Hi there! I’m Jeremy Krug, and welcome to my review of AP Chemistry Unit 3 – which covers Properties of Substances and Mixtures. If you find this video helpful, I’d really appreciate it if you could HIT that like button, subscribe, leave a comment, and share this video with the rest of your AP Chem community! And don’t forget that I have the FULL 30-minute review video, Unit 3 Key Concepts and guided notes, practice multiple choice AND free response questions, and a full unit study guide over at Ultimate Review Packet dot com. The link is in the description down below. And now, for my TEN MINUTE review of Unit 3…. All molecules exhibit London dispersion forces. Usually these are the weakest intermolecular forces, but in large molecules, they CAN be stronger than other forces. For nonpolar molecules, these are the only intermolecular force present. The more electrons a molecule has, the more polarizable it is, and the stronger its dispersion forces. Polar molecules exert dipole-dipole forces on each other. The positive pole of one molecule attracts the negative pole of its neighbors. These forces are USUALLY stronger than dispersion forces. Any molecule that contains a hydrogen atom bonded to an oxygen, fluorine, or nitrogen atom will exhibit hydrogen bonding. Hydrogen bonding is an especially strong type of intermolecular force. For all intermolecular forces, the stronger the attraction, the higher the boiling point and melting point of the compound. London dispersion forces are usually the weakest, then dipole-dipole, and hydrogen bonds are strongest. Another force can exist between polar molecules – like water – and ions in an ionic compound. This is called an ion-dipole force. In this example, the positive pole of water surrounds a negative ion like fluoride and literally drags it into solution. The negative pole of water surrounds positive ions like sodium and does the same thing. If the ion-dipole force is stronger than the ionic forces holding the compound together, the compound will dissolve in water quite easily. Different types of solids have different properties. Ionic solids have high melting points due to the strong attractions between oppositely-charged ions. They’re also brittle and conduct electricity when dissolved in water. Covalent network solids, such as diamond and silicon dioxide, have some of the strongest covalent forces out there. Each atom is bonded to multiple other atoms in multiple directions, making them extremely strong. Molecular solids have individual molecules as units, so that makes them different from ionic solids or covalent network. Sugar is a good example. These have relatively weak forces between molecules and relatively low melting points. Metallic solids include pure metals, as well as alloys. Metals are malleable and ductile, and they conduct electricity well, because of a metallic core surrounded by a sea of electrons. True solids are crystalline in structure. Some solids are not completely crystalline; we call these amorphous solids. These include materials like plastics and other semi-solid materials with a non-crystalline structure. The particles in a solid are usually very close to each other, and molecules have the least freedom of motion, with only vibrational motion. In a liquid, molecules are usually a little farther apart, so they can slip and slide around each other; that’s why a liquid can flow. In a gas, molecules are moving independently of each other and are very far apart from each other. Only gases can be truly compressed or can expand to fill a container. Be able to use the Ideal Gas Law, PV=nRT. Pressure is in atmospheres, volume is in liters, n is the number of moles of gas, and T is the temperature in Kelvins. R will be the Universal Gas Constant, which is equal to 0.08206 liter atmospheres per mole Kelvin. With the constant and three of the four variables, you can solve for the unknown. When there’s a mixture of gases, the partial pressures of the individual gases add up to the total pressure in the container. And to determine the partial pressure of any one gas, just multiply the mole fraction that gas occupies in the container by the total pressure. Be able to recognize the graphical relationships between pressure, volume, temperature, and number of moles of gas. Those graphs look like this. Temperature is a measure of the average kinetic energy of the molecules in a sample. So if two different materials have the same temperature, their molecules have the same average kinetic energy. If one material has a higher temperature, its molecules have a higher average kinetic energy. A graph like this Boltzmann distribution helps us visualize the range of motion in molecules at different temperatures. No matter the temperature, some molecules will be moving faster and some slower. At higher temperatures, a greater fraction of particles are moving faster. The Ideal Gas Law truly only works for ideal gases. Ideal gases have no intermolecular