Teaching High School Science

AP Chemistry Review in 30 Minutes

• Kesha "Doc" Williams • Episode 5

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Get ready to conquer the AP Chemistry exam with a comprehensive breakdown of the entire course material that will sharpen your skills and elevate your understanding to the next level. From the building blocks of atomic structures to the balancing act of equilibrium, we've got you covered. We begin with the mole concept, electron configurations, and Coulomb's Law, giving you the tools to unravel periodic trends and ionization energy. Then, we tackle the nature of ionic and covalent bonds, their distinct properties, and the reasons behind their melting and boiling points.

As we progress, you'll learn to navigate through the labyrinth of chemical reactions, stoichiometry, and the puzzle of reaction rates with ease. We'll decipher the theories behind acids and bases, and you'll learn the significance of catalysts and how to calculate enthalpy changes. Electrochemistry might seem like uncharted territory, but after this episode, concepts like the Nernst equation and salt bridges will be as familiar as your periodic table. Join me on this journey to demystify the intricacies of AP Chemistry and stride confidently into your exam. Check the show notes for exclusive study materials and resources to tailor your revision strategy, because we're not just preparing you for success—we're setting the stage for you to excel.

Helpful Resources - Practice Problems, Study Guides, Videos

AP Chemistry Course at a Glance (Units and Topics)

Fivable: AP Chemistry Study Guides by Unit

Khan Academy: AP Chemistry

Varsity Tutors: AP Chemistry Practice Tests

Tyler DeWitt Videos

Melissa Maribel (Videos) Cramming for a Chem Exam Playlist

Bozeman Science (Videos) AP Chemistry Video Essentials


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Speaker 1:

Okay y'all, ap Chemistry exam is right around the corner and several of my students have been working with me for a while. We're at the point where we are focusing in on key topics, on identified areas, and just practicing, practicing, practicing the entire exam. However, I have a few students that just joined me and are stressing the entire course, and that sparked the idea for this podcast, where I'm just going to talk about all of the key concepts for each of the nine units in the entire AP Chem course. I'm also going to provide links to some of the study materials and sites that students have found extremely helpful in the show notes. That way, you have one spot to get the topics and resources to help you organize and personalize your study plan. So this is going to be a little longer than my normal podcast. Make sure you check out the show notes so that you can look at the time stamps, and that's going to allow you to skip to the units that you need. So let's get started. Welcome to Teaching High School Science. I'm your host, doc, a former biochemist turned high school science teacher and private tutor. Whether you're homeschooling your team through high school science or teaching online, join me as I share tips and strategies I've learned over the years for at-home and online labs and activities, breaking down complex concepts and structuring learning in a way that makes sense. Now let's dive into today's topics.

Speaker 1:

Unit one is atomic structure and properties. The mole is a unit that we use to count for large numbers of atoms and molecules. An element is going to be equal to its atomic mass. From the periodic table, for example, one mole of carbon is going to be equal to 12.011 grams, and we say that as carbon is equal to 12.011 grams per mole. Now a molecule is going to be equal to the sum of the atomic masses for each element in a molecule. For example, a molecule of water contains one oxygen and two hydrogens. One oxygen is approximately 16 grams per mole, where each hydrogen is approximately one gram per mole. If we add 16 for the one oxygen and two, which is one each for both of the hydrogens, then we get a total of 18 grams per one mole for water.

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Electron configurations you definitely want to be able to write them and be able to identify elements from an electron configuration that is given to you, and you want to be able to look at that and tell if that element is stable or unstable, which is indicated by the number of valence electrons and whether or not there is a full octet. You need to relate Coulomb's law to the electrostatic attraction between ions. Basically, the greater the magnitude of the charges of the ions, the greater the attraction. But also the greater the distance between the charges, the weaker the attraction by squared. So electrons in the outermost shell are held less tightly to the nucleus because you have a greater distance between them and the nucleus. This concept is also applied to electron spectroscopy. The taller the peak, the more energy it is required to remove the electron. So peaks to the right of the graph require less energy to remove the electrons than peaks to the left, and each set of peaks represent a different energy level. So being able to identify the atom by the graph is a must. And you also want to be able to explain the energy required to remove electrons based on Coulomb's law.

