Chemistry Notes for NEET Chapter 3 PDF

60 E3 Chapter 3 Chemical Bonding The force which holds the atoms or ions together within the molecule is called a chemical bond and the process of their combination is called Chemical Bonding. It depends on the valency of atoms. 2 2 2 2 3 2 (1) Conditions for formation of electrovalent bond (i) The...

60 E3 Chapter 3 Chemical Bonding The force which holds the atoms or ions together within the molecule is called a chemical bond and the process of their combination is called Chemical Bonding. It depends on the valency of atoms. 2 2 2 2 3 2 (1) Conditions for formation of electrovalent bond (i) The atom which changes into cation (+ ive ion) should possess 1, 2 or 3 valency electrons. The other atom which changes into anion (–ve ion) should possess 5, 6 or 7 electrons in the valency shell. U Cause and Modes of chemical combination Some other examples are: MgCl , CaCl , MgO, Na S, CaH , AlF , NaH, KH, K 2O , KI, RbCl, NaBr, CaH etc. ID Atoms of different elements excepting noble gases donot have complete octet so they combine with other atoms to form chemical bond. D YG Chemical bonding takes place due to acquire a state of minimum energy and maximum stability and to convert atoms into molecule to acquire stable configuration of the nearest noble gas. We divide atoms into three classes, (1) Electropositive elements which give up one or more electrons easily. They have low ionisation potentials. (2) Electronegative elements, which can gain electrons. They have higher value of electronegativity. (3) Elements which have little tendency to lose or gain electrons. Different types of bonds are formed from these types of atoms. Type U Atoms involved ST A+B B+B A+A Electrovalent Covalent Coordinate H and electronegative element (F, N,O) Hydrogen       Na    r   r  is internuclear distance. The energy changes involved in the formation of ionic compounds from their constituent elements can be studied with the help of a thermochemical cycle called Born Haber cycle. H +IE 1/2Hdiss. sub       Cl        or Cl (g) Na  Cl  – EA – e– Na  (g)  1 Cl 2 (g) 2 + Na(g) An electrovalent bond is formed when a metal atom transfers one or more electrons to a non-metal atom.  Cl  (iv) Higher the lattice energy, greater will be the case of forming an ionic compound. The amount of energy released when free ions combine together to K form one mole of a crystal is called lattice energy (U). Lattice energy    ; r r Na(s) Electrovalent bond  (iii) There must be overall decrease in energy i.e., energy must be released. For this an atom should have low value of Ionisation potential and the other atom should have high value of electron affinity. Metallic Electrons deficient molecule or ion (Lewis acid) and electrons rich molecule or ion (Lewis base) Na  (ii) A high difference of electronegativity (about 2) of the two atoms is necessary for the formation of an electrovalent bond. Electrovalent bond is not possible between similar atoms. + +e – Cl  (g) H f Na Cl (s) + – (Crystal) –U (Lattice energy) (Born Haber Cycle) According to Hess's law of constant heat summation, heat of formation of an ionic solid is net resultant of the above changes. H f  H Subl.  1 H diss.  IE  EA  U 2 (2) Characteristics of electrovalent compounds nd NaF  NaCl  NaBr  NaI, MgO  CaO  BaO (iii) Electrovalent compounds are hard and brittle in nature. (iv) Electrovalent solids do not conduct electricity. While electrovalent compounds in the molten state or in solution conduct electricity. (v) Electrovalent compounds are fairly soluble in polar solvents and insoluble in non-polar solvents. (vi) The electrovalent bonds are non-rigid and non-directional. Thus these compound do not show space isomerism e.g. geometrical or optical isomerism. (vii) Electrovalent compounds furnish ions in solution. The chemical reaction of these compounds are known as ionic reactions, which are fast.      K  Cl   Ag NO 3  Ag Cl   K NO 3 E3 (ii) Electrovalent compounds possess high melting and boiling points. Order of melting and boiling points in halides of sodium and oxides of II group elements is as, 60 2 (i) Electrovalent compounds are generally crystalline is nature. The constituent ions are arranged in a regular way in their lattice.  (ii) Diamond, Carborandum (SiC), Silica (SiO ), AlN etc. have giant three dimensional network structures; therefore have exceptionally high melting points otherwise these compounds have relatively low melting and boiling points. (iii) In general covalent substances are bad conductor of electricity. Polar covalent compounds like HCl in solution conduct electricity. Graphite can conduct electricity in solid state since electrons can pass from one layer to the other. (iv) These compounds are generally insoluble in polar solvent like water but soluble in non-polar solvents like benzene etc. some covalent compounds like alcohol, dissolve in water due to hydrogen bonding. (v) The covalent bond is rigid and directional. These compounds, thus show isomerism (structural and space). (vi) Covalent substances show molecular reactions. The reaction rates are usually low. (vii) The number of electrons contributed by an atom of the element for sharing with other atoms is called covalency of the element. Covalency = 8 – [Number of the group to which element belongs]. The variable covalency of an element is equal to the total number of unpaired electrons in s, p and d-orbitals of its valency shell. The element such as P, S, Cl, Br, I have vacant d-orbitals in their valency shell. These elements show variable covalency by increasing the number of unpaired electrons under excited conditions. The electrons from paired orbitals get excited to vacant d-orbitals of the same shell. (Precipitate ) ID (viii) Electrovalent compounds show isomorphism. (ix) Cooling curve of an ionic compound is not smooth, it has two break points corresponding to time of solidification. 3 5 (3) The Lewis theory : The tendency of atoms to achieve eight electrons in their outermost shell is known as lewis octet rule. Lewis symbol for the representative elements are given in the following table, U (x) Ionic compounds show variable electrovalency due to unstability of core and inert pair effect. Four elements, H, N, O and F do not possess d-orbitals in their valency shell. Thus, such an excitation is not possible and variable valency is not shown by these elements. This is reason that NCl exists while NCl does not. Covalent bond D YG Covalent bond was first proposed by Lewis in 1916. The bond formed between the two atoms by mutual sharing of electrons so as to complete their octets or duplets (in case of elements having only one shell) is called covalent bond or covalent linkage. A covalent bond between two similar atoms is non-polar covalent bond while it is polar between two different atom having different electronegativities. Covalent bond may be single, double or a triple bond. We explain covalent bond formation by Lewis octet rule. Chlorine atom has seven electrons in the valency shell. In the formation of chlorine molecule, each chlorine atom contributes one electron and the pair of electrons is shared between two atoms. both the atoms acquire stable configuration of argon.       Cl * Cl    ( 2, 8 , 8 ) ( 2, 8 , 8 ) U **   *  Cl   Cl *  ** ( 2 , 8 , 7 ) ( 2, 8 , 7 ) ST Some other examples are : or Cl  Cl H 2 S , NH 3 , HCN , PCl3 , PH3, C2 H 2 , H 2 , C2 H 4 , SnCl 4 , FeCl3 , BH3 , graphite, BeCl 2 etc. (1) Conditions for formation of covalent bond (i) The combining atoms should be short by 1, 2 or 3 electrons in the valency shell in comparison to stable noble gas configuration. (ii) Electronegativity difference between the two atoms should be zero or very small. (iii) The approach of the atoms towards one another should be accompanied by decrease of energy. (2) Characteristics of covalent compounds (i) These exist as gases or liquids under the normal conditions of temperature and pressure. Some covalent compounds exist as soft solids. Group Lewis symbol 1 IA 2 IIA 13 IIIA X X  X 14 IVA 15 VA 16 VIA 17 VIIA  X    X    X   X    (4) Failure of octet rule : There are several stable molecules known in which the octet rule is violated i.e., atoms in these molecules have number of electrons in the valency shell either short of octet or more than octet. BeF2 , BF3 , AlH3 are electron- deficients (Octet incomplete) hence are Lewis