|Course Introduction||Course Syllabus|
|1.1: Matter||Extensive and Intensive Properties||
Read this text, which gives examples of each type of property.
|Physical and Chemical Properties||
Scientists classify properties of matter as physical or chemical. Physical properties are those we can observe without altering the identity of the substance. For example, the melting or boiling point of a substance is a physical property because they do not alter the identity of the substance.
We can only observe chemical properties when we alter the identity of the substance. Rusting is an example of a chemical property because it is a chemical reaction that changes the composition of the substance.
Read this text. Pay attention to the chart at the end of the section, which gives examples of some physical and chemical properties of the element sodium (Na).
|Solids, Liquids, and Gases||
The three phases of matter are solid, liquid, and gas. Note the differences between these three phases of matter on the microscopic and macroscopic levels.
Read this text. Pay attention to the first image which shows the microscopic differences between the phases of matter.
|Density and Its Uses||
Density is an intrinsic property of matter. We define density (d) as the mass or volume of a substance at a given temperature. We write d = m/v where d is density, m is mass, and v is volume. If we know two of the variables in this equation, we can solve for the third algebraically. The units for density are a mass unit divided by a volume unit. The units used to describe density often differ for the phases of matter: solids (g/cm3), liquids (g/mL), and gases (g/L).
After you read this section, try the practice problem examples 1 and 2.
|Calculations Using the Density of an Unknown Liquid||
Watch this video for additional practice with density problems as the instructor works through problems using the d = m/v formula.
|How Do We Classify Matter?||
Chemists classify matter as a pure substance or a mixture. A pure substance consists of only one type of matter, while a mixture consists of multiple types of matter. Pure substances are further categorized as single-element or compound. Mixtures are further categorized as homogeneous (single-phase) or heterogeneous (multiple phases). The distinction between homogeneous and heterogeneous mixtures presented here is dependent on phase, or physical, boundaries. Mixtures, whether homogeneous or heterogeneous, can be separated by physical means into pure substances.
Read this text. Pay attention to the flowcharts that classify types of matter.
|Physical and Chemical Changes||
In chemistry, we often study changes in matter. Two types of changes can occur in matter: physical and chemical.
To determine whether you are dealing with a physical or chemical change, ask yourself if you can reverse the process to recover the original material. Physical changes can be reversed, but chemical changes generally cannot. For example, ice melting is a physical change because you can re-freeze the water. However, cooking a steak is a chemical change because you cannot recover the raw meat. Note that we discuss the energetics of chemical change more thoroughly in Unit 6.
Watch this video to see examples of physical and chemical changes, and how we can observe a change to classify it.
|1.2: Measurement and Notation||Units of Measure||
Read this section. Be sure to memorize the SI units of measure listed in the text and the metric prefixes listed in the section "The SI Decimal Prefixes". We will use these units of measure throughout this course, so you need to commit them to memory.
|Practice: Converting Units||
Next, complete this practice set.
|Uncertainty Is Certain!||
Uncertainty exists in any measured quantity because measurements are always performed by a person or instrument. For example, if you are using a ruler to measure length, it is necessary to interpolate between gradations given on the ruler. This gives the uncertain digit in the measured length. While there may not be much deviation, what you estimate to be the last digit may not be the same as someone else's estimation. We need to account for this uncertainty when we report measured values.
When measurements are repeated, we can gauge their accuracy and precision. Accuracy tells us how close a measurement is to a known value. Precision tells us how close repeat measurements are to each other. Imagine accuracy as hitting the bullseye on a dartboard every time, while precision corresponds to hitting the "triple 20" consistently. Another example is to consider an analytical balance with a calibration error so that it reads 0.24 grams too high. Although measuring identical mass readings of a single sample would mean excellent precision, the accuracy of the measurement would be poor.
Read this text, which text describes how uncertainty comes about in measurements. It uses the example of a dartboard to differentiate between accuracy and precision.
To account for the uncertainty inherent in any measured quantity, we report measured quantities using significant figures or sig figs, which are the number of digits in a measurement you report based on how certain you are of your measurement. Reporting sig figs properly is important, and we need to account for sig figs when performing mathematical calculations using measured quantities. There are rules for determining the number of sig figs in a given measured quantity. There are also rules for carrying sig figs through mathematical calculations.
Watch these two videos to learn how to count sig figs for a given quantity.
|Using Significant Figures||
Next, watch these two videos. Note that the rules for sig figs for addition and subtraction are different from the rules for sig figs for multiplication and division.
|Significant Figure Practice||
After you watch the videos, complete this practice set.
|2.1: Atoms and Elements||Atoms, Elements, and the Nucleus||
Read this text, which gives some of the history of the development of atomic theory. In the section, "Atoms Become Real", pay close attention to the law of conservation of mass-energy and the law of definite proportions. These laws define how we describe chemical reactions. Also, pay close attention to Dalton's Atomic Theory. John Dalton was the first person to propose a cohesive theory for how atoms make up matter.
|The Nuclear Atom||
As stated above, we define the elements by their number of protons. We define the atomic number, Z, as the number of protons in an atom. For a neutral atom (not a charged ion), the number of electrons must equal the number of protons. However, the number of neutrons can vary within atoms of a given element.
