Chemical Bonding and Hybridization Notes

Introduction to Chemical Bonding

Chemical bonding refers to the process by which atoms combine to form compounds. The main types of bonds are ionic bonds, covalent bonds, and metallic bonds. Each bond type has unique characteristics, depending on how electrons are shared or transferred between atoms.

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Types of Chemical Bonds

Ionic Bond:

• Formed when electrons are transferred from one atom (usually a metal) to another atom (usually a nonmetal).
• Results in the formation of positively charged ions (cations) and negatively charged ions (anions) that are held together by electrostatic forces.
• Example: Sodium chloride (NaCl), where sodium donates an electron to chlorine.

Covalent Bond:

• Formed when two atoms (usually nonmetals) share electrons.
• Creates molecules with a stable electron configuration.
• Example: Water (H₂O), where oxygen shares electrons with hydrogen atoms.

Metallic Bond:

• Occurs in metals, where electrons are delocalized and free to move throughout the metal lattice.
• Gives metals properties like conductivity and malleability.

Coordinate (Dative Covalent) Bond:

• A type of covalent bond where both shared electrons come from the same atom.
• Commonly found in complex ions.
• Example: Ammonium ion (NH₄⁺), where a lone pair from nitrogen bonds with a hydrogen ion.

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Hybridization Concept

Hybridization is the mixing of atomic orbitals in an atom to form new hybrid orbitals, which have different shapes and energies. Hybrid orbitals allow atoms to form bonds in specific orientations, increasing molecular stability.

Importance of Hybridization

• Hybridization explains the shapes and bond angles in molecules, which are key to understanding molecular geometry and bonding behavior.
• By hybridizing orbitals, atoms can form bonds that have equal strength and are symmetrically arranged in space.

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Types of Hybridization and Examples

sp Hybridization

sp hybridization occurs when one s orbital and one p orbital from the same main energy level combine to create two equivalent orbitals. These newly formed orbitals are known as sp hybrid orbitals.

• Orbitals Involved: 1 s orbital and 1 p orbital
• Geometry: Linear
• Bond Angle: 180°
• Example: Beryllium chloride (BeCl₂), acetylene (C₂H₂)
• Explanation: In acetylene, each carbon undergoes sp hybridization, creating two sp hybrid orbitals that form a triple bond.

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sp² Hybridization

• Orbitals Involved: 1 s orbital and 2 p orbitals
• Geometry: Trigonal planar
• Bond Angle: 120°
• Example: Boron trifluoride (BF₃), ethene (C₂H₄)
• Explanation: In ethene, each carbon atom undergoes sp² hybridization, forming a double bond with planar geometry.

sp³ Hybridization

• Orbitals Involved: 1 s orbital and 3 p orbitals
• Geometry: Tetrahedral
• Bond Angle: 109.5°
• Example: Methane (CH₄), ammonia (NH₃), water (H₂O)
• Explanation: In methane, the carbon atom’s sp³ hybrid orbitals form a tetrahedral structure, making four sigma bonds with hydrogen atoms.

sp³d Hybridization

• Orbitals Involved: 1 s orbital, 3 p orbitals, and 1 d orbital
• Geometry: Trigonal bipyramidal
• Bond Angles: 90°, 120°
• Example: Phosphorus pentachloride (PCl₅)
• Explanation: In PCl₅, phosphorus uses sp³d hybrid orbitals to accommodate five bonds with a trigonal bipyramidal shape.

sp³d² Hybridization

• Orbitals Involved: 1 s orbital, 3 p orbitals, and 2 d orbitals
• Geometry: Octahedral
• Bond Angle: 90°
• Example: Sulfur hexafluoride (SF₆)
• Explanation: In SF₆, sulfur undergoes sp³d² hybridization, forming an octahedral structure with six fluorine atoms.

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Summary Table of Hybridization Types

Hybridization	Orbitals Involved	Geometry	Bond Angle	Example Compounds
sp	1 s, 1 p	Linear	180°	BeCl₂, C₂H₂
sp²	1 s, 2 p	Trigonal planar	120°	BF₃, C₂H₄
sp³	1 s, 3 p	Tetrahedral	109.5°	CH₄, NH₃, H₂O
sp³d	1 s, 3 p, 1 d	Trigonal bipyramidal	90°, 120°	PCl₅
sp³d²	1 s, 3 p, 2 d	Octahedral	90°	SF₆

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Molecular Geometry and Hybridization

Molecular geometry depends on the hybridization of the central atom and the arrangement of atoms around it. Here’s how hybridization affects geometry:

• Linear Geometry (sp): Molecules like BeCl₂ and CO₂ have a linear shape due to sp hybridization.
• Trigonal Planar Geometry (sp²): Molecules like BF₃ and NO₃⁻ adopt a trigonal planar shape due to sp² hybridization.
• Tetrahedral Geometry (sp³): Methane and ammonium ions have tetrahedral geometry, associated with sp³ hybridization.
• Trigonal Bipyramidal Geometry (sp³d): Molecules like PCl₅ exhibit a trigonal bipyramidal shape due to sp³d hybridization.
• Octahedral Geometry (sp³d²): Molecules like SF₆ have an octahedral shape, indicative of sp³d² hybridization.