acid. In PCl5 , P has 10 electrons in valency shell while in SF6 , S has 12 electrons in valence shell. Sugden introduced singlet linkage in which one electron is donated (Instead of one pair of electrons) to the electron deficient atom so that octet rule is not violated. This singlet is represented as (⇁). Thus, PCl5 and SF6 have structures as, Cl Cl P Cl Cl Cl F F S F F F F (5) Construction of structures for molecules and poly atomic ions : The following method is applicable to species in which the octet rule is not violated. (i) Determine the total number of valence electrons in all the atoms present, including the net charge on the species (n ). (ii) Determine n = [2 × (number of H atoms) + 8 × (number of other atoms)]. 1 2 (iii) Determine the number of bonding electrons, n , which equals n – n. No. of bonds equals n /2. (iv) Determine the number of non-bonding electrons, n , which equals n – n. No. of lone pairs equals n /2. (v) Knowing the central atom (you’ll need to know some chemistry here, math will not help!), arrange and distribute other atoms and n /2 bonds. Then complete octets using n /2 lone pairs. (vi) Determine the ‘formal charge’ on each atom. 3 1 Examples : CO, N O, H O , N O , N O , N O , HNO , NO 3 , SO , SO , 2 2 H SO , 2 2 2 2 3 2 4 2 5 3 SO 42  , SO 22  , 3 4 2 3 H 3 PO4 , H 4 P2O7 , 4 4 3 4 (vii) Formal Charge = [valence electrons in atom) – (no. of bonds) – (no. of unshared electrons)] (viii) Other aspects like resonance etc. can now be incorporated. Illustrative examples (i) CO 32  ; n1  4  (6  3)  2  24 [2 added for net charge] n 2 = (2 × 0) + (8 × 4) = 32 (no. H atom, 4 other atoms (1’C’ and 3 ‘O’) n 3 = 32 – 24 = 8, hence 8/2 = 4 bonds n 4 = 24 – 8 = 16, hence 8 lone pairs. Since carbon is the central atom, 3 oxygen atoms are to be arranged around it, thus, H 3 PO3 , Al2Cl6 (Anhydrous), O3 , SO 2Cl2 , SOCl 2 , HIO3 , HClO4 , HClO3 , CH 3 NC , N 2 H 5 , CH 3 NO 2 , NH 4 , [Cu(NH 3 )4 ]2 etc. Characteristics of co-ordinate covalent compound (1) Their melting and boiling points are higher than purely covalent compounds and lower than purely ionic compounds. (2) These are sparingly soluble in polar solvent like water but readily soluble in non-polar solvents. (3) Like covalent compounds, these are also bad conductors of electricity. Their solutions or fused masses do not allow the passage to electricity. (4) The bond is rigid and directional. Thus, coordinate compounds show isomerism. 60 3 Dipole moment E3 1 “The product of magnitude of negative or positive charge (q) and the distance (d) between the centres of positive and negative charges is called dipole moment”.  = Electric charge  bond length As q is in the order of 10 esu and d is in the order of 10 cm,  is in the order of 10 esu cm. Dipole moment is measured in “Debye” (D) unit. 1D  10 18 esu cm = 3.33  10 30 coulomb metre (In S.I. unit). –10 –8 –18 ID O | O  C  O , but total bonds are equal to 4. O | Hence, we get O  C  O. Now, arrange lone pairs to complete.O : |... octet : O  C  O : D YG.. (ii) CO 2 ; n = 4 + (6 × 2 ) = 16 1 n = (2 × 0) + (8 × 3) = 24 n = 24 – 16 = 8, hence 4 bonds n = 16 – 8 = 8, hence 4 lone-pairs Since C is the central atom, the two oxygen atoms are around to be arranged it thus the structure would be; O – C – O, but total no. of bonds 2 3 4 =4 Symmetrical polyatomic molecules are not polar so they do not have any value of dipole moment. H U.. Dipole moment is indicated by an arrow having a symbol ( ) pointing towards the negative end. Dipole moment has both magnitude and direction and therefore it is a vector quantity. O C.... U Co-ordinate covalent or Dative bond This is a special type of covalent bond where the shared pair of electrons are contributed by one species only but shared by both. The atom ST which contributes the electrons is called the donor (Lewis base) while the other which only shares the electron pair is known as acceptor (Lewis acid). This bond is usually represented by an arrow ( ) pointing from donor to the acceptor atom. BF molecule, boron is short of two electrons. So to complete its octet, it shares the lone pair of nitrogen in ammonia forming a dative bond. 3 H     F    H * N **  B   H     F   H     