Chapter 03: Chemical Bonding

(i) Dipole A molecule with St charge on one part and 8- change on other part is called dipole, e.g., HF, HCC, HBr (ii) Bond order Number of bonds formed between two atoms after the atomic orbitals overlaps, is called bond order. It is taken as half of the difference between the number of bonding electrons (say a) and antibonding electrons (say b) Bond order = 9-b (wii) Permanent Dipole-Permanent dipole force The negative end of one permanent dipole attracts the positive end of other permanent dipole molecule. This force of attraction is called permanent dipole-permanent dipole force e.g., St 8- CI Permanent dipole-permanent dipole force (iv) London dispersion force Momentary force of attraction between instantaneous dipole molecule and induced dipole molecule is called London dispersion force. It is present in all type of molecules whether polar or non-polar.

(i) Xenon trioxide has a trigonal pyramidal shape due to lone pair-bond pair repulsion. (ii) X e

See Q.8 of theory.

See Q.13,14,15 of theory. nucleus, which in turn exerts a stronger repulsion over bonding electrons. Thus, angle further shrinks to 102°. Moreover, the bond pairs N-F bonds are closer to F atom than N atom. The increased distances in these bond pairs makes their repulsion less operative. Intermolecular forces between

(i) 2p, sp sp' 2p, Figure: Hybridized structure of acetone CH3 CH3 (i) Yes propanone can make hydrogen bond with water when both are intermixed as shown below CH3 ••H CH3

See Q.9,10,11 of theory. the molecular orbital

SO-? According to VSEPR Theory, the shape of SO4? ion is tetrahedral. The central sulfur atom is surrounded by four oxygen atoms and there are no lone pairs on sulfur. 12- :0: BHA According to VSEPR Theory, the shape of ion is tetrahedral in which central Boron atom has four bond pair and no lone pair as shown below. H H H H Borate ion Tri-iodide ion (13) : I Tri iodide long [3] orbital

See Q.18 of theory.

Sideways overlap: When the axes of two p-orbitals (i.e. py or pz orbitals) are parallel to each other, they interact to form a molecular orbitals. The bonding molecular orbitals n(2py) or T(2pz) have żero electron density on the nuclear axis (called the nodal plane). On the other hand, anti-bonding molecular orbitals т"(2py) and т"(2pz) have the least electron density in the inter-nuclear region as in figure. П2p* Tі2p 2py+2p Bonding 2p, Antibonding т2p*. п2p Bonding 2p; + 2p, Antibonding 2p,- 2p,

See Q.19 of theory.

PCl3 It shows sp hybridization because phas three sigma bonds and one lone pair as show below :Cl Ö1= SiCl4: SiCla shows sp hybridization in which central silicon forms Four sigma bonds with four chlorin atoms as shown below. Cl Si Cl Cl. Cl [NH4|+ NH4 shows sp? hybridization in which central Nitrogen atom forms four sigma bonds with four hydrogen atoms. + + H H H The structures of PF3 and SiF4 are

PF3 St P 8- F F° (H = 0.985D) PF3 is a polar molecule its dipole moment is | (wii) NH4* 0.985D SiF4 F°- 8- F F™- SiF4 is a non-polar molecule. It has zero dipole moment.

(i) ClO4-: 0: In the above structure central Cl has expanded octets by eight electrons because. Cl belongs to 3rd period (ii) ICla In above structure I śÖ(: has expended octet by four electrons ￾Öc: Iodine because 5th to belongs :C!: period. + H 1 H H In above ions the central Nitrogen obey - octet rule. (iv) 15 : Central shows -I : iodine : I- expended octets by 2 electrons because iodine belongs to 5th period.

The bond between K and Cl is ionic because their electron affinity difference is 2.34 which is greater than 1.8. The bond between Si and Cl is polar because their electronegativity difference is 1.26 which fall between 0.4-1.8

SO2 is polar due to its bent molecular geometry which results in a net dipole moment 8+ 8+ S 8- P = 0.0 D P = 1.61 D SOs is non-polar because of its trigonal planer shape allowing individual bond dipoles to cancel each other

In the molecular orbital diagrams of O2, and 05. In the molecular orbital diagram of O2, there are two unpaired electrons in t*(2py) and т* (2pz) orbitals. So 02 is paramagnetic. In the molecular orbital diagram of O2, r(2py) So are and vacant. T* (2pz) orbitals paramagnetic behaviour vanishes and now it is diamagnetic. In case of Of two electrons are absorbed into O2. Therefore n" (2py) and т* (2pz) will contain four electrons. There is no electron. So, it is not unpaired paramagnetic.

Electronegativity Bone Polarity Difference 0.66 Slightly C-Cl Polar Si-F Ionic 2.1 Se-F Most Polar 1.45 SLO BASED SHORT QUESTION ANSWERS Chemical Bond

Triple bond is stronger than double and single bound it how following order. Triple bond > double bond > single bond Single bond means one shared pair electron Double bond means two shared pair electrons Triple bond means three shared pair electrons As the number of shared pair electrons increases it leads to stronger nuclear attraction and shorter bond lengths and higher bond energies so following trend observe. Triple bond > Double bond > Single bond electron to chlorine, forming Nat and Cl ions, which are held together by electrostatic attraction.

