Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → General chemistry → Atomic Structure ↓


Periodic trends


There are patterns observed in the properties of elements in different periods and groups of the periodic table. These patterns arise from regular and predictable changes in the electronic structure of atoms. Understanding these trends helps us predict the behavior of elements, and is a fundamental component of general chemistry. Let's embark on a subtle journey through the primary periodic trends that include atomic radius, ionization energy, electron affinity, and electronegativities.

1. Atomic radius

The atomic radius is defined as the distance from the nucleus to the outermost electron of the atom. As we move across the periodic table, especially from left to right in a period, the atomic radius decreases. Conversely, as we move down a group, the atomic radius increases. This pattern can be effectively visualized with a simple example:

Explanation:

  1. Across a period: from left to right, electrons pair up in the same shell or energy level, but the number of protons in the nucleus also increases, which brings the electrons closer together due to the increased positive charge. This results in a smaller atomic radius.
  2. Going down the group, electrons get added to a new outer shell, which is farther from the nucleus than in the previous period, resulting in larger atomic radius.

For example, trends in atomic radius can be represented by these changes: 

Li (Lithium) > Be (Beryllium) > B (Boron) in decreasing order across a period. Li (Lithium) < Na (Sodium) < K (Potassium) in increasing order down a group.

2. Ionization energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. Generally, ionization energy increases across a period and decreases down a group. This concept is important because it relates to the reactivity and bonding of an element. Consider the following example:

Explanation:

  1. Across a period: Ionization energy increases because the electrons are more strongly attracted to the increasingly positive nucleus, making them harder to remove.
  2. Going down the group: electrons are located farther from the nucleus and experience more shielding, making them easier to remove, thereby lowering the ionization energy.

For example, consider these trends: 

He (Helium) > Ne (Neon) > Ar (Argon) in increasing ionization energy across periods. Li (Lithium) < Na (Sodium) < K (Potassium) in decreasing ionization energy down a group.

3. Electron affinity

Electron affinity is the energy change that occurs when an electron is added to a neutral atom. Elements with higher electron affinities gain electrons more easily. This energy change provides insight into the formation of anions and is important in the context of ionic bonding.

Like ionization energy, electron affinity also generally becomes more negative across a period (reflecting higher affinity for electrons) and less negative going down a group.

For example: 

F (Fluorine) has a more negative electron affinity than O (Oxygen), while Cl (Chlorine) > F (Fluorine) in electron affinity but Cl is below F in group.

This trend is not as simple as the others, because there are larger numbers of electrons and electron-electron repulsion in specific subshell structures, but the general pattern remains.

4. Electronegativity

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. It plays an important role in determining the type of bond formed between atoms. Electronegativity increases across a period and decreases down a group.

Notable examples include: 

N (Nitrogen), O (Oxygen), and F (Fluorine) are highly electronegative elements. Electronegativity values: Li (0.98) < Be (1.57) < B (2.04) < C (2.55) < N (3.04) < O (3.44) < F (3.98).

Key concept summary

  • Atomic radius: Decreases across a period and increases down a group.
  • Ionization energy: Increases across a period and decreases down a group.
  • Electron affinity: Generally becomes more negative across a period and less negative down a group.
  • Electronegativity: Increases across a period and decreases down a group.

Understanding these trends is important for recognizing how the elements interact and bond with one another, which is central to the study of chemistry. These trends are not absolute, but exceptions usually occur due to unique electronic configurations or subtle inter-electronic interactions. By looking at these general patterns and exceptions, we form a broader understanding of the behavior and interactions of the elements.

Such essential knowledge not only provides insight into the fundamental nature of chemical reactions, but also supports advanced applications in materials science, biology, and physics. Trends in periodic properties reflect deeper principles of quantum mechanics and atomic theory, which are covered in detail in more advanced chemistry courses.


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Periodic trends

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