Speaker 1:

Another concept you want to be able to apply Coulomb's law is explaining the trends of the periodic table. These trends include ionization, energy electronegativity, atomic radius and ionicic radius. As you move across the periodic table in the same period you are gaining electrons and you are gaining protons, but the distance is not changing. So therefore, the atom becomes smaller because you have the stronger electrostatic force. However, as you move down a group, you are increasing energy levels, which increases that distance. So the atoms are becoming larger. The radius is becoming larger. However, with ions, as you are increasing the number of electrons to make it a negatively charged ion, then that atom is becoming larger and as you are decreasing the number of electrons, it is becoming smaller when it's forming a positively charged ion. Now you will apply the same logic when you're thinking about the first ionization energy or the trend in the first ionization energy. So, as you move across the same period, you're increasing the number of protons and increasing the number of electrons, so that electrostatic force is getting stronger because you're not increasing the distance between the nucleus and the electrons. This means that it's going to require more energy to remove that first electron. Now, as you move down a group, you are increasing the distance between the outermost electrons and the nucleus, so that means that it will require less energy to remove that first electron. This same logic will also be applied when you are explaining or discussing the trends in electronegativity as you move across a period and down a group.

Speaker 1:

In unit two we're talking about molecular and ionic compound structure and properties. Ionic bonds are formed between negatively charged non-metals, which are anions, and positively charged metals, which are cations. They are held together by an electrostatic force, which is that attraction between oppositely charged particles. Now covalent bonds are formed between non-metals where atoms share electrons but the electrons may not be equally shared. Covalent bonds are polar when the electrons are not shared equally and non-polar when they are shared equally. And remember, in most cases symmetric molecules tend to be non-polar. Elements in an ionic bond exist in 3D lattice structures. This is where you have an alternating pattern between an anion and a cation and they form extremely strong structures. So it requires a lot of energy to break that apart, which means they have a higher boiling and melting point. On the other hand, covalent bonds form basically independent molecules. They tend to have very low melting points and low boiling points because they do not require quite as much energy to break those bonds. Ionic bonds conduct electricity where covalent bonds generally do not. Now metallic bonds are formed between metals and alloys. This is where electrons move freely around the atoms in the bond and they are high conductors of electricity and thermal energy because of the sea of electrons.

Speaker 1:

We use Lewis dot structures to visualize the shapes of molecules using the valence electrons. The goal is to provide the most stable orientation of the atoms around a central atom, which will provide each atom with a full octet. So you want to be able to calculate formal charge and understand resonance and hybridization. Molecules form different shapes based on bonded electrons and long pair electrons. You want to be able to identify these shapes, the bond angles, and explain able to identify these shapes, the bond angles, and explain why they form these shapes.

Speaker 1:

Unit three is about intermolecular forces and solids, liquids and gases. So for this unit, you want to keep in mind that the force that holds atoms together within a compound are intramolecular forces, and these are the forces we talked about in unit two, with ionic bonds, covalent bonds and metallic bonds. Now, intramolecular forces are those that hold different compounds or molecules together. Dispersion forces is the weakest intramolecular forces and they exist among all molecules. Now, they're the weakest with the exception of molecules that are larger, and this is because they have more electrons and when we increase the number of electrons, we are making that molecule more polarizable, which leads to stronger dispersion forces between molecules. And anytime you're asked about forces that exist between molecules, always list dispersion forces, even if there is another stronger force that also exists between the molecules. Next is dipole-dipole forces, which exist between polar molecules, and this occurs when you have the positive end of one molecule attracting the negative end of another molecule. These are usually stronger than dispersion forces, but, but again, both dispersion and dipole forces are found in polar molecules, so be sure to list both. Hydrogen bonds are a type of dipole-dipole bond that exists only between hydrogen and oxygen, hydrogen and nitrogen and hydrogen and fluorine. These are usually the strongest of the three. So this means you want to check to see if there is a hydrogen in the bond and then if it's bonded to a fluorine, an oxygen or a nitrogen. The mnemonic FON fluorine, oxygen and nitrogen is a way to remember the three elements that form hydrogen bonds.

Speaker 1:

Now for solids, liquids and gases. Particles and solids are packed closely together and they vibrate in place. You have two categories of solids, which are crystalline and amorphous. Where crystalline has a definite internal structure and amorphous has an irregular internal structure, particles and liquids have more energy than the particles and solids, which means that they move a bit more, creating more space between them, allowing them to slip past one another. So, where solids have a definite shape and a definite volume, liquids would take the shape of their container, meaning they do not have a definite shape, but they do have a definite volume. The particles and gases have the most energy, the most kinetic energy, and they move independently. So they not only allow for the most compression and expansion, but they also do not have a definite shape and they do not have a definite volume. They will take the volume of any container.

