2026 AP Chemistry Exam Review (EVERYTHING YOU NEED TO KNOW!!)

Added: 2026-05-10

[00:00:01]We're about to review the entire AP Chemistry curriculum in this topic by topic guide made from the course and exam description provided by College Board. We'll cover atomic structure and properties, compound structure and properties, properties of substances and mixtures, chemical reactions, kinetics, thermmochemistry, equilibrium, acids and bases, and thermodynamics and electrochemistry.

[00:00:24]Let's get into it with unit one, atomic structure and properties.

[00:00:28]The mole is your conversion currency. 1 mole equals 6.022 * 10 23rd particles.

[00:00:35]The flowchart on the right shows how we can convert g to moles dividing by mass and moles to particles multiplying by Avagadro's number. Moles is always the middle. Percent composition is mass of element over total mass time 100.

[00:00:49]Empirical formula is the simplest whole number ratio. Molecular formula is a whole number multiple of the empirical.

[00:00:56]Divide actual molar mass by empirical mass and then multiply that by the empirical.

[00:01:02]On a mass spectrum, the x-axis is mass to charge ratio and the yaxis is relative abundance. Each peak is an isotope. Taller peak means more abundant.

[00:01:13]Average atomic mass is the sum of fractional abundance time isotopic mass.

[00:01:18]That weighted average is what you see on the periodic table, not the mass of the most common isotope.

[00:01:26]Kulum's law states that attraction grows with nuclear charge and drops with distance. This idea drives every periodic trend. Subshells S, P, D, and F hold 2, 6, 10, and 14 electrons respectively. There are three filling rules. Offbell, lowest energy first.

[00:01:45]Poly max two electrons per orbital with opposite spins. Hans one electron per orbital before pairing.

[00:01:53]Also remember that when forming ions, electrons leave from the highest principal quantum number first, meaning that 4s goes before 3d for transition metals.

[00:02:05]PES measures the energy needed to eject electrons from each subshell. Peak position on the right is binding energy with core electrons on the left at high binding energy and veence on the right at low. Peak height is the number of electrons in that subshell. Read off the configuration directly. The example here is sodium 1 s2 2 s2 2 p6 3s1.

[00:02:32]All five trends on the right come from one variable effective nuclear charge.

[00:02:37]It increases left to right across a period. Atomic radius shrinks left to right and grows top to bottom. Cations are smaller and annions are larger.

[00:02:47]Ionization energy and electro negativity both increase left to right and bottom to top with florine as the highest. Two exceptions to remember are that oxygen has lower ionization energy than nitrogen because of paired 2p repulsion and aluminum is lower than magnesium because 3p sits above 3s.

[00:03:09]Veence electrons live in the outermost principal energy level. For main group elements, group number gives the count.

[00:03:16]Ionic compounds form when metals hand electrons to non-metals. Lattice energy follows Kulum's law. So bigger charges and smaller radi mean a stronger lattice and higher melting point.

[00:03:30]Before we get into unit two, if you like what you're seeing so far, remember that we offer free AP study guides to go along with all our YouTube videos on our website, prep worksucation.org.

[00:03:40]Also, if you're watching this video, chances are you're applying for college soon. We offer college consulting from T10 admits and free college admissions tools. Our comprehensive and proprietary chancing calculator is a favorite.

[00:03:53]links in the description for all of those resources. Now for unit 2, compound structure and properties.

[00:04:00]There are three bond types. Ionic bonds involve electron transfer and have a big electro negativity gap. Covealent bonds share an electron and of similar electro negativity and metallic bonds are like a sea of deoized electrons around positive ports.

[00:04:16]On the potential energy curve to the right, the bottom of the well is the bond length and the depth of the well is the bond disassociation energy. When atoms get close, potential energy rises because of the nuclei repellent.

[00:04:29]Polarity roughly follows electro negativity difference. Above 1.7 is ionic, 0.4 to 1.7 is polar covealent and below 0.4 is non-polar coalent.

[00:04:44]Ionic solids form a regular lattice of alternating cations and annions packed to maximize attraction. Lattice energy U is proportional to Q Q Q1 * Q2 / R.

