Chemical Bonding

Synopsis

Introduction

• A chemical bond is the force which holds two or more atoms together in a stable molecule.
• An atom is the smallest unit of matter taking part in a chemical reaction. It is built up of subatomic particles—protons, neutrons and electrons.

Chemical Bonding

• Chemical Bond
  A chemical bond is defined as the force of attraction between any two atoms in a molecule to maintain stability.

      Noble Gases

• Have stable electronic configuration, i.e. their outermost shell is complete.
• They have 2 electrons in the outermost shell or 8 electrons in the outermost shell.
• They do not lose, gain or share electrons and are inert or unreactive.

      Atoms of Elements – Other than Noble Gases

• Have unstable electronic configuration, i.e. their outermost shell is incomplete.
• They can lose, gain or share electrons and are chemically reactive.

      Reasons for Chemical Bonding

• The driving force for atoms to combine is related to the tendency of each atom to attain stable electronic configuration of the nearest inert noble gas.
• For an atom to achieve stable electronic configuration, it must have
• Two electrons in the outermost shell (nearest noble gas He) – Duplet rule
• Eight electrons in the outermost shell (all noble gases other than He) – Octet rule

      Methods for Achieving Chemical Bonding

      There are three methods in which atoms can achieve a stable configuration.

• Transfer of one or more electrons from one atom to the other to form an electrovalent bond.
• Sharing of one, two or three pairs of electrons between two atoms to form a covalent bond.
• When the shared electron pairs are contributed by only one of the combining atoms, the bond formed is known as a coordinate bond.

Electrovalent (or Ionic) Bond

      Types of Elements

• Metallic elements have 1, 2 or 3 electrons in their valence shell. They lose 1, 2 or 3 electrons and become positively charged ions [cations].
• Non-metallic elements have 4, 5, 6 or 7 electrons in their valence shell. They gain (4), 3, 2 or 1 electrons and become negatively charged ions [anions].

• Ionic Bond
  The chemical bond formed between two atoms by transfer of one or more electrons from the atom of a metallic electropositive element to an atom of a non-metallic electronegative element.

• Ionic Compound
  The chemical compound formed as a result of transfer of one or more electrons from the atom of a metallic electropositive element to an atom of a non-metallic electronegative element.

• Electrovalency
  The number of electrons donated or accepted by the valence shell of an atom of an element so as to achieve stable electronic configuration is called electrovalency.

      Conditions for the Formation of an Ionic Bond

• Ionisation Potential (IP)
  Lower the value of IP of a metallic atom, greater the ease of formation of the cation.
• Electron affinity
  Higher the value of EA of a non-metallic atom, greater the ease of formation of the anion.
• Electronegativity
  Larger the electronegativity difference between combining atoms, greater the ease of electron transfer.

      Stability of Ionic Compounds

• Ionic compounds are formed by ions, but there also exists a repulsive force between ions for like charges.
• Since the electrostatic force of attraction between opposite charges is much higher, it makes ionic compounds stable.

            Examples: NaCl, MgCl₂, CaO, KBr, CaCl₂

•  Formation of Electrovalent Compounds
   Formation of Sodium Chloride:
  The electronic configuration of sodium atom is 2, 8, 1. Since its valency is 1, it loses one electron to attain the stable electronic configuration of neon and becomes a positively charged sodium ion (Na⁺) having a net charge +1.

                                          Na  − 1e−  →   Na⁺

                                      (2, 8, 1)                (2, 8)

The electronic configuration of chlorine atom is 2, 8, 7. As its valency is 1, it will accept one electron to complete its octet and attain stability.
So, it will accept the one electron lost by the sodium atom to become a chloride ion (Cl⁻) having −1 net charge.

                                   Cl  + 1e−   →      Cl⁻

                                     (2, 8, 7)          (2, 8, 8)

So, cation Na⁺ and anion Cl⁻ are attracted towards each other due to electrical charge and form an ionic compound, sodium chloride.

                               Na + Cl  →  Na⁺ Cl⁻   →  NaCl’

Electron Dot Structural Diagram

Covalent Bonding
Covalent Bond
The chemical bond formed due to mutual sharing of electrons between the given pairs of atoms of non-metallic elements.

• Covalent Compound
  The chemical compound formed due to mutual sharing of electrons between the given pairs of atoms thereby forming a covalent bond between them.

•  Covalency
  The number of electron pairs which an atom shares with one or more atoms of the same kind or different kind to achieve the stable electronic configuration is called covalency.

