Updated Unit 4 Chemical Bonding I sem (Basic Chemistry).pdf

Summary of Basic Chemistry Distribution of Theory Marks Unit Teaching Unit Title No. hours...

Summary of Basic Chemistry Distribution of Theory Marks Unit Teaching Unit Title No. hours R U A Total Level Level Level Marks IV Chemical bonding 06 02 03 04 09 V Electrochemistry and 12 03 04 05 12 Metal Corrosion, prevention Paints, Varnishes, Insulators, VI Polymer Adhesives Lubricants, 12 03 05 06 14 Catalysis TOTAL 30 08 12 15 35 UNIT 4: Chemical Bonding (09 Marks)  4.1 Indian Chemistry- Philosophy of atoms by Acharya Kanad  4.2 Electronic theory of valency: Assumptions, chemical bonds: Types and characteristics of electrovalent bond, covalent bond, coordinate bond, hydrogen bond, metallic bond and Intermolecular force of attraction.  4.3 Molecular arrangement in solid, liquid and gases.  4.4 Structure of solids: crystalline and amorphous solid, properties of metallic solids-, Unit cell- simple cubic, body centre cubic (BCC), face centre cubic (FCC), hexagonal close pack crystals. 4.1 Indian Chemistry- Philosophy of atoms by Acharya Kanad Dr S K Mandavgade  Maharshi Kanad was the founder of the school of Vedic philosophy known as Vaisheshika. The creation and existence of this universe was explained by this school by proposing an atomic theory which was discovered by Maharshi Kanada 2600 years ago.  However, not many people are aware that a theory of atoms was formulated approximately 2500 years before Dalton by an Indian sage and philosopher named Acharya Kanad.  Acharya Kanad was born in 600 BC in Prabhas Kshetra (near Dwaraka) in Gujrat, India. His real name is Kashyap.  He was a hindu sage and philosopher who founded the philosophical school of Vaisheshika and authored the text Vaisesika Sutra.  He also wrote a book on his research “Vaisheshik darshan” and known as “The Father of Atomic Theory” Structure of an atom The study of the fundamental particles of an element and the type of bonding between different elements help to understand the physical, mechanical and chemical properties of the elements.  Atom is the smallest particle of an elements which cannot be further subdivided and take part in all chemical changes.  It has three Fundamental particles (Subatomic particles) i. e. Proton, Neutron and Electron.  Molecule – Smallest particle of substance may be an element or compound. It is made up two or more either similar or different kinds of atoms.  Size of Atom: An atom of an element is extremely small with a spherical shape having 1 x 10 -8 cm as its radius. Characteristics of Fundamental Particles of an Atom Characteristics Proton Neutron Electron Discovery Rutherford (1911) Chadwick (1932) J.J.Thomson (1898) Symbol P n e- Nature Positively charged Neutral Negatively charged Location in Atom Inside nucleus Inside Nucleus Outside nucleus Relative charge +1 0 -1 Mass in amu 1.007825 amu 1.008665 amu 0.0005466 amu Dr S K Mandavgade  The electrostatic force of attraction between the nucleus and electrons is balanced by the centrifugal force. Hence electron does not fall into the nucleus and atom remain same.  Atom is electrical neutral because it consist of equal number of positive charge proton and negative charge electron.  Nucleus of an atom is electrically positively charge because it consist of positive charge proton and neutral charge of neutron. Representation of an element Atomic Number (Z) : Proton (P) = Electron (e) Atomic Mass Number A) : Proton (P) + Neutron (N) Atomic Number:  It is defined as the number of proton present in the nucleus is equal to number of electron outside the nucleus.  It is represented by Z.  Atomic Number = Number of Proton = Number of electron Z = P = E  Atomic number are different for different elements.  Atomic numbers fixes the position of an element in the periodic table.  It does not indicate the mass of the nucleus of an atom. Atomic Mass Number:  It is defined as the sum of the number of proton and neutron present in the nucleus of an atom.  It is represented by A.  