According to the Lewis theory, the ionic bond is formed by the complete transfer of electrons from an, atom with low ionization energy to another atom with high electron affinity.

A dative bond is formed between two atoms when the shared pair of electrons is donated by one of the bonding atoms.

An ionic bond is formed by the complete transfer of one or more electrons from one atom to another, resulting in the formation of oppositely charged ions. This bond occurs between metals and nonmetals. For example, in NaCl, sodium donates one Se-F will be most polar bond.

Ionic compounds consist of a regular lattice of alternating positive and negative ions. The strong electrostatic forces of attraction between oppositely charged ions require a large amount of energy to break. Therefore, ionic compounds are usually hard solids and have high melting and boiling points. why ionic compounds 221. Explain conduct electricity in molten or aqueous states but not in solid state. Ans. In solid form, ions in ionic compounds are fixed in position and cannot move freely. However, when melted or dissolved in water, the ions are free to move, allowing them to conduct electricity. The presence of mobile ions is essential for electrical conductivity in ionic substances.

Lattice energy is directly related to the strength of an ionic bond. The greater the lattice energy, the stronger the ionic bond. It depends on the size and charge of the ions; smaller and highly charged ions form lattices with higher lattice energies.

Covalent compounds do not contain free ions or electrons that can carry electric current. The atoms are held together by shared electron pairs in neutral molecules. Hence, in solid, liquid, or aqueous form, covalent compounds are poor conductors of electricity.

In a double bond, two electron pairs are shared between two atoms, while in a triple bond, three pairs are shared. Double and triple bonds are shorter and stronger than single bonds. For example, oxygen forms a double bond in O2, and nitrogen forms a triple bond in N2.

Bond order refers to the number of chemical bonds between a pair of atoms. A higher bond order means stronger and shorter bonds. For example, in N2, the bond order is 3 (a triple bond), making it stronger than O2, which has a bond order of 2:

Bond polarity arises when two atoms with different electronegativities share electrons more unequally. The electronegative atom attracts the shared electrons more strongly, developing a partial negative charge. For instance, in Hcl, chlorine is more electronegative than hydrogen, making the bond polar.

A coordinate covalent bond is formed when one atom donates both of the electrons shared in a bond. In contrast, in regular covalent bonds, each ätom contributes one electron: An example is the formation of NHa, where NH: donates a lone pair to Ht.

One atom must have a lone pair of electrons (a donor), and the other must have an empty orbital (an acceptor). The donor provides both electrons to form a shared pair, resulting in a stable coordinate covalent bond

Metallic bonding is the attraction between positively charged metal ions and a "sea" of delocalized electrons. Unlike ionic or covalent bonds, electrons are not shared or transferred between specific atoms. This delocalization gives metals their unique properties.

In metals, valence electrons are free to move throughout the structure. These mobile electrons carry electric current easily when a voltage is applied, making metals excellent conductors of electricity in solid and liquid states. Dipole moment

It is a quantitative measurement of the polarity of a bond or a molecule. The dipole moments of diatomic molecules like HF, Hcl, HBr, HI, NO, etc. are directed from the positive ends (0 t) to negative ends (0 -). Dipole moment is measured in Debye unit (D). Higher dipole moment indicates high polarity in a molecule.

CCI, has four C-Cl bonds which are expected to have high polarity due to a large electronegativity difference between C and Cl atoms. The bond moments associated to the four C-Cl bonds are directed in such a way that they cancel each other. The net dipole moment of the CCI, molecule is zero making it a non- polar molecule. The CCl4 molecule id perfectly tetrahedral.

BeClz is a linear molecule having two similar Cl atoms on both sides of the central atom at 180°. The individual Be- Cl bond moments are cancelled out as they are opposite in direction and equal in amount. So, it has zero dipole moment.

BF3 has a trigonal planner its dipole symmetrical structure and moment is also zero. As a rule, the molecules that have same ligands atoms or groups of atoms with the central zero in a regular geometry, the individual dipole moments may not be zero, but overall molecule has zero dipole moment. Such a molecule is said to be non-polar.

Bond polarity arises when atoms with different electronegativities form a covalent bond. The more electronegative atom attracts electrons more strongly, creating a partial negative charge, while the other atom becomes partially positive. The greater the electronegativity difference, the more polar the bond.

Dipole moment is a measure of the polarity of a molecule. It is the product of the charge difference and the distance between the charges. A molecule with a higher dipole moment is more polar. It is represented by the Greek letter l and measured in Debye units (D).

Water has a bent shape due to lone pairs on oxygen, so its bond dipoles do not cancel, making the molecule polar. In contrast, carbon dioxide is linear, and the dipoles of the two C=O bonds cancel each other out, resulting in a non-polar molecule despite having polar bonds. Expanded Octets

From Lewis structures of SO?, we can calculate the number of electrons around the central S atom. The number of electrons in the valence shell of S atom can be calculated as: No. of valence electrons = 2 x (double bond electrons) + 2 × (single bond electrons) = 2(4) + 2(2) = 12 Thus, S has 12 electrons in its valence shell and it exceeds the octet by 4 electrons.