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Bonding Theories and Hybridization

Valence Bond Theory (VBT)

• VBT explains bonding through the overlap of atomic orbitals.
• Hybridization occurs to maximize the overlap, resulting in stronger bonds.

Molecular Orbital Theory (MOT)

• MOT combines atomic orbitals into molecular orbitals that span the entire molecule.
• While MOT does not directly describe hybridization, it complements the understanding of bonding and electron distribution.

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Hybridization in carbon atom

The concept of hybridization in carbon atoms is fundamental to understanding how carbon forms various types of bonds in organic molecules. Carbon has four valence electrons and can form multiple bonding arrangements by hybridizing its orbitals.

Types of Hybridization in Carbon

Carbon can undergo three main types of hybridization, depending on the number of bonds it forms and the type of molecule it’s in. These are sp³ hybridization, sp² hybridization, and sp hybridization.

sp³ Hybridization (Tetrahedral Geometry)

• Formation: One 2s orbital mixes with all three 2p orbitals to create four equivalent sp³ hybrid orbitals.
• Bonding: Each sp³ orbital contains one electron and can form a single sigma (σ) bond with another atom.
• Geometry: Tetrahedral with a bond angle of 109.5°.
• Example: Methane (CH₄), where each hydrogen atom bonds with an sp³ hybrid orbital of carbon, resulting in a stable tetrahedral structure.

sp² Hybridization (Trigonal Planar Geometry)

• Formation: One 2s orbital mixes with two of the 2p orbitals to form three sp² hybrid orbitals, leaving one unhybridized p orbital.
• Bonding: The three sp² orbitals form sigma bonds in a planar arrangement, while the unhybridized p orbital can form a pi (π) bond with another atom.
• Geometry: Trigonal planar with a bond angle of 120°.
• Example: Ethene (C₂H₄), where each carbon forms three sigma bonds using sp² orbitals and a pi bond with each other using the unhybridized p orbitals.

sp Hybridization (Linear Geometry)

• Formation: One 2s orbital mixes with one 2p orbital to create two sp hybrid orbitals, leaving two unhybridized p orbitals.
• Bonding: The two sp orbitals form sigma bonds in a linear arrangement, and the two unhybridized p orbitals form two pi bonds.
• Geometry: Linear with a bond angle of 180°.
• Example: Acetylene (C₂H₂), where each carbon forms two sigma bonds in a straight line and two pi bonds with each other using the unhybridized p orbitals.

Summary Table of Carbon Hybridization

Hybridization	Orbitals Mixed	Geometry	Bond Angles	Example
sp³	1 s, 3 p	Tetrahedral	109.5°	CH₄
sp²	1 s, 2 p	Trigonal Planar	120°	C₂H₄
sp	1 s, 1 p	Linear	180°	C₂H₂

Why Hybridization Is Important for Carbon Compounds

Hybridization allows carbon to form diverse structures and bond geometries, which is why it can create such a wide variety of organic compounds. The hybrid orbitals maximize bond strength and stabilize the molecule, contributing to the versatility of carbon in organic chemistry.

How sp³d² Hybridization Forms in SF₆

• Electron Configuration of Sulfur:

The ground-state electron configuration of sulfur (S) is 1s² 2s² 2p⁶ 3s² 3p⁴. In SF₆, sulfur needs to form six bonds, which requires six hybrid orbitals.

• Excited State and Orbital Mixing:

To achieve six bonding sites, sulfur promotes two electrons from its 3p orbitals to vacant 3d orbitals, leading to the configuration 3s¹ 3p³ 3d² in an excited state.

• Hybridization Process:

• One 3s orbital, three 3p orbitals, and two 3d orbitals combine to form six equivalent sp³d² hybrid orbitals.
• Each of these six sp³d² hybrid orbitals holds one electron, allowing sulfur to form six sigma (σ) bonds with six fluorine atoms.