Atomic mass number = Number of Proton + Number of Neutron. A = P + N  Atoms of same or different elements may or may not have same atomic mass numbers. Ex. 8O16, 8O17, 8O18 18Ar 40, 40 19K , 20Ca 40  Atomic mass number does not fix the position of an element in the periodic table. Dr S K Mandavgade  It indicates the mass of the nucleus of an atoms of an elements.  Relation between Atomic Number and Atomic Mass number: N=A- Z  Number of neutrons = mass number – atomic number n = A - Z Ex. 1) Mass No=24 , Atomic No= 13 , Neutron No = ? 2) Proton No = 8 , electron No = ? 3 ) Proton No =10 , Atomic No = ? 4) Electron No = 14 , Atomic No = ? 5) Mass No = 36 , Neutron No = 12 , Atomic No =? 4.2 Electronic Theory of Valency Orbit – Electron revolved around the nucleus in fixed circular path. Max. number of electron in different orbit are calculated by the formula 2n2 Electronic Configuration- Distribution of electron in different orbit Ex Na 11 = ( 2, 8, 1 ) Mg 12 = (2,8,2) Al13 = ( 2,8,3) Ca20 = (2,8,8,2) Cl17 = (2,8,7) Si14= (2,8,4) Orbit – Electron revolved around the nucleus in fixed circular path. K, L, M, N Orbitals: The three dimensional region of space around the nucleus where the probability of finding the electron is maximum. S= 2, p= 6, d=10, f=14 But, each orbital contain maximum 2 electron with opposite spins. Dr S K Mandavgade  Energy level Total electrons 2n2 Total orbitals n2 K=1 2 1 L=2 8 4 M=3 18 9 N=4 32 16 Maximum electron capacity of different orbital's are s=2, p=6 , d=10, f=14  Valency electron- Electron present in the outermost orbit of an atom. Ex Na 11 = ( 2, 8, 1 ) , Ca 20 = ( 2, 8, 8,2 ) Valence electron = 01 Valence electron = 02  Kernel atom- If last orbit is removed then the rest of the atom is called Kernel of the atom. It possess the positive charge.  Stable configuration – The electronic arrangement of 8 e (octet) or 2e (duplet) in the last orbit of atom is known as stable configuration.  Element having less than 8e in the outermost orbit have tendency to take part in chemical combination. Dr S K Mandavgade  For Stability - 1. Element having less than 4 valence electron will loss of electron 2. Element having more than 4 valence electron will gain of electron. 3. Element having 4 valence electron will share of electron.  Valency – The number of electron loss, gain or share between the atoms in a molecule to complete its outermost shell.  Three modes of chemical combination- - By loss and gain of electrons. - By mutual sharing of electrons. - By one sided sharing of electrons. Valency – The valency obtained by the loss , gain or share of electrons between the atoms in a molecule to complete its outermost shell. Types of Valency - Electrovalency (Ionic valency) – ( loss and gain of electron) The charge on an ion is known as its electrovalency. The valency obtained by the loss , gain of electrons between the atoms in a molecule to complete its outermost shell. Ex. Na+ , Mg 2+ , Cl-, O2- Dr S K Mandavgade - Covalency - The valency obtained by the (sharing of electron) of an element to complete its last orbit. Types of electrovalency -Positive electrovalency - Number of electron of an element can loss to complete its last orbit. Na  Na+ + e [positive electrovalency = +1] (2,8,1) (2,8) Mg  Mg2+ + 2e [positive electrovalency = +2] (2,8,2) (2,8) Ca  Ca2+ + 2e [positive electrovalency = +2] (2,8,8,2) (2,8,8) -Negative electrovalency - Number of electron of an element can gain to complete its last orbit. Cl + e  Cl- [Negative electrovalency = -1] (2,8,7) (2,8,8) O + 2e  O2- [Negative electrovalency = -2] (2,6) (2,8) N + 3e  N3- [Negative electrovalency = -3] (2,5) (2,8) Chemical bond – The force of attraction between atoms in a molecule to hold them together is called chemical bond. Types of bond -  Ionic bond (electrovalent bond)  Covalent bond  Coordinate bond (Dative bond)  Hydrogen bond  Metallic bond Ionic bond (electrovalent bond) – The bond formed by loss and gain (transfer) of electron to complete its last orbit. Ionic Compound - The compound formed by loss and gain (transfer) of electron to complete its last orbit. Ex. NaCl, CaCl2, MgO, NaF, MgS, CaO, AlCl3, MgCl2 Dr S K Mandavgade  Characteristics