F      H * N   H     F  B     F    H | F    | | H F F H H 2 O H N Water H H 3 3 C H H  = 1.84D H H H Ammonia Methyl chloride  = 1.46D  = 1.86D   0 due to unsymmetry (1) Dipole moment is an important factor in determining the geometry of molecules. Table : 3.1 Molecular geometry and dipole moment General formula Molecular geometry AX Linear AX 2 AX 3 | F  H  N B  F H  = 0 due to symmetry Unsymmetrical polyatomic molecules always have net value of dipole moment, thus such molecules are polar in nature. H O, CH Cl, NH , etc are polar molecules as they have some positive values of dipole moments. Cl    C F.. octets, we get, : O  C  O : and thus final structure is : O  C  O :  O B Thus, O = C = O. After arrangement of lone pairs to complete.. F Dipole moment Example HF, HCl Linear Bent or V-shape May be non zero Zero Non zero Triangular planar Pyramidal T-shape Zero Non zero Non zero BF3 Tetrahedral Zero CO 2 ,CS 2 H 2O, NO 2 NH 3 , PCl3 ClF3 Formation of a co-ordinate bond between NH3 and BF3 AX 4 CH 4 ,CCl 4 Trigonal bipyramidal Square pyramidal Zero Non zero PCl5 AX 6 Octahedral Distorted octahedral Zero Non zero SF6 AX 7 Pentagonal bipyramidal Zero IF7 AX 5 XeF4 SF4 ,TeCl 4 (9) s -orbitals are spherically symmetrical and thus show only head on overlapping. On the other hand, p -orbitals are directionally concentrated and thus show either head on overlapping or lateral overlapping.Overlapping of different type gives sigma () and pi () bond. BrCl5 XeF6 (2) Every ionic compound having some percentage of covalent character according to Fajan's rule. The percentage of ionic character in compound having some covalent character can be calculated by the following equation. The % ionic character  Observed   100. Theoretica l  (3) The trans isomer usually possesses either zero dipole moment or very low value in comparison to cis–form H  C  Cl || H  C  Cl (8) Between two sub shells of same energy level, the sub shell more directionally concentrated shows more overlapping. Bond energy : 2s  2s < 2s  2 p < 2 p  2 p H  C  Cl || Cl  C  H Sigma () bond Pi () bond It results from the end to end overlapping of two s-orbitals or two p-orbitals or one s and one porbital. Stronger Bond energy 80 kcals More stable Less reactive Can exist independently It result from the sidewise (lateral) overlapping of two p-orbitals. 60 Zero Non zero The electron cloud is symmetrical about the internuclear axis. Fajan’s rule Always exist along with a -bond The electron cloud is above and below the plane of internuclear axis. Hybridization ID The magnitude of polarization or increased covalent character depends upon a number of factors. These factors are, (1) Small size of cation : Smaller size of cation greater is its polarizing power i.e. greater will be the covalent nature of the bond. (2) Large size of anion : Larger the size of anion greater is its polarizing power i.e. greater will be the covalent nature of the bond. (3) Large charge on either of the two ions : As the charge on the Less strong Bond energy 65 kcals Less stable More reactive E3 Square planar See saw U The concept of hybridization was introduced by Pauling and Slater. Hybridization is defined as the intermixing of dissimilar orbitals of the same atom but having slightly different energies to form same number of new orbitals of equal energies and identical shapes. The new orbitals so formed are known as hybrid orbitals. D YG ion increases, the electrostatic attraction of the cation for the outer electrons of the anion also increases with the result its ability for forming the covalent bond increases. (4) Electronic configuration of the cation : For the two ions of the same size and charge, one with a pseudo noble gas configuration (i.e. 18 electrons in the outermost shell) will be more polarizing than a cation with noble gas configuration (i.e., 8 electron in outer most shell). Valence bond theory or VBT ST U It was developed by Heitler and London in 1927 and modified by Pauling and Slater in 1931. (1) To form a covalent bond, two atoms must come close to each other so that orbitals of one overlaps with the other. (2) Orbitals having unpaired electrons of anti spin overlaps with each other. (3) After