• Bonding and Geometry:

• The six sp³d² hybrid orbitals arrange themselves in an octahedral geometry to minimize repulsion, with bond angles of 90° between each pair of bonds.
• Each fluorine atom uses a p orbital to overlap with an sp³d² hybrid orbital of sulfur, creating strong, stable sigma bonds.

Structure of SF₆

• Geometry: Octahedral, meaning each fluorine atom is symmetrically positioned around the sulfur atom.
• Bond Angles: 90° between each pair of adjacent bonds.
• Symmetry: High symmetry in the octahedral structure leads to equal bond lengths and bond energies for all six S-F bonds.

In SF₆, sp³d² hybridization enables sulfur to expand its octet and accommodate six fluorine atoms, resulting in an octahedral molecular structure. This type of hybridization is possible because sulfur can access its d orbitals, allowing it to form more than the typical four bonds found in sp³ hybridized atoms.

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Covalent Bonding

Covalent bonding involves the sharing of electron pairs between atoms, allowing them to achieve stable electron configurations. This bond formation occurs through the overlapping of atomic orbitals to create molecular orbitals, which hold the shared electrons between atoms. Covalent bonds can be classified into sigma (σ) bonds and pi (π) bonds, each formed through specific types of orbital overlaps.

Atomic Orbitals in Covalent Bonding

Atoms have specific orbitals (s, p, d, etc.) where electrons reside. When atoms approach each other to form covalent bonds, these orbitals overlap, creating new molecular orbitals.

Sigma (σ) Bonds:

• A sigma bond is the first bond formed between two atoms in a covalent bond.
• It results from head-on (end-to-end) overlap of atomic orbitals along the internuclear axis (the line between the nuclei of the two atoms).
• Sigma bonds can form between different types of orbitals:
  • s-s overlap: For example, in an H₂ molecule where two s orbitals overlap.
  • s-p overlap: Seen in HCl, where an s orbital from hydrogen overlaps with a p orbital from chlorine.
  • p-p overlap: Seen in F₂, where two p orbitals overlap.
• Characteristics:
  • Sigma bonds are strong and allow free rotation around the bond axis.
  • They form the primary structural framework of the molecule, creating single bonds or the first bond in double and triple bonds.

Pi (π) Bonds:

• Pi bonds are formed after a sigma bond when atoms form double or triple bonds.
• They result from the sideways (parallel) overlap of unhybridized p orbitals above and below the internuclear axis.
• Pi bonds form only between p orbitals and contribute to the electron density above and below the plane of the nuclei.
• Characteristics:
  • Pi bonds are weaker than sigma bonds because the orbital overlap is less direct.
  • They restrict rotation around the bond axis, leading to more rigid molecular structures.
  • In double bonds, there is one sigma bond and one pi bond, while in triple bonds, there is one sigma bond and two pi bonds.

Molecular Orbital Formation in Covalent Bonds

When atomic orbitals overlap, they create molecular orbitals that are shared by both atoms involved in the bond:

• Bonding Molecular Orbitals: Lower-energy orbitals that result from constructive interference of atomic orbitals; electrons in these orbitals stabilize the bond.
• Antibonding Molecular Orbitals: Higher-energy orbitals that result from destructive interference; these orbitals can destabilize a molecule if occupied.

Examples of Covalent Bonds with Sigma and Pi Bonds

Single Bonds (Sigma Bonds Only):

• Example: Methane (CH₄). In CH₄, each hydrogen forms a sigma bond with carbon, using sp³ hybridized orbitals.
• Hybridization: Carbon in methane undergoes sp³ hybridization, creating four equivalent sigma bonds with no pi bonds.

Double Bonds (One Sigma and One Pi Bond):

• Example: Ethene (C₂H₄). In C₂H₄, each carbon shares two electrons to form a double bond (one sigma and one pi bond).
• Hybridization: Each carbon undergoes sp² hybridization to form three sigma bonds, and the unhybridized p orbitals form one pi bond.

Triple Bonds (One Sigma and Two Pi Bonds):

• Example: Acetylene (C₂H₂). In C₂H₂, each carbon atom forms a triple bond (one sigma and two pi bonds).
• Hybridization: Carbon atoms in acetylene undergo sp hybridization, forming one sigma bond and two pi bonds, with the pi bonds providing additional bonding strength but restricting rotation.

