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


General chemistry is a fundamental science course taught primarily to undergraduate students. It serves as an introductory course that lays the foundation for understanding the principles of chemistry that apply in real-world contexts. The primary purpose is to create a framework of knowledge that can be expanded upon in more advanced courses. In this lesson, we will explore various concepts and principles central to general chemistry.

1. Introduction to matter and its states

Matter is defined as anything that occupies space and has mass. It exists in different states, mainly solid, liquid, and gas, depending on the arrangement and energy of its molecules.

  • Solids: The particles in solids are packed together in a fixed, rigid structure. This gives solids a definite shape and volume.
  • Liquids: Particles in liquids are close together but are not rigid like solids. Liquids have a fixed volume but they take the shape of their container.
  • Gas: The particles in gases are far away from each other and move around freely, so gases have neither a definite shape nor volume.

2. Atoms and elements

Everything in the universe is made up of atoms, which are the smallest units of elements. An element is a pure substance that contains only one type of atom, represented by a unique chemical symbol, such as H for hydrogen or O for oxygen.

The atomic structure includes the following:

  • Proton: Positively charged particle in the nucleus.
  • Neutrons: Neutral charged particles are also located in the nucleus.
  • Electrons: Negatively charged particles that orbit around the nucleus.

3. Periodic table of elements

The periodic table is a systematic chart of all known elements arranged according to their atomic numbers. This table helps in predicting the chemical properties and behaviours of the elements.

        H2O
     by Lee B. BCNOF
     NaMgAlSiPsClAr

This arrangement allows quick identification of the properties of the elements, with each row called a period and each column called a group.

4. Chemical bond

Chemical bonds are the forces that hold atoms together in compounds. The main types of chemical bonds include:

  • Ionic bond: Formed through the transfer of electrons from one atom to another, usually between a metal and a nonmetal.
  • Covalent bonds: bonds where atoms share electrons, usually between non-metals.
  • Metallic bond: attraction between free electrons and metal ions, common in metallic elements.

For example, in water (H2O), two hydrogen atoms form a covalent bond by sharing electrons with an oxygen atom.

5. Chemical reactions

In chemical reactions atoms are rearranged to form new substances. Reactions are represented by chemical equations, for example:

    2H2 + O2 → 2H2O

Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement reactions.

6. Mole concept

The mole is a standard unit of measurement in chemistry that represents a large number of molecules or atoms. Avogadro's number, 6.022 × 1023, is the number of units in a mole.

It is useful in stoichiometry for balancing chemical equations and calculating reactants/products in reactions.

7. Solutions and mixtures

A solution is a homogeneous mixture of two or more substances. The solvent (the dissolving medium) and the solute (the dissolved substance) are the components of a solution.

For example, a sugar solution is formed when sugar (the solute) is dissolved in water (the solvent).

Mixtures can be classified as homogeneous (same composition) or heterogeneous (different phases).

8. Acids and bases

Acids and bases are substances that can donate or accept protons (H+), respectively.

  • Acids: Substances that increase the concentration of hydrogen ions in a solution. Example: HCl (hydrochloric acid).
  • Bases: Substances that increase the concentration of hydroxide ions (OH-) in a solution. Example: NaOH (sodium hydroxide).

The pH scale from 0 to 14 is used to measure the acidity or alkalinity of a solution.

9. Thermodynamics

Thermodynamics involves the study of energy transformations, especially the transfer of heat and work in chemical processes.

Some of the key concepts are as follows:

  • First Law of Thermodynamics: Energy cannot be created or destroyed, it can only be transferred.
  • Enthalpy (H): represents the heat content of a system at constant pressure.

For example, during vaporization, a liquid absorbs heat to become a gas, causing a change in enthalpy.

10. Atomic structure and periodicity

The Bohr model of the atom helps explain atomic structure, where electrons orbit around the nucleus in fixed paths. Quantum mechanics later introduced the concept of electron clouds.

Periodicity refers to recurring trends in element properties within the periodic table, such as electronegativities, ionization energies, and atomic radii.

The properties of elements change in a predictable manner across different periods and groups.

11. Dynamics and balance

Kinetics focuses on the rate of chemical reactions. Factors that affect the rate include concentration, temperature, and the presence of a catalyst.

Equilibrium occurs when the rate of the forward reaction is equal to the rate of the reverse reaction. For a balanced chemical equation:

    a + b ⇌ c + d

It refers to a dynamic state where reactants and products exist together.

12. Redox reactions

Redox reactions involve the transfer of electrons between substances, which is reflected in oxidation and reduction processes. The main ones of importance are:

  • Oxidation: The loss of electrons by a molecule, atom, or ion.
  • Reduction: Gain of electrons by a molecule, atom, or ion.

A classic example of this is the rusting of iron:

    4Fe + 3O2 → 2Fe2O3

13. Basic organic chemistry concepts

Organic chemistry is the study of compounds containing carbon. Carbon can form a wide variety of compounds with chains and rings, and the simple ways these are structured give rise to a wide variety of organic chemicals.

Examples include hydrocarbons such as methane (CH4) and more complex molecules such as glucose (C6H12O6).

14. Environment and practical relations

Chemistry plays a vital role in solving environmental issues such as pollution, climate change, and sustainable resources.

Chemical principles are used to develop environmentally friendly materials, understand ozone layer depletion, and create energy-efficient processes.

General Chemistry encourages students to consider real-life applications and global implications of chemical phenomena, which fosters further scientific inquiry and innovation.

Conclusion

General chemistry serves as an essential framework for understanding the complex world of chemical substances and reactions. It forms the foundation upon which students can develop more complex knowledge in specialized areas of chemistry. The concepts learned help to appreciate the ubiquity of chemistry in everyday life and the influence of the molecular world on broader scientific discussions.


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General chemistry

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