Video Lesson

• Describe the valence shell electron pair repulsion theory (VSEPR).
• Distinguish between the bonding pairs and non-bonding pairs of electrons.
• Describe how electron pair arrangements and shapes of molecules can be predicted from the number of electron pairs
• Explain why double bonds and lone pairs cause deviations from ideal bond angles
• Explain the term dipole moment with the help of a diagram.
• Describe the relationship between dipole moment and molecular geometry.
• Describe how bond polarities and molecular shapes combine to give molecular polarity
• Predict the geometrical shapes of some simple molecules on the bases of hybridization and the nature of electron pairs
• Construct models to represent shapes of some simple molecules.
• Define intermolecular forces
• Name the different types of intermolecular forces.
• Explain dipole-dipole interactions.
• Dive examples of dipole-dipole interaction
• Define hydrogen bonding.
• Explain the effect of hydrogen bond on the properties of substances.
• Give reasons why hydrogen bonding is stronger than ordinary dipole-dipole interactions.
• Explain dispersion (London)forces.
• Give examples of dispersion forces.
• Predict the strength of intermolecular. forces for a given pair of molecules.

Brainstorming Questions

• How does the concept of electronegativity influence the type of covalent bond formed between two atoms?
• In what ways do single, double, and triple covalent bonds affect the properties of a molecule?
• How do lone pairs of electrons influence the bonding and geometry of molecules?
• Why are some covalent bonds polar while others are non-polar, and how is this related to molecular symmetry?

Key terms and Concepts

Covalent Bond theories

2.3.1. Representation of Covalent Bond or Drawing Lewis structure

Drawing Lewis structure Is very useful when predicting molecular shape. To predict arrangement of atoms within the molecule use the following rules.

1. H is always a terminal atom. Always connected to only one other atom.
2. Lowest electro negativity is central atom in molecule [not just the oddball element].
3. Find the total number of valence electrons by adding up group number of the element. For ions, add for negative and subtract for positive charges

4. Place one pair of electrons, a bond, between each pair of bonded atoms.
5. Subtract from the total number of bonds you just need.
6. Place lone-pairs about each terminal atom (except H) to satisfy the octet rule. Left over pairs are assigned to the central atom. If the central atom is from the 3rd or higher period, it can accommodate more than four electron pairs up to six pairs.

7. If the central atom is not yet surrounded by four electron pairs, convert one or more terminal atom lone pairs to Pi bonds. NOT ALL ELEMENTS FORM pi (𝝅 ) BONDS !! ONLY C, N, O, P & S!!!

EXAMPLES

1. Determine the total number of valence electrons for the following species.
   • a. SO₄^2-
   • b. NH₄⁺

Solution; a. SO₄^2-

S ⇀ 1 × 6 = 6
O ⇀ 4 × 6 = 24
2 ⟶ the charge
SO₄²⁻ = 32 valence electrons

b. NH₄⁺

N ⇀ 1 × 5 = 5
H ⇀ 4 × 1 = 4
−1 for the charge
NH₄⁺ has 8 valence electrons

2. predict the lewis structure for:

a. PH₄⁺

b. COCl₂

Solution;

a. PH₄⁺

b. COCl₂

P ⇀ 1 × 5 = 5
H ⇀ 4 × 1 = 4
−1 for the change
PH₄⁺ has 8 valence electrons

The central atom is P, thus the Lewis structure for PH₄⁺ is shown below;

b. COCl₂

C ⇀ 1 × 4 = 4
O ⇀ 1 × 6 = 6
Cl ⇀ 2 × 7 = 14
Total: COCl₂ has 24 valence electrons
C- is the central atom because it has the lowest electro negativity.

• The central atom has fewer than eight electrons, so one lone pair of electrons is removed from the terminal atom, for this case Oxygen, to make the central atom octet, then a double bond is formed.

2.3.2. Valence Shell Electron Pair Repulsion (VSEPR) Theory

• This theory was proposed for the first time by Sidewick and Powell in 1940 and developed by Gillespie and Nyholm in 1957.
  This theory is very useful in predicting the geometry or shape of polyatomic molecules or ions of non-transition elements.
  According to this theory, “the shape of a given species (molecule or ion) depends on the number and nature of electron pairs surrounding the central atom of the species.
• Molecular Geometry is the general shape of a molecule as determined by the relative position of atomic nuclei.
  The valence shell electron pairs are arranged about each atom so that electron pairs are kept as far away from one another as possible, thus minimizing the electron pair repulsion.