Differences Between Sigma (σ) Bond and Pi (π) Bond

Feature	Sigma (σ) Bond	Pi (π) Bond
Formation	Formed by the head-on (end-to-end) overlap of atomic orbitals	Formed by the sideways (parallel) overlap of atomic orbitals
Orbital Overlap	Stronger overlap, as orbitals directly face each other	Weaker overlap, as orbitals overlap sideways
Electron Density	Electron density is concentrated along the internuclear axis (line connecting nuclei)	Electron density is above and below the internuclear axis
Strength	Generally stronger due to greater orbital overlap	Generally weaker due to less orbital overlap
Bond Type	Found in single bonds and as the first bond in double or triple bonds	Found only in multiple bonds (second bond in a double bond, second and third in a triple bond)
Rotation	Allows free rotation around the bond axis	Restricts rotation, as rotating the atoms would break the bond
Shape of Orbitals	Formed by overlapping s-s, s-p, or p-p orbitals	Formed exclusively by p-p orbital overlap
Occurrence in Molecules	Present in all types of covalent bonds (single, double, triple)	Present only in multiple bonds (double and triple bonds)

Summary:

• Sigma bonds provide the main bond axis and can exist alone in single bonds.
• Pi bonds contribute to the bond strength in multiple bonds but restrict rotation and are less stable than sigma bonds.

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Example of Covalent Bond Formation 

Formation of the HCl Molecule

The molecule hydrogen chloride (HCl) is formed through covalent bonding between hydrogen (H) and chlorine (Cl) atoms. Here’s how the bonding process occurs, focusing on electron sharing and orbital overlap:

Electron Configuration of Hydrogen and Chlorine:

• Hydrogen (H): Has one electron in its 1s orbital (configuration: 1s¹).
• Chlorine (Cl): Has seven valence electrons, with the configuration 1s² 2s² 2p⁶ 3s² 3p⁵, needing one more electron to complete its outer shell and achieve the stable octet configuration.

Bond Formation through Electron Sharing:

• To achieve stability, the hydrogen atom needs one additional electron to fill its 1s orbital, while chlorine needs one electron to complete its 3p orbital.
• Hydrogen and chlorine atoms come together and share a pair of electrons: one from hydrogen and one from chlorine, forming a single covalent bond.

Orbital Overlap and Formation of a Sigma (σ) Bond:

• The 1s orbital of hydrogen overlaps with a 3p orbital of chlorine along the internuclear axis (the line connecting the two nuclei).
• This head-on overlap creates a sigma (σ) bond, which is the primary bond holding the H and Cl atoms together.
• The electron density of the sigma bond is concentrated along the internuclear axis, providing a strong, stable bond.

Polarity of the HCl Molecule:

• Chlorine is more electronegative than hydrogen, meaning it attracts the shared electrons more strongly.
• As a result, the shared electron pair spends more time closer to the chlorine atom, creating a polar covalent bond.
• This electron distribution gives chlorine a partial negative charge (δ-) and hydrogen a partial positive charge (δ+), making HCl a polar molecule.

Summary

The HCl molecule forms through covalent bonding with a sigma bond resulting from the overlap of hydrogen’s 1s orbital and chlorine’s 3p orbital. The bond is polar due to chlorine’s higher electronegativity, leading to a dipole moment in the molecule. This polarity is significant for HCl’s behavior in solutions and interactions with other molecules.

Formation of Methane (CH₄) and Hybridization

Methane (CH₄) is a simple molecule with a tetrahedral structure, where carbon forms single covalent bonds with four hydrogen atoms. To understand how methane forms, we need to look at the hybridization of carbon’s orbitals.

Step-by-Step Formation of CH₄

Electron Configuration of Carbon:

• Carbon (C) has an atomic number of 6, with the ground-state electron configuration: 1s² 2s² 2p².
• In the 2nd shell, carbon has four valence electrons (2s² 2p²), which it uses to form bonds.

Need for Hybridization:

• To form four equivalent bonds with hydrogen, carbon needs four orbitals with one unpaired electron each.
• However, in its ground state, carbon only has two unpaired electrons in the 2p orbitals and two paired electrons in the 2s orbital, which does not allow for four bonds.
• Solution: Carbon undergoes sp³ hybridization, where one 2s and three 2p orbitals combine to form four equivalent sp³ hybrid orbitals.

sp³ Hybridization:

• In sp³ hybridization, the 2s and three 2p orbitals mix to produce four new, equivalent sp³ hybrid orbitals.
• Each sp³ orbital has one electron, allowing carbon to form four sigma (σ) bonds.
• These sp³ hybrid orbitals arrange themselves in a tetrahedral geometry with bond angles of 109.5° to minimize electron repulsion.

Bond Formation with Hydrogen:

• Each of the four sp³ hybrid orbitals on carbon overlaps with the 1s orbital of a hydrogen atom, forming a sigma bond.
• These sigma bonds are single covalent bonds, where each shared electron pair resides in a molecular orbital between carbon and hydrogen.