In the same way, we can calculate the number of valence electrons around the central iodine atom in the try-iodide ion. The Lewis structure of the ion is given by: The number of valence electrons of the central atom can be calculated as follows, no of valence electrons = 2 x (no. of single bond electrons) + 2 x (no. of lone pair) =2(2) + 2(3) =10 The central iodine atom in try-iodide ion I3 has 10 electrons in the valence shell and the octet expands by two electrons with ions 040. Define polyatomic example. Ans. Definition: - Polyatomic ions are the ions composed of more than one type of atoms: There formal charge is the net charge on them which is calculated based on the number of electrons in their valence shells after the formation of bonds. With the exception of ammonium ion (NH4), these ions mostly carry a negative charge, for example the carbonate (CO;), sulfate (SO?), and nitrate (NO;) ions. Difference of Electronegativity can be

The atoms having electronegativity difference less than 0.4 are said to make a pure covalent bond. An electronegativity between 0.4 difference and 1.8 corresponds to a polar covalent bond, whereas, above this value the bond between two atoms will be ionic in nature Bond Energy and Bond Length

Bond length is the distance between the nuclei of two bonded atoms. It decreases as the number of shared electron pairs increases. A single bond is longer than a double bond, which in turn is longer than a carbon triple bond. For example, in compounds: C-C > C=C > C=C in bond length.

There is an relationship between bond energy and bond length. A shorter bond is generally stronger and has a higher bond energy. As bond length increases, the attraction between nuclei and bonding electrons decreases, leading to a weaker bond and lower bond energy

Valence Shell Electron Pair Repulsion (VSER) theory states that electron pairs around the central atom repel each other and arrange themselves to be as far apart as possible. This determines the shape of the molecule to minimize repulsion.

Methane has a tetrahedral shape with bond angles of approximately 109.5°. This is because the four bonding pairs of electrons repel each other equally and arrange themselves as far apart as possible around the central carbon atom.

Water has two lone pairs and two bonding pairs on the central oxygen atom. Lone pairs repel more strongly than bonding pairs, distorting the shape into a bent geometry with a bond angle of about 104.5°.

Carbon dioxide (COz) has two double bonds on opposite sides of the carbon atom and no lone pairs on the central atom. The repulsion between the two electron regions results in a linear arrangement with a bond angle of 180°

Hybridization is the mixing of atomic orbitals to form new orbitals (hybrids) that Molecular Orbital theory (MOT)

In NH3, the nitrogen atom is sp hybridized because it has three sigma bonds and one lone pair. The molecule has a trigonal pyramidal geometry, and the bond angle is about 107°, slightly less than tetrahedral due to lone pair repulsion. 2-2 = 0 A.B.M.O 0*(1s) • zHe 1s A.O 0(1s) B.M.O] 2-0 - = 1 2 2 Bond energy and bond length

The electron configuration of He is 1s2. For a successful formation of Hez molecule, 1s orbitals of two of Hez molecule, 1s orbitals of two He atoms must combine to form bonding (8 1s) and anti-bonding (o" 1s) orbitals as shown in Figure 3.36. Out of four electrons, two enter the bonding molecular orbital s (ls) and the remaining two occupy the antibonding s* (1s) molecular orbital. But on calculation we discover that the bond order for Hez is zero. Hence, Hez molecule is not formed Bond order of He2 = Energy гНе = 1s (A.O)

The number of bonds formed between two atoms after the atomic orbitals overlap, is called the bond order and is taken as half of the difference between the number of bonding electrons (say a) and antibonding electrons (say b). (No. of Electrons in B. M.O) - (No. of Electrons in A. B.M:O) Bond Order = are suitable for bonding. It explains the shape and bond angles of molecules better than using unhybridized orbitals. For sp? in CH4, carbon uses example, hybridization to form four equivalent bonds. o bonds + lone pairs =4 the geometry and

The energy is the average amount of energy required to break all bonds of a particular type in one mole of a substance. by It is determined experimentally measuring the heat involved in a chemical . Bond reaction and its unit is KJ mol. energy is a measure of the strength of a bond and its reactivity Bond length is the distance between the nuclei of two atoms forming a covalent bond are bond The lengths experimentally determined by physical techniques, such as electron diffraction, X-ray diffraction, or spectral studies. The bond length of a bond is governed by many factors including electronegativity, size, and the nature of the covalent bond (single, double, or triple). "

In NF3, the strong polarity of N-F bond pulls the lone pair of atom closer to its DESCRIPTIVE QUESTIONS (EXERCISE)

Chemical bond's, i.e. ionic, covalent, and metallic bond usually termed as true chemical bonds, as the affect the chemical properties of a substance. This is because of certain reasons. Firstly, chemical bonds result in the formation a new species through transfer and sharing of electrons. Whereas, intermolecular forces act to bring molecules closer and influence the physical properties. However, such a clear distinction between the chemical bonds and intermolecular forces is not possible.