[00:04:58]Bigger charges or smaller ions means stronger lattice and higher melting point. Metals are positive cores and a sea of veence electrons which explains conductivity, malleability and luster.

[00:05:11]Alloys are harder than pure metals and come in two types which we see to the right. Interstitial alloys like steel are when a small atom fits between larger ones and substitutional alloys like brass are when similarly sized atoms replace each other.

[00:05:27]To make a le structure count veence electrons connect atoms complete octets on terminal atoms and then form multiple bonds if you run short. Watch for exceptions to the octet rule shown below.

[00:05:40]Formal charge equals veence electrons minus lone pair electrons minus half the bonding electrons. If you have multiple options for a le structure, choose the one that minimizes formal charge.

[00:05:53]The resonance structures on the right share the same conductivity but different electron placements. The real molecule is an average of each resonance structure with equal bond orders. When you have different bond orders, a higher one means a shorter, stronger bond.

[00:06:11]Vesper rules tell us about the geometry of electron domains caused by the repelling of negative electrons. The chart to the right shows all the different shapes possible based on number of areas and lone pairs. Remember that lone pairs repel more than bonding pairs which compresses bond angles. Both bond polarity and geometry matter for molecular polarity. For example, carbon dioxide has polar bonds but is non-polar because it is linear. Water on the other hand is bent with polar bonds so is polar. Carbon tetrachloride is symmetric like carbon dioxide so is non-polar.

[00:06:48]Lastly, there are two types of bonds sigma and pi. Single bonds are sigma and the bonds after that in a double or triple bond each count as pi bonds. For example, a triple bond has one sigma and two pi bonds.

[00:07:03]Unit three is about properties of substances and mixtures. This is a heavily tested unit with 18 to 22% rating.

[00:07:11]IMFs control boiling point, melting point, viscosity, surface tension, and vapor pressure. There are four main ones to know for the AP. London dispersion forces are present in everything and stronger with more electrons and surface area. Dipole dipole attraction occurs between polar molecules and is stronger than London dispersion forces. Hydrogen bonding is a special intense dipole dipole when H is bonded to F, O, or N.

[00:07:39]That's why water, hydrogen fluoride, and ammonia have weirdly high boiling points. And the strongest IMF is ion dal attraction, responsible for salts dissolving in water.

[00:07:53]There are four main types of solids to know. Ionic solids are brittle with high melting points. Metallic solids are conductive and malleable. Molecular solids have low melting points and are held together by IMFs. And coalent network solids like diamond and silicon have extremely high melting points.

[00:08:13]The diagram on the right shows the bigger picture. In solids, particles are fixed and vibrate in place. In liquids, particles stay in close contact but are free to flow past each other. And in gases, particles have large spacing and move in constant random motion. Note that gas molar volumes are hundreds to thousands of times larger than solid or liquid.

[00:08:36]The ideal gas law is PV= NRT. Dalton's law tells us that the total pressure equals the sum of each gas's contribution, which is its mole fraction times the total pressure. Kinetic molecular theory makes four assumptions.

[00:08:52]Particle volume is negligible. There are no IMFs. Collisions are perfectly elastic and average kinetic energy is proportional to temperature in Kelvin.

[00:09:02]The dark panel on the right shows deviations from ideal behavior. At high pressure, particle volume becomes significant, so real pressure is greater than ideal. At low temperature, IMFs become significant, so real pressure is less than ideal. The Maxwell Boltzman distribution shows that at higher temperatures the curve becomes flatter, broader and shifted to higher speeds.

[00:09:27]Also, lighter gas moves faster than heavier gases at the same temperature.

[00:09:33]There are a few concentration units to know. Marity is moles of solute per liter of solution. Mole fraction is moles of one component over total moles.

[00:09:43]For dilutions, use M1 V1 equals M2 V2.

[00:09:48]The general rule is like dissolves like, meaning polar dissolves polar, and non-polar dissolves non-polar. The separation methods on the right work by different principles. Filtration separates solid from liquid, distillation separates by boiling point, chromatography separates by affinity for a stationary phase, and centrifugation separates by density.

