AP Chemistry Cheat Sheet PDF

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Shreya Agarwal Akshay Thakur AP Chemistry Cheat Sheet Unit 1: Atomic Structures and Properties Avogadro’s number — 6.022 × 1023  mol-1  ● Aka number of particles/atoms/molecules in one mole Atomic mass units (amu) : mass of an individual atom or molecule ● Average mass of 1 particle (atom or molecule) or formula unit of a substance = molar mass of that substance in grams. ● quantitative connection between mass and the # of particles that the substance contains. Mass Spectroscopy of Elements ● The mass spectrum of a sample of a single element can be used to determine the identity of the isotopes of that element and the relative abundance of each isotope in nature. ● An element’s amu can be estimated from the weighted average of the isotopic masses using the mass of each isotope and its relative abundance. Atomic Structure and Electron Configuration ● Aufbau Principle: electrons are added to the lowest subshells first ● Hund’s Rule: Each orbital should have one electron before any are doubled up ● Effective Nuclear Charge: net positive charge experienced by valence electrons. It can be calculated by Zeff = Z – S ○ Z : atomic #, S : # of shielding (core) electrons Photoelectron Spectroscopy ● energies of the electrons in a given shell can be measured ● position of each peak – Energy required to remove an electron from the corresponding subshell ● height of each peak – # of electrons in that subshell ● Photoelectric Effect: the emission, or ejection, of electron from the surface of, generally, a metal in response to incident light ● greater binding energy = electrons are closer to the nucleus Periodic Trends ● Zeff or effective nuclear charge is determined by the amount of protons and pulls electrons closer to the nucleus ● Ionization Energy: energy required to remove an electron ● Electron Affinity: The amount of energy released when an electron is added to a neutral atom ● Electronegativity: ability to attract electrons towards itself and to attract shared electrons to itself

Shreya Agarwal Akshay Thakur Unit 2: Molecular and Ionic Compound Structure and Properties Polar bonds: ● 2 different atoms have different electronegativities, so they have different pulls on the shared electrons, so the electrons are not shared equally → polar bond Non-polar bonds: ● Two same atoms will form a nonpolar bond since they have the same pull on shared electrons → electrons are shared equally → nonpolar bond Polar Covalent bonds: ● In a polar covalent bond, the atom that has a greater electronegativity value will have a partial negative charge ● Greater the difference in electronegativity between the 2 atoms bonded → the more polar the bond is Covalent Bonds: ● Two nonmetal atoms bonded ● Not a good conductor ○ Low melting point Ionic Bonds: ● Metal and nonmetal ○ Not a conductor as a solid, but is great conductor when either melted or dissolved in water ○ High melting point ○ Brittle, crystalline Metallic Bonds: ● Metal atoms form a crystal lattice structure with positive nuclei and core electrons, valence electrons are delocalized so they are free to move → good conductor ○ Electron sea model and delocalized electrons ● Alloys: ○ Substitutional: one element takes another's place in the lattice structure ○ Interstitial: one element fits within the spaces between another element in the lattice structure Intramolecular Forces: forces that hold individual molecules together Coulomb's Law: ● Calculates force between 2 charged particles ● Higher charge and smaller radius means stronger force ● Closer the charges are to each other, the stronger the force

Shreya Agarwal Akshay Thakur Potential Energy: ● When bonded atoms separate, PE ↑ because energy must be added to overcome the coulombic attraction between each nucleus and shared e– s Bond Length: ● Increases as you go down a group because atomic radius increases, so the distance between the 2 atoms is greater Bond Energy: ● Decreases down a column because atomic radius increases, the distance between 2 atoms increases, so coulombic attractions are weaker → less energy required to break the bond Bond Order: ● As bond order increases, bond length decreases and bond energy increases ● To calculate, add up total number of bonding pairs and divide by number of bonds Lattice Energy: ● How much energy is required to completely separate a mole of solid ionic compound into its gaseous ions ○ Lattice energy increases when the charge on the ions increases and when the size of the ions decreases Formal charge of atom: (# of valence electrons) - (# of lone electrons) - (# of bonds) ● Formal charges closer to 0 are more stable and show a preferred distribution of e-s Hybridization: ● Occurs when atomic orbitals mix to form “hybrid” orbitals to produce a set of equivalent hybrid orbitals ● Hybridization only depends on lone pairs and bonds on central atom ● Central atom is least electronegative ○ A double bond or triple bond still only counts as one bond ● To determine hybridization of an atom in a molecule: Count each pairing bond(if its double/triple bond, it still counts as one) and nonpairs ● sp hybridization - 1 s and 1 p orbital hybridizes ○ Present in atoms with only 2 electron regions ○ 180°-Bond Angle ● sp2 hybridization - 1 s and 2 p orbitals hybridize ○ Present in atoms with 3 electron regions ○ 120°-Bond Angle ● sp3 hybridization - 1 s and 3 p orbitals hybridize ○ Present in atoms with 4 electron regions ○ 109.5°-Bond Angle

