Video Lesson

Lesson Objects

At the end of the unit , you will be able to:

• List the favourable conditions for the formation of ionic bonds.
• Explain the formation of ionic bonding
• Give examples of ionic compounds
• Define lattice energy.
• Calculate lattice energy of ionic crystals from given data using the Born-Haber cycle
• Discuss the exceptions to the octet rule.
• Describe the properties of ionic bonding
• Carryout an activity to demonstrate the effect of electricity on ioniv compounds (PbI2 and NaCl).
• Carryout an investigation into the melting point and solubility of some ionic compounds (NaCl and CuCl2.
• List types of chemical Bonding.
• Define covalent bond.
• Understand the metallic bonding
• Lewis structures or electrondotformulasof some covalent moleculecule.
• Illustrate the formation of coordinate covalent bonding using example draw resonance structures of some covalent molecules and polyatomic ion.
• Discuss the exceptions to the octet rule in covalent bond in distinguish between polar and non polar covalent molecule describe the properties of covalent molecules ,electricity and some solvents on covalent compounds (naphthalene, graphite, iodine and ethanol ).

Brainstorming questions

• why do atoms combine ?
• why molecules are stable than atoms ?

key terms and Concepts

2.1. Introduction

What is chemical bonding?

• Chemical bond is a strong force of binding between two or many atoms in compounds, molecules or ions.

Why do atoms combine? Atoms are bonded together to:

• Attain noble gas electronic configuration.
• Attain stability.
• have lowest possible energy.

2.2. Types of Chemical Bonding

Chemical bonds are basically classified into three types consisting of

• Ionic or electrovalent bond.
• Covalent bond and.
• Metallic bond.

Mostly, valence electrons in the outer energy level of an atom take part in the chemical bonding.

2.2.1. Ionic Bonding

• The electrostatic force of attraction between oppositely charged ions results in the formation of ionic bond or electrovalent bond.
  • Example: Consider the formation of some ionic bandings b/n Na⁺ and Cl^– ions:

• Examples of ionic compounds include: NaCl, MgO, MgBr₂, CaC₂, HCl, Li₃N…

Factors Affecting The Formation Of Ionic Bonds are:

• Ionization energy(IE)
• Electron Affinity (EA)
• Lattice Energy (U)

Lattice Energy formation of Ionic bond

Finding the lattice energy of anion in solid by experiment is difficult. However, this quantity can be found indirectly using the Born–Haber cycle. The reasoning is based on Hess’s law, which states that an overall reaction’s enthalpy change is the sum of the enthalpy changes for the individual reactions that make it up:
∆H _total= ∆H1+ ∆H2+∆H3+ …
Consider the Born-Haber cycle for the formation of NaCl. Solid sodium chloride can be formed from the elements by two different routes, as shown in the following figure. In one route, NaCl(s) is formed directly from ; Na(s) + ½Cl (g) → NaCl(s) ∆H =−411 kJmol^-1

Figure :Born–Haber cycle for NaCl

The second route consists of the following five steps , along with the enthalpy change for each.
Step1: Metallic sodium is vaporized to a gas of sodium atom:
Na(s)→ Na(g) ……………………………….∆H_sublimation = + 108 kJ mol^–1
Step 2: Chlorine molecules are dissociated to atoms:
½Cl₂(g)→Cl(g) …………………………..∆H°_dissociation =½ bond energy of Cl₂ =½ (240 KJ) = + 120 kJ mol^–1
Step3:Sodium atoms are ionized to Na+ ions:
Na(g)→Na⁺(g)+e– ……………………∆Η°_ionization =IE1=+ 496 kJ mol^–1
Step 4:Formation of chloride ion:
Cl(g)+e–→Cl^–(g) ………………………∆Η°electron afinity =EA=–349 kJ mol^–1

Step 5:Formation of NaCl(s) from ions. The ions Na+ and Cl−combine to give solid
sodium chloride whose enthalpy changes (the lattice energy)is unknown:
Na⁺(g) + Cl^–(g) → NaCl (s).…………………….∆H°_{lattice energy} =U(lattice energy)=?
We know the enthalpy formation (∆H°_formation) of NaCl (Direct route) and equals −411 kJmol-1. therefore, we can calculate the lattice energy using Hess’s low: Solving for U NaCl gives:

∆H°_formation = ∆H_sublimation + ∆H°_dissociation + ∆Η°_ionization + ∆Η°_{electron afinity} + ∆H°_{lattice energy}

∆H°_{lattice energy} =∆H°_formation – [∆H_sublimation + ∆H°_dissociation + ∆Η°_ionization + ∆Η°_{electron afinity} ]
∆H°_{lattice energy} =–411kJmol^–1 – [108kJmol^–1+120kJmol^–1+496kJmol^–1+(–349kJmol^–1)]
∆H°_{lattice energy} =–786 kJ mol^–1
Note: Ionic solids exist only because the lattice energy exceeds the energy required for the electron transfer.

