The universe is a vast soup of interacting particles and energy. The ways in which those interactions take place, as well as the structure and composition of matter, is the main focus of the field of chemistry. Our chemistry learning modules introduce you to the world of chemistry, exploring current research and scientific findings on concepts like the structure and function of atoms, forms of energy and its transfer, chemical bonding and reactions, and more.


Air contains matter, even if you cannot see it.

Tracking the development of our understanding of the atomic structure of matter, this module begins with the contributions of ancient Greeks, who proposed that matter is made up of small particles. The module then describes how Lavoisier's Law of Conservation of Mass and Proust's Law of Definite Proportions contributed to Dalton's modern atomic theory.

Modern atomic theory has evolved dramatically from the 19th century view of the atom as a small, solid sphere resembling a billiard ball. This module explores that story: from the discovery of electrons and protons in the late 19th century to the planetary model of the atom in the early 20th century. The module explains the function of subatomic particles as well as their relative size and weight. The concepts of atomic number and atomic mass are introduced.

The 20th century brought a major shift in our understanding of the atom, from the planetary model that Ernest Rutherford proposed to Niels Bohr’s application of quantum theory and waves to the behavior of electrons. With a focus on Bohr’s work, the developments explored in this module were based on the advancements of many scientists over time and laid the groundwork for future scientists to build upon further. The module also describes James Chadwick’s discovery of the neutron. Among other topics are anions, cations, and isotopes.

  • Drawing on experimental and theoretical evidence, Niels Bohr changed the paradigm of modern atomic theory from one that was based on physical particles and classical physics, to one based in quantum principles.
  • Under Bohr’s model of the atom, electrons cannot rotate freely around the atom, but are bound to certain atomic orbitals that both constrain and define an atom's electronic behavior.
  • Atoms can gain or lose electrons to become electrically charged ions.
  • James Chadwick completed the early picture of the atom with his discovery of the neutron, a neutral, nuclear particle that affects an atom’s mass and the different physical properties of atomic isotopes.

The 20th century was a period rich in advancing our knowledge of quantum mechanics, shaping modern physics. Tracing developments during this time, this module covers ideas and refinements that built on Bohr’s groundbreaking work in quantum theory. Contributions by many scientists highlight how theoretical insights and experimental results revolutionized our understanding of the atom. Concepts include the Schrödinger equation, Born’s three-dimensional probability maps, the Heisenberg uncertainty principle, and electron spin.

  • Electrons, like light, have been shown to be wave-particles, exhibiting the behavior of both waves and particles.
  • The Schrödinger equation describes how the wave function of a wave-particle changes with time in a similar fashion to the way Newton’s second law describes the motion of a classic particle. Using quantum numbers, one can write the wave function, and find a solution to the equation that helps to define the most likely position of an electron within an atom.
  • Max Born’s interpretation of the Schrödinger equation allows for the construction of three-dimensional probability maps of where electrons may be found around an atom. These ‘maps’ have come to be known as the s, p, d, and f orbitals.
  • The Heisenberg Uncertainty Principle establishes that an electron’s position and momentum cannot be precisely known together, instead we can only calculate statistical likelihood of an electron’s location.
  • The discovery of electron spin defines a fourth quantum number independent of the electron orbital but unique to an electron. The Pauli exclusion principle states that no two electrons with the same spin can occupy the same orbital.

Our Atomic Theory series continues, exploring the quantum model of the atom in greater detail. This module takes a closer look at the Schrödinger equation that defines the energies and probable positions of electrons within atoms. Using the hydrogen atom as an example, the module explains how orbitals can be described by type of wave function. Evidence for orbitals and the quantum model is provided by the absorption and emission spectra of hydrogen. Other concepts include multi-electron atoms, the Aufbau Principle, and Hund’s Rule.

  • The wave-particle nature of electrons means that their position and momentum cannot be described in simple physical terms but must be described by wave functions.
  • The Schrödinger equation describes how the wave function of a wave-particle changes with time in a similar fashion to the way Newton’s second law describes the motion of a classical particle. The equation allows the calculation of each of the three quantum numbers related to individual atomic orbitals (principal, azimuthal, and magnetic).
  • The Heisenberg uncertainty principle establishes that an electron’s position and momentum cannot be precisely known together; instead we can only calculate statistical likelihood of an electron’s location.
  • The discovery of electron spin defines a fourth quantum number independent of the electron orbital but unique to an electron. The Pauli exclusion principle states that no two electrons with the same spin can occupy the same orbital.
  • Quantum numbers, when taken as a set of four (principal, azimuthal, magnetic and spin) describe acceptable solutions to the Schrödinger equation, and as such, describe the most probable positions of electrons within atoms.
  • Orbitals can be thought of as the three dimensional areas of space, defined by the quantum numbers, that describe the most probable position and energy of an electron within an atom.

The modern periodic table is based on Dmitri Mendeleev’s 1896 observations that chemical elements can be grouped according to chemical properties they exhibit. This module explains the arrangement of elements in the period table. It defines periods and groups and describes how various electron configurations affect the properties of the atom.