Atoms of the same element with different numbers of neutrons are called isotopes. Most elements have multiple isotopes. For a given isotope, we define the mass number, A, as the atomic number plus the number of protons. We can write this as A = Z + N where A is mass number, Z is atomic number, and N is number of neutrons.
Read this section, which explains how we write atomic symbols with atomic numbers and mass numbers.
When we see the mass of an element on the periodic table, we are seeing the weighted average of the masses of all isotopes of that element. Then read the section "Isotopic Mixtures and Abundances" near the bottom of the page. This section describes how to determine the average atomic mass of an element if we know the isotope masses and their relative abundance. Try the practice problems to do these types of calculations yourself.
|2.2: Avogadro's Number and Moles||Avogadro's Number and the Mole||
Read this text, which defines Avogadro's Number and moles. It also shows examples of conversions between particles and moles, and conversions between moles and mass.
|Atoms to Moles||
Watch these two videos, which provide step-by-step tutorials on how to convert between the number of atoms and moles using Avogadro's Number as a conversion factor, and how to convert between mass and moles using molar mass as a conversion factor. After you watch one example, you may want to pause the video and try an example on your own. We will be using these types of calculations later in this course.
|2.3: Atomic Theories||Quanta: A New View of the World||
Read this text, which outlines many of the experiments done in the early part of the twentieth century that led to quantum theory. Pay attention to the equation E = hv, which relates frequency (v) to energy (E) of a wave by Plank's Constant (h).
|Light, Particles, and Waves||
Read this text. Section 1 describes Young's Double Slit Experiment, which showed that atom-sized particles exhibit wave-like properties by producing wave-like interference patterns. The illustrations in this section show the differences between macroscopic particles and atom-sized particles.
Section 2 provides an overview of the properties of waves, including amplitude, frequency, wavelength, and energy. Pay attention to the examples where energy, frequency, and wavelength are calculated for a given wave. Also, note the relative energies in the electromagnetic spectrum.
Section 3 shows that different atom types emit different frequencies of energy, or spectra, which are characteristic of that element. We can use these atomic spectra to identify elements in a sample.
Finally, section 4 describes how we can calculate the wavelength of a moving atom-sized particle using the deBroglie relation. It also outlines the Heisenberg Uncertainty Principle, which states that it is impossible to know both the position and momentum of a moving atom-sized particle with any precision.
|The Bohr Model||
The first model of the atom to use quantum mechanics was the Bohr Model of the hydrogen atom. You have no doubt seen this model, as it is frequently used in logos. The Bohr model, often called the planetary model, consists of the nucleus surrounded by electrons in fixed orbits circling the nucleus, much like planets orbit the sun. This model properly accounted for the hydrogen atomic spectrum and introduced the idea of quantum numbers – numbers that define the energy of an electron. However, as we learn in our next resource, this model is incorrect.
|The Quantum Mechanical Model of the Atom||
The current view of atomic structure is that electrons exist in a cloud surrounding the nucleus, rather than in fixed orbits. The electrons exist in orbitals, areas of high probability of finding the electron. Different orbitals have different shapes and sizes, which correspond to different energy levels.
Read this text, which reviews the Bohr model and then shows how more advanced quantum mechanics gave rise to the orbital model of the atom.
|2.4: The Periodic Table of the Elements||Shells, Subshells, and Orbitals||
This video reviews the different orbital shapes and names. The electrons in an atom are arranged in shells that surround the nucleus, with each successive shell being farther from the nucleus. Electron shells consist of one or more subshells, and subshells consist of one or more atomic orbitals. Electrons in the same subshell have the same energy, while electrons in different shells or subshells have different energies.
This video reviews the types of orbitals and shows how we can count electrons and place them in orbitals using the periodic table.
Electron configurations describe where electrons are located around the nucleus of an atom. For example, the electron configuration of lithium, 1s²2s¹, tells us that lithium has two electrons in the 1s subshell and one electron in the 2s subshell.
|The Aufbau Principle||
Our third video shows how to use the Aufbau Principle when filling electrons in orbitals for larger elements. This principle tells us that we fill electrons from the lowest energy to the highest energy.
The Aufbau principle states that electrons fill lower-energy atomic orbitals before filling higher-energy ones (Aufbau is German for 'building-up'). By following this rule, we can predict the electron configurations for atoms or ions. The Aufbau principle is most useful for the first 20 elements: from Sc on, the Aufbau principle does not accurately predict the order of electron filling in atoms.
Our final video in this sequence introduces the idea of an outer shell, or valence electrons, in an electron configuration. Because valence electrons are in orbitals in the outermost shell, they are the electrons that can be involved in chemical reactions.