overlapping a new localized bond orbital is formed which has maximum probability of finding electrons. (4) Covalent bond is formed due to electrostatic attraction between radii and the accumulated electrons cloud and by attraction between spins of anti spin electrons. (5) Greater is the overlapping, lesser will be the bond length, more will be attraction and more will be bond energy and the stability of bond will also be high. (6) The extent of overlapping depends upon: Nature of orbitals involved in overlapping, and nature of overlapping. (7) More closer the valence shells are to the nucleus, more will be the overlapping and the bond energy will also be high. Characteristics of hybridization (1) Only orbitals of almost similar energies and belonging to the same atom or ion undergoes hybridization. (2) Hybridization takes place only in orbitals, electrons are not involved in it. (3) The number of hybrid orbitals produced is equal to the number of pure orbitals, mixed during hybridization. (4) In the excited state, the number of unpaired electrons must correspond to the oxidation state of the central atom in the molecule. (5) Both half filled orbitals or fully filled orbitals of equivalent energy can involve in hybridization. (6) Hybrid orbitals form only sigma bonds. (7) Orbitals involved in  bond formation do not participate in hybridization. (8) Hybridization never takes place in an isolated atom but it occurs only at the time of bond formation. (9) The hybrid orbitals are distributed in space as apart as possible resulting in a definite geometry of molecule. (10) Hybridized orbitals provide efficient overlapping than overlapping by pure s, p and d-orbitals. (11) Hybridized orbitals possess lower energy. How to determine type of hybridization : The structure of any molecule can be predicted on the basis of hybridization which in turn can be known by the following general formulation, H  1 (V  M  C  A) 2 Where H = Number of orbitals involved in hybridization viz. 2, 3, 4, 5, 6 and 7, hence nature of hybridization will be sp, sp , sp , sp d, sp d , sp d respectively. 2 3 3 3 2 3 3 V = Number electrons in valence shell of the central atom, M = Number of monovalent atom C = Charge on cation, O | C // \ O O Bond order  A = Charge on anion O || C   / \ O O O |  C / \\  O O 2 11  1.33 3 Bond characteristics In case of certain molecules, a single Lewis structure cannot explain all the properties of the molecule. The molecule is then supposed to have many structures, each of which can explain most of the properties of the molecule but none can explain all the properties of the molecule. The actual structure is in between of all these contributing structures and is called resonance hybrid and the different individual structures are called resonating structures or canonical forms. This phenomenon is called resonance. “The average distance between the centre of the nuclei of the two bonded atoms is called bond length”. It is expressed in terms of Angstrom (1 Å = 10 10 m) or picometer (1pm = 10 12 m). In an ionic compound, the bond length is the sum of their ionic radii ( d  r  r ) and in a covalent compound, it is the sum of their covalent radii (e.g., for HCl, d  rH  rCl ). Factors affecting bond length (i) The bond length increases with increase in the size of the atoms. For example, bond length of H  X are in the order, HI  HBr  HCl  HF. ID To illustrate this, consider a molecule of ozone O3. Its structure can be written as...... O O.O.......... O..O O. O.O......O...... (c )... (a ) (b ) (1) Bond length 60 The phenomenon of resonance was put forward by Heisenberg to explain the properties of certain molecules. E3 Resonance Resonance is shown by benzene, toluene, O , allenes 3 U As a resonance hybrid of above two structures (a) and (b. For simplicity, ozone may be represented by structure (c), which shows the resonance hybrid having equal bonds between single and double. (>C = C = 4 D YG C F – F > Br – Br > I – I, (Decreasing order of bond energy) Resonance increases bond energy. In case the central atom remains the same, bond angle increases with the decrease in electronegativity of the surrounding atom. (3) Bond angle In case of molecules