Atomic Structure and Periodicity of Elements - Chemistry (Undergraduate Foundation)
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[NUC Core] CHM 101: General Chemistry IThis 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. Introduction1
This chapter provides the roadmap for the course. It outlines the progression from early atomic theories to the modern understanding of atomic structure and its direct relationship to the periodic table. Key learning objectives include: understanding the overall course structure and appreciating the historical and experimental development of atomic theory.
Chapter lessons
2. Early Atomic Models6
This chapter traces the origin of atomic theory, starting with the empirical laws that led to Dalton's foundational model. It then covers the pivotal discovery of the electron, the first experimental proof that the atom is divisible, which forced the initial theory to evolve. Key topics include the laws of chemical combination, Dalton's atomic postulates and their critical shortcomings, and the experiments by Thomson and Millikan that discovered and characterised the electron, leading to the 'plum pudding' model.
Chapter lessons
2-1. Laws of chemical combination11:05
2-2. Dalton's atomic theory9:58
2-3. Shortcomings of Dalton's theory7:23
2-4. Thomson's cathode ray experiment10:00
2-5. Milikan's Oil drop experiment11:18
3. The Nuclear Model52
This chapter covers the development of the nuclear atom, from Rutherford's definitive gold foil experiment to Chadwick's discovery of the neutron. It establishes the modern planetary model of the atom and immediately addresses its critical failure, demonstrating why classical physics is insufficient. Key topics include the experimental evidence for the nucleus, the discovery of the neutron, the shortcomings of Rutherford's classical model, and the fundamental properties of protons, neutrons, and electrons.
Chapter lessons
3-1. Rutherford's gold foil experiment11:48
This lesson examines Rutherford's pivotal gold foil experiment. The surprising deflection of a few alpha particles at large angles invalidated the Thomson model. This observation was only explainable by concentrating the atom's mass and positive charge into a tiny nucleus.
3-2. Shortcomings of Rutherford's nuclear model6:40
Rutherford's model is fundamentally unstable according to classical physics. An orbiting electron must radiate energy, causing it to rapidly spiral into the nucleus. This lesson explains why this contradiction forced the development of a quantum model of the atom.
3-3. Chadwick's nuclear bombardment4:57
This lesson details Chadwick's bombardment experiment, which provided the evidence for the neutron. By observing the neutral radiation from alpha particle bombardment of beryllium, he identified a new nuclear particle. This completed the proton-neutron atomic model.
3-4. Sub-atomic particles9:42
This lesson defines the fundamental properties of the three subatomic particles. We will detail the relative mass and charge of the proton, neutron, and electron. A firm command of these values is required for all subsequent atomic calculations.
3-5. Isotopy6:06
This lesson defines isotopy, where atoms of the same element possess different numbers of neutrons. We will examine how this affects mass number whilst the atomic number remains constant, using key examples like the isotopes of hydrogen and carbon.
4. Bohr's Model31
This chapter details Niels Bohr's model, the first attempt to solve the failures of the classical nuclear atom using quantum theory. Understanding this transitional model is critical to grasping the full conceptual leap to modern quantum mechanics. Topics include: Bohr's postulates, explaining hydrogen's emission spectrum, energy level calculations, and the model's ultimate shortcomings.
Chapter lessons
4-1. Introduction8:16
This lesson reviews the critical failures of Rutherford's classical atomic model, specifically its inability to account for atomic stability and line spectra. We establish the context for Bohr's revolutionary quantum hypothesis, which directly addresses these shortcomings.
4-2. Energy levels8:07
This lesson explains how Bohr's model of quantized energy levels accounts for the hydrogen emission spectrum. We will demonstrate how electron transitions between these discrete energy levels produce the specific lines observed, and define the principal quantum number, n.
4-3. Shortcomings12:31
This lesson details the critical failures of the Bohr model, focusing on its inability to describe multi-electron atoms and its violation of the Heisenberg Uncertainty Principle. Understanding these shortcomings is essential to appreciate the necessity of the modern quantum mechanical model.
