Chemical Bonding and Shapes of Molecules - Chemistry (Undergraduate Foundation)

This course covers the principles that govern how atoms connect to form molecules. It provides a full treatment of ionic and covalent bonding, valence forces, and the various types of intermolecular forces, including hydrogen bonding. The core of the course is a practical guide to predicting the three-dimensional shape of molecules using VSEPR theory and the concept of orbital hybridization. A command of chemical bonding and molecular geometry is essential for understanding the properties and reactivity of all chemical substances. The shape of a molecule dictates its function, from the specificity of drug-receptor interactions in medicine to the properties of polymers in materials science. This is the foundational knowledge for designing new molecules and materials. By the end of this course, you will be able to draw Lewis structures for any molecule, predict its three-dimensional geometry and bond angles using VSEPR theory, and determine its polarity. You will also be able to explain the formation of sigma and pi bonds using the concept of orbital hybridization and identify the types of intermolecular forces present in a substance. This course is for students who have a complete understanding of atomic theory and electronic configuration. It is a mandatory course for all students of chemistry, biochemistry, and materials science, and a direct prerequisite for the study of organic chemistry.

Enrolment valid for 12 months
This course is also part of the following learning track. You may join the track to gain comprehensive knowledge across related courses.
[NUC Core] CHM 101: General Chemistry I
[NUC Core] CHM 101: General Chemistry I
This learning track delivers the complete NUC CCMAS curriculum for General Chemistry I. It is a comprehensive programme designed to build a robust, university-level foundation in modern chemistry. The track systematically covers all essential topics, from atomic theory, chemical bonding, and the states of matter, to the quantitative principles of stoichiometry, equilibrium, thermodynamics, and kinetics. This programme is for first-year undergraduates in science, technology, engineering, and mathematics (STEM) faculties who are required to take CHM 101. It is also essential for any student or professional globally who needs a rigorous and complete foundation in first-year university chemistry for further study or career development. This track delivers a full skill set in chemical theory and quantitative problem-solving. Graduates will be able to determine molecular structures, calculate reaction quantities, analyse the energetics and rates of reactions, and solve complex equilibrium problems. This programme provides the non-negotiable prerequisite knowledge for all subsequent chemistry courses and for any degree in the physical sciences, engineering, or medicine.

This learning track delivers the complete NUC CCMAS curriculum for General Chemistry I. It is a comprehensive programme designed to build a robust, university-level foundation in modern chemistry. The track systematically covers all essential topics, from atomic theory, chemical bonding, and the states of matter, to the quantitative principles of stoichiometry, equilibrium, thermodynamics, and kinetics. This programme is for first-year undergraduates in science, technology, engineering, and mathematics (STEM) faculties who are required to take CHM 101. It is also essential for any student or professional globally who needs a rigorous and complete foundation in first-year university chemistry for further study or career development. This track delivers a full skill set in chemical theory and quantitative problem-solving. Graduates will be able to determine molecular structures, calculate reaction quantities, analyse the energetics and rates of reactions, and solve complex equilibrium problems. This programme provides the non-negotiable prerequisite knowledge for all subsequent chemistry courses and for any degree in the physical sciences, engineering, or medicine.

Course Chapters

1. Introduction
1

This chapter provides the roadmap for the course. It outlines the progression from the fundamental nature of chemical bonds to the theories that predict the three-dimensional shapes of molecules. Key learning objectives include: understanding the overall course structure and appreciating why the shape of a molecule is critical to its function.

Chapter lessons

1-1. Welcome
6:11

This lesson provides a brief overview of the course, outlining the key topics of chemical bonding, molecular geometry, and intermolecular forces.

2. Types of Bonds
9

This chapter establishes the foundation by classifying and detailing the strong intramolecular forces that hold atoms together. Mastery of ionic and covalent bonding is the prerequisite for subsequent analysis of molecular structure. We also introduce the critical hydrogen bond. Key objectives: Define ionic, covalent, and coordinate covalent bonds; contrast the physical properties of ionic and covalent substances; and understand hydrogen bonding.

Chapter lessons

2-1. Ionic bonds
11:05

This lesson defines the ionic bond. We explain electron transfer from a metal to a non-metal - using Sodium chloride (NaCl) as a primary example. The resulting electrostatic attraction between the ions is the fundamental bonding force.

