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Chemistry 1220 exam 3 study guide

by: Sean Bhatnagar

Chemistry 1220 exam 3 study guide CHEM 1220

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Sean Bhatnagar
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Chapters 11-20 supplemental information for Exam 3 for Chemistry 1220
General Chemistry 1220
Dr. Fus
Study Guide
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This 27 page Study Guide was uploaded by Sean Bhatnagar on Sunday March 13, 2016. The Study Guide belongs to CHEM 1220 at Ohio State University taught by Dr. Fus in Spring 2016. Since its upload, it has received 127 views. For similar materials see General Chemistry 1220 in Chemistry at Ohio State University.


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Date Created: 03/13/16
Sean Bhatnagar Chemistry 1220 Exam 3 study guide Chapter 11: Liquids and Intermolecular Forces 11.1  Intermolecular forces in liquids and solids are much stronger than those in gases Gases: Assumes both volume and shape of its container Expands to fill its container Is compressible Flows readily Diffusion within a gas occurs rapidly Liquid: Assumes shape of portion of container it occupies Does not expand to fill its container Is virtually incompressible Flows readily Diffusion within a liquid occurs slowly Solid: Retains own shape and volume Does not expand to fill its container Is virtually incompressible Does not flow Diffusion within a solid occurs extremely slowly 11.2  Intermolecular forces are weaker than intramolecular forces  3 types of intramolecular forces: ionic, metallic, and covalent bonds.  Boiling points go up as intermolecular forces go up  Melting points go up as intermolecular forces go up  3 types of intermolecular forces: dispersion forces, dipole-dipole, and hydrogen bonding (collectively called the “van der Waals forces”)  All intermolecular forces are electrostatic, involving positive and negative species like ionic bonds  Interactions get stronger the as the magnitude of the charges increases, and weaker as the distance between charges increases  Dispersion forces are based on instantaneous dipole moments, also called “London dispersion forces”  Instantaneous dipole moments are based on electrons moving in the electron cloud. At some point in time the distribution of electrons on each side of the nucleus may be lopsided, creating a small dispersion charge.  Polarizability – the ease at which the charge distribution is distorted (the “squishiness of the electron cloud”)  Polarizability increases as the number of electrons increases. So in general, the dispersion forces increase as the size of the molecule increases.  The shape of the molecule also changes the magnitude of the dispersion forces  The permanent presence of dipole moments are called “dipole-dipole forces”, based on if a molecule is polar or not (draw a lewis structure)  Dipole-Dipole forces are stronger than dispersion forces  Hydrogen Bonds > Dipole-Dipole Bonds > Dispersion Forces  Hydrogen bonding occurs between hydrogen and fluorine/oxygen/nitrogen  Hydrogen bonds are the strongest because there is only one electron, so the dipole is very very strong.  Ion-dipole forces occur between an ion and a polar molecule  When the molecular weights differ greatly the attractive forces are generally stronger in the substance with the higher weight 11.3  Viscosity – resistance of a liquid to flow (molasses)  Viscosity can be measured by the time it takes a liquid to flow through a thin vertical tube, or at which a steel ball falls through liquid.  Viscosity decreases at higher temperatures  Surface tension – the energy required to increase the surface area of a liquid by a unit amount  Surface area is based on the net interior pull of molecules, making the liquid at the surface pack closely together  Cohesive forces – intermolecular forces that bind similar molecules to one another  Adhesive forces – intermolecular forces that bind molecules to a surface  Capillary action – when liquid rises up a very narrow tube, based on adhesive forces to the tube wall and surface area reducing surface area, pulling the liquid up the tube (used for nutrients in plants) 11.4  Phase changes – changes of state  Phase changes are accompanied by a change of energy in the system  Fusion and melting are the same, “heat of fusion” is in kJ/mol.  Sublimation – solid to gas  Vaporization – liquid to gas  Temperature remains constant until the END of the phase change  Heating curves look like stairs, the flat parts are the phase changes and the inclines are the heating portions  Heating curve – graph of temperature vs. heat added  Critical point – the highest temperature at which a distinct liquid phase can form  Critical pressure – pressure required to bring about liquefaction at this critical temperature  Supercritical fluid – when temperature exceeds the critical temperature and pressure exceeds the critical pressure, the gas and liquid phases are indistinguishable 11.5  As molecules vaporize from liquid to gas the gas exerts a pressure on the closed container  Vapor pressure – pressure of the vapor on the container once it has reached a constant value (when the L2G/G2L transitions are in equilibrium)  Volatile – liquids that evaporate readily  Substances with high vapor pressures means the equilibrium has higher gas content, because equilibrium only refers to the equality in RATES  Temperature increase will increase vapor pressure because the increased kinetic energy given to molecules allows more of them to break out of the liquid phase.  Normal boiling point – boiling point of a liquid at 1atm  Boiling point – point at which the vapor pressure is equal to that of the external pressure.  