made up of three or more atoms, the average Bond angle angle between the bonded orbitals (i.e., between the two covalent bonds) is known as bond angle . Valence shell electron pair repulsion theory (VSEPR ) The basic concept of the theory was suggested by Sidgwick and Powell (1940). It provides useful idea for predicting shapes and geometries of molecules. The concept tells that, the arrangement of bonds around the central atom depends upon the repulsion’s operating between electron pairs(bonded or non bonded) around the central atom. Gillespie and Nyholm developed this concept as VSEPR theory. Factors affecting bond angle (i) Repulsion between atoms or groups attached to the central atom may increase or decrease the bond angle. sp sp Bond angle 109º28 120° 3 60 (ii) In hybridisation as the s character of the s hybrid bond increases, the bond angle increases. Bond type The main postulates of VSEPR theory are sp 2 PCl3 PBr3 PI3 , AsCl3 AsBr3 AsI3 100 o 101.5 o 102 o 98.4 o 100.5 o 101o (1) For polyatomic molecules containing 3 or more atoms, one of the 180° E3 atoms is called the central atom to which other atoms are linked. (2) The geometry of a molecule depends upon the total number of valence shell electron pairs (bonded or not bonded) present around the central atom and their repulsion due to relative sizes and shapes. (3) If the central atom is surrounded by bond pairs only. It gives the symmetrical shape to the molecule. (4) If the central atom is surrounded by lone pairs (lp) as well as (iii) By increasing lone pair of electron, bond angle decreases approximately by 2.5%. Bond angle NH 4 109º 107 HO 3 2 105 o o bond pairs (bp) of e  then the molecule has a distorted geometry. (5) The relative order of repulsion between electron pairs is as ID CH (iv) If the electronegativity of the central atom decreases, bond angle decreases.  H 2S 92.2 o  H 2 Se 91.2 o  H 2 Te 89.5 o greater when a lone pair is involved. U H 2O 104.5 o Bond angle follows : lp – lp > lp – bp > bp – bp. A lone pair is concentrated around the central atom while a bond pair is pulled out between two bonded atoms. As such repulsion becomes Table : 3.2 Geometry of Molecules/Ions having bond pair as well as lone pair of electrons No. of lone pairs of electrons Hybridization Bond angle Expected geometry AX 3 2 1 sp 2 < 120 Trigonal planar AX 4 2 2 sp 3 < 109 28 Tetrahedral AX 4 3 1 sp 3 < 109 28 AX 5 4 1 sp 3 d < 109 28 AX 5 3 2 sp 3 d 90 AX 6 AX 7 o o o o o Actual geometry Examples V-shape, Bent, Angular V-shape, Angular H O, H S, SCl , OF , NH , ClO Tetrahedral Pyramidal NH , NF , PCl, PH, AsH, ClO , H O Trigonal bipyramidal Trigonal bipyramidal Trigonal bipyramidal Irregular tetrahedron SF , SCl , TeCl T-shaped ICl , IF , ClF Linear XeF , I , ICl SO , SnCl , NO 2 2 – 2 2 2 2 2 3 3 3 4 3 4 3 3 5 1 sp 3 d 2 < 90 Octahedral Square pyramidal ICl , BrF , IF 2 3 2 – Octahedral Square planar XeF , ICl 3 3 – Pentagonal pyramidal Distorted octahedral XeF 4 6 1 sp d sp d Molecular orbital theory or MOT Molecular orbital theory was given by Hund and Mulliken in 1932. The main ideas of this theory are, (1) When two atomic orbitals combine or overlap, they lose their identity and form new orbitals. The new orbitals thus formed are called molecular orbitals. + 3 3 180 – 2 3 4 sp 3 d o – 2 – 3 3 o – 2 2 ST AX 6 U AX 5 D YG Type of mole-cule No. of bond pairs of electron 3 – 2 5 5 4 5 – 4 6 (2) Molecular orbitals are the energy states of a molecule in which the electrons of the molecule are filled just as atomic orbitals are the energy states of an atom in which the electrons of the atom are filled. (3) In terms of probability distribution, a molecular orbital gives the electron probability distribution around a group of nuclei just as an atomic orbital gives the electron probability distribution around the single nucleus. (4) Only those atomic orbitals can combine to form molecular orbitals which have comparable energies and proper orientation. (5) The number of molecular orbitals formed is equal to the number of combining atomic orbitals. (6) When two atomic orbitals combine, they form two new orbitals called bonding molecular orbital and antibonding molecular orbital. (7) The