[00:10:12]A quick rule to remember. Solid solubility goes up with temperature while gas solubility goes down. The solubility table on the right is worth memorizing. Always soluble includes salts of sodium, potassium, ammonium, nitrate, and acetate. Mostly soluble are halides except those with silver, lead, and mercury, and sulfates except those with berium, serrantium, and lead.

[00:10:41]The electromagnetic spectrum on the right shows the different regions from radio waves up to gamma rays. Energy equals plank's constant time frequency or plank's constant times the speed of light over wavelength. This means gamma rays have the highest energy and radio waves have the lowest.

[00:10:59]Each region on the spectrum matches a different type of transition. Microwaves cause molecular rotation. Infrared causes vibration. visible and UV light cause electronic transitions where electrons jump between orbitals.

[00:11:14]Lastly, the beer Lambert law states A= epsilon BC, meaning absorbance is linear with concentration.

[00:11:23]Now, let's look at unit 4 chemical reactions.

[00:11:28]There are five clues that tell you a chemical change has happened. color change, gas production, precipitate formation, temperature change, and light emission.

[00:11:39]Physical changes like phase changes and dissolving keep composition intact.

[00:11:45]Chemical reactions create new substances and obey conservation of mass and charge. So, always balance for both.

[00:11:51]There are three forms of equations to know. Molecular equations use full formulas. Complete ionic equations split strong electrolytes into ions. And net ionic equations cancel out the spectator ions.

[00:12:05]There are three steps to writing a net ionic equation. First, write the balanced molecular equation. Second, split all strong electrolytes, meaning strong acids, strong bases, and soluble salts into their ions. Third, cancel out anything identical on both sides, which are the spectator ions. Look at the example on the right. silver silver chloride precipitation. The nitrate and sodium ions are left out of the net ionic equation because they cancel out.

[00:12:34]The stochometry workflow on the right has four steps. Balance the equation.

[00:12:39]Convert your given quantity to moles.

[00:12:41]Apply the mole ratio from the balanced equation. Then convert to your desired units.

[00:12:47]The limiting reagent is the reactant that gives the smaller theoretical yield. Percent yield is actual yield over theoretical yield times 100.

[00:12:57]Titration involves adding a titrant of known concentration to your analyte until you reach the equivalence point.

[00:13:04]For 1:1 reactions, the equation is m acid * v acid equals m base * vb.

[00:13:13]There are six reaction types all shown on the table to the right. In synthesis, A plus B gives AB. In decomposition, AB breaks apart into A and B. In single displacement, A plus BC gives A C plus B driven by the activity series. In double displacement, AB plus CD gives A D plus BC, which happens when the precipitate, gas, or water forms.

[00:13:41]By the Bronstead Lowry definition, an acid donates a proton while a base accepts one. On the diagram to the right, water donates a proton to the bicarbonate ion, turning water into hydroxide and bicarbonate into carbonic acid. In this case, water is a Bronst Lowry acid and bicarbonate acts as a base. Water is amphoteric, meaning it can play either role depending on what it's reacting with.

[00:14:07]Now, let's talk about redux reactions starting with oxidation states. A free element has an oxidation state of zero.

[00:14:15]A monotomic ion equals its charge.

[00:14:18]Oxygen is usually -2 except in peroxides where it's1.

[00:14:23]Hydrogen is usually + one except in hydrides where it's negative 1. And the sum of all oxidation states equals the total charge.

[00:14:32]Remember the pneummonic oil rig.

[00:14:35]Oxidation is loss, reduction is gain of electrons.

[00:14:39]The example on the right shows iron 2+ losing an electron to become iron 3+ which which is oxidation.

[00:14:48]To balance a half reaction in acidic solution, balance non oxygen and non-hydrogen atoms first then balance oxygen with water, hydrogen with H+ and charge with electrons. Then equalize and add the half reactions. Per basic solution, convert each H+ to water by adding O minus to both sides.

[00:15:10]Kinetics is about how fast reactions occur and the mechanisms involved. Let's start off with reaction rates.