Valence electrons are the electrons in the outermost shell, or energy level, of an atom. For example, oxygen has six valence electrons, two in the 2s subshell and four in the 2p subshell. We can write the configuration of oxygen's valence electrons as 2s²2p⁴.
|Periodic Properties of the Elements||
Because the elements in the periodic table are listed in a way that relates to their valence electrons, we can infer many properties of elements based on their relative position on the periodic table. Many of these properties explain the reactivity of the elements.
Read this text, which defines properties including atomic radius (size), ionization energy (energy required to remove an electron from an atom), electron affinity (tendency of an atom to gain electrons), and electronegativity (tendency of an atom to pull electrons toward itself in a bond).
|3.1: Chemical Bonds||What are Covalent and Ionic Bonds?||
Watch these videos for an overview of the differences between the two types of bonding. Note the types of elements that are involved in each type of bonding.
|Ionic, Covalent, and Metallic Bonds||
Watch this video, which reviews ionic and covalent bonds, and introduces another common type of bonding called metallic bonds. Metallic bonds occur between metal atoms.
|Predicting Bond Type – Electronegativity||
When discussing covalent bonds, we need to make one further definition. Nonpolar covalent bonds are covalent bonds where the electrons are equally shared in the bond. Polar covalent bonds are covalent bonds where the electrons are not equally shared in the bond. This occurs when the bonding atoms have different electronegativities. Watch this video to learn how to identify polar covalent bonds.
|Models of Chemical Bonding||
Now that we have an overview of covalent and ionic bonding, we can see how quantum mechanics informs our understanding of bonding.
As you read this text, note that the section "Classical Models of the Chemical Bond" describes ionic bonding, which is largely based on the attractive forces among charged particles. This section also discusses models of covalent bonding that describe bonds as areas of shared electrons between atoms. While this is simplistic, it does accurately describe the three-dimensional shape of molecules.
The section "Quantum Models of Chemical Bonding" briefly outlines hybrid orbital theory, which is also often used to describe the three-dimensional shape of molecules. We will explore these models further in later sections of this unit.
|Drawing Dot Structures and Lewis Diagrams||
Now that you understand the types of chemical bonds, we explore Lewis Dot Diagrams, a common way to draw covalently-bonded molecules. In Lewis dot diagrams, we use small dots to show the valence (outer shell) electrons of an atom. We can then combine the valence electrons of different atoms together to form covalent bonds, using a set of rules.
Watch the following videos in order. These four videos guide you through the steps for making Lewis dot diagrams for small molecules, and show worked examples.
|Exceptions to the Octet Rule||
Most Lewis dot structures are guided by the Octet Rule, which states that bonded atoms must have eight valence electrons. However, there are exceptions to the octet rule. Some atoms, such as hydrogen, helium, and boron, have less than an octet. Many atoms, such as those past sulfur on the periodic table, can have an expanded octet, meaning more than eight valence electrons. Watch this video, which outlines these cases.
|Resonance Hybrids and Dot Structures||
In some cases, we can draw more than one equivalent Lewis structure for a given molecule or polyatomic ion. These equivalent structures are called resonance structures. The true structure of these types of molecules or polyatomic ions are actually a hybrid, or mixture, of the resonance structures. This is called a resonance hybrid.
Watch these videos, which introduce resonance structures with the example of the nitrate, Na3- ion, which has three resonance structures. All three structures are equivalent – they only differ by the placement of the double bond.
|Formal Charge and Dot Structures||
Sometimes we can draw more than one Lewis structure for a given molecule or polyatomic ion that are not equivalent. In this case, we use a concept called formal charge to determine, which Lewis structure is best. Formal charge is not actually a charge; rather, it is just a system to keep track of electrons in a given Lewis structure.
Watch these two videos, which teach you the rules for assigning formal charge to atoms in a Lewis structure, and show examples of using formal charge to determine the best possible Lewis structure for a given molecule.
|Interactions Between Molecular Units||
Now that we have studied bonding and how to draw chemical structures, we can investigate how molecules interact with each other. Molecules interact with each other through intermolecular forces; forces that hold molecules together, but are not covalent or ionic bonds. As we will learn, there are different types of intermolecular forces that occur between different types of molecules.
Read this text. The first two sections describe the differences between bonding interactions and intermolecular force interactions in terms of electrostatics, or charge interactions. You should focus on section 3, which outlines the different types of intermolecular forces and the types of molecules that experience these different forces.
|3.2: Molecular Structure||Molecular Geometry||
Read this text, which outlines the different geometries predicted by VSEPR theory. Pay attention to the pictures that show the shapes of these different geometries.
Next complete this VSEPR practice assessment, which gives you the opportunity to practice applying your understanding of Lewis structures and VSEPR.
|The Hybrid Orbital Model||
While VSEPR works well to describe molecular geometry, it does not take into account what we know about atomic structure from quantum mechanics. How do electrons in atomic valence orbitals form bonds? The hybrid orbital model helps to explain this.