•  Non-polar Covalent Compounds
  Covalent compounds are said to be non-polar when the shared pair of electrons are equally distributed between the two atoms.
  Examples: H₂, Cl₂, O₂, N₂, CH₄, CCl₄

•  Polar Covalent Compounds
  Covalent compounds are said to be polar when the shared pair of electrons are unequally distributed between the two atoms.
  Examples: H₂O, NH₃, HCl

      Conditions for Formation of Covalent Compounds

• Both atoms should have high electronegativity, electron affinity and ionisation energy.
• The electronegative difference between the two combining atoms should be negligible.
• Both atoms should have four or more electrons in their outermost shell.
• ­The percent covalent character of ionic bond is determined by the factors such as polarizing power of the cation, polarizability of the anion and degree of polarization of anion. ­erefore, we shall give some idea of the polarizing power of cation and polarizability of anion.

1. Polarizing power of a cation is its power to distort the electronic charge of the anion. In general, the polarizing power increases as cations become smaller and more highly charged.
2. Polarizability of an anion is the extent to which its electronic charge can be distorted. In general, anions with larger size and large negative charge are more polarizable than smaller ones.

Fajan’s Rules

1.  When the size of the cation is small and the size of the anion is large, the covalent character of an ionic bond is large.
2.  Reason: In small cation, the positive charge is concentrated over a small area. ­is makes the cation highly polarizing and more effective in distorting the negative charge of the large anions.
3. When the charge on cation or on anion, or on both ions is large, the covalent character of an ionic bond is large. Reason: A high charge increases the extent of polarization.
   When the cation does not have a noble gas configuration, the covalent character of an ionic bond is large.
   Explanation: A noble gas configuration has closed outermost shell and it shields the nuclear charge most effectively. ­us, a cation with noble gas electronic configuration has less power to polarize an anion.
   A cation without the noble gas configuration has high positive charge at its surfaces, and thus it has high power to polarize an anion.
   Some cations which do not have a noble gas configuration and favour covalent character:

1. Cations of transition metals such as Ti³⁺, V³⁺, Cr²⁺, Mn²⁺, and Cu⁺, Ti⁺ , Pb²⁺, and Bi³⁺ (b) Cations of some lanthanide metals such as Ce³⁺ and Eu²⁺

      Formation of covalent molecules:

1. Hydrogen molecule (Non-polar molecule):
   A hydrogen atom has only one electron in its K shell. It needs one more electron to complete its duplet or to attain the stable electronic configuration of the nearest rare gas helium. To meet this need, the hydrogen atom shares its electron with another hydrogen atom to form a hydrogen molecule.

      Formation of hydrogen molecule:

    2. Chlorine molecule:

 

     3. Formation of Water – Polar Covalent Compound

    4. Formation of an ammonia molecule:

  

   5. Carbon tetrachloride molecule (Non-polar molecule):

Limitations of the Octet Rule

• Incomplete octet of the central atom in some compounds:
  The number of electrons surrounding the central atom is less than eight. This is especially the case with elements with less than four valence electrons. Examples: LiCl, BeH₂, BCl₃

• Odd-electrons molecule:
  In molecules with an odd number of electrons such as nitric oxide (NO) and nitrogen dioxide (NO₂), the octet rule is not satisfied for all the atoms.

• Expanded octet:
  Elements belonging to groups 15, 16 and 17 have more than four electrons in their valence (outermost) shell. The elements of these groups form stable compounds in which there are more than eight electrons around the central atom. Examples: PF₅, PCl₅, IF₆, SF₆

Drawbacks of the Octet Theory

• It is clear that the octet rule is based on the chemical inertness of noble gases. However, some noble gases (e.g. xenon and krypton) also combine with oxygen and fluorine to form compounds such as XeF₂, KrF₂ and XeOF₂.
• This theory does not account for the shape of molecules.
• It does not explain the relative stability of the molecules being totally silent about the energy of a molecule.

• A covalent bond is formed because of electrostatic attraction between radii and the accumulated electrons cloud and by attraction between spins of anti-spin electrons.
• Greater is the overlapping, lesser will be the bond length, more will be the attraction and more will be the bond energy and the stability of the bond will also be high.
• The extent of overlapping depends on the nature of orbitals involved in overlapping and the nature of overlapping.
• Closer the valence shells to the nucleus, more will be the overlapping and the bond energy will also be high.
• Between two subshells of the same energy level, the subshell more directionally concentrated shows more overlapping. Bond energy: 2s–2s < 2s–2p < 2p–2p
• 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 types gives sigma (σ) and π (p) bonds.