Clausius-clapeyron equation describes the relationship of vapor pressure and temperature 11.6  Phase diagram – (three branched chart) used to describe the equilibrium conditions between states (think SLG in a clockwise fashion starting with the leftmost state)  Vapor pressure curve – the curve between gas and liquid on the phase diagram, ends at the critical point, beyond which is the supercritical fluid  Sublimation curve – curve between the solid and gas phase  Melting curve – curve between the liquid and solid phase, usually slopes towards the right as pressure increases because in most cases solid is more dense than liquid  Normal melting point – melting point at 1atm  Triple point – the intersection of the vapor-pressure, sublimation, and melting curves  Melting curve for water slants backwards, because of the rare liquid is denser than solid characteristic for water 11.7  Liquid crystal – the viscous milky state of some organic compounds  Liquid phase – molecules arranged randomly (think scattered rattlesnake egg magnets (REMs))  Nematic liquid crystalline phase – long axes of molecules aligned but not aligned at ends (think REMs all standing up but at different heights)  Smectic A liquid crystalline phase – Molecules aligned in layers, long axes of molecules perpendicular t layer planes (remember perpendicular)  Smectic C liquid crystalline phase – molecules aligned in layers, long axes of molecules inclined with respect to layer planes (think slightly skewed, but similar to smectic A)  Certain liquid crystals stack better due to benzene rings  Cholisteric liquic crystal – the molecules pack into layers; the long axis of each molecule is oriented parallel to its neighbors within the same layer (think single slices of the other phases stacked horizontally atop one another)  Cholisteric liquid crystals produce unusual color patterns with visible light, used to monitor temperature changes when conventional methods are unfeasible Chapter 13: Properties of Solutions 13.1  Solutions are formed when one substance disperses uniformly throughout another  The mixing of gases is a spontaneous process, associated with an entropy increase  Gases spontaneously mix unless restrained by their containers  Solutions made from solutes of solid or liquid state may or may not mix based on the intermolecular forces (think NaCl and H O) 2  Dispersion forces dominate in nonpolar (NP) – nonpolar solutions  Ion-dipole forces dominate in ionic substances in water  Solute-solute interactions – between solute particles must overcome in order to disperse the solute particles through the solvent  Solvent-solvent interactions – between solvent particles must be overcome to make room for the solute particles in the solvent  Solvent-solute interactions – between solvent and solute particles occur as the particles mix  Solvation – when solute molecules are surrounded by solvent molecules  Hydration – solvation with water as the solvent  ΔH(soln) = ΔH(solute) + ΔH(solvent) + ΔH(mix)  Separation of solvent and solute molecules requires energy, so the enthalpy is positive, then it drops down largely when mixed, as this releases energy. The difference between the mixed energy released and the energy required to break the bonds is the ΔH(soln)  This chart of the aforementioned points is the Hess’s Law one  Ionic solutes do not dissolve in nonpolar solvents  Polar liquid solutes (water) will not dissolve in a nonpolar liquid solvent (octane)  Chemical reactions cannot be recovered when dehydrated, unlike physical reactions (salt in water when dehydrated will recover the salt) 13.2  Crystallization – as more particles dissolve there is a higher chance of them colliding with the surface and reattaching to the surface of the original solid (this is the opposite of the solution process)  Equilibrium is reached when the rate of crystallization and the rate of solution are equal  Saturated – a solution in equilibrium with undissolved solute, additional solute will not dissolve  Solubility – the amount of solute needed to form a saturated solution in a given quantity of solvent  Unsaturated – when less solute is dissolved than the amount needed to form a saturated solution  Supersaturated – when, under suitable conditions, more solute is dissolved than that of a saturated solution.  Usually heat is used to make a solution supersaturated, when the solution is cooled down particles will fall out of solution and crystallize because a supersaturated solution is considered unstable  However, for crystallization to occur the particles must align themselves to properly form crystals 13.3  The stronger the attractions between the solute and the solvent the greater the solubility of solute in that solvent  POLAR DISSOLVES POLAR (due to favorable dipole-dipole interactions)  Miscible – pairs of liquids that mix in all proportions (acetone and water)  Immiscible – pairs of liquids that do not dissolve one another (gasoline and water – gasoline is NP and water is P)  Hydrocarbons (molecules just containing C and H atoms are NP)  NONPOLAR DISSOLVES NONPOLAR  Alcohols – Organic compounds with a framework of hydrocarbons with a polar OH group attached (these groups are able to form hydrogen bonds)  For alcohols, as the number of carbon atoms increases the OH group becomes an ever smaller part of the molecule making it act more nonpolar, decreasing its solubility in water and increasing its solubility in NP solvents  To increase its solubility the molecules should increase OH groups, why glucose which has 5 OH groups, is very soluble in water  Finally: LIKE DISSOLVES LIKE  The solubility of any gas is increased as the partial pressure of the gas above the solvent increases, pressure increases increase solubility  