[00:15:18]Reaction rate is measured in marity per second and is always recorded as a positive number. You can use stochometric coefficients to relate the rate of one species to another. Rate is influenced by concentration, temperature, surface area, and catalysts.

[00:15:35]The rate law is rate= K * A to the M * B to the N, where M and N are rate orders.

[00:15:43]It's important to know that these orders are determined experimentally. They're not the coefficients from the balanced equation.

[00:15:51]To find orders from initial rate data, double the concentration of one reactant and watch what happens to the rate. If the rate doubles, that reactant is first order. If it quadruples, second order.

[00:16:03]And if it stays the same, zero order.

[00:16:08]The table on the right is one you should memorize. To identify reaction order, look for which plot is linear. If A versus T is linear, the reaction is zero order. If natural log of A versus T is linear, first order. And if 1 / A versus T is linear, second order. First order halfife is constant meaning it doesn't depend on starting concentration. That's why radioactive decay is first order.

[00:16:32]After n half- livives the concentration is the original * 1/2 to the^ of n.

[00:16:38]An elementary reaction is a single molecular scale step. For elementary steps only the rate law follows directly from stochometry.

[00:16:47]Un molecular reactions have a rate of K * A, while biomolecular reactions have a rate of K * A * B. The collision model says reactions need both enough energy and the right orientation to proceed.

[00:17:01]Higher temperature shifts the Maxwell Boltzman distribution so more moleculars clear the activation barrier. The reaction coordinate diagram on the right shows reactants on the left, the transition state at the peak, and products on the right.

[00:17:15]Delta H equals products minus reactants, which is just the difference between the start and end energy levels.

[00:17:23]A mechanism is a sequence of elementary steps that adds up to the overall balanced equation. Intermediates show up in one step and get consumed in another.

[00:17:32]So they don't appear in the overall equation and they don't show up in the rate law. The slowest step is called the rate determining step or RDS and it dictates the overall rate law.

[00:17:44]On the energy profile to the right, the bigger hump is the RDS. To validate any mechanism, every step must be elementary. The steps must sum to the overall reaction and the derived rate law must match experiment.

[00:18:00]A catalyst gives the reaction an alternative pathway with lower activation energy. Look at the energy diagram on the right where the catalyzed curve has a smaller hump. The lower the hump, the faster the reaction.

[00:18:13]A catalyst speeds the forward and reverse reactions equally, so it doesn't shift equilibrium. It also isn't consumed overall. There are two types of catalysts. Homogeneous catalysts are in the same phase as the reactants, while heterogeneous catalysts are in a different phase, like a solid catalyst working with gas reactants.

[00:18:34]Now for unit 6, thermmochemistry, we'll start off with heat change and energy diagrams.

[00:18:41]There are two key terms here. Exothermic means the system releases heat, so delta H is negative and products sit lower than reactants like diagram B on the right. Endothermic means the system absorbs heat. So delta H is positive and the products sit higher than reactants like diagram A. The sign of delta H reverses if you reverse the reaction.

[00:19:03]Also, adding a catalyst lowers activation energy in both directions, but does not change delta H. The start and end levels stay the same. You just get a shorter hump in between.

[00:19:16]Heat flows from hot to cold until temperatures equalize. Q lost equals Q gained because energy is conserved. The core equation is Q= MC delta T where M is mass, C is specific heat capacity and delta T is the temperature change in Kelvin or Celsius.

[00:19:36]The picture on the right is a coffee cup calerimeter which runs at constant pressure. So Q at constant P equals delta H giving you enthaly directly. A bomb calerimeter on the other hand runs at constant volume. So Q there equals delta E. The connection between them is delta H= delta E plus delta N * R * T where delta N is the change in moles of gas.

[00:20:04]Phase changes happen at constant temperatures. All the energy goes into breaking IMFs, not into raising T.

[00:20:11]That's why you see flat plateaus on the heating curve to the right. There are two enalpies to note. Enthalpy of fusion is for melting and enaly of vaporization is for boiling. Vaporization is always larger because you have to fully separate the molecules. For the flat regions on the curve, which are phase changes, use Q= N * delta H. And for the sloped regions, which are heating a single phase, use Q= MC delta T.