2-2. Examples of ionic bonds (1)
12:11

This lesson reviews the formation of simple ionic compounds via electron transfer. We demonstrate bond formation for Sodium chloride (NaCl), Calcium chloride (CaCl2), and Lithium oxide (Li2O). Focus on calculating the correct number of electrons transferred to achieve stability.

2-3. Examples of ionic bonds (2)
12:42

This second lesson reinforces ionic bond formation via electron transfer. We demonstrate bond formation for Calcium chloride (CaCl2) and Lithium oxide (Li2O). Master how to balance the number of atoms to achieve a neutral compound.

2-4. Properties of ionic compounds
4:40

This lesson details the physical properties of ionic compounds. We explain how strong electrostatic forces cause high melting points, brittleness, and electrical conductivity when molten or in solution.

2-5. Covalent bonds
16:41

This lesson defines the covalent bond as the sharing of electron pairs between non-metal atoms. We examine single, double, and triple bonds using Chlorine (Cl2), Disulfur (S2), and Nitrogen (N2). Master the concept of shared electrons creating molecular stability.

2-6. Covalent bonds in heteronuclear molecules
12:18

This lesson covers covalent bonding between different atoms - heteronuclear molecules. We use Hydrogen fluoride (HF), Carbon dioxide (CO2), and Carbon monoxide (CO) to explain how differences in electronegativity lead to unequal electron sharing and polar bonds. Understand the resulting bond dipole moment.

2-7. Coordinate covalent bonds
13:46

This lesson defines the coordinate covalent bond. We use the Ammonium ion (NH4+) and the Hydronium ion (H3O+) to demonstrate bond formation where one atom supplies both shared electrons. Understand that the resulting bond is structurally identical to a standard covalent bond.

2-8. Polar covalent bonds
12:05

This lesson defines the polar covalent bond. We link differences in atomic electronegativity to unequal electron sharing - which creates a bond dipole. Bonds analysed include H-F, P-H, H-O, and C-Cl. Understand how bond polarity lies between pure covalent and ionic bonding.

2-9. Hydrogen bonding
12:55

This lesson defines hydrogen bonding - the strong intermolecular force. We explain its formation between hydrogen and N, O, or F. We contrast Water (H2O), Ethanol (C2H5OH), and Hydrogen fluoride (HF) with Hydrogen chloride (HCl) to establish the critical requirements.

3. Lewis Structures
7

This chapter establishes the Lewis structure framework, the fundamental method for mapping valence electron arrangement in molecules and ions. Mastering this structured approach is the prerequisite skill for predicting molecular geometry and understanding bond characteristics later in the course. We cover systematic rules for neutral species and polyatomic ions, including the essential concept of resonance. By the end of this chapter, you will master: drawing valid Lewis structures for molecules and ions; identifying and drawing resonance structures; applying the three main octet rule exceptions; and calculating and using formal charges to select the most plausible structure.

Chapter lessons

3-1. Neutral molecules
19:35

This lesson details rules for drawing Lewis structures for neutral molecules like Sulfur Dichloride (SCl2), Dichlorodifluoromethane (CCl2F2), Phosgene (COCl2), and Carbon Dioxide (CO2). We apply the systematic method: count valence electrons, arrange bonds, and place lone pairs.

3-2. Polyatomic ions
9:23

This lesson applies Lewis structure rules to polyatomic ions. We will draw the Tetrafluoroborate (BF4-), Hydronium (H3O+), and Disulfide (S22-) ions. The key step is adjusting the total valence electron count based on the overall charge.

3-3. Resonance structures
10:24

This lesson defines resonance for when a single Lewis structure fails. We examine Ozone (O3) and the Carbonate ion (CO32-). Learn to draw all valid contributing structures. Understand these average into a resonance hybrid - which represents the true electron delocalisation.

3-4. Exceptions to the octet rule
7:33

The octet rule is not absolute. This lesson details its three exceptions - incomplete octets, odd-electron species, and expanded octets. We examine Nitrogen monoxide (NO) and Boron trifluoride (BF3) to master the odd-electron and incomplete octet cases.

