Has AP Chemistry got you down? Does it seem like TOO MUCH material? Hi there! My name is Jeremy Krug, and in this AP Chemistry speed review, we’re going to hit ALL the major topics in the AP Chemistry course in less than 20 minutes. Now this video can’t replace a full AP course, but if you want to review, this is a good place to start. But there’s a good chance you might want an even better, more complete, maybe even the ULTIMATE review packet, for AP Chem. I’m proud to be joining the AP content area experts at Ultimate Review Packet dot com to provide the absolute best AP Chem review out there. If you’ve seen my YouTube review videos and my full 101-video AP Chem course, then you’re going to LOVE the Ultimate Review Packet, because I’m in the process of preparing study guides for every unit (and I’ll even give you the answers), longer review videos that cover tips and tricks for every unit, plus a full length exam with answers. It’s just $24.99 (and if you can get your teacher to pay for the whole class, they’ll even get a FORTY PERCENT DISCOUNT). So head over to Ultimate Review Packet dot com and check it out! Now, buckle up, here goes the Speed Review! Unit 1 covers atoms. We use a unit called the MOLE to count large numbers of atoms and molecules. When you weigh out a mole of an element, it’s equal to its atomic mass, or for a compound – the sum of those atomic masses, expressed in grams. So one mole of iron is about 55.85 grams. One mole of water, H2O, is about 18.02 grams. Both of these quantities have the same amount of fundamental particles, 6.022 times 10 to the 23rd particles. You have to know electron configurations. For neon, that’s 1s2, 2s2, 2p6. Atoms are usually most stable when they have 8 electrons (or an octet) in their outermost, or valence shell. We all know that opposite charges attract. Well, Coulomb’s Law says that the greater the magnitude of charge, the stronger that attractive force will be. And the closer together those charged particles are, the stronger the attractive force will be. This is why valence electrons are held less tightly to the nucleus; they’re literally farther away. And this is how photoelectron spectroscopy works. Each peak in this diagram represents a sublevel, and the taller the peak, the more electrons it has. The sublevels on the far left have electrons that require more energy to be stripped away, while the sublevels on the far right need way less energy. So this diagram represents calcium. The Periodic Table has several patterns. Like atomic radius: the bigger atoms are at the bottom of the table, and also toward the left. Atoms at the top and right of the table have the highest first ionization energy. When an atom gains electrons, it becomes a negatively-charged anion, and it gets larger. When an atom loses electrons, it becomes a positively-charged cation, and it gets smaller. Unit 2 covers chemical compounds. Metals and nonmetals are held together by ionic bonds, which are electrostatic forces: positives and negatives attracting. Nonmetals are held together by covalent bonds, where atoms share electrons. Covalent bonds can be polar, where they share electrons unequally, or they can be nonpolar, where they’re sharing electrons fairly equally. Covalent bonds form molecules, which can usually be individual units. However, ionic compounds exist in a three-dimensional lattice, where cations alternate with anions. Metals, and metal alloys, exhibit metallic bonding, where electrons can move freely, consisting of positive metal ions surrounded by a sea of electrons. Lewis electron-dot diagrams help us visualize the shapes of molecules. Each atom has a certain number of valence electrons, and we have to draw every one of those in the diagram. There are exceptions, but we try to arrange these so every atom has eight valence electrons, creating double or even triple bonds if we have to. These shapes have names. This molecule has a tetrahedral geometry and a bond angle of 109.5 degrees. This geometry would be called linear, with a bond angle of 180 degrees, and this one is trigonal planar, 120 degrees. Unit 3 covers Intermolecular Forces. Dispersion forces are usually weak interactions, but they get stronger as molecules get larger and have more electrons. The more electrons a molecule has, the more polarizable it is. Dispersion forces are the principal intermolecular force between nonpolar molecules. Polar molecules also have dipole-dipole forces. This is where the positive pole of one molecule attracts the negative pole of a neighboring molecule. These are usually stronger than dispersion forces. Some polar molecules, like water, have an especially strong force, called hydrogen bonding, found between molecules with an oxygen-hydrogen bond, a nitrogen-hydrogen bond, or a fluorine-hydrogen bond. Solids are usually crystalline, with molecules packed tightly together, having a fixed shape and volume. Liquids usually have a little more space between the molecules, so they can slip and slide around each other, which is why a liquid flows. Gases have molecules that are basically independent of each other. So gases can expand or be compressed more easily. The Ideal Gas Law, PV=nRT, shows relationships among pressure, volume, number of moles of gas, and the temperature. However, the Ideal Gas Law is just that, IDEAL. Gases in the real world aren’t ideal, but they sometimes approximate ideal conditions when we’re working with especially small molecules, very weak attractions, or even at a high temperature or low pressure. At higher temperatures, molecules have a higher average kinetic energy. This Maxwell-Boltzmann distribution shows that at higher temperatures, more molecules are moving faster, and at lower temperatures, most molecules are moving more slowly. We measure solution concentration with