Chemistry

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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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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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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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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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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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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Did you know?

Did you know that just as one dozen equals 12 of something, one mole equals 602,000,000,000,000,000,000,000 of something? This huge number, written as 6.02 x 10²³, is used by scientists to describe the amount of extraordinarily small things like atoms and molecules. The mole is a standard unit of measure in the metric system, and it is useful for converting the number of particles in a substance into its mass, and vice versa.

Summary

The mole is an important concept for talking about a very large number of things — 6.02 x 10²³ of them to be exact. This module shows how the mole, known as Avogadro’s number, is key to calculating quantities of atoms and molecules. It describes 19th-century developments that led to the concept of the mole, Topics include atomic weight, molecular weight, and molar mass. Sample equations illustrate how molar mass and Avogadro’s number act as conversion factors to determine the amount of a substance and its mass.

Key Concepts

• The mole is a term for a very large number, 6.02 x 10²³, known as Avogadro’s number.
• Avogadro’s number is the experimentally determined number of carbon-12 atoms in 12 grams of carbon-12.
• The numerical value for the mass of one mole (molar mass) of an element’s atoms is the same as the value for the mass of an individual atom of that element (atomic mass); however, the units on these mass numbers are different.
• An element’s molar mass can be used to calculate the number of atoms in a sample with a known mass, and to calculate the mass of a sample with a known number of atoms.

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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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Did you know?

Did you know that you encounter and interact with a variety of substances on a daily basis? Substances are everywhere and can vary significantly, from the water you drink to your favorite breakfast meal. Given that there are so many, how do scientists keep all of these substances straight? Let’s find out.

Summary

This module explores substances through hypothetical and real-world examples. Substances are broadly classified as pure substances, such as elements and compounds, or mixtures, such as rainwater, but the classification system goes further. How a substance is classified depends on its makeup and properties, and understanding the differences helps scientists solve major issues, such as creating clean drinking water.

• NGSS

• MS-PS1.A1, MS-PS1.A2, MS-PS1.A3, MS-PS1.A4, MS-PS1.A5, MS-PS1.A6

Key Concepts

• Substances can be classified as pure substances or mixtures. This classification helps scientists understand what particular substances are made up of and their properties.
• Experiments over many years have helped scientists recognize that pure substances include elements, which cannot be broken down, and compounds, which can be broken down chemically into the elements that make them up. Compounds have a definite composition and are chemically formed.
• Mixtures are physical combinations of pure substances that can be homogeneous or heterogeneous. Homogeneous mixtures are also called solutions and look the same throughout. Heterogeneous mixtures have clearly distinguished parts.
• Mixtures do not have a definite composition and may be separated physically, such as the distillation of water to separate salts and impurities from pure water.
• Elements are made up of atoms and compounds are made up of molecules. Mixtures may be any combination of atoms, molecules, or a combination of the two.

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Did you know?

Did you know that if you took a helium balloon to the top of Mount Everest, it would get bigger and might even pop? Conversely, if you took a helium balloon deep enough under the ocean, it would shrivel up. This is because of the basic properties of gases, which in addition to explaining the behavior of balloons are key to critical functions like breathing and lifesaving technology like automobile airbags.

Summary

This module describes the properties of gases and explores how these properties relate to a common set of behaviors called the gas laws. With a focus on Boyle’s Law, Charles’s Law, and Avogadro’s Law, an overview of 400 years of research shows the development of our understanding of gas behavior. The module presents the ideal gas equation and explains when this equation can—and cannot—be used to predict the behavior of real gases.

Key Concepts

• Unlike solids or liquids, the molecules in a gas are very far apart and rarely interact with each other, which is why gases made out of different molecules share similar behaviors.
• The gas laws describe the relationships between a gas's temperature, pressure, volume, and amount. These laws were identified in experiments performed by multiple scientists over four centuries.
• Because gases share common behaviors, the behavior of a real gas at a given pressure (P), absolute temperature (T), volume (V), and amount (n, in moles) can often be predicted by the ideal gas equation, PV = nRT, which perfectly describes the behavior of an idealized gas.
• The behavior of real gases deviates from ideal gases at very low temperatures and high pressures.