Read these sections, which describe how atomic orbitals combine to form hybrid orbitals in molecules. The molecular shapes and the bond angles of hybrid orbitals match the molecular shapes and bond angles predicted by VSEPR theory.
|4.1: Chemical Formulas||Chemical Formulas and Their Arithmetic||
Read this text, which reviews how to write chemical formulas, how to determine information about the compound such as molar mass, mole ratios, and percent composition.
|Empirical and Molecular Formulas||
Watch this video, which explains the difference between empirical and molecular formulas with multiple examples.
|Formula from Mass Composition||
We can determine the empirical formula of an unknown compound from experimental data. Using instrumentation in the lab, we can determine the elemental mass composition of a compound. In other words, we can determine the percent of each element present in the compound. From that information, we can backtrack to determine the empirical formula of an unknown compound.
Watch these videos, which go through these types of problems step by step.
|Molecular and Empirical Formulas from Percent Composition||
If we know the molecular mass or molar mass of the unknown compound, we can determine the molecular formula of the unknown compound.
Watch this video, which shows an example of using percent composition data and a known molar mass to determine the molecular formula of an unknown compound.
Nomenclature is a fancy word for the process of naming compounds. In chemistry, we have a set of specific rules used to name compounds so scientists can communicate effectively. We will use this nomenclature for the remainder of the course, so you must be comfortable naming a compound correctly from its formula, and become comfortable determining the formula of a compound from its name. Here, primarily inorganic substance nomenclature is covered.
Read this text. The first section covers rules for naming inorganic molecular compounds. The second and third sections cover naming rules for ionic compounds. Be sure to focus on the systematic name for compounds that have a systematic name and a common name.
|4.2: Stoichiometry and Limiting Reagents||Balancing Chemical Reactions||
Watch this video to see examples of how to systematically balance a chemical equation.
Watch these videos, which show all the types of stoichiometry problems we see in chemistry.
|Stoichiometry Example Problems||
Pause each video after the problem is presented and try the problem yourself. Then, watch the rest of the video to check your work.
|5.1: Gases and Gas Laws||Observable Properties of Gases||
Read this text, which introduces the important gas properties of pressure, temperature, and volume.
|The Ideal Gas Equation||
We use the Ideal Gas Law PV = nRT – where P is pressure, V is volume, n is the number of moles, T is temperature, and R is the gas constant – to describe the relationship among pressure, volume, temperature, and the number of moles of a gas.
Watch these videos, which show how to solve for the different variables in an Ideal Gas Law problem. Pay attention to the units used for the variables in the Ideal Gas Law. It is necessary to use the proper units to ensure your units cancel when you complete an Ideal Gas Law problem.
|Charles' Law, Boyle's Law, and Avogadro's Law||
The scientists who formulated the Ideal Gas Law based their theory on Charles' Law, Boyle's Law, and Avogadro's Law, which showed the relationship between just two variables in a gas sample.
Watch these videos to see simulations of what happens when you manipulate one gas variable in a container – pressure, volume, temperature, number of moles – and how the modification affects the other variables.
In many cases different types of gases are mixed. This is true of the air around us. Air consists of nitrogen, oxygen, and a variety of other gases. We can use Dalton's Law of Partial Pressures to determine the pressure exerted by each gas in a given mixture, and the total pressure exerted by a mixture of gases. Dalton's Law of Partial Pressure states that a mixture has a total pressure that is the sum of its component, or partial, pressures.
Watch this video, which introduces Dalton's Law of Partial Pressures, and shows how to use it and the Ideal Gas Law to calculate the partial pressures of gases in a mixture.
Kinetic molecular theory (KMT) explains the properties of gas molecules and their motion. KMT helps explain how radio waves are transmitted and why the sky is blue.
Read this text. The first section defines the main points of KMT. The second section details how KMT explains the gas laws. The third section, "Some Important Practical Applications of KMT," describes many interesting natural phenomena, such as the color of the sky and how lightbulbs work, that come from KMT.
|5.2: Phase Changes||States of Matter||
Watch this video, which uses the example of water to remind us of the properties of solids, liquids, and gases. It also introduces a graph called a heating curve. Heating curves show us how the temperature changes as we add heat (q or △H) to a substance. Notice that while a phase change occurs, the temperature stays constant. This is why we can define melting points and boiling points for substances.
|Specific Heat, Heat of Fusion and Vaporization||
We will investigate thermodynamics, the study of heat transfer, in detail in Unit 6. However, we need to introduce some thermodynamics terminology here to understand how phase changes occur. We can define the specific heat of a substance as the energy it takes to raise 1 g of a substance by 1 degree Celsius. The unit used for heat is the joule (J), or kilojoule (kJ). We define heat of fusion as the energy or heat necessary to change a given amount of a substance (often one mole) from solid to liquid. Heat of fusion uses the unit of J or kJ. Heat of vaporization is the amount of heat needed to change a given amount of a substance (often one mole) from liquid to gas, with the unit of J or kJ.