Speaker 1:

Now this brings me to the ideal gas law, which is PV equals NRT. This describes how changes in pressure, volume, temperature and the number of moles of a gas will affect the behavior of ideal gases. Now, I say ideal because we're looking at behavior gases in a perfect situation, an ideal situation where there is absolutely no conditions that will allow attraction between molecules. In the real world we have very small molecules, particularly helium, hydrogen and neon, and then we have gases at low pressures and high temperatures that can cause them to behave a little differently, and the Maxwell-Boltzmann distribution curve for gases shows the velocity of gases at higher temperatures. So basically, at very high temperatures, gas particles have higher kinetic energy, which causes them to have a higher velocity, and the opposite is true at lower temperatures.

Speaker 1:

Now, moving on to solubility, like dissolves like. This means that polar molecules dissolve in polar solvents and nonpolar molecules dissolve in nonpolar solvents and we can describe the concentration of solutions using molarity, which is the number of moles of solute divided by the liters of a solution. Another way we work with concentrations of solution is with spectrum photometry, which analyzes the concentration of an unknown solution. So this uses how light interacts with matter and there's two basic equations that we use with this part. We have the speed of light equation, which is the speed of light equals lambda, the wavelength times, nu, the frequency, and once we have the frequency of that light from that equation, we can use that to find the energy of a single photon of light using Planck's equation, which is energy equals Planck's constant times. That frequency, the higher the absorbance of energy, the higher the concentration. So if we were to graph this information, we can determine the concentration of an unknown solution using a line of best fit, and that's the key. When you're asked to graph this type of information, do not connect the dots. Always use a line of best fit, because that is what's going to give you the proper reading and you will probably be asked to discuss why you have some outliers.

Speaker 1:

Unit four is all about chemical reactions. The first thing to make sure is that all of your chemical equations are balanced and remember, when you're balancing them, only change the coefficients and never ever the subscripts. Ionic equation shows the separate ions of dissolved solids in a solution. These are written by breaking down all of the aqueous components into their ions so that we can remove the spectator ions and leave the net ionic equation. Balanced chemical equations are essential for performance stoichiometry, where we change from one substance to another using mole ratio. All units of the given must be converted into moles, because we cannot use mole ratio if the units are not in moles. So you may have to use mole mass, avogadro's number or some other conversion factor in order to convert the units of the given into moles. Then you use mole ratio to convert from substance A to substance B. Make sure you're checking the units of substance B, because you may have to go back to mole ratio at Avogadro's number to change the units of the substance you were looking for into the units that the question is asking you to find.

Speaker 1:

There are three types of chemical reactions that we focus on in AP chemistry. One is redox reactions. This is where one element loses electrons while another one gains electrons. We write the oxidation and reduction components in half reactions, then we balance the electrons to gain the net reaction. Another chemical reaction type are precipitation reactions. This is where two solutions are mixed together. The two solutions are your reactants and you will form a precipitate that will be your product. The third chemical reaction are acid-based reactions, where acids and bases react to form salt and water. There are three types of acids and bases. One are aranias acids and bases where, when dissolved in water, the acid will produce a proton and the base will produce a hydroxide ion. Then we have Bronsted-Lowry acids and bases, where the acid is a proton donor and forms a conjugate base and the base is a proton acceptor and forms a conjugate acid. And with Lewis acids and bases, the acid are electron acceptors and the bases are electron donors. Unit five is kinetics.

Speaker 1:

The balanced chemical equation, which you know to make sure is balanced, helps us describe relative rates of reaction. For example, in a chemical reaction of one mole of nitrogen plus three moles of hydrogen will yield two moles of ammonia. Then the rate at which ammonia is produced is twice as fast as the rate at which nitrogen is consumed, and we can tell that by looking at the chemical equation, because we have one mole of nitrogen to every two moles of ammonia. Each reaction will have its own rate law which we determine experimentally. That's key. Remember that rate laws are determined experimentally and the equation we use to determine the rate of reaction is the reaction rate is equal to the rate constant times the concentrations of each reactant multiplied by each other, and they're going to be raised to the power of this order. This means that you're going to have to look at your experimental data, then determine the rate order and then use that to calculate your rate constant.

Speaker 1:

Tips for determining rate order is that it is a zero order if you see an increase in the concentration of reactants but that does not cause a change in the rate of the reaction. It is first order when the multiple by which the concentration of the reactants increase is the same as the increase in the rate of the reaction. So if we see that the concentration of the reactants increase by two and the rate of the reaction increased by two, then that reaction is a first order. Now second order is when the multiple by which the concentration of the reactants increase is the square of the multiple of the increase in the rate. For example, if the concentration of the reactants tripled and then you look at the rate and that increased by nine, then that will be a second order, because three squared equals nine. Knowing the order for the reaction allows us to use the integrated rate laws. Since most reactions occur in multiple steps, it can only go as fast as the slowest step. So determine the rate law for the slowest steps gives us the rate law for the entire reaction.