3-5. Formal charges
12:42

Learn to calculate formal charge for each atom. This tool is essential for evaluating competing Lewis structures - such as those for Phosgene (COCl2). We apply rules to select the most stable structure - which is the one that minimises formal charge.

3-6. Examples using formal charges (1)
26:21

This lesson demonstrates the practical application of formal charge. We review examples using the Chlorite ion (ClO2-) and Sulfuric acid (H2SO4). Master the rules for selecting the most plausible Lewis structure by minimising formal charges.

3-7. Examples using formal charges (2)
5:58

This lesson applies formal charge rules to structures with expanded octets. We use Phosphorus pentachloride (PCl5), Sulfur hexafluoride (SF6), and Sulfur tetrafluoride (SF4). Master how minimising formal charge determines the most plausible structure.

4. Shapes of Molecules
3
3

This chapter introduces Valence Shell Electron Pair Repulsion (VSEPR) theory, the primary model for predicting the three-dimensional geometry of molecules. The focus is on applying the theory to determine electron and molecular shapes. Key learning objectives include: using VSEPR theory to determine the electron geometry around a central atom; and predicting the final molecular geometry and bond angles for any simple molecule.

Chapter lessons

4-1. VSEPR theory
37:58

This lesson introduces Valence Shell Electron Pair Repulsion (VSEPR) theory. We use ABn notation - covering AB2, AB3, and AB4 domains - to predict electron and molecular geometry. Examples include Water (H2O), Beryllium chloride (BeCl2), Boron trifluoride (BF3), Methane (CH4), and Ammonia (NH3).

4-2. Examples (1)
12:30

This lesson is the first application of VSEPR theory. We demonstrate predicting molecular shapes for Tin(II) chloride (SnCl2), Sulfur tetrafluoride (SF4), Phosphorus pentafluoride (PF5), and Chlorine trifluoride (ClF3). Master the procedure for determining both electron and molecular geometry.

4-3. Electron vs molecular geometry

This lesson explains the crucial distinction between the electron geometry (the arrangement of all electron pairs) and the molecular geometry (the arrangement of only the atoms).

5. Hybridization Theory
3
4

This chapter covers the concept of orbital hybridization, a model used to explain the observed shapes of molecules that VSEPR theory alone cannot. It connects bonding to the mixing of atomic orbitals. Key learning objectives include: defining orbital hybridization; and identifying the correct hybridization (sp, sp², sp³) for the central atom in a molecule.

Chapter lessons

5-1. Valence bond theory

This lesson introduces valence bond theory, which describes a covalent bond as the overlap of atomic orbitals.

5-2. Hybridization of orbitals

This lesson defines hybridization as the concept of mixing atomic orbitals to form new, hybrid orbitals with different shapes and energies, suitable for bonding.

5-3. Sigma and pi bonds

This lesson explains the two main types of covalent bonds based on orbital overlap: sigma (σ) bonds (end-to-end overlap) and pi (π) bonds (side-to-side overlap).

6. Intermolecular Forces
2
1

This chapter covers intermolecular forces, the attractive forces that exist between molecules. These forces are responsible for the physical properties of substances, such as melting and boiling points. Key learning objectives include: defining the different types of intermolecular forces; and predicting the relative strength of these forces to explain the physical properties of a substance.

Chapter lessons

6-1. Dispersion and dipole forces

This lesson covers the two main types of van der Waals forces: London dispersion forces, which exist in all molecules, and dipole-dipole forces, which exist only in polar molecules.

6-2. Hydrogen bonding

This lesson defines hydrogen bonding as an especially strong type of dipole-dipole interaction that occurs when hydrogen is bonded to a highly electronegative atom (N, O, or F).

7. Conclusion
2

This concluding chapter summarises the key concepts of chemical bonding and molecular geometry. It reinforces the connection between electronic structure, molecular shape, and the physical properties of matter. This summary prepares the student for the next course, 'Kinetic Theory of Matter and Gas Laws', where the interactions between molecules are explored further.

Chapter lessons

7-1. Course summary

This lesson consolidates knowledge by reviewing the progression from simple bonding theories to the prediction of 3D molecular shapes and the analysis of intermolecular forces.

7-2. Next steps

This final lesson looks ahead, explaining how a command of molecular shape and intermolecular forces is a direct prerequisite for understanding the bulk properties of matter, which is the focus of the next course.