Watch these three videos in order. Our first video uses the heating curve of water to define specific heat, heat of fusion, and heat of vaporization.
|Chilling Water Problem||
Our second video is a worked example showing how to calculate the heat needed to change an ice cube at a given temperature to liquid water at a warmer temperature.
|Change of State Example||
This video is a similar example. You can pause the third video after it presents the problem and try it yourself using the same technique as seen in the "chilling water problem" video.
As we saw above, heating curves show us how the temperature of a substance changes as we apply heat to it. We can use heating curves to determine the phase of a substance at a given temperature.
Another graph used to determine the phase of a substance is a phase diagram. Phase diagrams show the effect of pressure and temperature on the phase of a given substance. In a phase diagram, the temperature is shown on the x-axis and the pressure is shown on the y-axis. Curved lines show where phase changes occur. Areas between the curved lines show the conditions where the substance is a solid, liquid, and gas. All phase diagrams have a similar form, making them relatively straightforward to interpret.
Watch this video to see examples of phase diagrams, and how we can determine the phase of a substance at a given temperature and pressure. Also note the triple point, melting point, boiling point, and supercritical point on the phase diagram.
|Liquids and Their Interfaces||
Now, let's explore the properties of liquids, and what happens at the liquid-gas interface. Have you ever wondered why all soap bubbles are round? This is because of the interfacial properties of water!
Read this text, which highlights many important properties of liquids including viscosity and surface tension. The text explains how surface tension leads to the formation of soap bubbles.
Directly above the surface of a liquid, there is a small amount of vapor that comes from liquid particles that gain enough energy to escape into the gas phase. The pressure exerted by this vapor is called vapor pressure.
Watch this video to learn how this phenomenon occurs and how the strength of intermolecular forces affects vapor pressure.
|Vapor Pressure Example||
Next, watch this video to learn how to perform calculations involving vapor pressure that use the ideal gas law.
|Cubic Crystal Lattices and Close-Packing||
Our last topic in this unit is the solid phase of matter. There are two types of solids: crystals and amorphous solids. Crystalline solids have a regular, repeating pattern. The repeating unit is called the unit cell. Ionic compounds are examples of crystalline solids. Amorphous solids are things like glass that lack a repeating structure. Here we will focus on crystal structures.
Read this text. Pay attention to the three types of unit cells (cubic, body-centric cubic, and face-centered cubic) and how many atoms are contained in each type of unit cell.
|6.1: Energy||Energy, Heat, and Work||
Read this text, which offers more detailed definitions of these topics. Do not focus on the chart that shows different energy units. In chemistry, we almost exclusively use the joule as our energy unit. The other somewhat commonly used energy unit is the calorie.
|Molecules as Energy Carriers and Converters||
Read this text, which provides more detail about chemical energy. Here, we learn that the total energy of a system is the sum of all of the kinetic energy and all of the potential energy of that system. The energy of a molecular system is complex because molecules are constantly in many different forms of motion contributing to kinetic energy and many different interactions contributing to potential energy. Therefore, we can only measure the change in energy of a process rather than absolute energy.
This text also introduces the term enthalpy (H), meaning energy at constant pressure. If a process produces heat, we call it an exothermic process. If a process uses heat, we call it an endothermic process. Pay attention to the last section of the text, which discusses the phase changes we learned about in Unit 5.
Watch this video, which goes into more detail about the difference between internal energy (U) and enthalpy (H). The presenter uses a pressure-volume graph to show why internal energy is defined at constant volume and enthalpy is defined at constant pressure.
|Heat of Formation||
Watch this video, which defines the heat of formation (△Hf) as the energy change involved in forming one mole of a compound from its elements in their natural form. The presenter tabulates the heat of formations for most common compounds, which can be used to calculate the enthalpy change of a reaction if you have a balanced chemical equation. You can use the sign of the enthalpy of reaction to determine if the reaction is endothermic (positive enthalpy) or exothermic (negative enthalpy).
|Calculate Standard Enthalpy of Reaction||
Watch this video, which shows another worked example of how to use heats of formation to calculate the enthalpy of reaction. Note that many chemists often use the terms "heat" and "enthalpy" interchangeably.
|Thermochemistry and Calorimetry||
Scientists use calorimetry to measure the change in enthalpy of reactions in a laboratory setting. In calorimetry, we conduct the reaction in an isolated setting and measure the temperature change. We can then use the equation– where q is heat, m is mass, s is specific heat of the substance, and △T is change in temperature – to find the heat of the reaction.
There are two main types of calorimetry: constant pressure, or coffee cup calorimetry, and constant volume, or bomb calorimetry.