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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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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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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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Did you know that during the 18th century scientists theorized that particles were engulfed in a heat substance called “caloric” which imparted temperature to matter and caused gas molecules to be repelled from one another? This idea was rejected by the scientist Rudolph Clausius who proposed that heat is a form of energy that affects the temperature of matter by changing the motion of molecules in matter. This kinetic theory of heat enabled Clausius to study and predict the flow of heat—a field we now call thermodynamics and key to the development of kinetic-molecular theory.

Summary

Over four hundred years, scientists including Rudolf Clausius and James Clerk Maxwell developed the kinetic-molecular theory (KMT) of gases, which describes how molecule properties relate to the macroscopic behaviors of an ideal gas—a theoretical gas that always obeys the ideal gas equation. KMT provides assumptions about molecule behavior that can be used both as the basis for other theories about molecules and to solve real-world problems.

Key Concepts

• Kinetic-molecular theory states that molecules have an energy of motion (kinetic energy) that depends on temperature.
• Rudolf Clausius developed the kinetic theory of heat, which relates energy in the form of heat to the kinetic energy of molecules.
• Over four hundred years, scientists have developed the kinetic-molecular theory of gases, which describes how molecule properties relate to the macroscopic behaviors of an ideal gas—a theoretical gas that always obeys the ideal gas equation.
• The kinetic-molecular theory of gases assumes that ideal gas molecules (1) are constantly moving; (2) have negligible volume; (3) have negligible intermolecular forces; (4) undergo perfectly elastic collisions; and (5) have an average kinetic energy proportional to the ideal gas’s absolute temperature.

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Did you know?

Did you know that the more concentrated a solution is, the lower its freezing point and the higher its boiling point? This is why antifreeze keeps your car engine from freezing in frigid weather or overheating on very hot days. Forces at work on a molecular level determine what happens when a solution is formed. A look at the chemistry of solutions reveals why some substances dissolve more easily than others and why some compounds don’t dissolve at all.

Summary

Aqueous solutions are found throughout our world, and their chemistry depends in part on how much of a dissolved substance is in them. This module explores how substances dissolve, why some substances don’t dissolve, and how we express the concentration of a solution. The module describes the forces that hold particles together and interactions that keep dissolved particles apart. It also examines how concentration affects freezing point, boiling point, and vapor pressure.

Key Concepts

• A solution is formed when solute particles are randomly distributed and dissolved in a solvent.
• Molarity is a measure of the solute concentration in a solution, and remains consistent when a fraction of the solution is poured off.
• In polar solutions, the charges on both the solute and solvent particles keep the solute dissolved, as the polar solvent molecules surround the solute particles and keep them apart.
• The relative solubility of a salt or polar compound in water is a balance of two forces: the attraction between atoms of the salt molecule, and the attraction between the ions and the water molecules.
• Solutions of non-polar solutes in non-polar solvent are driven by London dispersion forces, another type of attraction between molecules.
• Colligative properties of solutions—freezing point depression, boiling point elevation, and vapor pressure lowering—are related to the concentration of solute molecules but independent of the specific solute type.

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Did you know?

Did you know that Galileo and his chief rival, Ludovico delle Colombe, had a famous debate on why ice floats on water? Delle Colombe claimed it was the broad and flat shape of ice, whereas Galileo claimed it was the lower density of ice that allowed it to float.

Summary

Water is a truly unusual and important substance. The unique chemical properties of water that give rise to surface tension, capillary action, and the low density of ice play vital roles in life as we know it. Floating ice protects aquatic organisms and keeps them from being frozen in the winter. Capillary action keeps plants alive. Surface tension allows lily pads to stay on the surface of a lake. In fact, water’s chemistry is so complex and important that scientists today are still striving to understand all the feats this simple substance can perform.