Speaker 1:

And in order for reactions to occur, reactants must collide in the correct orientation and with sufficient energy to react, which is the collision theory. The energy that is required to start the reaction is called the activation energy, and that energy can be lowered using a catalyst. We can speed up reactions by increasing the temperature, increasing the concentration of the reactants or decreasing the size of the particles. We're moving on to unit six, which is thermodynamics. Now some reactions will lose energy to the environment. These are called exothermic reactions, and other reactions will absorb energy from the surroundings. These are called endothermic reactions. The amount of heat absorbed or transferred can be calculated using the equation Q equals MC delta T, where Q is the energy transfer or absorbed. M is the mass, c is the specific heat capacity of the substance and T is the change in temperature, which is final temperature minus the initial temperature.

Speaker 1:

Enthalpy, or delta H, describes the heat change for a reaction, and we can determine this in a few ways. One is through bond enthalpies, which is the sum of the energy of the bonds broken minus the sum of the energies of the bonds form. The next way is through the enthalpies of formation, and this is the sum of the enthalpies of the formation of the products minus the sum of the enthalpies of the formation of the products minus the sum of the enthalpies of formation of the reactants. And then the last way we can determine enthalpy is with Hess's law, which states that, no matter the number of steps of a reaction, the total change in enthalpy is the sum of all of its steps, even if the reaction took place in a single step. Unit seven is equilibrium.

Speaker 1:

A reaction in equilibrium does not mean that the reaction stops. It simply means that the forward reaction is at the same rate as the reverse, so products are being formed at the same rate as they are breaking down to form reactants. And we can determine if a reaction is in equilibrium using the reaction quotient Q, and that equation is the concentration of the products multiplied together, divided by the concentration of the reactants. Now both the concentration of the products and the concentration of the reactants are raised to the power of the coefficients. Also, you do not include the concentrations of solids or liquids in this equation. This equation is also used when determining partial pressures.

Speaker 1:

When we calculate Q and Q equals the equilibrium constant K, then a reaction is at equilibrium. If K is greater than one, that means that the reaction is in the forward direction and producing more product. This will cause the reaction to shift and try to make more reactants to obtain equilibrium. Now, if K is small, meaning it's less than one, then that means that the reaction is in the reverse and it's forming more reactants, and this will cause the reaction to shift in the opposite direction to form more products in order to achieve equilibrium. And the shifting of the reaction to obtain equilibrium. All of the different factors that affect this can be described using Le Chatelier's principle. For example, if we increase the concentration of products, as we just said, it causes the reaction to shift in the reverse, to form more reactants, and vice versa. If we change the temperature, or the volume or the pressure, the reaction will shift to obtain equilibrium. Now, one thing to keep in mind with this is that the only way to change your actual value for your equilibrium constant K is to change the temperature.

Speaker 1:

Now some of the problems. For this, you will be given values and you'll be asked to determine the final concentrations when provided initial concentrations. To accomplish this, we need to set up ICE tables. These type of calculations can be done for pressures. We can also apply these concepts for solutions. Now, when we're talking about solutions, we're talking about the desolation process, which is KSP, the rate at which a salt is broken down. You need to be able to relate solubility rules to KSP. For example, a KSP greater than one tells us that the salt is soluble. But if the solution already contains ions from the salt, then this is called the common ion effect and we can use Le Chatelier's principle to describe what happened to this reaction. So if you already have ions in the solution before you add the solid, then you already have a high concentration of products before you even add more of the salt. So that means that the reaction will shift toward the reactants until equilibrium is reached.

Speaker 1:

Unit eight covers acids and bases. We already touched on Bronsted-Lowry acids and bases earlier, where you have your acid and conjugate base or base and conjugate acid pairs. In addition to this, before we even get even deeper, I want you to understand and be able to use three main equations. The first two deals with pH and pOH. For pH, we would take the negative log of the concentration of the hydronium ion. For pOH, to calculate the pOH of a solution, you would take the negative log of the hydroxide concentration. The third equation is taking the concentration of the hydronium ion multiplied by the concentration of the hydroxide ion and setting that equal to 1.0 times 10 to the negative 14. This equation allows you to solve for either the concentration of the hydroxide ion or the concentration of the hydronium ion.