Read this text, which describes the two types of calorimetry, and shows worked examples of how to calculate heat of reaction from calorimetry data. Pay attention to the sign conventions here. In calorimetry, we directly measure the temperature change of the surroundings, not the system, and first calculate the heat of the surroundings. Then, we need to convert this to heat of the system.
|Coffee Cup Calorimetry||
Watch this video to see another worked example of a coffee cup calorimetry problem.
|Hess' Law Example||
Enthalpy is a state function. This means that enthalpy is path independent: it does not matter how you determine enthalpy experimentally or through calculation, you always get the same answer. This means you can use whatever information you have about a reaction to solve for enthalpy change. This is called Hess' Law.
One way to apply Hess' Law is to break a complex reaction up into the sum of smaller reactions. If you know the enthalpies of the smaller reactions, you can simply sum them to determine the enthalpy of the larger reaction.
Watch this video to see how to apply Hess' law. Pay close attention to what happens when a small reaction is flipped or multiplied by a factor. Changing a reaction in any way changes the enthalpy of that reaction.
|Calculate Standard Enthalpy of Reaction from Bond Dissociation Energies||
Another way to determine the enthalpy of reaction is by using tabulated bond dissociation energies. You can apply this in a similar manner to how we used standard enthalpies of formation. Using bond energy to calculate the enthalpy of reaction is not particularly accurate because energies are affected by each molecule's unique surroundings, such as intermolecular forces. For this reason, energies are only an average across many different reactions.
Watch this video, which shows a worked example of this type of problem.
|6.2: Thermodynamics||The First Law of Thermodynamics||
Read this text. Note the highlighted equation that defines the first law of thermodynamics. This equation shows that the change of internal energy of a system is the sum of heat and work. Pay attention to the sign convention box. This shows that when heat or work is transferred from system to surroundings, it has a negative sign. When heat or work is transferred from the surroundings to the system, it has a positive sign.
This text also introduces pressure-volume work, which refers to the work involved in changing the pressure or volume of a system. Problem Example 1 shows a typical pressure-volume work problem. Later parts of this text use calculus to derive useful equations. If you have not learned calculus, you can skip the calculus parts and still learn the important material. Focus on the worked problems and the problem examples.
|What is Entropy?||
The second law of thermodynamics involves a thermodynamic quantity we call entropy (S). Entropy is a measure of the disorder of a system, measured in joules (J). The second law of thermodynamics states that the entropy of the universe is always increasing.
One consequence of the second law of thermodynamics is that in any engine there will be some energy lost as heat that cannot be harnessed to do work. We observe this in our everyday lives. If you touch the hood of your car while the engine is running, the hood of the car will feel hot. This is because some of the energy from your car engine is lost as heat. Because of this, the second law of thermodynamics explains why a perpetual motion machine can never exist.
Read this text. The first section explains the difference between reversible and irreversible processes. A reversible process can be modeled as a series of tiny steps, while an irreversible process must be modeled as a single large change. The second section discusses the meaning of entropy, and what disorder means on a microscopic level.
Entropy is a state function, which means we can apply Hess' Law to it. Absolute entropies of most common substances are tabulated, allowing us to calculate the entropy of a reaction in the same way we can calculate enthalpy of reaction from standard enthalpies of formation.
|The Availability of Energy||
Read this text, which introduces the second law of thermodynamics. Pay close attention to the green box, which shows entropy calculations for the process of water freezing. It shows that while the system (water becoming ice) decreases in entropy, the entropy change of the universe is still positive.
The third law of thermodynamics states that the entropy of a perfect crystal at absolute zero (0K) is zero. This means that all molecular motion ceases in a perfect crystal at absolute zero. Pay attention to the definition of the third law of thermodynamics.
|Gibbs Free Energy and Spontaneity||
Gibbs Free Energy uses enthalpy, temperature, and entropy to predict whether a reaction or process will happen spontaneously or not. Recall that a spontaneous process is one that happens on its own, without needing "help" to go. A non-spontaneous process needs constant "help" to occur. If the Gibbs Free Energy for a process is negative, the process is spontaneous. If the Gibbs Free Energy is positive, the process is non-spontaneous.
It is important to point out that spontaneous does not mean fast; in fact, many spontaneous reactions are extremely slow. For example, the reaction of diamond turning to graphite is a spontaneous reaction, but it is extremely slow. So, we do not need to worry about diamond jewelry turning into graphite!
Watch this video, which describes the Gibbs Free Energy equation and shows how enthalpy, temperature, and entropy work together to determine if a process is spontaneous.
|Gibbs Free Energy Example||
This video shows a worked example of using the Gibbs Free Energy equation to determine the spontaneity of a chemical reaction.
|7.1: Acids and Bases||What are Acids and Bases?||
Read this text. Pay attention to the dissociation reactions for acids and bases when dissolved in water. Section 4 introduces neutralization reactions, where an acid and a base react together to form a salt (ionic compound) and water. Pay attention to the form of the neutralization reactions, and make sure you can identify the acid, base, and salt in each reaction.
|Arrhenius' Definition of Acids and Bases||
Watch this video for more discussion of the Arrhenius definition. Notice that Arrhenius acids increase the concentration of hydrogen ions in aqueous solutions, while Arrhenius bases increase the concentration of hydroxide ions.
|Aqueous Solutions, pH, and Titration||
Read this text, which introduces the pH scale that scientists use to determine whether a substance is acidic, basic, or neutral. The pH scale ranges from 0 to 14. A solution with a pH of 7 is neutral. If the pH is less than 7, the solution is acidic. If the pH is greater than 7, the solution is basic. Since the pH scale is logarithmic, each whole pH value below 7 is ten times more acidic than the next higher value. We determine pH based on the concentration of H+ ions in the solution.
