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ISC Class XI Notes : Chemistry (J.S.S. International School (JSS IS), Dubai) Chemical Bonding

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JSS INTERNATIONAL SCHOOL, DUBAI CHEMISTRY NOTES - GRADE 11 CHEMICAL BONDING WHY DO ATOMS BOND? Tendency to acquire noble gas configuration. An octet in the outermost shell/ noble gas configuration gives stability. Therefore elements that do not have noble gas configuration tend to acquire it by losing, gaining or sharing electrons. Tendency to acquire minimum energy There is a force of attraction between the nucleus and the electrons and repulsive forces between electrons. When there is a net increase attractive force, the energy decreases. Decrease in energy makes molecules more stable. TYPES OF CHEMICAL BONDS ELECTROVALENT BOND The attractive force that holds oppositely charged ions together is called an ionic bond or electrovalent bond. They are formed by the complete transfer of one or more electrons from a metal to a nonmetal. Electrovalency The number of electrons an atom loses or gains while forming an ionic bond is called electrovalency. Formation of NaCl: Sodium [Ne]3s1 has one electron in its outermost shell, which it loses to form Na+ ion with a noble gas configuration. Cl [Ne]3s23p5 has five electrons in its outermost shell. Gaining the electron from sodium, it acquires the configuration of its nearest noble gas, Argon. The positively charged Na+ and the negatively charged Cl- ions are held together by electrostatic forces. The electron dot structure can be represented as [Ne]3s1 [Ne]3s23p5 Electron dot structures of other compounds: Magnesium Oxide: [Ne]3s2 [He]2s22p4 Sodium sulphide (Na2S) Draw the electron dot structures of Li2O, CaO, and MgF2 Factors influencing the formation of ionic bond: 1. Low ionization energy of the metal Lesser the ionization energy, greater is its tendency to lose one or more electrons and form a cation. Alkali metals form cations most readily due to their low ionization potentials. 2. High electron affinity of the nonmetal Atoms having high electron affinity accept electrons more readily. Greater the electron affinity, greater is the ease of formation of anion. Hence, halogens form ionic compounds more readily than other nonmetals. 3. High lattice energy Lattice energy is the amount of energy released when one mole of a crystalline ionic compound is formed by the close packing of gaseous anions and cations . When the attraction between the cations and anions is greater, the lattice energy is greater. This in turn results in greater stability of the ionic compound. 4. Electronegativity difference between the metal and nonmetal A large electronegativity difference between the combining atoms favours ionic bond formation. Generally, an electronegativity difference of about 2 favours ionic bond formation. 5. Size of ions and the charge on them Smaller sizes and higher charges form stronger ionic bonds. Relation between periodic table and formation of ionic bonds Elements of groups 1 and 2 form cations more readily due to their low ionization energies and elements of groups 16 and 17 form anions more easily due to their high electron affinity. Hence, compounds between elements of groups 1 and 2 with groups 16 and 17 will be ionic in nature. Variable electrovalency: Some elements exhibit different electrovalencies in different compounds. This is called variable electrovalency. Reasons for variable electrovalency: 1. Inert pair effect: In elements like Ga, Sn, Pb etc, the inner ns2 electrons donot participate in bonding. They remain inert. This is called Inert Pair Effect . Eg. a. The electronic configuration of Gallium (Ga) is [Ar]3d10 4s2 4p1 . It loses the 4p electron to form Ga+. It can also lose the 4s electrons, if sufficient energy is available forming Ga3+. b. The electronic configuration of Lead (Pb) is [Xe] 4f14 5d10 6s2 6p2. It loses the 6p electron to form Pb2+. It can also lose the 6s electrons, if sufficient energy is available forming Pb4+. However, Pb2+ is more predominant, due to inert pair effect. 2. Unstable electronic configuration of the core When electrons are lost from the outermost shells of atoms, sometimes, unstable core is formed. To gain more stability, these ions lose more electrons. Eg: a. Electronic configuration of Fe is [Ar]3d64s2. When it forms Fe2+, its electronic configuration becomes [Ar]3d6. Losing another electron results in d5 configuration, which is more stable (half-filled orbitals give extra stability). Therefore, Iron exhibits +2 and +3 electrovalencies. b. Electronic configuration of Cu is [Ar]3d104s1. When it forms Cu+, its electronic configuration becomes [Ar]3d10. This d10 configuration is more stable. However if sufficient energy is available, another electron can be lost forming Cu+2. Therefore, copper exhibits +1 and +2 electrovalencies. COVALENT BOND Bond formed by mutual sharing of electrons between atoms is called covalent bond. Conditions favouring formation of covalent bond: 1. High ionization energy Elements with high ionization energy cannot lose electrons easily. Hence, they show greater tendency to share electrons to attain octet. 