This module introduces the mole, a unit of measurement for quantifying atoms and molecules. The module describes 19th century developments that led to the concept of the mole, also known as Avogadro’s number, or 6.02 x 1023. Molar mass is explained, and examples of mole/weight relationships are presented. In addition, the module shows how to calculate molecular weight.

There are many states of matter beyond solids, liquids, and gases, including plasmas, condensates, superfluids, supersolids, and strange matter. This module introduces Kinetic Molecular Theory, which explains how the energy of atoms and molecules results in different states of matter. The module also explains the process of phase transitions in matter.

When it comes to different liquids, some mix well while others don’t; some pour quickly while others flow slowly. This module provides a foundation for considering states of matter in all their complexity. It explains the basic properties of liquids, and explores how intermolecular forces determine their behavior. The concepts of cohesion, adhesion, and viscosity are defined. The module also examines how temperature and molecule size and type affect the properties of liquids.

  • Liquids share some properties with solids—both are considered condensed matter and are relatively incompressible—and some with gases, such as their ability to flow and take the shape of their container.
  • A number of properties of liquids, such as cohesion and adhesion, are influenced by the intermolecular forces within the liquid itself.
  • Viscosity is influenced by both the intermolecular forces and molecular size of a compound.
  • Most liquids we encounter in everyday life are in fact solutions, mixtures of a solid, liquid or gas solute within a liquid solvent.

Solids are formed when the forces holding atoms or molecules together are stronger than the energy moving them apart. This module shows how the structure and composition of various solids determine their properties, including conductivity, solubility, density, and melting point. The module distinguishes the two main categories of solids: crystalline and amorphous. It then describes the four types of crystalline solids: molecular, network, ionic, and metallic. A look at different solids makes clear how atomic and molecular structure drives function.

  • A solid is a collection of atoms or molecules that are held together so that, under constant conditions, they maintain a defined shape and size.
  • There are two main categories of solids: crystalline and amorphous. Crystalline solids are well ordered at the atomic level, and amorphous solids are disordered.
  • There are four different types of crystalline solids: molecular solids, network solids, ionic solids, and metallic solids. A solid's atomic-level structure and composition determine many of its macroscopic properties, including, for example, electrical and heat conductivity, density, and solubility.

The process of diffusion is critical to life, as it is necessary when our lungs exchange gas during breathing and when our cells take in nutrients. This module explains diffusion and describes factors that influence the process. The module looks at historical developments in our understanding of diffusion, from observations of “dancing” particles in the first century BCE to the discovery of Brownian motion to more recent experiments. Topics include concentration gradients, the diffusion coefficient, and advection.

  • Diffusion is the process by which molecules move through a substance, seemingly down a concentration gradient, because of the random molecular motion and collision between particles.
  • Many factors influence the rate at which diffusion takes place, including the medium through with a substance is diffusing, the size of molecules diffusing, the temperature of the materials, and the distance molecules travel between collisions.
  • The diffusion coefficient, or diffusivity, provides a relative measure at specific conditions of the speed at which two substances will diffuse into one another.

Water, critical to our survival, behaves differently from any other substance on Earth. The unique chemical properties of water are presented in this module. The module explains how the dipole across the water molecule leads to hydrogen bonding, making water molecules act like little magnets. Also explored are surface tension and water’s properties as a solvent.

Chemical bonding between atoms results in compounds that can be very different from the parent atoms. This module, the second in a series on chemical reactions, describes how atoms gain, lose, or share electrons to form ionic or covalent bonds. The module lists features of ionic and covalent compounds. Lewis dot structures and dipoles are introduced.

When scientists want to describe a chemical reaction in writing, they use precise chemical equations. This module, the third in a series on chemical reactions, explains chemical equations in a step-by-step process. The module shows how chemical equations are balanced in the context of the chemical changes that take place during a reaction. The Law of Conservation of Matter is introduced.

Since acids and bases were first labeled and described in the 17th century, their definition has been refined over the centuries to reflect an increased understanding of their chemical properties. This module introduces the fundamentals of acid/base chemistry, including neutralization reactions. The relationship between hydrogen ion concentration [H+] and pH is shown alongside everyday examples of acids and bases.

Exploring the how and why of chemical reactions, this module describes properties of the inert – or "noble" – gases. It also explains how different elements achieve a stable configuration by bonding to fill their valence shells with electrons, resulting in a chemical compound. The difference between exothermic and endothermic reactions is discussed.

Beginning with the work of Marie Curie and others, this module traces the development of nuclear chemistry. It describes different types of radiation: alpha, beta, and gamma. The module then applies the principle of half-life to radioactive decay and explains the difference between nuclear fission and nuclear fusion.

The chemical basis of all living organisms is linked to the way that carbon bonds with other atoms. This introduction to organic chemistry explains the many ways that carbon and hydrogen form bonds. Basic hydrocarbon nomenclature is described, including alkanes, alkenes, alkynes, and isomers. Functional groups of atoms within organic molecules are discussed.

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