of Ionic Compound Soluble in water, polar compound Conduct electricity High melting and boiling point Non directional Ionic reaction Covalency (sharing of electron) Dr S K Mandavgade Covalency (share of electron) - Number of electron of an element can share to complete its last orbit. Single covalency - Covalency obtained by sharing of one pair of electron (Two electron) Ex. H2, Cl2, H2O, NH3, CH4, HCl etc Double covalency - Covalency obtained by sharing of two pair of electron. (Four electron) Triple covalency - Covalency obtained by sharing of three pair of electron. (Six electron) Dr S K Mandavgade   Covalent bond - The bond formed by sharing of electron to complete its last orbit.  Covalent compound - The compounds formed by sharing of electron to complete its last orbit. Ex. Cl2, H2O, CO2, O2, N2, C2H2  Characteristics: Non polar compounds, Do not conduct electricity, Low Melting point and boiling point, Insoluble in polar solvent, Take part in molecular reaction Coordinate Bond (Dative bond) – Share electron pair are contributed by any one atoms. It is denoted by arrow  Ex. SO2 , NH4+ Ammonium ion , SO3, CO, N2O, HNO3, H2SO4 Properties of Coordinate Bond –  They have lone pair of electron of any one atom.  M.P. and B.P. are higher than purely covalent compounds and lower than purely ionic compounds.  These are sparingly soluble in polar solvent like water but readily soluble in non-polar solvents.  Like covalent compounds, these are also bad conductor of electricity. Hydrogen Bond -  Bond is formed by sharing of electron between hydrogen atom and electronegative atoms (F, N, O)  The shared electron pair lies more nearer to the electronegative atom.  So the hydrogen atoms possess fractional positive charge and electro native atom possess fractional negative charge. Ex. HF Dr S K Mandavgade Properties of Hydrogen Bond –  Hydrogen bond is weak than ionic bond and covalent bond.  Compound containing hydrogen bonding are partially polar.  They possess high melting and boiling points.  Compound contains always hydrogen atom. Types of Hydrogen Bonding – a) Intramolecular hydrogen bonding:  If positive and negative end develop within the same molecule, electrostatic forces of attraction results intra molecular hydrogen bonding.  In this bond hydrogen and electronegative elements both present in the same molecule.  The forces of attraction between the molecules situated at a distance less than molecular diameter are known as intermolecular forces of attraction. b) Intermolecular hydrogen bonding:  In this bond hydrogen atom and other electronegative atoms are present in different molecules of the same substance.  Ex. Water, HF,, NH3, alcohol and acids Dr S K Mandavgade Significance of Hydrogen Bonding –  The hydrogen bond form electrostatic forces between the positive end of one molecule and negative end of another molecule of the same substance.  It is a very weak bond.  Generally OoC to 4oC , hydrogen bonds continue to be broken and the molecules come closer and closer. This leads to contraction.  Above 4oC however the normal expansion effect take place due to rise in temperature. Hence the volume increases as the temperature rises.  Water exist in liquid state because of hydrogen bonding in it.  Hydrogen bonding play an important role in making wood fibers more rigid.  The stickiness of glue or honey is due to hydrogen bonding. Vander Waals –  The attraction and repulsion of intermolecular forces between molecules, surfaces and atoms are called as vander walls. Dipole –Dipole Interactions-  Dipole dipole interaction occurs only between polar molecules. They are caused by permanent dipoles between opposite charges on neighboring molecules.  Partially positive charges in one molecule lies up close together with the partial negative charges in another molecule forming a dipole dipole interaction.  These forces are weaker than covalent bonds. Vander Waals –  The attraction and repulsion of intermolecular forces between molecules, surfaces and atoms are called as vander walls. Dipole –Dipole Interactions-  Dipole dipole interaction occurs only between polar molecules. They are caused by permanent dipoles between opposite charges on neighboring molecules.  