• Atomic Theory and Structure

  • Early Ideas about Matter
  • Did you know?

    Did you know that some ancient Greeks believed that all matter was made up of four substances: fire, air, water, and earth? They believed that rabbits were soft because they had more water than earth. Although this idea seems silly now, it contains a fundamental principle of atomic theory: that matter is made up of a small number of fundamental elements.

    • NGSS
    • HS-C5.1, HS-PS1.A3
    Summary

    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.

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  • Atomic Theory I
  • Did you know?

    Did you know that for 100 years scientists believed that atoms were the smallest particles that existed? It took many scientists and numerous experiments to show that atoms were made up of smaller particles with very different properties. Over a 75-year period beginning in the first part of the 19th century, two subatomic particles were discovered: the electron and the nucleus. In addition, by means of a clever experiment the negative charge of a single electron was calculated.

    Summary

    The 19th and early 20th centuries saw great advances in our understanding of the atom. This module takes readers through experiments with cathode ray tubes that led to the discovery of the first subatomic particle: the electron. The module then describes Thomson’s plum pudding model of the atom along with Rutherford’s gold foil experiment that resulted in the nuclear model of the atom. Also explained is Millikan’s oil drop experiment, which allowed him to determine an electron’s charge. Readers will see how the work of many scientists was critical in this period of rapid development in atomic theory.

    • NGSS
    • HS-C4.4, HS-C6.2, HS-PS1.A1, HS-PS1.A3
    Key Concepts
    • Atoms are not dense spheres but consist of smaller particles including the negatively charged electron.
    • The research on passing electrical currents through vacuum tubes by Faraday, Geissler, Crookes, and others laid the groundwork for discovery of the first subatomic particle.
    • J.J. Thomson’s observations of cathode rays provide the basis for the discovery of the electron.
    • Rutherford, Geiger, and Marsden performed a series of gold foil experiments that indicated that atoms have small, dense, positively-charged centers – later named the nucleus.
    • Millikan’s oil drop experiment determines the fundamental charge on the electron as 1.60 x 10-19 coulombs.

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  • Atomic Theory II
  • Did you know?

    Did you know that energy is not released in a continuous flow, but rather is released in “packets”? This discovery, known as quantum theory, changed the way we understand the basic properties of the atom. Many other advances in atomic theory were made in the 20th century, including the discovery of the neutron, which made the atom bomb possible.

    Summary

    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.

    • NGSS
    • HS-C4.4, HS-C6.2, HS-PS1.A1, HS-PS1.A3
    Key Concepts
    • 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.

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  • Atomic Theory III
  • Did you know?

    Did you know that atoms could not be described accurately until quantum theory as developed? Quantum theory offered a fresh way of thinking about the universe at the atomic level. After tremendous advances in quantum mechanics in the last century, the position of electrons and other infinitesimal particles can be predicted with confidence.

    Summary

    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.

    • NGSS
    • HS-C1.4, HS-C4.4, HS-PS1.A2, HS-PS2.B3
    Key Concepts
    • 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.

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  • Atomic Theory IV
  • Did you know?

    Did you know that electrons are so tiny that when you shine light on them, the light itself changes the electron’s path? Because of this, we can’t know exactly where an electron is within an atom. Rather, it necessary to describe the position of an electron in terms of probability. Thus, scientists use a mathematical equation to describe how electrons are most likely distributed around the atom's nucleus.

    Summary

    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.

    • NGSS
    • HS-C1.4, HS-C4.4, HS-PS1.A2, HS-PS2.B3
    Key Concepts
    • 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.

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  • The Periodic Table of Elements
  • Did you know?

    Did you know that although electrons are minuscule compared to other parts of an atom, the way they are arranged around the nucleus is the biggest factor in determining the chemical properties of an element? The periodic chart is ordered by atomic number, but drastic shifts in chemical properties can occur from one element to the next. These shifts are explained by how the elements are displayed on the periodic table.

    • NGSS
    • HS-C1.1, HS-PS1.A2
    Summary

    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.

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  • The Mole
  • Did you know?

    Did you know that just as a “pair” equals two of something and a “dozen” equals 12 of something, a “mole” equals 602,000,000,000,000,000,000,000 of something? This huge number, written as 6.02 x 1023, comes in very handy for describing amounts of extraordinarily small things like atoms and molecules.

    Summary

    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.

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  • Physical States and Properties

  • States of Matter
  • Did you know?

    Did you know that solids, liquids, and gases are not the only states of matter? Among others are plasmas, which have such high energy that molecules are ripped apart. And Bose-Einstein Condensates, seen for the first time in 1995, are a weird state of matter that can actually trap light.

    • NGSS
    • HS-C5.2, HS-PS1.A3, HS-PS1.A4, HS-PS2.B3
    Summary

    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.

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  • Properties of Liquids
  • Did you know?

    Did you know that various liquids behave differently because of how the tiny molecules of which they are composed interact with each other? This is why gasoline flows more quickly than syrup and why certain insects can walk across the surface of water without falling in. In fact, pitch, a liquid that comes from plants and petroleum, flows so slowly that when placed in a funnel, an entire decade can pass between each drop!

    Summary

    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.