[00:20:41]Standard enthalpy of reaction equals the sum of delta H of formation for products minus the sum of delta H of formation for reactants each weighed by a stochometric coefficient. Standard conditions are 25° C and one atmosphere.

[00:20:56]Remember that delta H of formation for any element in the standard state is zero by definition. Enaly is a state function meaning only initial and final states matter which is exactly why Hessa's law works.

[00:21:10]Bond breaking is endothermic because energy goes in while bond forming is exothermic because energy comes out. As an approximation, delta H of reaction equals the sum of bond enalpies broken minus the sum of bond enalpies formed.

[00:21:26]This is just an approximation because bond enalpies are averages and it works best for gasphase reactions.

[00:21:34]If a reaction is the sum of multiple steps, then delta H of the overall reaction equals the sum of delta H of each step. There are two manipulations we'll use constantly. If you reverse a reaction, flip the sign of delta H. And if you multiply by a coefficient, multiply delta H by the same factor. The example shown combines three reactions to give nitrogen plus oxygen forming NO2. We can find the enalpy of this reaction by looking at the three enalpies and reactions given. If we add the first two reactions, we cancel out the ammonia. Then we're left with four water and hydrogen molecules that we don't want. So we can reverse the third reaction and multiply by four. So that when we combine this, we're left with only nitrogen, oxygen, and nitrogen dioxide. To get the enalpy, we add the first two enalpies and then subtract 4 * the third reaction's enalpy to get 82.2 2 K.

[00:22:31]Now on to unit 7 equilibrium.

[00:22:34]For a generic reaction A plus BB going to CC plus DD, KC equals product over reactants with each species raised to its coefficient. The equation on the right makes this visual. One important rule is that pure solids and pure liquids are left out of the expression.

[00:22:54]KP is the equivalent expression using partial pressures. KP= KC * RT to the delta N where delta N is the moles of gas products minus moles of gas reactants.

[00:23:06]Q has the same expression as K but uses none equilibrium concentrations. So it's basically a snapshot of where the reaction currently is. If Q is less than K, the reaction goes forward and if Q is greater than K, it reverses. If Q equals K, then the system is already at equilibrium.

[00:23:26]K is a temperature dependent quantity only. Concentration changes don't affect it and catalyst don't affect it either.

[00:23:36]There are three manipulations to know.

[00:23:39]If you reverse a reaction, K becomes 1 / K. If you multiply coefficients by N, K becomes K to the^ of N. And if you add reactions, you multiply their K values together.

[00:23:51]The magnitude of K tells you which side is favored. If K is much greater than one, products are favored. And if K is much less than one, reactants are favored. If K is around one, you have significant amounts of both.

[00:24:07]Look at the ice table on the right. Ice stands for initial concentrations, the change in terms of X, and equilibrium expressions. In the example, HCN reacts with water with an initial concentration of 3.5, a change of minus X and plus X for the products. Then you plug the equilibrium expressions into KC and solve for X. There's a very useful shortcut called the small X approximation. When K is small, typically less than 10 -3rd, you can assume initial minus X is approximately equal to the initial concentration.

[00:24:41]That kills the quadratic. But verify your answer because X must be 5% or less of the initial concentration and if it's more you have to redo without the approximation.

[00:24:54]The shot's principle says that when you disturb a system at equilibrium it shifts to minimize that disturbance. The chart on the right summarizes each case.

[00:25:04]For concentration changes adding a reactant shifts the equilibrium right while adding a product shifts it left. K stays unchanged. For pressure changes with gases, increasing pressure or decreasing volume shifts towards the side with fewer moles of gas. Adding inert gas at constant volume causes no shift.

[00:25:25]Temperature is the one factor that actually changes K. For exothermic reactions, raising T causes K to go down. And for endothermic reactions, raising T causes K to go up. The trick is to treat heat like a product or a reactant.

[00:25:41]Ksp is the equilibrium expression for sparingly soluble salts. It's just the ions raised to their coefficients.