Molecular Geometry of CH₄:

• The resulting structure of CH₄ is tetrahedral, with four C-H bonds evenly spaced at 109.5°.
• This geometry gives methane its stability and symmetrical shape, with no lone pairs on the central carbon atom.

Summary

The formation of methane (CH₄) involves:

• sp³ hybridization of carbon’s orbitals to create four equivalent sp³ hybrid orbitals.
• Each sp³ orbital overlaps with a hydrogen 1s orbital, forming four sigma bonds in a tetrahedral geometry with 109.5° bond angles.
• This bonding arrangement makes methane a stable, nonpolar molecule.

Formation of Ammonia, (NH₃)

Ammonia (NH₃) forms through covalent bonding, where nitrogen (N) and hydrogen (H) atoms share electrons to achieve stable electron configurations. Here's a detailed explanation of how ammonia forms:

Electron Configuration of Nitrogen and Hydrogen:

• Nitrogen has an atomic number of 7, meaning it has 7 electrons. Its electron configuration is 1s² 2s² 2p³. To achieve a stable configuration (similar to the nearest noble gas, neon), nitrogen needs 3 more electrons to complete its outer shell (which can hold 8 electrons).
• Hydrogen has an atomic number of 1, with an electron configuration of 1s¹. Each hydrogen atom needs one more electron to achieve the stable configuration of helium.

Bond Formation:

• In the ammonia molecule, one nitrogen atom bonds with three hydrogen atoms.
• Nitrogen shares one electron with each hydrogen atom, forming three covalent bonds. Each of the three hydrogen atoms shares its one electron with nitrogen. This sharing allows both nitrogen and hydrogen to achieve a more stable electron configuration.

Lewis Structure of Ammonia (NH₃):

• The nitrogen atom has three single bonds with hydrogen atoms, and it also has one lone pair of electrons (two electrons) remaining on it.
• The nitrogen atom in NH₃ uses three of its electrons to form covalent bonds with the three hydrogen atoms, while the remaining two electrons stay as a lone pair on nitrogen.

Properties of Covalent Bonding in Ammonia:

• Ammonia has a polar covalent bond, where the nitrogen atom is more electronegative than the hydrogen atoms. This creates a slight negative charge on the nitrogen and a slight positive charge on the hydrogen atoms, giving the molecule a polar nature.
• The molecule has a trigonal pyramidal shape due to the lone pair of electrons on nitrogen, which repels the bonding pairs of electrons.

This formation of ammonia via covalent bonding results in a stable molecule where nitrogen and hydrogen share electrons to complete their valence electron shells.

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WASSCE Questions and Answers on Chemical Bonding and Hybridization 

Question 1 

Explain why CO2 is linear molecule but H20 is not 

Solution 

The difference in molecular shape between carbon dioxide (CO₂) and water (H₂O) is due to the arrangement of electron pairs around the central atom, as described by the Valence Shell Electron Pair Repulsion (VSEPR) theory. Here’s an explanation of why CO₂ is linear, while H₂O is bent:

a. Structure of CO₂

• Molecular Formula: CO₂
• Central Atom: Carbon (C)
• Electron Configuration: Carbon has four valence electrons, while each oxygen has six.
• Bonding: In CO₂, the carbon atom forms two double bonds with two oxygen atoms—one on each side.
• Electron Geometry: The two double bonds around the central carbon atom create a situation where there are only two regions of electron density (each double bond counts as one region of electron density in VSEPR theory).
• Molecular Geometry: According to VSEPR theory, to minimize repulsion, these two regions of electron density will orient themselves 180° apart. This results in a linear geometry for CO₂, where the bond angle between the oxygens is 180°, giving it the linear shape: $$\text{O} = \text{C} = \text{O}$$
  
• Non-Polar Nature: Additionally, CO₂ is non-polar because the bond dipoles (from carbon to oxygen) are equal in magnitude and opposite in direction, canceling each other out.

b. Structure of H₂O

• Molecular Formula: H₂O
• Central Atom: Oxygen (O)
• Electron Configuration: Oxygen has six valence electrons, needing two more to complete its octet.
• Bonding: In H₂O, the oxygen atom forms two single bonds with two hydrogen atoms.
• Electron Geometry: Oxygen has four regions of electron density around it: two bonding pairs (with hydrogen atoms) and two lone pairs of electrons.
• Molecular Geometry: According to VSEPR theory, these four regions of electron density around the oxygen atom adopt a tetrahedral electron geometry to minimize repulsion. However, because only two of these regions are bonding pairs, the molecular geometry of H₂O is bent or angular. The lone pairs exert slightly more repulsion than bonding pairs, resulting in a bond angle of approximately 104.5° between the hydrogen atoms: $$\text{H} - \text{O} - \text{H}$$
  
• Polar Nature: This bent shape makes H₂O a polar molecule because the bond dipoles do not cancel out, resulting in a partial negative charge on the oxygen atom and a partial positive charge on the hydrogen atoms.

