode to positive electrode. This constitutes electrical conductivity.
5. Thermal conductivity:
When a part of the metal is heated, the kinetic energy of the electrons in that region
increases.
Since the electrons are free and mobile, these energetic electrons move rapidly to the
cooler parts and transfer their kinetic energy by means of collisions with other
electrons. 
Therefore, the heat travels from hotter to cooler parts of the metals.

9. Malleability and ductility:
Metals can be beaten into sheets (malleability) and drawn into wires (ductility). Metallic bonds are non-directional in nature.
Whenever any stress is applied on metals, the position of adjacent layers of metallic kernels is altered without destroying the crystal. 
The metallic lattice gets deformed but the environment of kernels does not change and remains the same as before.
The deforming forces simply move the kernels from one lattice site to another.
7. High tensile strength:
Metals have high tensile strength. Metals can resist stretching without breaking. 
A strong electrostatic attraction between the positively charged kernels and the mobile electrons surrounding them is the reason for tensile strength.

10. 8. Hardness of metals:
The hardness of metals is due to the strength of the metallic bond. In general, the strength of a metallic bond depends upon
The greater the number of valence electrons for delocalisation the stronger is the metallic bond.
Smaller the size of the kernel of metal atom, greater is the attraction for the delocalised electrons. Consequently, stronger is the metallic bond.
For example, alkali metals have only one valence electron and larger atomic kernels, which makes the metallic bonds weak. Consequently these metals are soft metals.
Examples; for materials having metallic bonds are most metals such as Cu, Al, Au, Ag etc. Transition metals (Fe, Ni etc) form mixed bonds that are comprised of covalent bonds (involving their 3d-electrons) and metallic bonds. This is one of the reasons why they are less ductile than Cu, Ag and Au.

11. Band Theory of Metals:
Band Theory was developed with some help from the knowledge gained during the quantum revolution in science. In 1928, Felix Bloch had the idea to take the quantum theory and apply it to solids. In 1927, Walter Heitler and Fritz London discovered bands- very closely spaced orbitals with not much difference in energy.
Molecular orbital theory extended to solids is referred as Band Theory.
Example : Construction of a crystal of a sodium metal by adding Na atoms one at a time forming first Na2 than Na3, Na4 ,Nan respectively.
In Na2 molecule, each Na-atom has electronic configuration [Ne] 3s1, with a single 3s valence electron. Two 3s-atom orbitals, one from each Na-atom, overlap to form two molecular orbitals σ (3s) and σ *(3s). There are just two valence electrons, which will occupy lower energy bonding molecular orbital σ (3s). The antibonding molecular orbital σ* (3s) is vacant.

13. Explanation:
Three Na atoms joined to form Na3. Three 3s atomic orbitals would combine to form three molecular orbital one bonding, one non­bonding and one antibonding. The energy of the non-bonding MO is between that for the bonding and antibonding orbitals. The three valency electrons from the three sodium atoms would occupy the bonding and non-bonding molecular orbitals .
In Na4, the four atomic orbitals would form four molecular orbitals-two bonding, and two anti-bonding. The four valence electrons would occupy the two lowest energy bonding molecular orbitals, half of the total number of molecular orbitals are vacant.
As the number of atoms in the cluster increases, the spacing between the energy levels of the various orbitals decreases further, and when there are a large number of atoms, the energy levels of the orbitals are so close that a band of closely spaced molecular orbital is formed. The band is half-full because each molecular orbital can hold two electrons, and there are N valence electrons. This band is also known as valence band as it contains the outer or valence electrons. The empty band is known as conduction band. Since only half the molecular orbitals in the 3s valence band are filled in the bonding molecular orbitals. It requires very small amount of energy to excite an electron to an unoccupied molecular orbital. The molecular orbitals extend in there dimensions over all the atoms in the crystal. So electrons have a high degree of mobility. That is why metals have high thermal and electrical conductivities.

14. Usefulness of band Theory Of Metals:
With the help of Band theory, we can classify materials into three categories viz. conductors, insulators and semi-conductors, depending on the energy gap between the valence and conduction bands.
In conductors (metallic), either the valence and conduction bands overlap or the valence band is only partly full .

15. In Insulators (non-metallic elements), there is large band gap between the filled valence band and empty conduction band as in diamond. Therefore Electrons cannot qe promoted from the valence band to conduction band where they could move freely.
Semiconductors are of the types, intrinsic and extrinsic semiconductors. Intrinsic semiconductors (like Si or Ge) are having small energy gap between the filled valence band (VB) and empty conduction band (CB) sufficient to promote an electron from VB to CB. The hole left in the VB and the promoted electron in the CB both contribute towards conductivity.

16. Hydrogen Bonding: Types of hydrogen bonding:
A hydrogen bond is the electromagnetic attraction between polar molecules in which hydrogen is bound to a larger atom, such as oxygen or nitrogen. This is not a sharing of electrons, as in a covalent bond. Instead, this is an attraction between the positive and negative poles of charged atoms.
Types of hydrogen bonding:
Hydrogen bonds can be classified into two types;
Intermolecular hydrogen bond and
Intramolecular hydrogen bond.
1. Intermolecular hydrogen bond:
This occurs when the hydrogen bonding is between H-atom of one molecule and an atom of the electronegative element of another molecule. For example (i) hydrogen bond between the molecules of hydrogen fluoride. (ii) hydrogen bond in alcohol or water molecules.
Intermolecular hydrogen bond results into association of molecules. Hence, it usually increases the melting point, boiling point, viscosity, surface tension, solubility, etc.

17. Intermolecular Hydrogen Bonding Intramolecular Hydrogen Bonding
This bond is formed between the hydrogen atom and an atom of the electronegative element (F,O,N) , of the same molecule. Intramolecular hydrogen bond results in the cyclization of the molecules and prevents their association. Consequently, the effect of this bond on the physical properties is negligible. For example, intramolecular hydrogen bonds are present in molecules such aso-nitrophenol, o-nitrobenzoic acid, etc.

18. Cause of hydrogen bond formation:
Hydrogen atom is bonded with a highly electronegative element, and therefore the shared pair of electrons move away from the hydrogen atom towards the electronegative atom. Hydrogen atom becomes electropositive with respect to the electronegative element. This results in the development of positive charge over hydrogen atom and partial negative charge over the electronegative element. This further leads to the formation of a polar molecule with electrostatic force of attraction. The magnitude of H-bond depends on the physical state of the compounds. It reaches a maximum value in solid state and minimum in gaseous state.

19. References:
“Solids and Surfaces”: A chemists view of bonding ìn Solids and Surfaces: A chemists view of bonding in extended structures in extended structures; Roald Hoffmann, VCH Publishers, (1988).
“The Electronic Structure and Chemistry of Solids” P.A. Cox, Oxford University Press, Oxford (1987). P.A. Cox, Oxford University Press, Oxford (1987).
“Chemical Bonding in Solids” Jeremy K. Burdett, Oxford University Press, Oxford (1995).
George A. Jeffrey. An Introduction to Hydrogen Bonding (Topics in Physical Chemistry). Oxford University Press, USA (March 13, 1997).