attractions, and their particles take up no space. Of course, there’s no such thing as an ideal gas in the real world, but some gases do get fairly close. Gases with very little interparticle attraction – and a very small molecular size – like helium, approximate ideal gases. Gases normally approximate ideal conditions the best when they’re at high temperatures and low pressures, which are conditions that bring about the least interactions between particles. There are two main types of mixtures. Heterogeneous mixtures are those where you can see the different components with your eyes. Homogeneous mixtures are those where all the components are distributed uniformly, and they’re more commonly called solutions. Make sure you can comfortably work with molarity, which is the most common way to talk about solution concentration. Molarity equals moles of solute divided by liters of solution. To calculate moles of solute in the solution, just rearrange the equation and multiply molarity by liters. Diagrams representing the relative concentrations of components in a solution can be helpful for describing what’s in the solution. When drawing a diagram like this, first of all pay close attention to mole ratios. In this diagram, two of the aluminum atoms disappeared, so that means they reacted. Following the mole ratios, that means we should lose six silver ions, leaving just two on the product side. This tells us that we gain two aluminum ions, along with six silver atoms. Finally, check the states. Anything that’s a solid should be drawn clumped together, while ions are aqueous and should be swimming around in the solution. There are two primary ways to separate solutions into components in AP Chem. Distillation takes advantage of the fact that different substances have different boiling points. So you can boil the mixture, and the different components will boil away at different temperatures. You can collect the steam and condense the components into different containers. Chromatography involves a mixture passing through a column. Some of the components adhere more strongly to the column, so they pass through it more slowly. Other components don’t adhere as well to the column, so they pass through more quickly. The column itself is sometimes called the stationary phase, and the substances that pass through the column are called the mobile phase. Sometimes you’re asked if a compound will dissolve into another compound. If you don’t have a straightforward solubility rule to follow, just remember the rule of thumb “Like dissolves like.” Polar molecules tend to dissolve into polar solvents, often through hydrogen bonding or dipole-dipole forces. On the other hand, nonpolar molecules tend to dissolve into nonpolar solvents. Different parts of the electromagnetic spectrum do different things to molecules. For example, ultraviolet or visible light can cause an atom’s electrons to transition to different energy levels. If you see the phrase infrared radiation in the question, you should realize that the molecule is undergoing vibrations at those frequencies. And finally, if the question mentions microwave radiation, you can associate that with rotation of molecules. Light has a dual nature. It can act as a wave, and it can act as a packet of light, which we call a photon. We can calculate the energy of these photons of light. Use the equation c equals lambda nu, where c is the speed of light, around 3 x 108 meters per second, lambda is the wavelength in meters, and nu is the frequency in Hertz. If you know the wavelength or the frequency, you can figure out the other one. We can use the equation E equals h nu to calculate the energy of a photon. E is the energy of the photon, h is Planck’s constant, which is 6.626 x 10-34 Joule seconds, and nu is the frequency. If you’re given wavelength, you’ll need to use both equations, but it’s a fairly simple calculation. When we work with spectrophotometry in the lab, we use the Beer-Lambert Law. Often written A = epsilon b c, the A is absorbance, which is a value that we normally read right off the spectrophotometer. Epsilon is the molar absorptivity, which is dependent on the substance and the wavelength. B is the path length, or the width of the cuvette, usually that’s 1 centimeter. And c is the concentration of the sample in moles per liter. Since epsilon and B are usually held constant in an experiment, we can plot concentration on the x axis and absorbance on the y axis for our known concentrations. Build a calibration curve, and use that to determine your unknown concentration. If you have outliers on your calibration curve, there’s usually a reason. If you have a dot that’s too high, it probably got contaminated with a more concentrated solution. If a dot is too low, it probably got diluted with a more dilute solution, or very likely with water. That’s Unit 3! I’m Jeremy Krug, join me next time for my 10 minute review of Unit 4. Keep up the good work, and keep learning chemistry!