molarity, which is equal to moles of solute, divided by liters of solution. The rule “like dissolves like” helps us decide if a solute will dissolve in a particular solvent. Polar solutes dissolve in polar solvents (like water). Nonpolar solutes dissolve in nonpolar solvents. Light interacts with matter. The wavelength of light, lambda, multiplied by it frequency, nu, equals the speed of light. And if you multiply that frequency by Planck’s constant, you find the energy of a single photon of that light. We often use spectrophotometry to analyze the concentration of a solution. The higher the absorbance of the solution by the instrument, the higher the solution’s concentration. We use this data to build a graph to determine the concentration of an unknown. Unit 4 covers chemical reactions. When we write equations for reactions in solution, we usually omit so-called ‘spectator’ ions that don’t actually participate in the reaction, like sodium or potassium cations, or nitrate anions. The result is a net ionic equation. When you write an equation, check that the number of atoms of an element is the same on the reactant side as it is on the product side. We use coefficients to do this, and it’s called balancing the equation. A balanced equation is basically a recipe for how the reaction works. The coefficients in an equation form a MOLE RATIO. When you make a stoichiometric calculation, your first step should be to convert the quantity to moles, if it’s not already in moles. Then, use the coefficients of the balanced equation to form a mole ratio to determine the moles of the substance you’re converting to. Finally, if you’re asked to find a quantity in some other unit, like grams, you’ll convert to that final unit. AP Chem focuses on three types of reactions. In precipitation reactions, two solutions are mixed, and a solid precipitate is formed. In oxidation-reduction reactions, one element loses electrons, a process we call oxidation. At the same time, another element gains those electrons, which is called reduction. In a third type, acid-base reactions, an acid reacts with a base to form a conjugate acid and a conjugate base. Remember, an acid is a proton donor, while a base is the proton acceptor. Since a proton is just an H+ ion, an acid always has exactly one more H+ than its conjugate base. Unit 5 covers Kinetics. Balanced equations help us describe relative rates. For example, in this reaction, since the coefficient of NH3 is twice that of N2, the rate of appearance of ammonia will be twice as fast as the rate of disappearance of nitrogen. Each reaction has its own rate law, and these are determined experimentally. The rate law is always written as: Rate equals k, which is the rate constant, times the concentration of the first reactant, raised to the power of its order, times the next reactant, raised to its order, and so on. If we double the concentration, and the rate quadruples, that’s a second-order relationship. On the other hand, if we double a reactant’s concentration, and the rate doubles, that’s a first-order relationship. If we double the concentration, and the rate doesn’t change at all, that’s zero-order. Once we know the order for a reactant, we can use an integrated rate law equation to calculate the amount that will be left over after a certain amount of time. For each integrated rate law, we have the rate constant k, time elapsed t, initial concentration A subzero, and elapsed concentration A sub T. If we know any three values, we can calculate the fourth. Most reactions take place in multiple steps, forming a reaction mechanism. One step is slower than the others, and that slow step determines the rate of the whole reaction. If we can determine the rate law of the slow step, we’ll know the rate law for the whole reaction. For molecules to react, they have to collide with enough energy and in the right orientation. When they do, a high-energy transition state can form, which is at the peak on this graph. Then, the products can be formed. The energy required to start the reaction is called the activation energy. This reaction has a net loss of heat to the surroundings, which makes it an exothermic reaction. To speed up a reaction, we can raise the temperature, use a smaller particle size, or raise the concentration of the reactants. We can also add a catalyst, which actually provides a completely different reaction mechanism and lowers the activation energy required for the reaction to take place. Unit 6 covers Thermodynamics. Endothermic reactions absorb heat from the surroundings, while exothermic reactions release heat into the surroundings. We calculate heat transfer with the equation Q equals M C delta T. Q represents heat in Joules, M is the mass in grams, C is the specific heat capacity of the material, and delta T is the change in temperature. The heat change for a reaction is called change in enthalpy, or delta H, measured in kilojoules per mole. We can estimate it using bond enthalpies, adding all the enthalpies for the bonds broken in a reaction, minus the total enthalpies for the bonds formed. Or we can use enthalpy of formation, where the sum of the enthalpies of formation of the products, minus the total of the enthalpies of formation of the reactants, equals delta H. Or we can use Hess’s Law, which says that if several individual reactions add up to give a new reaction, the delta H values of those individual reactions can be added to give us the delta H of the new reaction. Unit 7 discusses equilibrium. When a reaction reaches equilibrium, it doesn’t stop. Instead, the forward reaction has the same rate as the reverse reaction, so the overall concentrations of the substances stop changing. The expression for the reaction quotient is abbreviated Q: the concentrations of the products over the concentrations of the