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

The shape of a molecule dictates its function. This course provides a practical, foundational treatment of chemical bonding and molecular geometry - the principles governing how atoms connect to form substances. It covers ionic and covalent bonding, the drawing of Lewis structures, and methods for determining molecular shape using VSEPR theory and orbital hybridisation. Finally, we examine intermolecular forces that control a substance's physical properties. A command of this material is essential for molecular design. Knowledge of bonding and geometry dictates the properties of everything from pharmaceuticals, where shape governs drug-receptor specificity, to new materials like polymers and catalysts. Mastery of this foundational knowledge is the starting point for developing or synthesising any new chemical compound. Our system uses comprehensive videos and allows you to ask questions directly, making it fully online and easy to master at your pace. By the end of this course, you will be able to draw correct Lewis structures for any molecule, predict its three-dimensional geometry and polarity using VSEPR theory, determine the central atom's hybridisation and bond types, and identify the various intermolecular forces present in a substance. This course is mandatory for all undergraduate students of chemistry, biochemistry, and materials science, and is a direct prerequisite for studying organic chemistry. It benefits anyone needing a rigorous, structured, student-centred refresher or initial exposure to this critical subject, assuming a complete understanding of atomic theory and electronic configuration.

4

9 hrs

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.
CHM 101: General Chemistry I
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.

See more

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

This chapter focuses on mastering Valence Shell Electron Pair Repulsion (VSEPR) theory to accurately predict the three-dimensional geometry of molecules. Understanding molecular shape is non-negotiable; it governs all chemical properties, reactivity, and function, providing the framework for molecular design. Upon completion, you will be able to: predict electron and molecular geometry for molecules with two to six electron domains, identify geometries in examples with multiple lone pairs, and master the step-by-step procedure for determining complex molecular shapes.

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. Examples (2)
32:17

This second lesson provides advanced VSEPR examples involving six or more electron domains. We determine the geometry of Xenon difluoride (XeF2), Sulphur hexafluoride (SF6), Bromine pentafluoride (BrF5), and Xenon tetrafluoride (XeF4). Master the prediction of complex geometries with multiple lone pairs.

5. Hybridization Theory
4

This chapter introduces Valence Bond Theory and the crucial concept of orbital hybridization. Hybridization is a necessary model for explaining molecular shapes and bond angles that cannot be accounted for by simple VSEPR theory alone, bridging the gap between bonding theory and observed molecular geometry. Upon completion, you will be able to: define bonding as the overlap of atomic orbitals, explain the formation of sigma and pi bonds, and accurately determine the hybridization (from sp to sp3d2) for the central atom in any molecule.

Chapter lessons

5-1. Valence bond theory
17:18

This lesson introduces Valence Bond Theory, defining a chemical bond as the overlap of atomic orbitals. We explain the formation of sigma and pi bonds. This foundational theory is essential for understanding hybridization.

5-2. Hybridization of orbitals
19:08

This lesson explains orbital hybridization - the mixing of atomic orbitals to form new, equivalent hybrid orbitals. Learn why this model is necessary to explain observed molecular geometries and bond angles.

5-3. Examples (1)
21:49

This lesson provides initial examples of determining the hybridization (sp, sp2, sp3) of the central atom. We analyze the bonding in Methane (CH4), Ammonia (NH3), Ethene (C2H4), and Beryllium chloride (BeCl2). Master the method for identifying the correct hybridization from the Lewis structure.

5-4. Examples (2)
23:43

This second lesson reinforces hybridization theory - focusing on molecules with expanded octets. We determine the hybridization of Boron trichloride (BCl3), Phosphorus pentachloride (PCl5), and Sulphur hexafluoride (SF6). Master identifying hybrid orbitals including those that use d-orbitals (sp3d and sp3d2).

6. Intermolecular Forces
3

This chapter details intermolecular forces (IMFs) - the attractive forces between molecules. These non-covalent interactions dictate all physical properties, including melting points, boiling points, and solubility, which is critical for predicting substance behaviour. Upon completion, you will be able to: define and identify London dispersion and dipole-dipole forces, explain the unique strength of hydrogen bonding, and use dipole moment and molecular symmetry to determine a molecule's overall polarity.

Chapter lessons

6-1. Dispersion and dipole forces
10:54

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
21:09

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).

6-3. Dipole moment
20:07

This lesson defines the dipole moment as the vector sum of all bond dipoles in a molecule. We use 1,4-dichlorobenzene (C6H4Cl2) and 1,2-dichlorobenzene (C6H4Cl2) to show how molecular geometry determines if the net dipole is zero or non-zero. Master how symmetry impacts overall molecular polarity.

7. Conclusion
1

This final chapter is dedicated to comprehensive application and synthesis of all course material. It reinforces the critical connections between Lewis structures, molecular geometry, and physical properties, ensuring foundational knowledge is secure for advanced topics. Upon completion, you will be able to demonstrate mastery across: drawing Lewis structures, predicting VSEPR geometry, determining orbital hybridization and bond types, and identifying intermolecular forces in any given molecule.

Chapter lessons

7-1. Practice questions
3:48

This lesson features comprehensive practice questions covering all course topics: Lewis structures, VSEPR geometry, hybridisation, and intermolecular forces. Use this session to test your mastery of the connection between electronic structure and molecular properties.