Shreya Agarwal Akshay Thakur Unit 3: Intermolecular Forces and Properties Intermolecular Forces ● Determines the physical properties(g, l, s) of molecular(covalent) liquids and solids ● “Forces”, NOT “bonds” ; WEAKER than ionic/covalent bonds Types of IMFs ● Ion-Induced Dipole: force of attraction between charged ion and nonpolar molecule ● Dipole-induced dipole: force of attraction between polar and nonpolar molecule ● Ion-Dipole Forces (IDF): Between an ion and a partial charge on the end of a polar molecule ● Dipole-Dipole Forces (DDF): Medium force, Between neutral polar molecules, as dipole moment (polarity) increases, DDF increases ● London Dispersion Forces (LDF) (induced dipole-induced dipole): Weak force, between ALL molecules and the ONLY forces between nonpolar molecules, as molar mass increases, LDF increases. This because they have a larger electron cloud which increases the amount of forces in the atom, so they are more polarizable ● Hydrogen bonding forces: Strong force, special type of DDF (stronger than regular DDF), between H (bonded with F/O/N) and F/O/N (of the other molecule) Properties of Solids ● Covalent Network solids: high melting points, insoluble, hard to break or deform, held together by covalent bonds, usually do not conduct electricity as a solid ● Equilibrium Vapor pressure: when the rate of evaporation equals rate of condensation in a closed container Ideal Gas Law ● PV = nRT ● Gas Stoichiometry ○ Molar volume: volume/mole gas at STP; 22.4 L/mol ○ STP: 273 K (0.0°C and 1.0 atm); conditions of an ideal gas ○ Density: DRT/P=Molar Mass ● Dalton’s Law of Partial Pressures: The pressure of a mixture of gases is the sum of the pressures of the different components of the mixture Kinetic Molecular Theory ● Postulates ○ Gases are composed of a large number of particles that behave like hard, spherical objects in a state of constant, random motion. ○ These particles move in a straight line until they collide with another particle or the walls of the container. ○ These particles are much smaller than the distance between particles. Most of the volume of a gas is therefore empty space.

Shreya Agarwal Akshay Thakur There is no force of attraction between gas particles or between the particles and the walls of the container. ○ Collisions between gas particles or collisions with the walls of the container are perfectly elastic. None of the energy of a gas particle is lost when it collides with another particle or with the walls of the container. ○ The average kinetic energy of a collection of gas particles depends on the temperature of the gas and nothing else. Statements 2,3,6 are untrue in real gases as molecules have finite volume and do exert forces on each other Temperature: measure of the average KE of the particles of matter; can also be thought of as an index of random motions of the particles Diffusion: mixing of gases Effusion: passage of a gas through a tiny orifice into an evacuated chamber ○

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Solubility ● Solute: the compound that dissolves in solution ● Solvent: the compound that dissolves the Solute ● Miscible = soluble ● Polarity: “Like dissolves Like” → Polar dissolves Polar best, nonpolar with nonpolar. ○ *This does not mean that polar cannot dissolve in nonpolar