bonding molecular orbital has lower energy and hence greater stability than the corresponding antibonding molecular orbital. (8) The bonding molecular orbitals are represented by  ,  etc, whereas the corresponding antibonding molecular orbitals are represented Increasing energy (for electrons > 14) (b)  1s,  1s,  2 s,  * 2 s,  2 p y ,  2 p x , *  * 2pz Increasing energy (for electrons  14) (12) Number of bonds between two atoms is called order and is given by  N  NA  Bond order   B  2   by  * ,  * etc. (11) Electrons are filled in the increasing energy of the MO which is in order 60 where N B  number of electrons in bonding MO. N A  number of electrons in antibonding MO. For a stable molecule/ion, N B  N A (13) Bond order  Stability of molecule  Dissociation energy  1. Bond length E3 (9) The shapes of the molecular orbitals formed depend upon the type of combining atomic orbitals. (10) The filling of molecular orbitals in a molecule takes place in accordance with Aufbau principle, Pauli's exclusion principle and Hund's rule. The general order of increasing energy among the molecular orbitals formed by the elements of second period and hydrogen and their general electronic configurations are given below. (14) If all the electrons in a molecule are paired then the substance is a diamagnetic on the other hand if there are unpaired electrons in the molecule, then the substance is paramagnetic. More the number of unpaired electron in the molecule greater is the paramagnetism of the substance. * (a)  1s,  *1s,  2 s,  * 2 s,  2 p x ,  2 p y  * 2 py ,  * 2 p x  2 pz ,  * 2 pz  (2pz) ID *(2pz) 2p *(2py) 2p *(2px) U  (2pz) 2s D YG Increasing energy  (2py)  (2px) 2s *(2s) Increasing energy 2p 2s U  (1s) Atomic orbitals Molecular orbitals Atomic orbitals ST Molecular orbital energy level diagram (Applicable for elements with Z > 7) Hydrogen bonding In 1920, Latimer and Rodebush introduced the idea of “hydrogen bond”. For the formation of H-bonding the molecule should contain an atom of high electronegativity such as F, O or N bonded to hydrogen atom and the size of the electronegative atom should be quite small. Types of hydrogen bonding (1) Intermolecular hydrogen bond : Intermolecular hydrogen bond is formed between two different molecules of the same or different substances. (i) Hydrogen bond between the molecules of hydrogen fluoride. 2p *(2px)  (2pz) 2s *(2s)  (2s) 1s 1s *(1s) *(2py)  (2py)  (2px)  (2s) 1s bond 1s *(1s)  (1s) Atomic orbitals Molecular orbitals Atomic orbitals Molecular orbital energy level diagram obtained by the overlap of 2s and 2pz atomic orbitals after mixing (Applicable for elements with Z < 7) (ii) Hydrogen bond in alcohol and water molecules (2) Intramolecular hydrogen bond (Chelation) Intramolecular hydrogen bond is formed between the hydrogen atom and the highly electronegative atom (F, O or N) present in the same molecule. Intramolecular hydrogen bond results in the cyclisation of the molecules and prevents their association. Consequently, the effect of intramolecular hydrogen bond on the physical properties is negligible. For example : Intramolecular hydrogen bonds are present in molecules such as o-nitrophenol, o-nitrobenzoic acid, etc. O H | | C C O O H O H O O H N N O | | Salicyldehyde O O (o-hydroxy benzaldehyde) Ortho nitrophenol Ortho nitrobenzoic acid Hydrogen bond helps in explaining the abnormal physical properties in several cases. Some of the properties affected by H-bond are given below, (1) Dissociation : In aqueous solution, hydrogen fluoride dissociates and gives the difluoride ion (HF2 ) instead of fluoride ion (F  ). This is due to H-bonding in HF. This explains the existence of KHF2. H-bond formed is usually longer than the covalent bond present in the molecule (e.g. in H O, O–H bond = 0.99 Å but H-bond = 1.77 Å). 