Summary

• CO₂ has a linear shape because it has two double bonds that arrange themselves 180° apart, with no lone pairs on the central carbon atom.
• H₂O has a bent shape due to two lone pairs on the central oxygen atom, which push the two hydrogen atoms closer together, resulting in an angular shape.

Question 2

(a) Give the type of hybridization shown by the central atoms in the following compounds: NH₃, H₂O, BCl₃.
(b) Deduce the shape of the molecules.

Solution:

Part (a): Hybridization of the Central Atoms

To determine the hybridization of the central atom in each molecule, we consider the number of regions of electron density (bonds and lone pairs) around it.

1. NH₃ (Ammonia)

• Central Atom: Nitrogen (N)
• Electron Regions: 3 bonding pairs (with H atoms) + 1 lone pair = 4 regions
• Hybridization: The four regions of electron density around nitrogen require sp³ hybridization.

2. H₂O (Water)

• Central Atom: Oxygen (O)
• Electron Regions: 2 bonding pairs (with H atoms) + 2 lone pairs = 4 regions
• Hybridization: The four regions of electron density around oxygen also require sp³ hybridization.

3. BCl₃ (Boron Trichloride)

• Central Atom: Boron (B)
• Electron Regions: 3 bonding pairs (with Cl atoms) + 0 lone pairs = 3 regions
• Hybridization: The three regions of electron density around boron require sp² hybridization.

Part (b): Molecular Shapes of NH₃, H₂O, and BCl₃

Using the VSEPR theory, we can deduce the shapes of the molecules based on their hybridization and electron pairs:

1. NH₃ (Ammonia)

• Hybridization: sp³
• Shape: The sp³ hybridization suggests a tetrahedral electron geometry. However, due to the lone pair on nitrogen, the molecular geometry becomes trigonal pyramidal.
• Bond Angle: Approximately 107°.

2. H₂O (Water)

• Hybridization: sp³
• Shape: The sp³ hybridization indicates a tetrahedral electron geometry. With two lone pairs on oxygen, the molecular geometry is bent or angular.
• Bond Angle: Approximately 104.5°.

3. BCl₃ (Boron Trichloride)

• Hybridization: sp²
• Shape: The sp² hybridization leads to a trigonal planar molecular geometry, as there are no lone pairs on boron.
• Bond Angle: 120°.

Question 3

a) Draw the complete structural formula for the compound CH₃C≡CCH₂CN.

(b) Indicate the type of hybridization for each carbon atom in the compound.

(c) For each carbon atom in the compound, state the geometry of the hybridized orbitals.

Solution:

To analyze the compound CH₃C≡CCH₂CN, we will:

1. Draw the complete structural formula.
2. Determine the hybridization of each carbon atom.
3. State the geometry associated with each carbon atom based on its hybridization.

(a) Complete Structural Formula of CH₃C≡CCH₂CN

The compound CH₃C≡CCH₂CN consists of different groups:

• CH₃ (methyl group) attached to a C≡C (carbon-carbon triple bond).
• CH₂ attached to CN (nitrile group).

The complete structural formula is:

    H   H
    |   |
H—C—C≡C—C—C≡N
    |   |
    H   H

In condensed notation, it can be represented as: CH₃—C≡C—CH₂—C≡N

(b) Type of Hybridization for Each Carbon Atom

We determine the hybridization based on the bonding and electron density regions around each carbon:

1. Carbon 1 (CH₃ group):
   • This carbon forms three single bonds with three hydrogens and one single bond with the adjacent carbon.
   • Hybridization: sp³ (4 bonding regions)
2. Carbon 2 (C≡C):
   • This carbon is involved in a triple bond with the next carbon and a single bond with Carbon 1.
   • Hybridization: sp (2 bonding regions)
3. Carbon 3 (C≡C):
   • Similar to Carbon 2, this carbon forms a triple bond with Carbon 2 and a single bond with Carbon 4.
   • Hybridization: sp (2 bonding regions)
4. Carbon 4 (CH₂ group):
   • This carbon is bonded to two hydrogens and has single bonds to Carbon 3 and Carbon 5.
   • Hybridization: sp³ (4 bonding regions)
5. Carbon 5 (C≡N):
   • This carbon forms a triple bond with the nitrogen in the nitrile group and a single bond with Carbon 4.
   • Hybridization: sp (2 bonding regions)