Key Concepts

• Water has a number of unique properties that make it important in both the chemical and biological worlds.
• The polarity of water molecules allows liquid water to act as a "universal solvent," able to dissolve many ionic and polar covalent compounds.
• The polarity of water molecules also results in strong hydrogen bonds that give rise to phenomena such as surface tension, adhesion, and cohesion.

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Did you know that you don’t need to be a professional scientist to make use of acid-base chemistry? When you eat too much of a spicy food, it is acid that puts the “burn” in heartburn. And when you seek relief through medication or baking soda in water, your goal is to bring about a chemical reaction called neutralization.

Summary

We experience the effects of acids and bases in our everyday lives, which is why these types of molecules were first described so many centuries ago. The definition of acids and bases has been refined over time to reflect an increased understanding of their chemical properties. We now know that the properties that accompany pH arise because of the concentration of hydrogen ions. A major step in our ability to analyze and use acids and bases was the creation of the pH scale, which gave scientists a numerical way to describe H+ concentration. This understanding of pH allows us to determine pH when we don’t know it, and adjust pH value as needed. Thus, we can wield acid-base chemistry as a tool with wide-ranging uses, from cleaning up industrial waste to calming our overactive stomachs and keeping our pet fish healthy and happy. Acids and bases also play a crucial role in the body, maintaining a blood pH of around 7.4 (and saving us from nausea, headaches, and heart problems). To understand how that happens, we need a more accurate definition of acids and bases.

Key Concepts

• The Arrhenius definition of an acid is a substance that loses a proton (H+) when it dissolves in water. Conversely, a base is a substance that releases an OH- ion when dissolved in water.
• The Brønsted -Lowry definition of an acid is a substance that donates a proton (H+). A base is then a substance that accepts a proton.
• pH measures the acidity of a solution, and is determined by the concentration of hydrogen ions
• A pH of 7 is considered neutral, a pH below that is acidic, above that is basic.
• Titration is a neutralization technique that can be used to determine the pH of the substance being neutralized.

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Did you know?

We think of water as neutral, but did you know that water can act as both an acid and a base? This is one of water’s most important roles in our body chemistry. Key to the body’s pH-control system is this property of water that allows it to donate or accept protons to keep our body fluids and tissues in a state of equilibrium. Acid-base chemistry is also important to many other common processes such as dying fabrics, producing everyday grooming products, and manufacturing life-saving medications.

Summary

Our bodies constantly produce carbon dioxide, some of which stays in the blood and becomes carbonic acid. So what keeps blood from becoming too acidic? Our blood has a built-in buffering system that works to maintain a neutral pH. This module explores how our body achieves this remarkable balancing act, and traces scientific discoveries that contributed to our current understanding of acid-base chemistry.

Key Concepts

• The Brønsted-Lowry system defines a conjugate base as an acid without its proton, and a conjugate acid as a base with an additional proton.
• Water is amphoteric, meaning it can act as both an acid and a base, accepting a proton to become a hydronium ion (H₃O+), or dissociating into a proton and a hydroxide ion (OH-).
• The strength of an acid is based on its tendency to ionize in water, not on the concentration of the acid in a solution
• Buffers are solutions containing a weak acid and its conjugate base, allowing them to absorb a strong acid or base without much change in pH.

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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, a highly reactive 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 chemical bonding on a sliding scale 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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Did you know?

Chemical equations are an efficient way to describe chemical reactions. This module explains the shorthand notation used to express how atoms are rearranged to make new compounds during a chemical reaction. It shows how balanced chemical equations convey proportions of each reactant and product involved. The module traces the development of chemical equations over the past four centuries as our understanding of chemical processes grew. A look at chemical equations reveals that nothing is lost and nothing is gained in a typical chemical reaction–matter simply changes form.