2. Comparable or similar electron affinities 3. Equal electronegativity or less electronegativity difference For electrons to be shared, the combining atoms should have almost the same electronegativity so that electrons don t get transferred from one atom to another. The elements having high electronegativity form covalent bonds. 4. Number of valence electrons Elements with 4 to 7 valence electrons form covalent bonds with each other or themselves. Relation between periodic table and formation of covalent bonds Elements of groups 14 to 17 form covalent bonds with each other. They are placed towards the right of the periodic table. Covalency: The number of electrons an atom shares with other atoms in a covalent compound is called covalency. TYPES OF COVALENT BONDS Single bond: It involves one pair (two) bonding electrons between atoms. A single bond is represented by one line between the atoms. Eg: In hydrogen molecule, the two hydrogen atoms are bonded by a single bond. Double bond: The two atoms are bonded by two pairs of electrons. A double bond is represented by two lines joining the two atoms. Eg: In oxygen molecule, the two oxygen atoms are held by two pairs of electrons, resulting in a double bond between the two atoms. Triple bond: The two atoms are bonded by three pairs of electrons. A triple bond is represented by three lines joining the two atoms. Eg: In a nitrogen molecule, the two nitrogen atoms are held by a triple bond. SIGMA AND PI BONDS Sigma bond Pi bond Formed by axial overlap of atomic orbitals. Formed by lateral or side-ways overlap of orbitals. Stronger due to maximum overlap. Weaker as extent of overlap is less. Formed between s-s, s-p or p-p orbitals. Formed between p-p orbitals. Contains one electron cloud symmetrical Contains two electron clouds, one above about the internuclear axis. and one below the plane of the nuclei. POLAR AND NONPOLAR COVALENT BONDS When two atoms with equal electronegativity bond with each other, the shared electron pair is attracted equally by the two atoms. Hence, the electron pair remains in the centre. Such a bond is nonpolar and the compounds are called nonpolar covalent compounds. Eg: H2, N2 etc When atoms sharing electrons have different electronegativities, the element with greater electronegativity attracts the shared pair of electrons more strongly than the other. The bond pair of electron move towards the greater electronegative atom, resulting in it having a slight negative charge and the other, a slight positive charge. Such a compound is called a polar covalent compound. Eg: HCl Note: Compounds like CO2, CCl4, CH4 etc are nonpolar because the polarity of one bond is cancelled by the other. DIPOLE MOMENT The product of the magnitude of charge and the distance between the positive ad negative centres in a polar molecule is called dipole moment. Dipole moment, = q x d Unit: D (Debye) Dipole moment is a vector. When there is more than one bond in a molecule, the net dipole moment is the sum of the dipole moments of the individual bonds. In compounds having regular geometry, the dipole moment due to one bond cancels the effect of the others. Therefore, net dipole moment is zero. ie., the molecule is nonpolar. ELECTRON DOT STRUCTURES FOR COVALENT COMPOUNDS: 1. Methane (CH4) 2. Water (H2O) 3. Ammonia (NH3) 4. Ethyne Draw electron dot structures for ethane, ethyne and carbon dioxide following the same method. Resonance structures of ozone, carbon dioxide, carbonate ion and nitrate ion Ozone Carbon dioxide Carbonate ion Nitrate ion Variable Covalency: Some elements exhibit different covalency in different compounds. This is called variable covalency. Eg: Phosphorous exhibits covalency of 3 and 5, Chlorine exhibits covalency of 1, 3, 5 and 7. Reason for variable covalency: It is due to the increase in the number of unpaired electrons in the excited states of atoms. Explanation of variable covalency of Chlorine: The electronic configuration of chlorine is 2 2 6 2 5 1s 2s 2p 3s 3p (2.8.7) [Ne]3s 3p There is one electron in the outermost shell of Cl. Therefore it exhibits a covalency of 1.In the second excited state, one of the 3p electrons moves to the 3d orbital and the number of unpaired electrons becomes 3 ([Ne]3s23px2py1pz1 3d1), resulting in a covalency of 3. In the next two excited states, the configuration becomes [Ne]3s23px1py1pz1 3d2 (both d electrons are unpaired) and ([Ne]3s13px1py1pz1 3d3), resulting in covalencies of 5 and 7 respectively. The variable covalencies of S and P can be explained in a similar way. Deviation