Speaker 1:

Now let's talk about some properties of acids and bases. Strong acids and strong bases ionize completely when dissolved in water. That means that they will completely break down, but weak acids and weak bases do not. So if we're going to determine concentrations of these, then we have to treat them as equilibrium problems and use ice tables For our constants. You should be using Ka for weak acid and Kb for weak bases. Another thing you need to be very familiar with are titrations acid-based titrations. You always add the base to the acid, so the acid will be in your flask, you will have an indicator in the acid and base will be in your pipette. You add the base to the very moment the color changes to let you know that the reaction is complete. If you go past that, if there's too much of a change, then you have to start over, because we're going to graph this information and determine the concentration of the acid. Buffers are usually mixtures of a weak conjugate acid-base pair where the conjugate base reacts with an added acid and the conjugate acid reacts with the added base. The significance of these solutions is that they resist changes in pH. So the Henderson-Hasselbalch equation is what we use to calculate the pH of buffers.

Speaker 1:

Last unit, unit nine, is the application of thermodynamics. Entropy, or S, is the disorder or chaos of a system. In nature, our natural world, things spontaneously move towards high entropy, which means higher chaos and higher disorder. So that means moving from a solid to a gas. We are increasing entropy because gas particles exhibit a higher level of disorder and chaos than solid particles. Acreous solutions fall between liquids and gases, which means that aqueous solutions have a higher disorder or entropy than liquids, but a lower entropy or disorder than gases or entropy than liquids, but a lower entropy or disorder than gases. This leads to the role that temperature plays in entropy. As we increase the temperature of solids and causes that to change into liquid, we are increasing that kinetic energy, increasing motion and increasing entropy. A positive delta S indicates an increase in entropy and delta S can be calculated by the sum of entropy of the products minus the sum of the entropy of the reactants.

Speaker 1:

Gibbs free energy tells us how much energy is free to do work and we can use this to determine if a reaction is thermodynamically favorable, which means if it will spontaneously occur. A negative delta G is a thermodynamically favored reaction, meaning that it will spontaneously occur since energy is being released into the environment, and a positive delta G is not thermodynamically favored, which means it will not spontaneously occur. So sometimes we have to consider enthalpy, which is H, and entropy, which is is S, when determining spontaneity and there are different conditions where you have a positive S, a negative G, a positive G, a negative S and so forth. So you want to take the time to really look at that and determine how delta S and delta H impact Gibbs, free energy and spontaneity.

Speaker 1:

Now let's talk about galvanic or voltaic cells and electrochemistry. We use galvanic cells to convert the chemical energy from redox reactions into electrical energy. You want to be familiar with how to set up this type of cell, how to complete a drawing or a diagram. If we're using two beakers, then the two beakers are connected with a salt bridge and they're also connected with two metals or electrodes that are connected with a wire and a voltmeter. Both of the beakers contains an electrolyte.

Speaker 1:

Oxidation takes place on the anode side, which is by the negative metal or the metal with the lower energy, and reduction takes place on the cathode side, which is the positive metal or the metal with the higher energy. Electrons move from the anode through the wire into the cathode, and we can measure this flow with the voltmeter that's connected to the wire. Now the salt bridge allows the charges to move freely between the two beakers, where each particle is attracted to the oppositely charged side. The positive particles are attracted to the anode and the negative particles are attracted to the cathode. Now, as these charges continue to move and it gets closer to equilibrium, the voltage begins to drop. So the further away this reaction is from equilibrium, the higher the voltage, and we can calculate the impact of various conditions that are not standard meaning not at 25 degrees Celsius, not at one molar, not at one atmosphere using Nernst equation. And we can determine the amount of charge flow or the number of electrons transferred using the Nernst equation and we can determine the amount of charge flow or the number of electrons transferred using the equation. I equals Q over T, which is current, equals coulombs divided by time.

Speaker 1:

Now that is AP Chem the entire course in a nutshell, really highlighting those major ideas and concepts. I hope this helped to provide an outline for a study guide that will help you with each unit. Follow me on social media or visit my blog for study tips and just let me know if you think this helped. Let me know if you have any questions, ideas or other experiences that you'd like to share. Head on over to my podcast page, which you can access by visiting my website at thesciencementorcom. Then select podcast from the menu and subscribe now to the Teaching High School Science podcast for your regular dose of motivation and just-in-time science ideas, and together let's make high school science a journey of exploration and achievement. Until next time, remember curiosity leads to endless possibilities.