2 (7) Explanation of lower density of ice than water and maximum density of water at 277K : In case of solid ice, the hydrogen bonding gives rise to a cage like structure of water molecules as shown in following figure. As a matter of fact, each water molecule is linked tetrahedrally to four other water molecules. Due to this structure ice has lower density than water at 273K. That is why ice floats on water. On heating, the hydrogen bonds start collapsing, obviously the molecules are not so closely packed as they are in the liquid state and thus the molecules start coming together resulting in the decrease of volume and hence increase of density. This goes on upto 277K. After 277 K, the increase in volume due to expansion of the ID (2) Association : The molecules of carboxylic acids exist as dimers because of the hydrogen bonding. The molecular masses of such compounds are found to be double than those calculated from their simple formulae. For example, molecular mass of acetic acid is found to be 120. (6) The substances which contain hydrogen bonding have higher viscosity and high surface tension. 60 Effects of hydrogen bonding (5) As the compounds involving hydrogen bonding between different molecules (intermolecular hydrogen bonding) have higher boiling points, so they are less volatile. E3 The extent of both intramolecular and intermolecular hydrogen bonding depends on temperature. The high melting points and boiling points of the compounds (H 2 O, HF and NH 3 ) containing hydrogen bonds is due to the fact that D YG some extra energy is needed to break these bonds. (4) Solubility : The compound which can form hydrogen bonds with the covalent molecules are soluble in such solvents. For example, lower alcohols are soluble in water because of the hydrogen bonding which can take place between water and alcohol molecules as shown below,     liquid water becomes much more than the decrease in volume due to breaking of H-bonds. Thus, after 277 K , there is net increase of volume on heating which means decrease in density. Hence density of water U (3) High melting and boiling point : The compounds having hydrogen bonding show abnormally high melting and boiling points.  is maximum 277 K. H 0.90 Å (99 pm) H H H U The intermolecular hydrogen bonding increases solubility of the compound in water while, the intramolecular hydrogen bonding decreases. N | O O o- Nitrophenol Due to chelation, – OH group is not available to form hydrogen bond with water hence it is sparingly soluble in water. H H O H O H H O H H O H H H O H – O …… H – O – H H ST O Vacant spaces O O  C2 H 5 H H H H  O................ H  O............... H  O C2 H 5 1.77 Å (177 pm) O H O H Cage like structure of H2O in the ice ON=O p- Nitrophenol – OH group available to form hydrogen bond with water, hence it is completely soluble in water.  A chemical bond is expected to be formed when the energy of the aggregate formed is about 40 kJ mole lower than the separate particles. –1 60  Formation of a chemical bond is always an exothermic process.  Lattice energies of bi-bivalent solids > bi-univalent solids > uniunivalent 2 solids. 2 For example, 1 2 lattice energy  of 1 Mg O (3932 kJ mole )  Ca (F )2 (2581 kJ mole )   When co-ordination number increases, the coulombic forces of attraction increases and hence stability increases.  Ionic solids have negative vapour pressure. ID  As a general rule, atomic crystals are formed by the lighter E3 Li  F  (1034 kJ mole 1 ). elements of the middle columns of the periodic table.  FeCl3 is more covalent than FeCl2 because polarising power of covalent than SnCl 2. U Fe 3  is more than that of Fe 2 . Similarly SnCl 4 is more  Boron forms the maximum number of electron deficient D YG compounds than any other elements in the periodic table.  Roughly each lone pair decreases the bond angle by 2.5°.  Greater the number of the lone pairs at the two bonding atoms, greater is the repulsion between them and weaker is the bond.  The actual number of s- and p-electrons present in the outermost shell of the element is called maximum covalency of that atom.  The hydrogen bonds are tetrahedral in their directions and not planar. U  The hydrogen bond is stronger in HF and persists even in vapour ST state. Such bonds account for the fact that gaseous hydrogen fluoride is largely polymerised into the molecular species H 2 F2 , H 3 F3 , H 4 F4 , H 5 F5 and H 6 F6.  Hydrogen bonding is strongest when the bonded structure is stabilised by resonance.  Critical temperature of water is higher than that of O 2 because H 2 O molecule has dipole moment.