Types of overlap

• Sigma (σ) bond:
  The covalent bond formed because of overlapping of atomic orbitals along the internuclear axis is called the σ-bond. It is a stronger bond and cylindrically symmetrical. All sigma bonds have axial symmetry. A sigma bond is formed by s–s, s–p_z and p_z–p_z overlap.
  
    

 

• Pi (π) bond:
  The covalent bond formed by sidewise overlapping of atomic orbitals is called π-bond. In this bond, the electron density is present above and below the internuclear axis. It is relatively a weaker bond because the electrons are not strongly attracted by the nuclei of bonding atoms. If the Z axis is assumed to be the molecular axis, then the π–bond is given by p_x–p_x and p_y–p_y.
  
   

Differences between sigma and pi bonds

• Covalent bond may be single, double or a triple bond. Double and triple covalent bonds are called multiple covalent bonds.
• Single covalent bond is formed by sharing of only one electron pair ­is bond is represented by single dash.
• Double and triple covalent bonds are formed when atoms bonded together share two or three electron pairs, respectively. ­These bonds are represented by double dash ( = ) and triple dash ( ≡ ) respectively. Some examples of covalent bonding are given below:

 Sharing of one pair of electrons (single bond):
 Two hydrogen atoms share one pair of electrons to form a hydrogen molecule.

 

Sharing of two pairs of electrons (double bond):
Two oxygen atoms share two pairs of electrons to form one oxygen molecule.

 

Sharing of three pairs of electrons (triple bond)
Two nitrogen atoms share three pairs of electrons to form one nitrogen molecule.

 

 

Hybridisation
Pauling introduced the concept of hybridisation. According to him, atomic orbitals combine to form a new set of equivalent orbitals known as hybrid orbitals. The hybrid orbitals are used in bond formation.

 

 

Hybridisation: Mixing of two atomic orbitals with the same energy level to give a degenerated new type of orbitals.
This intermixing is based on quantum mechanics.

 Some Important Features of Hybridisation

• The number of hybrid orbitals is equal to the number of the atomic orbitals which get hybridised.
• Hybridised orbitals are always equivalent in energy and shape.
• Hybrid orbitals form more stable bonds than pure atomic orbitals.
• Hybrid orbitals are directed in space in some preferred direction. So, they show minimum repulsion between electron pairs and result in a stable arrangement.

Conditions for Hybridisation

 

Types of Hybridisation

 

sp Hybridisation

• Intermixing of one 's' and one 'p' orbital of almost equal energy to give two identical and degenerate hybrid orbitals is called 'sp' hybridisation.
• These sp-hybrid orbitals are arranged linearly by making a 180° angle.
• They possess 50% 's' and 50% 'p' character.
  
  

Examples:

Structure of BeCl₂:

• In BeCl₂, the two singly occupied orbitals (2s and 2p) hybridise to give two sp-hybrid orbitals. These hybrid orbitals lie along the z-direction and point in opposite directions.
• The ground state electronic configuration of Be is 1s²2s². In the exited state, one of the 2s-electrons is promoted to the vacant 2p orbital to account for its bivalency.
• One 2s and one 2p-orbital hybridise to form two sp hybridised orbitals. These two sp hybrid orbitals are oriented in opposite directions forming an angle of 180°.
  
   
  
  

 

 

Structure of C₂H₂:

• The ground state configuration of ‘C’ being 1s² 2s² 2p_x¹ 2p_y¹ has only two unassociated electrons. Carbon makes four bonds because its valency is four.
• For this, 4 unpaired electrons are required. Hence, it promotes its 2s electron to the empty 2p_z orbital in the excited state.
• So, the excited state electronic configuration of carbon is 1s² 2s¹ 2p_x¹ 2p_y¹ 2p_z¹. Each carbon atom undergoes ‘sp’ hybridisation by using 2s and 2p orbitals in the excited state to give two half-filled ‘sp’ orbitals, which are arranged linearly.
  
   

 

sp² Hybridisation

• Intermixing of one 's' and two 'p' orbitals of almost equal energy to give three identical and degenerate hybrid orbitals is known as sp² hybridisation.
• The three sp² hybrid orbitals are oriented in trigonal planar symmetry at angles of 120° to each other.
• The sp² hybrid orbitals have 33.3% 's' character and 66.6% 'p' character.
  