5. Wave-Particle Duality12
This chapter details wave-particle duality, the principle that all matter has wave properties. This concept explains the failure of Bohr's fixed orbits and establishes the necessary theoretical foundation for the modern quantum mechanical model of the atom. Key topics include: De Broglie’s wave hypothesis, applying the de Broglie relation (λ = h/mv), and calculating the wavelength of moving particles based on their mass and velocity.
Chapter lessons
5-1. De Broglie’s Theory4:04
This lesson explains de Broglie’s proposal that all matter exhibits wave properties. It introduces the relation λ = h/mv and shows how this idea connects particle motion to quantised electron states.
6. The Quantum Model61
This chapter introduces the modern quantum mechanical model, replacing Bohr's flawed orbits with probabilistic orbitals. Mastery of this model is non-negotiable, as it provides the definitive framework for describing electron behaviour and predicting all chemical properties. Key topics include: the four quantum numbers, defining orbital shapes (s, p, d, f) and their spatial orientation, and applying the Aufbau principle, Pauli exclusion, and Hund's rule to write correct electronic configurations.
Chapter lessons
6-1. Fundamental quantum numbers13:00
This lesson introduces Schrödinger’s equation as the foundation of atomic structure. It explains how its solutions give the principal and azimuthal quantum numbers.
6-2. Shapes of orbitals4:59
This lesson describes the spatial shapes of s, p, d, and f orbitals. It explains how quantum numbers determine these forms and their significance in atomic structure.
6-3. Magnetic quantum number8:16
This lesson defines the magnetic quantum number and its role in atomic structure. It explains how it determines the spatial orientation and distinction of orbitals within a subshell.
6-4. Orientation of orbitals8:50
This lesson explains how orbitals orient themselves in three-dimensional space. It relates each orientation to the magnetic quantum number and the distinct shapes of atomic orbitals.
6-5. Spin quantum number4:23
This lesson defines the spin quantum number and explains how electron spin distinguishes paired and unpaired electrons within an orbital.
6-6. Electronic configuration30:31
This lesson explains how electrons occupy orbitals using the Aufbau principle, Pauli exclusion, and Hund’s rule to determine the electronic configuration of any element.
7. Periodicity43
This chapter explains how the electronic structure of atoms leads to predictable, periodic trends in their properties across the periodic table. The focus is on justifying these trends based on concepts like shielding and effective nuclear charge. Key learning objectives include: defining and explaining the trends in atomic radii, ionization energies, and electronegativity across a period and down a group.
Chapter lessons
7-1. The periodic table24:45
This lesson reviews the structure of the modern periodic table, defining periods and groups. We establish the direct link between an element's position and its electronic configuration, which is the foundation for understanding periodicity.
7-2. Atomic and ionic radii14:11
This lesson defines atomic and ionic radii, explaining the trends across periods and down groups. We will justify these size changes using the concepts of effective nuclear charge and electron shielding.
7-3. First ionisation energy
This lesson explains the periodic trend of first ionisation energy across periods and down groups. We will justify these variations using effective nuclear charge, electron shielding, and orbital stability.
7-4. Electronegativity
This lesson defines electronegativity as the ability of an atom to attract electrons in a chemical bond and explains its periodic trends.
8. Conclusion2
This concluding chapter summarises the key concepts of atomic theory. It reinforces the understanding of atomic structure and its connection to the periodic properties of the elements. This summary prepares the student for the next course, 'Chemical Bonding and Molecular Geometry', where electronic structure is used to predict how atoms form molecules.
Chapter lessons
8-1. Course summary
This lesson consolidates knowledge by reviewing the progression from early atomic models to the modern understanding of electronic structure and periodic trends.
8-2. Next steps
This final lesson looks ahead, explaining how the principles of electronic configuration and periodicity are the direct prerequisites for understanding chemical bonding.