from octet rule: Deviation from octet rule is observed when a. Atoms have less electrons than inert gas configuration (incomplete octet) Elements like Be and B have two and three electrons in their outermost shells respectively. Even after sharing their valence electrons, they cannot attain an octet. Eg: BeCl2 and BF3 b. Atoms have more electrons than inert gas configuration (exceeding octet) Elements like S, P, Cl etc have more than eight electrons around them in their compounds like SF6, PCl5 etc. This is due to the availability of d-orbitals so that their s and p electrons get excited, resulting in more number of unpaired electrons. Fajan s rule: When cations and anions approach each other, the valence shell of anions are pulled towards the cation s nucleus due to electrostatic attraction and thus shape of the anion is deformed. This phenomenon of deformation of anion by a cation is known as polarization and the ability of cation to polarize a nearby anion is called as polarizing power of cation. Fajan points out that greater is the polarization of anion in a molecule, more is covalent character in it. This is Fajan's rule. According to Fajan s rule, 1. maximum polarization of anions takes place when a. Cation has a small size and anion has a large size b. Charge on either the anion or cation is large Greater polarization results in greater covalent character. 2. A cation with 18 valence electrons gives covalent compounds while that with 8 valence electrons gives ionic compounds. COORDINATE OR DATIVE BOND Bond formed when both the shared electrons are donated by one atom. Formation of ammonium ion: Formation of hydronium ion: Structure of nitric acid: O H O N=O Structure of ozone: Structure of hypochlorous acid (HClO): Structure of chlorous acid (HClO2): Or Structure of chloric acid (HClO3): Or Structure of perchloric acid (HClO4): Or Hydrogen Bonding: A weak bond formed between an electronegative atom having a partial negative charge and hydrogen having a partial positive charge is called hydrogen bond. Types of H- bonding: Intermolecular hydrogen bonding: Intermolecular hydrogen bond is formed between the hydrogen of one molecule and the highly electronegative atom of another molecule of the same or different substance. Eg: HF, H2O Ice Cage-like structure Intramolecular H-bond: If the hydrogen bond exists between the hydrogen atom and highly electronegative atom of the same molecule, the hydrogen bond is called an intramolecular hydrogen bond. Metallic bond: Metals have a low ionization energy. The outermost electrons are loosely bound. They can move freely in the metal. The positive metal ions called kernels are arranged in a regular order and the electrons move around the metal ions like sea water. This is hence called the electron sea model. Explanation of properties of metals using electron sea model: Thermal and electrical conductivity: When a metal is heated, the free electrons get extra energy which is passed to the neighbouring electrons. This process continues from the heated end of the metal to the other. When a potential difference is created across the ends of a metal, the electrons move towards the positive electrode. This flow of electrons results in electrical conductivity. Metallic lustre: When light is incident on a metal, the free electrons absorb the energy and then transmit it. This is why a metal has metallic lustre. Since the incident light is reflected and does not pass through the metal, a metal is opaque. Malleability and ductility: When force is applied on a metal, the positively charged metal ion layers slide on each other. Band model Metallic bond: The atomic orbitals of a metal combine to form molecular orbitals which are so close to each other in energy that they form a band The valence electrons of the metal are in the valence band. If the band is partially filled or is overlapping with a higher energy conduction band, electrons can easily flow easily. This results in conduction. In insulators, there is a gap between the valency and conduction bands. Hence the electrons cannot jump to it. In semiconductors, there is a very small gap between the two bands. Therefore they show conductivity under certain conditions. HYBRIDIZATION The combination of two or more atomic orbitals to form an equal number of hybrid orbital having the same energy is called hybridization. sp3 hybridization: one s-orbital and three p-orbitals combine to form 4 sp3 hybrid orbitals of the same energy. Eg: Ethane The electronic configuration of C in ground state is [He] 2s2 2px1 2py1. In the excited state, it becomes [He] 2s1 2px1 2py12pz1. The singly occupied 2s and three 2p orbitals combine to form four sp3 hybrid orbitals of equal energy. These orbitals are oriented towards the corners of a regular tetrahedron. The sp3 hybrid orbital of one C overlaps with the other carbon in ethane