Shreya Agarwal Akshay Thakur Unit 4: Chemical Reactions Physical Change: ● Occurs when a substance undergoes a change in properties but not a change in composition, ie. changes in the phase of a substance, formation/separation of mixtures of substances Chemical Change: ● Occurs when a substance is transformed into new substances with different compositions, ie. introduction of heat/light, formation of gas, precipitate is formed when mixing 2 aqueous substances, and color change ● Dissolution can be considered physical and/or chemical Percent Yield: actual yield/theoretical yield x 100% Titration: ● Titrant - Solution of known concentration ● Analyte - Solution of unknown concentration ● Equivalence point - Point at which the amount of titrant added to analyte results in perfect ● neutralization ● Indicator - a substance added at the beginning of the titration that changes color at the ● equivalence point ● Endpoint - the point at which the indicator changes color Oxidation-Reduction (redox) reactions: ● 1 or more electrons are transferred in a reaction ● Oxidation: increased oxidation states (lost electrons, charge increased) ● Reduction: decreased oxidation states (gained electrons, charge decreased) ● In general, metals tend to get oxidized because they “want” to lose electrons to get octets ● Nonmetals tend to get reduced as they “want” to gain electrons to get octets Oxidizing Agent: is reduced Reducing Agent: is oxidized OIL RIG: oxidation is loss of electrons, reduction is gain of electrons

Shreya Agarwal Akshay Thakur Unit 5: Kinetics Reaction rate: ● change in concentration of a reactant/product over time: Δ[A]/Δt Reactions are dependent upon: ● Temperature ○ As temp ↑, rate ↑ ● Concentration ○ as reactants’ concentration ↑, rate↑ ● Surface Area ○ As reactant surface area ↑, rate ↑ ● Catalyst ○ With a catalyst, rate ↑ ● Chemical nature of reactants Zero Order Reactions:

First Order Reactions:

Second Order Reactions:

Shreya Agarwal Akshay Thakur

Effective Collisions: ● A chemical reaction can only occur if reactant molecules, atoms or ions collide with more than a certain amount of kinetic energy and in the proper orientation. ● When a system is heated up, the reactants will not only collide with more energy, but they will also collide more frequently. The increase in frequency increases the probability of having a collision in the proper orientation. Thus, increasing the rate of the reaction. Reaction Intermediate: ● substance that is produced and used in the reaction, but doesn’t appear in the net equation. A reaction intermediate could be a catalyst. Rate Determining Step: ● The rate determining step is the slowest step in a reaction mechanism and determines the rate or speed of the reaction.

Shreya Agarwal Akshay Thakur Unit 6: Thermodynamics Endothermic and Exothermic Processes ● Endothermic: substance absorbs energy from the surroundings, causing vessel to become colder ● Exothermic: substance releases energy into its surroundings, causing vessel to become warmer Energy Diagrams

Heat Capacity and Calorimetry ● q = mcΔT  ● q = heat, m = mass, C = specific heat, ΔT = change in temperature ● Heat capacity is the energy needed to raise the temperature by one degree J/K or J/C (specific = raise one gram of substance) ● 1st law of thermodynamics = conservation of energy Enthalpy of Reaction ● Enthalpy change represents amount of heat released/absorbed at constant pressure ● Bond energy/bond enthalpy are interchangeable Bond Enthalpies ● The average energy required to break all of the bonds in the reactant molecules can be estimated by adding up the average bond energies of all the bonds in the reactant molecules. Likewise, the average energy released in forming the bonds in the product molecules can be estimated. If the energy released is greater than the energy required, the reaction is exothermic. If the energy required is greater than the energy released, the reaction is endothermic. Enthalpy of Formation ● Enthalpy of Formation (ΔHf): enthalpy change for the formation of 1 mol of a compound from its component elements ● Hess’s Law: states that if a reaction can be described as a series of steps, then ΔH for the overall reaction is simply the sum of the ΔH values for all the steps

Shreya Agarwal Akshay Thakur Unit 7: Equilibrium Introduction to Equilibrium ● Chemical Equilibrium: the state where the concentrations of all reactants and products remain constant with time ○ Forward and reverse reactions are still occurring, just at equal rates Reaction Quotient and Equilibrium Constant

for reaction aA + bB → cC + dD ● K is the value the equation will have at equilibrium ● Q is the value of the equation at any given time, no matter if the reaction is at equilibrium or not ● If Q > K: the system will shift left to attain equilibrium by decreasing product concentration while increasing reactant concentration ● If Q = K: the system is at equilibrium ● If Q < K: the system will shift right to attain equilibrium by increasing product concentration while decreasing reactant concentration Calculating the Equilibrium Constant ● Calculating Kc involves concentration [molarity] ● Calculating Kp involves pressure [atm] ● Do not include solids and liquids when they are the solvent bc densities stay same ● Include liquids when they are the solute and gases Magnitude of the Equilibrium Constant (K) ● K is a measure of the extent of the reaction ● If K>1, products are favored ● If K1. If the Ksp...