In Part 1, pay close attention to the ion product of water (Kw) – this is an important value used in chemistry. In Part 2, you should know the equations that define pH. Try Problem Example 2.
This text also introduces titrations, an important analytical technique chemists use to determine the concentration of an unknown substance, called the analyte. When we do a titration, we slowly add a substance of known concentration, called the titrant, to the analyte and allow them to react until the reaction is complete. The point when the reaction is complete is called the equivalence point. Color-changing indicators are often used to visually show the equivalence point. If we know the amount of titrant used, we can use stoichiometry to determine the amount of analyte.
There are many types of titrations, but the most well-known are acid-base titrations, in which a neutralization reaction takes place during the titration. In Part 3, pay close attention to Problem Example 3, since this is a typical titration problem. You should also note the shape of a titration curve, and how to determine the equivalence point from a titration curve.
|Definition of pH||
Now, we'll take a more detailed look at pH and titration calculations. This video defines pH and the pH scale and shows examples of how to conduct pH calculations.
|Strong Acids and Strong Bases||
This video shows how to do pH calculations for solutions of strong acids or bases.
This video demonstrates an example of a strong acid-strong base titration. It covers indicators, endpoint, equivalence point, and calculating the unknown analyte's concentration.
|Bronsted-Lowry Definition of Acids and Bases||
Now, we are ready to learn about a second definition of acids and bases, the Brønsted-Lowry definition. According to the Brønsted-Lowry definition, an acid is a substance that can donate a proton (H+ ion) to another molecule. A base is a substance that can accept that donated H+.
After the Brønsted-Lowry acid donates its proton, it becomes the conjugate base of the acid. After the Brønsted-Lowry base accepts a proton, it becomes the conjugate acid of the base. We think of Brønsted-Lowry acids and bases in terms of their reactions with other molecules rather than just their structure. This is a more general definition and broadens the compounds that can be considered acids or bases.
Watch this video, which shows examples of how we can interpret a chemical reaction to determine, which substance is the acid, base, conjugate acid, and conjugate base. Water acts as a base because it is reacting with a strong acid. Water is an amphiprotic compound, which means it can act as an acid or a base. Hydrogen ions do not exist in water on their own, but immediately get grabbed by water molecules to form hydronium ions. A conjugate pair is always one acid and one base.
|Autoionization of Water||
Water acts as a base when reacted with a strong acid. This seems odd, since we know water is neutral. Water is an amphiprotic compound, meaning it can act as either an acid or a base. When an acid reacts with water, the water will act as a base. When a base reacts with water, the water will act as an acid. Water also autoionizes, meaning it can form H+ and OH-. This video shows how two water molecules react to form hydronium ion (H3O+, an acid), and OH-, but is still overall pH-neutral.
|Conjugate Acids and Bases||
Watch this video, which reviews how to determine the conjugate acid-base pairs in an acid-base neutralization reaction. You should be able to recognize conjugate acid-base pairs for a given reaction.
|7.2: Oxidation-Reduction Reactions||Redox Reactions||
Oxidation-reduction reactions are commonly called redox reactions. When a redox reaction occurs, two half-reactions occur simultaneously. One of the reactants undergoes oxidation, meaning that it loses electron(s). The other reactant undergoes reduction, meaning it gains electron(s). Oxidation and reduction must occur together – one half-reaction cannot happen alone. The reactant that undergoes oxidation is called the reducing agent, while the reactant that undergoes reduction is called the oxidizing agent.
Read this text. The author breaks the redox reaction up into the two separate half-reactions. Each half-reaction is either oxidation or reduction. In the oxidation half-reaction, electrons are lost. In the reduction half-reaction, electrons are being gained.
|Oxidation Numbers and Redox Reactions||
In the example in the previous text, we could tell, which species were gaining and losing electrons because the species were single elements and elemental ions. However, most redox chemistry is much more complex and cannot be determined by quickly looking at the reaction.
To determine if a substance is being oxidized or reduced in a reaction, we use oxidation numbers. Oxidation numbers are simply a record-keeping tool. We can calculate the oxidation numbers for all elements in a reaction. If an element's oxidation number increases from reactant to product, it is being oxidized. If an element's oxidation number decreases from reactant to product, it is being reduced.