(c) Geometry of the Hybridized Orbitals for Each Carbon Atom

The geometry of each carbon atom is based on its hybridization type:

1. Carbon 1 (CH₃ group):

• Hybridization: sp³
• Geometry: Tetrahedral (bond angles approximately 109.5°)

2. Carbon 2 (C≡C):

• Hybridization: sp
• Geometry: Linear (bond angle 180°)

3. Carbon 3 (C≡C):

• Hybridization: sp
• Geometry: Linear (bond angle 180°)

4. Carbon 4 (CH₂ group):

• Hybridization: sp³
• Geometry: Tetrahedral (bond angles approximately 109.5°)

5. Carbon 5 (C≡N):

• Hybridization: sp
• Geometry: Linear (bond angle 180°)

Summary Table

Carbon Atom	Group	Hybridization	Geometry	Bond Angle
Carbon 1	CH₃	sp³	Tetrahedral	~109.5°
Carbon 2	C≡C	sp	Linear	180°
Carbon 3	C≡C	sp	Linear	180°
Carbon 4	CH₂	sp³	Tetrahedral	~109.5°
Carbon 5	C≡N	sp	Linear	180°

 

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Dipole Moments and Intermolecular Forces

Dipole Moments

A dipole moment is a measure of the separation of positive and negative electrical charges within a molecule, indicating the overall polarity of the molecule. Dipole moments are vector quantities, meaning they have both direction and magnitude, represented by the symbol μ. Molecules with a dipole moment are referred to as polar molecules, while those without a dipole moment are nonpolar.

Formation of Dipole Moments

• Electronegativity Difference: When two atoms in a bond have different electronegativities (the ability of an atom to attract electrons), the electrons are not shared equally. The more electronegative atom pulls electron density toward itself, creating a partial negative charge (δ⁻), while the less electronegative atom takes on a partial positive charge (δ⁺).
• Bond Polarity: This unequal sharing of electrons results in a polar covalent bond. The separation of charges within the bond creates an electric dipole. The bond dipole points from the positive pole to the negative pole, representing the direction of electron density shift.
• Symmetry and Dipole Cancellation: In molecules with symmetric structures, individual bond dipoles may cancel each other out, resulting in a nonpolar molecule overall. For example, in carbon dioxide (CO₂), the linear geometry causes the dipoles from the C=O bonds to cancel, leading to a net dipole moment of zero, even though each C=O bond is polar.

Calculating Dipole Moment

• The dipole moment (μ) is calculated using the formula: $$\mu = q \times d$$where:
  • $q$ is the magnitude of the charge separation.
  • $d$ is the distance between the positive and negative charges.
• Units: Dipole moments are measured in debye units (D)

Molecular Dipole Moment and Molecular Shape

• Molecular Shape Influence: The overall dipole moment of a molecule depends on both the individual bond dipoles and the 3D structure (geometry) of the molecule. Molecular geometry determines whether the bond dipoles reinforce each other or cancel out.
  • Bent Shape (e.g., water, H₂O): Water has an angular structure with a bond angle of about 104.5°, and the individual dipoles from the O–H bonds do not cancel out, resulting in a strong net dipole moment.
  • Linear Shape (e.g., carbon dioxide, CO₂): CO₂ has a linear shape with two C=O bonds in opposite directions, causing their dipoles to cancel out. As a result, CO₂ has no net dipole moment.
• Polarity of Molecules: Molecules with a net dipole moment are polar, while those with zero net dipole moment are nonpolar. Polarity influences many physical and chemical properties, such as solubility, boiling points, and interactions with other molecules.

Significance of Dipole Moments

• Predicting Molecular Polarity: Dipole moments are a quantitative indicator of molecular polarity. The higher the dipole moment, the more polar the molecule is.
• Influence on Intermolecular Forces: Polar molecules (with dipole moments) tend to interact more strongly with each other than nonpolar molecules, affecting physical properties like melting and boiling points.
• Applications: Dipole moments are crucial in molecular modeling, predicting molecular interactions, understanding reactivity patterns, and designing materials with specific polar or nonpolar characteristics.

Intermolecular Forces

Intermolecular forces (IMFs) are attractive forces that occur between molecules or ions. While they are weaker than covalent or ionic bonds within molecules (intramolecular forces), IMFs significantly influence the physical properties of substances, such as boiling points, melting points, viscosity, and solubility.