Summary

Chemical equations are an efficient way to describe chemical reactions. This module explains the shorthand notation used to express how atoms are rearranged to make new compounds during a chemical reaction. It shows how balanced chemical equations convey proportions of each reactant and product involved. The module traces the development of chemical equations over the past four centuries as our understanding of chemical processes grew. A look at chemical equations reveals that nothing is lost and nothing is gained in a typical chemical reaction–matter simply changes form.

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Summary

Stoichiometry is the mathematics of chemistry. Starting with a balanced chemical equation, we make use of the proportional nature of chemical reactions to calculate the amount of reactant needed at the start or predict the amount of product that will be produced. While it may not seem all that “chemical,” stoichiometry is a concept that underlies our ability to understand the impact and implications of many chemical processes. A bandage manufacturer may use mole ratios to determine how much silver is required (and therefor the cost) to treat a batch of bandages with silver nitrate. A fertilizer company might apply the concept of limiting reactant to figure out how much product they can produce with a given amount of hydrogen gas. And so on. Stoichiometry, mole ratios, and limiting reactants are indispensable concepts for fully understanding any chemical process.

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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? But no matter the type of reaction, in every case a new substance is produced and is often accompanied by an energy and/or an observable change.

Summary

This modules explores the variety of chemical reactions by grouping them into general types. We look at synthesis, decomposition, single replacement, double replacement, REDOX (including combustion), and acid-base reactions, with examples of each.

• NGSS

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

Key Concepts

• The steps from a qualitative science to quantitative one, were crucial in understanding chemistry and chemical reactions more completely.
• When a substance or substances (the reactants), undergo a change that results in the formation of a new substance or substances (the products), then a chemical reaction is said to have taken place.
• Mass and energy are conserved in chemical reactions. Matter is neither created or destroyed, rather it is conserved but rearranged to create new substances. No energy is created or destroyed, it is conserved but often converted to a different form.
• Chemical reactions can be classified into different types depending on their nature. Each type has its own defining characteristics in terms of reactants and products.
• Chemical reactions are often accompanied by observable changes such as energy changes, color changes, the release of gas or the formation of a solid.
• Energy plays a crucial role in chemical reactions. When energy is released into the surroundings the reaction is said to be exothermic; when energy is absorbed from the surroundings the reaction is said to be endothermic

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Did you know?

Did you know that temperature affects the chemical reactions that form air pollution? That’s why air quality tends to be better in the winter than in the summer. Changes in energy drive chemical reactions, so the hotter the weather, the faster air pollutants form. Other factors that can change the speed of chemical reactions include concentration, temperature, and surface area.

Summary

This module explores the study of reaction rates and the variables that can speed them up or slow them down, also known as chemical reaction kinetics. Through real-life examples, we examine our understanding of chemical reaction products, and the developments scientists have made in manipulating them.

• NGSS

• HS-PS1.B1, HS-PS1.B3

Key Concepts

• Current scientific theory suggests that chemical reactions are driven by collisions between molecules - and factors that affect those collisions affect chemical reactions.
• The study of chemical reaction rates, and the variables affecting those rates, is referred to as Chemical Reaction Kinetics. And the rate of a chemical reaction is the speed at which reactant molecules convert into product molecules.
• By collecting data about reaction rates during experiments, scientists have recognized that most reactions do not happen spontaneously, but rather need a certain amount of energy to get them started. This minimum energy needed to “activate” the reactants is called a reaction’s activation energy.
• Because reactions are driven by collisions, and the speed at which reactants are moving affect those collisions, chemical reactions are driven by changes in energy, and changing the energy of the reactants changes the rate of a chemical reaction.
• Experimentation conducted by scientists have shown that several factors influence reaction rates, including temperature, surface area, concentration, pressure, mixing, and sunlight (for photochemical reactions only).
• Many scientists have observed that certain substances can help speed up chemical reactions although they are not chemically changed in the process. These substances are called catalysts, and they work by reducing a chemical reaction’s activation energy.

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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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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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