Partially positive charges in one molecule lies up close together with the partial negative charges in another molecule forming a dipole dipole interaction.  These forces are weaker than covalent bonds. Dr S K Mandavgade Metallic Bond:  The metallic bond is a type of chemical bond that occurs between atoms of metallic element.  A metallic bond is a sharing of delocalized electron between the atoms of a metal element.  Thus, metallic bonding is the electrostatic attractive force between valence electron and positive charged metal ions (kernel) in a metal element.  Ex. Metallic bond are seen in pure metals like sodium, aluminum, copper, silver, gold etc., Properties of metallic Bond –  Metallic bonds are non-directional.  A metallic bonds show the metallic properties.  They are mostly hard.  They possess high melting and boiling points.  They are very good conductor of both heat and electricity. Metallic Properties –  Metallic luster: The bright luster of metal is due to the presence of delocalized mobile electron.  Electrical conductivity: The presence of mobile electrons causes conductivity of a metal when a potential difference is applied across the metal sheet.  Thermal conductivity: When a part of the metal is heated, the kinetic energy of the electrons in that region increases. Dr S K Mandavgade  Malleability: It is property of metal that convert the metal into thin sheet without breaking.  Ductility: It is property of metal that convert the metal into thin wire without breaking.  (The position of adjacent layers of metallic kernels is altered without destroying the crystal.)  High tensile strength: Metal can resist stretching without breaking. ( Due to strong electrostatic attraction )  Hardness of metals: Hardness of metal is due to the strength of the metallic bond.  Opaqueness: The light that fall on metal is either reflected or completely absorbed by the delocalized electrons. 4.3 Molecular Arrangement  Matter has the unique property of existing in different states at different temperature and pressure. Solid, liquid and gas are the three states of matter. Water present in its three states as ice, water and vapors.  Arrangement in Solids,  Arrangement in liquids,  Arrangement in Gaseous. Arrangement in Solids –  Particles are closely packed close to each other.  Particles attract each other very strongly.  Particles of solids have fixed positions.  They have negligible intermolecular spaces. Arrangement in Liquids –  Particles are less tightly packed than in solids.  Particles in liquids can move about a little.  Particles of liquids do not have fixed positions.  They have bigger intermolecular spaces. Arrangement in Gases –  Particles are far apart from one another.  Particles collide with each other and move in any direction.  Particles have no intermolecular force of attraction.  They have large intermolecular spaces between their particles. 4.4 Structure of Solids Solids are held together by strong forces of attraction. Solids are substances characterized by definite shape, definite volume, non-compressibility, very slow diffusion, rigidity and mechanical strength. They exist in two forms - Dr S K Mandavgade Crystalline Solids: o It is formed by arrangement of atom, ion or molecules in three dimensional network or geometrical pattern. o Total intermolecular force of attraction is maximum. o They are in crystalline form and homogeneous. o The forces are responsible for stability of solid crystal. o They have sharp melting point. Ex. Sugar, NaCl, Sulphur Amorphous Solids: Amorphous solids are having random particle arrangement. They appears like solid but do not have perfectly crystalline structure. They are super cooled liquids having small structural units. They have stiffness due to their highly increased viscosities. They do not have sharp melting point. Ex. Rubber, plastic, glass Crystal Lattice and Unit Cell:  Crystal lattice is a highly ordered three dimensional structure formed by its constituent atoms or molecules or ions.  