[00:25:50]Molar solubility S relates directly to Ksp. For MX type salts, Ksp= S^2. For MX2 type salts, Ksp= 4S cubed. The common ion effect on the right shows that when you add a common ion equilibrium shifts left, less solid dissolves and solubility goes down.

[00:26:13]Now for unit 8, acids and bases. This unit is pretty heavily tested at 11 to 15% waiting. Remember that an acid donates a proton and a base accepts one.

[00:26:24]The diagram on the right shows the conjugate pair concept. HCl gives the proton to water becoming chloride which is the conjugate base. Water becomes hydrronium which is the conjugate acid.

[00:26:35]The general rule is that stronger acids have weaker conjugate bases.

[00:26:40]Water also selfionizes with KW = to H+ * O minus which equals 10 to the -14 at 25°.

[00:26:53]PH equals the negative log of H+ concentration and PO equals the negative log of O minus concentration. At 25° C, pH plus P always equals 14.

[00:27:08]Strong acids disassociate completely.

[00:27:11]The list on the right is one to memorize. HCl, HBr, HI, nitric acid, sulfuric acid, and perchloric acid. For these, H+ just equals the initial acid concentration.

[00:27:23]Strong bases work the same way. The common ones are sodium hydroxide, potassium hydroxide, and berium hydroxide. For O minus, multiply the initial concentration by the number of hydroxides per formula unit.

[00:27:38]Weak acids only partially dissolve, disassociate. So we use equilibrium arrows instead of full arrows. Ka equals H+ * A minus over HA.

[00:27:50]For weak bases, KB= BH+ * O minus over B.

[00:27:56]For a conjugate pair, K a * KB equals KW. pKa is just the negative log of Ka.

[00:28:03]And a smaller pKa means a stronger weak acid.

[00:28:07]There's a fast approximation where H+ is roughly the square root of Ka* initial concentration. This is valid when X is much less than C by the 5% rule. One counterintuitive point to remember is that diluting a weak acid actually increases its percent ionization.

[00:28:26]A buffer resists pH change. The recipe is a weak acid plus its conjugate base both in significant amounts. Any added H+ gets eaten by A minus and any added O minus gets eaten by HA. The Henderson household equation gives you pH equals pKa plus the log of A minus concentration over HA concentration.

[00:28:50]When the concentrations are equal, pH equals pKa.

[00:28:55]The effective range of a buffer is pH within plus or minus one unit of pKa.

[00:29:01]Buffer capacity is greatest when HA equals A minus and both concentrations are high.

[00:29:08]The titration curve on the right has four regions. The pre-equivalence region acts as a buffer for weak acid plus strong base titrations. At half equivalence, pH equals pKa. At the equivalence point, moles of titrate equals moles of analyte. And in the post equivalence region, you have excess titrant.

[00:29:28]The pH at the equivalence point depends on the type of titration. Strong acid plus strong base gives a pH of 7. Weak acid plus strong base gives a pH greater than 7. Strong acid plus weak base gives a pH less than 7.

[00:29:46]For weak acid plus strong base at equivalence, all HA has converted to A minus. So use KB= KW over KA to find the pH.

[00:29:56]When choosing an indicator, pick one with pKa close to the pH at equivalence.

[00:30:01]The AP will give you the pKa or information you can use to find them. So no need to memorize these.

[00:30:09]Molecular structure affects how strong an acid is. There are two cases to consider. The first is binary acids where H is bonded to some element X.

[00:30:20]Strength of the acid rises as the H to X bond weakens.

[00:30:26]Going down a group, bond strength drops.

[00:30:28]So HI is stronger than HBr, which is stronger than HCl, which is stronger than HF. That's why HF is actually a weak acid despite florine's high electro negativity because the H to F bond is just too strong to break.

[00:30:44]The second case is oxy acids shown on the right. More oxygens means more electron withdrawal which weakens the OD H bond and makes the acid stronger. So perchloric acid is stronger than chloric acid which is stronger than chloris acid which is stronger than hypocchloric acid.