forming a sigma bond. The remaining three sp3 hybrid orbitals of each carbon overlap with the 1s atomic orbital of hydrogen to form sigma bonds. The structure of the molecule is shown below. Bond angle = 190 28 Geometry tetrahedral sp2 hybridization: One s-orbital and two p-orbitals combine to form 3 sp2 hybrid orbitals of the same energy. Eg: Ethene The electronic configuration of C in ground state is [He] 2s2 2px1 2py1. In the excited state, it becomes [He] 2s1 2px1 2py12pz1. The singly occupied 2s and two 2p orbitals combine to form three sp2 hybrid orbitals of equal energy. These orbitals form a trigonal structure. The sp2 hybrid orbital of one C overlaps with the other carbon in ethene forming a sigma bond. The remaining two sp2 hybrid orbitals of each carbon overlap with the 1s atomic orbital of hydrogen to form sigma bonds. The unhybridized p-orbital on the two carbon atoms overlap side-ways forming a pi bond. The structure of the molecule is shown below. Bond angle: 120 Structure: Trigonal planar sp hybridization: One s-orbital and one p-orbital combine to form 2 sp hybrid orbitals of the same energy. Eg: Ethyne The electronic configuration of C in ground state is [He] 2s2 2px1 2py1. In the excited state, it becomes [He] 2s1 2px1 2py12pz1. The singly occupied 2s and one 2p orbitals combine to form two sp hybrid orbitals of equal energy. These orbitals form a linear structure. The sp hybrid orbital of one C overlaps with the other carbon in ethyne forming a sigma bond. The remaining sp hybrid orbitals of each carbon overlap with the 1s atomic orbital of hydrogen to form sigma bonds. The unhybridized p-orbitals on the two carbon atoms overlap side-ways forming two pi bonds. The structure of the molecule is shown below. Bond angle: 180 Structure: linear sp3d hybridization: Eg: PCl5 The electronic configuration of P is [Ne]3s2 3px1 3py1 3pz1. In the excited state, one of the s-electrons moves to the 3d orbital. The configuration becomes The five singly occupied orbitals (one s, three p and one d) orbitals combine to form 5 sp3d hybrid orbitals oriented along the corners of a triagonal bipyramid. Each of these 5 hybrid orbitals overlaps with a p-orbital of chlorine to form five sigma bonds. Bond angle: The orbitals along the triangle form an angle of 120 and the orbitals above and below the plane form an angle of 90 with the plane. Structure: trigonal bipyramidal sp3d2 hybridization: Eg: SF6 The electronic configuration of S atom is [Ne]3s2 3px2 3py1 3pz1. In the excited state, the paired electrons from the 3s and 3p orbitals move to the vacant 3d orbitals. The configuration becomes The six singly occupied orbitals (one 3s, three 3p and two 3d orbitals) hybridize to form six sp3d2 hybrid orbitals. The structure is octahedral. Each of these hybrid orbitals overlaps with a p-orbital of F forming 6 sigma bonds. The bond angle is 90 . VALENCE SHELL ELECTRON PAIR REPULSION THEORY (VSEPR THEORY) The electrons in the valence shell of atoms repel each other. A molecule assumes certain geometry to minimize these repulsions. The repulsion increases in the order given below. Bond pair bond pair < lone pair bond pair < lone pair lone pair These repulsions produce distortion in structures where lone pairs are involved. Examples: 1. Ammonia The nitrogen atom in ammonia undergoes sp3 hybridization. Hence, ammonia molecule should have a tetrahedral structure with a bond angle of 109 28 . But ammonia molecule has a pyramidal structure with a bond angle of 107 . In ammonia molecule, one of the corners of the tetrahedron is occupied by a lone pair. In the molecule, there will be lone pair bond pair and bond pair bond pair repulsions. The lone pair bond pair repulsions are greater. Therefore, to minimize the repulsion, the bond pairs are pushed closer to each other reducing the bond angle to 107 , resulting in a pyramidal shape. 2. Water Similar to nitrogen atom in ammonia, the oxygen atom in water undergoes sp3 hybridization. Hence, water molecule should have a tetrahedral structure with a bond angle of 109 28 . But water molecule has an angular structure with a bond angle of 104.5 . In water molecule, two of the corners of the tetrahedron are occupied by lone pairs. In the molecule, there will be lone pair lone pair, lone pair bond pair and bond pair bond pair repulsions. The lone pair lone pair repulsions are greatest. Therefore, to minimize the repulsion, the lone pairs move away from each other. This brings the lone pairs closer to the bond pairs. To minimize the lone pair bond pair repulsions, the bond pairs are pushed closer to each other reducing the bond angle further to 104.5 and an angular shape.

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Additional Info : ISC Class XI Notes 2020 : Physics (Dhirubhai Ambani International School (DAIS), Mumbai)

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