Unit cell is the smallest building unit of a crystal lattice which when repeated in different direction generate crystal substance.  Unit cell is simply a box with an atom at each corner.  Crystal lattice is actually a array of points.  Crystal structure may be conveniently specified by describing the arrangement within the solid of a small representative group of atoms or molecules called as unit cell. Dr S K Mandavgade Crystal lattice are described as- 1. Simple cubic or primitive structure 2. Body centered cubic (BCC) structure 3. Face centered cubic (FCC) structure 4. Hexagonal close packed (HCP) structure  Coordination Number: The Coordination number of an atom in a given molecule or a crystal refers to the total number of atoms, ion, or molecules bonded to the other.  Contribution of each atom 1. Simple cubic or primitive crystal lattice-  Simple cubic unit cell has atoms at each of the eight corners of a cube.  It has all sides equal, a=b=c and bond angles are α = β = γ = 90O  Each atom in this structure can form bond to its six nearest neighbors.  Therefore it is said to have a coordination number is 6.  The each atom is situated at the corner of each unit is shared by a total of 8 unit cells.  Each unit cell has 1/8 share of every corner atom.  So total contribution of all the eight corner atoms to each cell  = 1/8×8 = 1 atom/unit cell of SC.  The bond angles are α = β = γ = 90O and all sides equal a=b=c  Simple cubic structure is found in crystal lattice of Polonium metal NaCl and KCl molecules. 2. Body centered Cubic (BCC) structure - Dr S K Mandavgade  The BCC unit cell has atoms at each of the eight corners of a cube plus one atom in the centre of the cube.  Each particle is in directly in contact with 4 other particle in the layer above and 4 particle in the layer below.  So the coordination number of body centered cubic structure is 4+4 = 8.  It has all sides equal, a=b=c and bond angles are α = β = γ = 90O  The each atom is situated at the corner of each unit is shared by a total of 8 unit cells.  Each unit cell has 1/8 share of every corner atom.  So total contribution of all the eight corner atoms to each cell and centre atom = 1/8×8 + 1 = 2 atoms.  Such metals are usually harder and less malleable.  Metals having BCC structure are Li, K, Na, Cr, Ba, V, W. 3. Face Centered Cubic (FCC) structure -  The FCC structure has atoms located at each of the corners and the centers of all cubic faces.  Each of the corner atom is the corner of another cube, so the corner atoms are shared among eight unit cells.  Each of its six face centered atoms is shared with an adjacent four atom.  Since 12 of its atoms are shared, So coordination number is 12.  The atoms of FCC structure can pack closer together than in BCC structure.  Ex. Al, Cu, Ag, Pb, Ni, Pt and Silver.  The FCC structure has atoms located at each of the corners and the centres of all cubic faces.  The each atom is situated at the corner of each unit is shared by a total of 8 unit cells..  Each unit cell has 1/8 share of every corner atom.  Each face centre atom share ½ of every face atom.  So total contribution of all the eight corner atoms to each cell and each face centre atom  = 1/8×8 + ½×6 = 1+3 = 4 atoms. Dr S K Mandavgade 4. Hexagonal Close Packed (HCP) structure -  The HCP structure has three layers of atoms.  It has two sides of equal lengths and third side of different length.  There are six atoms that arrange themselves in the shape of a hexagon and a seventh atom that in the middle of the hexagon.  Each atom of HCP is shared with six another unit cells and three in the layer above and three in the layer below.  So the coordination number of HCP structure is 12.  Ex Beryllium, Cd, Mg, Titanium, Zn and Zirconium. Number of Number of Number of Type of Unit Cell atoms at atoms on atoms in Total corners faces center Simple Cubic 8 × 1/8 = 1 0 0 1 Body Centred 8 × 1/8 = 1 0 1 2 Cubic Face Centred 8 × 1/8 = 1 6 × 1/2 = 3 0 4 Cubic HCP 6 Coordination Type of Unit Cell Number Simple Cubic 6 Body Centred Cubic 8 Face Centred Cubic 12 Hexagonal structure 12 The End Dr S K Mandavgade