[00:31:01]A higher electro negativity central atom also makes the acid stronger.

[00:31:08]Salts of weak acids like hydroxides and carbonates are more soluble in acidic solutions because the acid removes the basic annion. Take a metal hydroxide M O2 in equilibrium with M2+ and 2 O minus. When you add H+ it consumes the O minus so the equilibrium shifts right and more solid dissolves. Conversely, a basic pH precipitates many metal hydroxides out of solution.

[00:31:36]Now for the final unit thermodynamics and electrochemistry entropy S equals Boltzman Boltzman's constant time the natural log of W where W is the number of microates.

[00:31:49]You don't need to memorize this but you should conceptually understand that entropy means disorder. The diagram on the right illustrates this. Solids have low entropy, liquids have more and gases have much more. In general, entropy increases with more particles, higher temperature, the gas state, larger molecules, and dissolution.

[00:32:11]Delta S of reaction equals the sum of S of products minus the sum of S of reactants. Just like with enaly, reactions that produce more moles of gas typically have a positive delta S. And the third law of thermodynamics states that the entropy of a perfect crystal at 0 Kelvin is zero.

[00:32:31]The big equation on the right is delta G= delta H minus T delta S with T measured in Kelvin. This is what determines spontaneity. If delta G is negative the reaction is spontaneous. If positive nonspontaneous and if zero the system is at equilibrium.

[00:32:51]There's also a special temperature T transition equal to delta H over delta S where delta G changes time and the system is at equilibrium.

[00:33:01]Also worth knowing is thermodynamic versus kinetic control. A reaction can be favorable but slow because of high activation energy and low temperature tends to favor the kinetic product while high temperature tends to favor the thermodynamic product.

[00:33:17]The triangle on the right ties three quantities together.

[00:33:21]Delta G standard equals RT natural log of K which links which links free energy to the equilibrium constant.

[00:33:31]Delta G standard also equals NF E standard cell which links it to electrochemistry which we'll talk more about later.

[00:33:41]Under non-standard conditions, delta G equals delta G standard plus RT natural log Q. At equilibrium, Q= K and delta G equals Z.

[00:33:55]A non-spontaneous reaction can be driven forward by coupling it with a spontaneous one. The reactions add together and the coupling only works if the total delta G is negative. The classic example on the right is ATP hydraysis, which releases about -30 kJ per mole.

[00:34:16]There are two cell types, both shown on the right. Galvanic cells, also called voltaic cells, use spontaneous redux to generate electrical energy like a battery.

[00:34:28]Electrolytic cells push electrical energy in to drive non-spontaneous redux like an electro plate. In a galvanic cell, the anode is where oxygenation happens and is the negative terminal.

[00:34:40]The cathode is where reduction happens and is the positive terminal. Electrons flow externally from anode to cathode and the salt bridge keeps each half cell electrically neutral. Estandard cell equals E standard cathode minus E standard anode. A positive E standard cell means the reaction is spontaneous.

[00:35:01]The link back to thermodynamics is delta G standard= N * F * E standard where F is farad's constant. 96,485 kums per mole.

[00:35:15]The nurse equation boxed in yellow on the right says EC cell equals E standard cell minus RT / NF * natural log of Q at 25° C. This simplifies to E standard cell minus 0.0592 / N * natural log Q. If Q is less than one, products are depleted and E cell is greater than E standard. If Q is greater than one, products have built up and E cell is less than E standard. At equilibrium, E cell equals 0 and Q= K.

[00:35:51]So E standard equals 0.0592 / N * log K.

[00:35:59]In electrolysis shown in the picture to the right, the polarities flip from a galvanic cell. The anode is now positive and the cathode is now negative because you're forcing a non-spontaneous reaction to run. Faraday's law states that moles of substance equals Q / N * F. Here Q is the total charge in N is the electrons transferred per atom and F is Faraday's constant. The key relationship is that Q= I * T or current time. So with amps, seconds, and the number of electrons, you can calculate the grams plated out.

[00:36:37]And that's all nine units of AP count.

[00:36:39]Thanks for watching and good luck on your exam.