General Properties of Ionic Compounds

• Ionic compounds usually exist in the form of crystalline solids.
• Ionic compounds have high melting point and high boiling point due to strong attraction of force between ions.
• Ionic compounds are generally soluble in water and other polar solvents having high dielectric constants.
• Ionic compounds are good conductors of electricity in solution or in their molten states.
• In ionic – compounds, each ion is surrounded by oppositely charged ions uniformly distributed all around the ion. I.e. ionic compounds are non directional.

2.2.2. The Covalent Bond

• Covalent bonds involve sharing of a pair of valence electrons by two atoms.
• When two atoms have a small difference in their tendencies to lose or gain electrons, we observe electron sharing and covalent bonding.
• This type of bonding most commonly occurs between non metal atoms

Multiple covalent bonds are common for certain atoms depending upon their valence configuration.
For example, in carbon dioxide (CO₂), a double covalent bond results from the sharing of two sets of valence electrons.

Atomic nitrogen (N₂) is an example of a triple covalent bond. :N≡N:

The polarity of a covalent bond is defined by any difference in electro-negativity between the two atoms participating in the covalent bond formation. Bond polarity describes the distribution of electron density around two bonded atoms. For two bonded atoms with similar electronegativities, the electron density of the bond is equally distributed (shared) between the two atom, this is a non-polar covalent .bond. eg, H₂,

The electron density of a covalent bond is shifted towards the atom with the largest electro-negativity.

Polar covalent molecule.

https://phet.colorado.edu/sims/html/molecule-polarity/latest/molecule-polarity_all.html

https://phet.colorado.edu/sims/html/molecule-polarity/latest/molecule-polarity_all.html

Coordinate covalent Bond

A coordinate covalent bond (also called a dative bond) is formed when one atom donates (contribute) both of the electrons from single atom. The atom which gives both electron is called donor atom and the atom which accepts the electrons is called Acceptor atom.

Properties of covalent compounds

•  The covalent compounds do not exist as ions but they exist as molecules.
• The melting point. and boiling point of covalent compounds are generally low
• Covalent compounds are generally insoluble or less soluble in water and other polar solvents. However, these are soluble in non-polar solvents.
• They are poor conductors of electricity in the fused or dissolved state.
• Molecular reactions are quite slow because energy is required to break covalent bonds.
• Since the covalent bond is localized in between the nuclei of atoms, it is directional in nature

The Octet Rule

States that an atom tends to gain, lose or share electrons until there are eight electrons in its valence shell (the nearest noble gas configuration at its row). The octet rule guides us in allotting electrons to the atoms in a Lewis structure; in a few cases, however, we set the rule aside.
Example: ₁₁Na⁺ = 1s²2s²2p⁶ is the sub shell Configuration of Ne. Its main shell configuration is 2,8
₁₇Cl^– = 1s²2s²2p⁶3s²3p⁶ is the sub shell Configuration of Ar. Its main shell configuration is 2,8,8

Exceptions to the Octet Rule

1) Electron deficient molecules
Compounds containing either Beryllium(Be) or Boron(B) are often electron deficient, i.e., There are fewer than 8 electrons around Be or B atom.
Example: BH₃ , BeCl₂, BF₃,

2) More than Octet: Some molecules may contain more than 8 electrons on their central atoms, the following Lewis structure shows that;

2.2.3. Metallic Bonding

The atoms in a pure metal are in tightly-packed layers, which form a regular lattice structure.
Electrostatic forces of attraction between the positively charged nuclei and the negatively charged electrons hold the lattice together.
A metal is therefore seen as a rigid framework of cations immersed in a ‘sea’ of electrons that serve as the element holding the three-dimensional cationic network together – Metallic bonding.

The metal atoms become positively charged ions and are attracted to the sea of electrons. This attraction is called metallic bonding.

Properties of Metals

• Conduct heat and electricity
• High Melting and boiling Point
• Malleable and ductile
• Strong, not brittle