    • NGSS
    • HS-C6.2, HS-PS1.A3, HS-PS1.A4
    Key Concepts
    • 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.

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  • Properties of Solids
  • Did you know?

    Did you know that the melting point of solids can be as low as -38°C (or -36°F) for mercury and as high as 4,489°C (or 8,112°F) for graphite? This is because differences in the composition, bonding, and structure of various solids determine how they behave. The way that different solids are formed also determines which ones conduct heat and electricity and which dissolve easily when stirred into a beverage.

    Summary

    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.

    • NGSS
    • HS-C6.2, HS-PS1.A3
    Key Concepts
    • 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.

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  • Diffusion I
  • Did you know?

    Did you know that the process of diffusion is responsible for the way smells travel from the kitchen throughout the house? In diffusion, particles move randomly, beginning in an area of higher concentration and ending in an area of lower concentration. This principle is fundamental throughout science and is very important to how the human body and other living things function.

    Summary

    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.

    • NGSS
    • HS-C5.4, HS-PS3.A3, HS-PS3.B5
    Key Concepts
    • 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.

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  • Water
  • Did you know?

    Did you know that the way water molecules interact with each other caused the sinking of the Titanic? And the unique chemical properties of this miracle liquid that cause icebergs to float also cause the surface of water to stick together with enough force that the Basilisk or “Jesus” lizard can walk on water.

    • NGSS
    • HS-C6.2, HS-PS1.A3
    Summary

    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.

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  • Chemical Relationships

  • Chemical Bonding
  • Did you know?

    Did you know that the 118 elements on the periodic table combine to make millions and millions of chemical compounds? This is because chemical bonds between atoms result in new substances that are very different from the elements they are made of. For example, chlorine can be used as a chemical weapon and yet it combines with sodium, an explosive element, to make common table salt.

    Summary

    The millions of different chemical compounds that make up everything on Earth are composed of 118 elements that bond together in different ways. This module explores two common types of chemical bonds: covalent and ionic. The module presents a sliding scale of chemical bonding from pure covalent to pure ionic, depending on differences in the electronegativity of the bonding atoms. Highlights from three centuries of scientific inquiry into chemical bonding include Isaac Newton’s ‘forces’, Gilbert Lewis’s dot structures, and Linus Pauling’s application of the principles of quantum mechanics.

    • NGSS
    • HS-C4.3, HS-C6.2, HS-PS1.A3, HS-PS1.B1
    Key Concepts
    • When a force holds atoms together long enough to create a stable, independent entity, that force can be described as a chemical bond.
    • The 118 known chemical elements interact with one another via chemical bonds, to create brand new, unique compounds that have entirely different chemical and physical properties than the elements that make them up.
    • It is helpful to think of chemical bonding as being on a sliding scale, where at one extreme there is pure covalent bonding, and at the other there is pure ionic bonding. Most chemical bonds lie somewhere between those two extremes.
    • When a chemical bond is formed between two elements, the differences in the electronegativity of the atoms determine where on the sliding scale the bond falls. Large differences in electronegativity favor ionic bonds, no difference creates non-polar covalent bonds, and relatively small differences cause the formation of polar-covalent bonds.

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  • Chemical Equations
  • Did you know?

    Did you know that when methane and oxygen are combined, they produce carbon dioxide and water vapor? This is because the chemical bonds that hold molecules together are broken during chemical reactions, and when molecules re-form they can produce very different substances.

    • NGSS
    • HS-C3.5, HS-C5.1, HS-PS1.B1, HS-PS1.B3
    Summary

    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.

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  • Acids and Bases
  • Did you know?

    Did you know that some juices and vinegar taste sour because of the chemical properties of the acid in those liquids? And that when acids and bases are mixed together, they always counteract each other, producing water and a salt?

    • NGSS
    • HS-C5.2, HS-PS1.A3, HS-PS1.B1
    Summary

    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.

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  • Reactions and Changes

  • Chemical Reactions
  • Did you know?

    Did you know that chemical reactions happen all around us, such as when you light a match, start a car, or even take in a breath of air? To accomplish stability among atoms, sometimes electrons are taken from or shared with another atom. This need to balance electrons with protons causes chemical reactions and can create compounds.

    • NGSS
    • HS-C5.4, HS-PS1.A2, HS-PS1.A3, HS-PS1.B3
    Summary

    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.

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  • Nuclear Chemistry
  • Did you know?

    Did you know that the sun and stars are actually enormous thermonuclear fusion reactors? And that atoms can be split artificially, releasing energy that can be harnessed to generate electrical power? Thanks to pioneers in nuclear chemistry like Marie Curie, we have come to understand different types of radiation and nuclear reactions.

    • NGSS
    • HS-C5.5, HS-PS1.C1, HS-PS3.A1
    Summary

    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.

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  • Carbon Chemistry
  • Did you know?

    Did you know that organic chemicals make up all the life forms we know of? Organic chemistry, defined by the carbon-hydrogen bond, is at the foundation of life. Because of the unique properties of the carbon atom, it can bond with other atoms in many different ways, resulting in millions of different organic molecules.

    • NGSS
    • HS-C6.2, HS-PS1.A3
    Summary

    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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