Types of Intermolecular Forces

1. Dipole-Dipole Interactions

• Definition: Dipole-dipole interactions are attractive forces that occur between polar molecules with permanent dipole moments.
• Mechanism: The partially positive end (δ⁺) of one molecule attracts the partially negative end (δ⁻) of another. This alignment creates an attractive force, stabilizing the molecules in a particular orientation.
• Examples: Hydrogen chloride (HCl) has a dipole-dipole interaction where the δ⁺ end of hydrogen in one HCl molecule attracts the δ⁻ end of chlorine in another HCl molecule.
• Strength: Dipole-dipole forces are generally stronger than London dispersion forces but weaker than hydrogen bonds and covalent or ionic bonds.

2. Hydrogen Bonding

• Definition: Hydrogen bonding is a special, stronger type of dipole-dipole interaction occurring when hydrogen is covalently bonded to highly electronegative atoms (such as nitrogen, oxygen, or fluorine).
• Mechanism: The highly electronegative atom pulls electron density away from hydrogen, leaving it with a substantial partial positive charge. This hydrogen is then attracted to the lone pairs on electronegative atoms in neighboring molecules.
• Examples: In water (H₂O), each hydrogen atom is attracted to the lone pairs on oxygen atoms in neighboring water molecules, creating extensive hydrogen bonding.
• Strength: Hydrogen bonds are the strongest type of dipole-dipole interaction, significantly raising the boiling and melting points of compounds with hydrogen bonding.

3. London Dispersion Forces (Van der Waals Forces)

• Definition: Also known as dispersion forces, these temporary attractive forces arise due to momentary shifts in electron density within molecules.
• Mechanism: Even nonpolar molecules can exhibit momentary dipoles due to random movement of electrons. These temporary dipoles induce similar dipoles in nearby molecules, leading to an attractive force.
• Examples: Noble gases like argon (Ar) and nonpolar molecules like methane (CH₄) primarily exhibit London dispersion forces.
• Strength: London dispersion forces are generally the weakest type of intermolecular force but increase in strength with molecular size and mass.

4. Ion-Dipole Forces

• Definition: Ion-dipole forces are the attraction between an ion and a polar molecule, commonly seen when salts dissolve in polar solvents.
• Mechanism: The ion interacts with the partially charged ends of polar molecules, creating a strong attraction. For instance, a cation (positive ion) will attract the negative end (δ⁻) of the dipole, while an anion (negative ion) will attract the positive end (δ⁺).
• Examples: When sodium chloride (NaCl) dissolves in water, the Na⁺ ions are attracted to the oxygen atoms in water (δ⁻), and the Cl⁻ ions are attracted to the hydrogen atoms in water (δ⁺).
• Strength: Ion-dipole forces are generally stronger than dipole-dipole interactions and are a primary factor in the solubility of ionic compounds in polar solvents.

Impact of Intermolecular Forces on Physical Properties

1. Boiling and Melting Points: Stronger intermolecular forces result in higher boiling and melting points since more energy is needed to separate the molecules. Water has a high boiling point for its molecular weight due to hydrogen bonding, while methane, with weaker London forces, boils at a much lower temperature.
2. Solubility: "Like dissolves like" is a rule based on intermolecular forces. Polar substances dissolve well in polar solvents (e.g., salt in water) because of dipole interactions, whereas nonpolar substances dissolve in nonpolar solvents (e.g., oil in hexane) due to dispersion forces.
3. Surface Tension: Molecules at the surface of a liquid experience a net inward force, creating surface tension. Liquids with strong intermolecular forces, like water (hydrogen bonding), have high surface tension, allowing objects to float despite being denser than water.
4. Viscosity: Viscosity, the resistance to flow, is higher in substances with strong intermolecular forces. Glycerol, with multiple hydroxyl (OH) groups, has higher viscosity due to hydrogen bonding compared to water.
5. Vapor Pressure: Substances with strong intermolecular forces have lower vapor pressures because fewer molecules can escape into the vapor phase. For instance, ethanol has lower vapor pressure than diethyl ether due to hydrogen bonding.

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Summary Table: Types of Intermolecular Forces

Type of Force	Occurs Between	Relative Strength	Examples
Dipole-Dipole	Polar molecules	Moderate	HCl, SO₂
Hydrogen Bonding	Molecules with H–N, H–O, or H–F bonds	Strong	H₂O, NH₃
London Dispersion Forces	All molecules (stronger in larger/heavier atoms)	Weak	CH₄, Ar
Ion-Dipole	Ions and polar molecules	Very Strong	NaCl in water

 

Understanding dipole moments and intermolecular forces is crucial in chemistry, as they determine molecular interactions, solubility, and physical states, and are foundational for materials science, biology, and environmental science.