Read this text, which lists the rules we use to determine oxidation numbers. You need to know these rules and be able to apply them to determine the oxidation numbers of elements in compounds. After you study the rules, try the example problems.
|Balancing Redox Equations||
Balancing redox reactions is more difficult than regular reaction balancing because we need to account for the electron transfer. We also need to balance the reaction differently depending on whether the solution is acidic or basic. This text reviews the steps used to balance redox reactions. You should know these steps for balancing a redox reaction and be able to balance a redox reaction using these steps.
|Write a Balanced Redox Reaction||
Watch these videos, which provide worked examples of how to balance a redox reaction in an acidic and in a basic solution. Be sure you are able to follow all of the steps in this balancing process.
|Common Oxidizing and Reducing Agents||
Read these sections, which look at some common oxidizing and reducing agents, and substances that can act as an oxidizing agent or reducing agent depending on what they react with. Water is one of the compounds that can act as an oxidizing or reducing agent.
|8.1: Types of Nuclear Decay||Radioactive Decay||
Read this text for a brief introduction to nuclear decay. Look at the reaction in the text. Recall from Unit 1 that we can write elements with the mass number (A) as a superscript and the atomic number (Z) as a subscript before the element symbol. The atomic number is the number of protons, while the mass number is the sum of the number of protons and neutrons. In the reaction, the element changes because the atomic number changes. The other product is called an alpha particle, which we will discuss in our next reading.
Read this text, which introduces alpha decay and provides an example of an alpha-decay equation. Notice that an alpha particle is helium. Take note of the conditions listed that must be met when writing an alpha decay reaction.
Read this text, which introduces beta decay and provides an example of a beta-decay equation. A beta particle is a high energy electron. In a beta decay reaction, there is no change in the mass number in beta decay, but the atomic number increases by one.
|Types of Decay||
There are four different types of nuclear decay. In addition to alpha and beta decay, gamma and positron emission can occur. Gamma emission does not have mass and often accompanies other types of nuclear emission. In positron decay, there is no change in mass number, but the atomic number decreases by one.
|Writing Nuclear Equations for Alpha, Beta, and Gamma Decay||
This video shows examples of how to write nuclear equations for alpha, beta, and gamma decay.
|8.2: The Half-Life of Radioactive Isotopes||Half-Life||
Read this text. Pay attention to the figure with the red and blue circles, which illustrates how samples shrink by half during each successive half-life.
Note that Equation 1 shows the equation for the radioactive decay rate constant.
If you know the rate constant for the decay of a given material, you can determine the half-life of the material using this equation. Also pay attention to Equation 2, which shows how to determine the initial concentration of the material if you know the half-life and the rate constant for the material. Carefully read the step-by-step example of this type of calculation in the text.
Read this text, which explains how we can use carbon dating to determine the age of an object. Scientists use these calculations extensively in archaeology, anthropology, and paleontology. Since the half-life for C-14 is approximately 5,700 years, carbon dating is only useful for fossils that are younger than approximately 50,000 years old. For older fossils, scientists use isotopes with longer half-lives.
|8.3: Nuclear Fusion, Nuclear Fission, and Energy Production||Transmutation of the Elements||
Read this text to see examples of transmutation using the types of nuclear decay you have already studied in this course.
|The Mass Defect||
Read this text, which introduces the concept of mass defect. Mass defect refers to the difference observed in the atomic mass of an atom, and the sum of the masses of the protons, neutrons, and electrons that make up the atom. The unaccounted-for mass is converted to energy in nuclear reactions, which is why nuclear reactions produce an enormous amount of energy.
As we just learned, the mass defect can potentially lead to an enormous release of energy. Read this text to learn about this type of nuclear reaction. Fission means to break up. This text shows the chain reaction that can occur when a nuclear fission reaction is started. In the example in the text, each reaction between a neutron and a uranium isotope produces three more neutrons, which can each react in turn with another uranium. A certain amount of fission material must be available to sustain this reaction. We call this the critical mass.
|How Does a Nuclear Reactor Work?||
Read this text, which explains how a nuclear reactor works. It also covers the different parts of the nuclear reactor (the core, coolant, turbine, containment, and cooling towers) and their functions. While nuclear reactors are generally clean and efficient at producing power and accidents are rare, there is an issue of storing radioactive waste. There are extreme dangers if a nuclear reactor malfunctions, as radiation poisoning can be deadly, lead to cancer, and cause birth defects.
Read this text. Fusion means merging together, which is what occurs during fusion reactions when two atoms fuse to form an atom of a larger element. The fusion of two hydrogen isotopes into helium is the process that powers the stars in the universe, such as our sun. There have been attempts to harness nuclear fusion to produce power, but so far these attempts have been unsuccessful.
|Study Guide||CHEM101 Study Guide|
|Course Feedback Survey||Course Feedback Survey|
|Archived Materials||Formula Sheet||
This is the formula sheet that may be used while taking the CHEM101 final exam.