Basic principle of VSEPR Theory

• Electron pairs will arrange themselves to be as far apart as possible so that the repulsion between them is at minimum.
• Lone pair of electrons takes more room on the surface of atom than that of bonding pairs.
• According to the VSEPR theory, the repulsion of electron pair is in the following order: Lp-Lp >Lp -Bp > Bp -Bp
• The repulsive force decreases sharply with increasing bond angle.

Molecular Geometry

Molecular geometry refers to the spatial arrangement of atoms in a molecule. It describes the actual three-dimensional shape of the molecule, taking into account the positions of all atoms relative to each other. Molecular geometry is determined primarily by the positions of atoms, regardless of whether they are bonded directly to the central atom or not.

Example: For example, in a molecule like methane (CH₄), the molecular geometry is tetrahedral, meaning the carbon atom is at the center with four hydrogen atoms positioned at the vertices of a tetrahedron.

Electron Set Arrangement (Electron Pair Geometry)

Electron set arrangement (or electron pair geometry) refers to the spatial arrangement of all electron pairs (both bonding and non-bonding) around the central atom in a molecule.

Electron set arrangement is determined by the total number of electron pairs (bonding pairs + lone pairs) around the central atom.

Example: Continuing with the example of methane (CH₄), the electron set arrangement around the central carbon atom is tetrahedral because there are four regions of electron density (four bonding pairs of electrons).

Differences:

Molecular geometry focuses on the actual positions of atoms in space, providing information about the shape of the molecule. Where as Electron set arrangement considers all electron pairs around the central atom, including both bonding pairs and lone pairs, providing a broader perspective on electron distribution.

Relationship:

The molecular geometry is often a subset or specific manifestation of the electron set arrangement. For example, in molecules with no lone pairs on the central atom (like methane), the molecular geometry and electron set arrangement are the same (tetrahedral). However, in molecules with lone pairs, the molecular geometry describes the actual shape of the molecule, while the electron set arrangement describes the overall arrangement of electron pairs around the central atom.

• To classify molecular shapes, the AX_mE_n designation is assigned
  • Were A- is the central atom
  • x- is the terminal atom
  • E- is the lone pair of electron
  • m and n are positive integers.

Predicting the shapes of Molecules

1. Molecular shapes with two electron sets
    Molecular formula − Ax₂
    Number of electron pair− two bonding pairs
    Ideal Bond angle – 180⁰

• Molecular geometry- Linear
  Example:- BeCl₂

2. Molecular shapes with three Electron Sets
   a . AX₃ Type
   Molecular formula- Ax₃
   Number of electron pair- 3 bonding pairs
   Ideal bond Angle- 120⁰
   Molecular Geometry- Trigonal planar
   Example BF₃

Lewis structure is;

b. Ax₂E Type
 Molecular formula- Ax2E
 Number of electron pairs- 2 bonding and 1-lone pair electron
 Ideal bond angle- 95⁰

• Molecular Geometry- V- shaped or Angular
• Example:- SO₂

Lewis structure is;

3. Molecular shapes with four electrons set
   a. Ax₄ Type
   Molecular formula- Ax₄
   Number of electron pairs- 4 bonding electrons
   Ideal bond angle- 109.5⁰
   Molecular geometry- Tetrahedral
   Example:- CH₄. NH₄⁺, SO₄^2-

b. Ax₃∈ Type
Molecular formula- Ax₃ E
No of electrons pair- 3 bonding and one lone- pair of electrons
Ideal bond angle- 107.3⁰
Molecular geometry- Trigonal pyramidal
Example:- NH₃ , PCl₃ ,

c. Ax₂E₂ Type
Molecular formula- Ax₂E₂
Number of electrons pair- 2 bonding and 2 lone- pair of electrons
Ideal bond angle- 104.5⁰
Molecular geometry- V- shaped, Angular or Bent.

Example:- H₂0

4. Molecular shapes with five electron sets
   a. Ax₅ Type
   Molecular formula- Ax₅
   Number of electrons pair- 5 bonding electrons
   Ideal bond Angle-120⁰ and 90⁰
   .
   Molecular Geometry- Trigonal bi pyramidal
   Example:- PCl₅

b. Ax₄E Type
Molecular formula- Ax₄E
Number of electron pairs- 4 bonding and one lone pair of electrons.
Molecular Geometry- sea saw shape
Example:- SF₄,

C. Ax₃E₂ Type
Molecular formula- Ax₃E₂
Number of electron pairs- 3 bonding and 2 lone-pair of electrons
Molecular geometry- T- shape
Example:- ClF₃ , ICl₃