reactants, raised to the power of the coefficients, omitting any liquids or solids. If the reaction is at equilibrium, then the reaction quotient Q is equal to K, the equilibrium constant. If we calculate Q and it’s not equal to K, the reaction will proceed until it attains equilibrium. When the equilibrium constant K is very large, much greater than one, we’ll have lots of product and very little reactant. However, when K is very small, we’ll have lots of reactant, and very little product. We’re often given initial concentrations or pressures for a reaction and asked to calculate the final concentrations or pressures. The best way to do this is to organize the data into a chart; I call it an ICE box – initial, change, equilibrium. Plug in the numbers you know and solve for the numbers you don’t know. Using algebra, you can solve these problems like a puzzle. We can apply equilibrium in many ways. Le Chatelier’s Principle says that if a reaction is at equilibrium, and we add a component, the reaction will shift toward the other side of the reaction. And if we take away a product, the reaction shifts so that product is replenished, at the expense of the reactants. We can shift the direction of the reaction, but the ONLY WAY to change the actual value of the equilibrium constant K is to change the temperature. Unit 8 is Acids and Bases. There are a few essential equations in acid-base chemistry. pH equals negative log of the hydrogen ion concentration. pOH equals negative log of the hydroxide ion concentration. And at 25 degrees Celsius, pH plus pOH of any solution equals 14. And at 25 degrees Celsius, the hydrogen ion concentration times the hydroxide ion concentration equals 1 times ten to the negative 14th power, a constant we call Kw. Strong acids and strong bases ionize completely, so for example, in 0.50 molar nitric acid, the concentration of the hydrogen ions is 0.50 molar, so to find the pH you just take negative log of 0.50. For weak acids and bases, less than 100 percent of the molecules dissociate, so these are actually equilibrium problems. Set up an ICE box, where the initial concentration of the acid or base is written in the appropriate spot, the initial concentrations of the products are zero, and you can solve the ICE box by plugging into the equilibrium constant expression, using Ka for a weak acid or Kb for a weak base. In acid-base titrations, we’re trying to find the concentration of an acid or a base. We add base to an acid until we reach the endpoint – the moment where an indicator changes color, signaling the reaction is complete. When you plot pH versus volume of base added, you get a titration curve that looks like this. The inflection point represents the equivalence point. And at pH 8.7, this means the titration was between a strong base and a weak acid. Halfway between the starting point and the equivalence point, the pH is 4.8, which tells us the Ka of the acid is 10 to the negative 4.8 power. Buffers are mixtures of a weak acid and its conjugate base that resist changes in pH. We can calculate a buffer’s pH with the Henderson-Hasselbalch equation. Unit 9 covers applications of thermodynamics. Entropy, abbreviated S, is the disorder present in matter. Solids have the least entropy, since they are highly ordered. Pure liquids have more entropy, solutions have even more, and gases have the most entropy. Systems at higher temperatures have more entropy than those at lower temperatures. We can usually predict if entropy is increasing or decreasing by looking at a reaction. For example, when water melts, solid is transitioning to liquid; entropy is increasing, so delta S is positive. Gibbs Free Energy, delta G, is a measure of the thermodynamic favorability of a process. Delta G equals delta H minus temperature in Kelvins times delta S. If a reaction is thermodynamically favored at a certain temperature, delta G is negative. If it’s NOT favored, delta G is positive. Thermodynamic favorability is related to the equilibrium constant by the equation Delta G equals negative R times T times natural log of K. In electrochemistry, every galvanic cell has two half-reactions, one is an oxidation and one is a reduction. The side where reduction takes place is called the CATHODE, and the side where oxidation takes place is called the ANODE. Electrons move through the wire from the anode to the cathode. The salt bridge allows ions to flow freely through the cell. At the salt bridge, anions flow toward the anode, and cations flow toward the cathode. Every galvanic cell has a voltage that we can calculate, using a list of standard reduction potentials. As a galvanic cell runs, the voltage slowly drops until it reaches zero volts, when we say the cell is at equilibrium. We assume galvanic cells are at standard conditions, which is 25 degrees Celsius, and a concentration of one molar for solutions. For any other conditions, we use the Nernst Equation to calculate the actual voltage. Galvanic cells are thermodynamically favored, and its Delta G is equal to negative n times Faraday’s Constant times E, where n is the number of electrons transferred and E is the overall voltage of the cell. An external electricity source passed through a solution will power an electrolysis process. To determine the amount of an element plated out, we use this equation: Electrical current in amps, I, equals electrical charge in Coulombs, Q, divided by time in seconds. Once you know the number of coulombs, you can calculate the amount of metal plated out. There it is: The major points of the entire AP Chemistry course in about 19 minutes. Don’t forget to Like, Subscribe, and watch my other review videos and entire AP Chemistry Course. If you’re taking AP Exams in May, check out ultimate review packet dot com. See you next time, and thanks for watching!