  

Examples:

Structure of BCl₃:

• In the BCl₃ molecule, the ground state electronic configuration of the central boron atom is 1s²2s²2p¹.
• In the excited state, one of the 2s electrons is promoted to the vacant 2p orbital. As a result, boron has three unpaired electrons. These three orbitals (one 2s and two 2p) hybridise to form three sp² hybrid orbitals.
• The three hybrid orbitals so formed are oriented in a trigonal planar arrangement and overlap with 2p orbitals of chlorine to form three B–Cl bonds. The geometry is trigonal planar with a Cl–B–Cl bond angle of 120°.
  
   
  
  

Structure of C₂H₄:

• In the excited state, the carbon atom experiences sp² hybridisation by mixing 2s and two 2p orbitals to give three half-filled sp² hybrid orbitals oriented in trigonal planar symmetry.
• Another unhybridised 2p_z orbital is also present on each carbon perpendicular to the plane of sp² hybrid orbitals in the half-filled state. The carbon atoms form a σ-bond by the overlap of sp²–sp² hybrid orbitals.
• Lateral overlapping between the unhybridised 2p_z orbitals also leads to the formation of a π-bond (p–p bond). Thus, there is a double bond (σ_sp2–sp2 and π_p–p) between two carbon atoms. Each carbon atom also forms two σ_sp2–s bonds with two hydrogen atoms.
  
  

 

sp³ Hybridisation

• In sp³ hybridisation, one 's' and three 'p' orbitals of almost equal energy intermix to give four identical and degenerate hybrid orbitals.
• These four sp³ hybrid orbitals are oriented in tetrahedral symmetry with 109°28' angle with each other.
• The sp³ hybrid orbitals have 25% ‘s’ character and 75% 'p' character.
   

 

Example:

Structure of CH₄:

• The CH₄ molecule in which there is mixing of one s-orbital and three p-orbitals of the valence shell to form four sp³ hybrid orbitals of equivalent energies and shape.
• There is 25% s-character and 75% p-character in each sp³ hybrid orbital. The four sp³ hybrid orbitals so formed are directed towards the four corners of the tetrahedron.
• The angle between sp³ hybrid orbitals is 109.5°. Hence, it is planar and has ∠HCH and ∠HCC bond angles equal to 120°.
  
  
   

Structure of C₂H₆:

• C₂H₆ shows sp³ hybridisation in the excited state which yields four sp³ hybrid orbitals in tetrahedral geometry.
• Overlapping of sp³ hybrid orbitals along the internuclear axis leads to the formation of a σ_sp3–sp3 bond. Each carbon atom also forms three σ_sp3–s bonds with hydrogen atoms.
• Hence, ethane has tetrahedral symmetry around each carbon with ∠HCH and ∠HCC bond angles equal to 109^o28′.
  
   

VSEPR Theory
The Lewis concept is unable to explain the shapes of molecules. Sidgwick and Powell (1940) provided a useful idea for predicting shapes and geometries of molecules. The theory was based on the repulsions between electron pairs, known as valence shell electron pair repulsion (VSEPR) theory.

Postulates of this theory:

• In polyatomic molecules which contain three or more atoms, one of the atoms is called the central atom to which the other atoms are linked.
• The geometry of a molecule is dependent on the total number of valence shell electron pairs (bonded or not bonded) present around the central atom and their repulsion because of relative sizes and shapes.
• If the central atom is surrounded by bond pairs only, then it gives a symmetrical shape to the molecule.
• If the central atom is surrounded by a lone pair (lp) and bond pair (bp), then the molecule has a distorted geometry.
• The relative order of repulsion between electron pairs is 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 greater when a lone pair is involved.
  
  
  
  

Coordinate Bond
The bond formed between two atoms by sharing a pair of electrons provided entirely by one of the combining atoms but shared by both is called a coordinate bond or dative bond.
Examples: Ammonium ion (NH₄⁺), hydronium ion (H₃O⁺)
A coordinate bond has properties of both covalent and ionic bonds. So, it is also called a co-ionic bond.

• Lone Pair of Electrons
  A pair of electrons which is not shared with any other atom is known as a lone pair of electrons. It is provided to the other atom for the formation of a coordinate bond.

• Conditions for the Formation of a Coordinate Bond

1. One of the two atoms must have at least one lone pair of electrons. Examples: Ammonia (NH₃), water (H₂O)
2. Another atom should be short of at least one lone pair of electrons. Example: Hydrogen ion (H⁺)

        Formation of hydronium ion [H₃O⁺]

  

Formation of ammonium ion (NH₄⁺) :

 

• Properties and comparison of electrovalent and covalent compounds