D. Ax₂E₃ Type
 Molecular formula- Ax₂E₃
 Number of electron pairs- 2 bonding and 3 lone-pair of electrons.
 Molecular Geometry- Linear
Example:- XeF₂, I₃

5. Molecular shapes with six sets of electrons
A. Ax₆ Type
 Molecular formula- Ax₆
 Number of electron pairs- 6 bonding pairs
 Ideal bond angle-90⁰
 Molecular geometry- octahedral

Example:- SF₆

B. Ax₅E Type
 Molecular formula- Ax₅E
 Number of electron pairs- 5 bonding and one lone of e-pair of electrons.
 Molecular geometry- square pyramidal

• Example:- BrF₅
  

C. Ax₄E₂ Type
 Molecular formula – Ax₄E₂
 Number of electron pairs- 4 bonding and 2 lone- pair of electrons
 Molecular geometry- square planar

• Example:- XeF₄

https://phet.colorado.edu/sims/html/molecule-shapes/latest/molecule-shapes_all.html

Merits of Knowing Shape of Covalent Compounds

• Knowing the shape (geometry) of covalent molecules is a key to understand its physical and chemical behavior.
  One of the most important and far-reaching effects of molecular shape is molecular polarity, which can influence melting and boiling points, solubility, chemical reactivity, and even biological function.
• Bond Polarity In general, a covalent bond is:
  polar if it occurs between two different atoms.
• non polar if it occurs between two identical atoms

Dipole Moments

• – It is a measure of the polarity of a bond.- Is often represented by a special arrow.

Diatomic Molecules – molecules made of only two atoms.
If atoms are the same, molecule is nonpolar.
If atoms are diff., molecule is polar.
NOTE: Polar does not mean charged.
Is Cl2 polar or nonpolar?
Is CO polar or nonpolar?

Molecules With 3 or More Atoms

• A molecule with 3 or more atoms is:
  Polar if its central atom has lone pairs OR
  If the outer atoms are not all the same.
  Non-polar
• if its central atom has no lone pairs and All the outer atoms are identical.
• Example: CO₂ vs. H₂O
  i) Consider the Lewis structure of CO₂:
  This molecule is non-polar, and water is polar

Nonpolar and polar molecules

Intermolecular Force in the covalent bond

• There are two types of force that holds matter together. These are
• Intramolecular forces
• Intermolecular forces.
• Intramolecular force is a chemical bond (ionic, covalent or metallic)that exists within a particle (molecule or polyatomic ion) and affects the chemical property of the species.
• Intermolecular force are those bonds that hold particles (ions or molecules) together.
• A glass of water for example, contains many molecules of water.
• These molecules are held together by intermolecular forces, whereas the intramolecular forces hold the two hydrogen atoms to the oxygen atom in each molecule of water.
• Intermolecular forces are relatively weak as compared to intramolecular forces, because they typically involve lower charges that are farther apart. However, the strength of the intermolecular forces is important because they affect physical properties of the species such as melting point and boiling point.
• Three types of attractive force are known to exist between neutral molecules:
• Dipole–dipole forces
• London(or dispersion)forces,
• Hydrogen bonding

The term Vander Waals Forces area general term for those intermolecular forces that include dipole–dipole and London forces.

• Van der Waals forces are the weak attractive forces in a large number of substances, including Cl2, and Br2.

Dipole-Dipole forces

• Dipole-dipole forces act between the molecules possessing permanent dipole.
• When polar molecules are brought near one another, their partial charges act as tiny electric fields that orient them and give rise to dipole-dipole forces; the partially positive end of one molecule attracts the partially negative end of another.
• Ends of the dipoles possess “partial charges” and these charges are shown by Greek letter delta (δ).

Dispersion or London Forces

• Usually, the electrons in anon-polar covalent molecule are distributed symmetrically.
• However, the movement of the electrons may place more of them in one part of the molecule than another, which forms a temporary dipole.
• These momentary dipoles align the molecules so that the positive end of one molecule is attracted to the negative end of another molecule.
• This interaction is stronger than the London forces but is weaker than ion-ion interaction because only partial charges are involved.
• The attractive force decreases with the increase of distance between the dipoles.

Hydrogen Bonding

• Polar molecules containing hydrogen atoms bonded to highly electronegative atoms of nitrogen, oxygen, or fluorine form especially strong dipole–dipole attractions
• Hydrogen bonds are the strongest type of attractive forces between polar covalent molecules

https://phet.colorado.edu/en/simulations/molecule-polarity