Henry’s law (S =gkP ) gelates partial pressure of gases and solubility of the gases  Solubility of solid solutes in water tends to increase as temperature increases (as seen by solubility curves)  Solubility of gaseous solutes in water decreases with increasing temperature (think of when you warm a glass of cold tap water, bubbles rise because the air dissolved comes out of solution)  Thermal pollution – warm water is less dense than cold water, so it sits at the surface. As it is warmer less oxygen can dissolve in it, which means less oxygen can get to deeper water and life can die 13.4  Mass percentage = (mass of component in solution)/(total mass of solution) X 100  Very dilute solutions are measured in “parts per million (ppm)” and “parts per billion (ppb)” 6  Ppm = (mass of component)/(total mass) X 10  Ppb is the same as ppm except with 10 9  Mole fraction of component = (moles of component)/(total moles of all components) (X is typically used as a symbol for mole fraction)  Molarity (M) = moles of solute/liters of soln  Molality (m) = moles of solute/kg of solvent  KNOW HOW TO CONVERT BETWEEN CONCENTRATION UNITS 13.5  Colligative properties – depend on the quantity (concentration) but not the identity of solute particles  Examples of colligative properties – freezing-point lowering, boiling-point raising, vapor-pressure lowering, and osmotic pressure  Nonvolatile – a substance with no measurable vapor pressure  Adding a nonvolatile solute to a volatile solvent decreases the overall vapor pressure  The vapor pressure of a volatile solvent above a solution containing a nonvolatile solute is proportional to the solvent’s concentration in the solution o  Raoult’s Law – P soln (X solventP solvent  ΔP = (X solutePosolvent  Ideal solution – one that oberys Raoult’s Law (there are many cases where the solution is not ideal)  Addition of a nonvolatile solute brings down the vapor pressure, which brings down the vapor-pressure curve (L-G curve), making a higher temperature for the normal boiling point. Hence increasing the boiling point  Change in boiling point is directly related to molality (m)  ΔTb= K mb  K – “molal boiling-point-elevation constant” b  Molality is based on the TOTAL moles of solute particles 1m of NaCl is 2m – (1m Na+, 1m Cl-)  So consider whether the solute is an electrolyte or non-electrolyte when calculating boiling-point elevation  Based on the previous note of how the vapor-pressure curve moves down, the triple point is moved to the left, thus decreasing the normal freezing point  ΔTf= K mf  Osmosis – the net movement of solvent toward the solution of higher solute concentration, as if the solutions were driven to attain equal concentrations  Osmotic Pressure (OP) oberys a law similar to the ideal gas law  (OP)(V)=(n)(R)(T) – when the V is divided, n/V is M, so OP = MRT  Cell in hypertonic environment (higher solute concentration outside), causes solvent to leave the cell, the cell will shrivel, this is called crenation  Cell in a hypotonic environment (higher solute concentration inside) causes solvent to enter the cell, the cell will swell or even rupture, this is called hemolysis  Isotonic solution – one that matches the concentration of the cell, preventing crenation and hemolysis 13.6  Particles 5nm to 1000nm are considered “colloidal particles”, they are small enough to remain suspended in solution, but large enough to be seen or have an effect on the aesthetic  Colloids – solutions containing colloidal particles or dealing with colloidal dispersions  Tyndall effect – the scattering of light by a colloid  Examples of colloids – fog, smoke (aerosol), whipped cream (foam), milk (emulsion), paint (sol), marshmallow (solid foam), butter (solid emulsion), ruby glass (solid sol)  Most important colloids are those where the solvent is water, they are either hydrophilic (water loving) or hydrophobic (water fearing)  Typically, a colloid will have hydrophobic groups on the interior of the molecule with hydrophilic, polar groups on the surface interacting with the water (enzymes, antibodies)  Adsorption – adhering to a surface, ions adsorb to hydrophobic molecules to stabilize them.  To remove colloidal particles, coagulation must be used  Heating a colloidal dispersion, or adding an electrolyte to a dispersion brings about coagulation, increasing the size of these colloidal particles and allows them to be filtered out  In dialysis, a semipermeable membrane is used to allow the passage of ions and prevent the passage of colloidal particles Chapter 14: Chemical Kinetics  Kinetics is based on the rates at which reactions occur, called “reaction rates”  Reaction mechanism – step-by-step, molecular-level view of the pathway from reactants to prducts 14.1  Factors that affect reaction rates are as follows:  Physical state of the reactants – for example reactions involving solids will occur faster if the solid has greater surface area… hence why many medicines are in the form of a powder  Reactant concentrations – many reactions will occur faster if the concentrations of one or more reactants is greater (steel wool bursting into flames in pure oxygen, but burning slowly in air which is only 20% O ) 2  Reaction temperature – reaction rates generally increase as the temperature increases (as T goes up, KE goes up, increasing collisions, increasing the likelihood of reactions occurring)  The presence of a catalyst – catalysts speed up reactions  Catalyst – agents that increase reaction rates without themselves being used up  On a molecular level, reactions are based on collisions. For a collision to lead to a reaction, however, it must be with enough energy to break the bonds and the right orientation to form the new bonds 14.2  Reaction rate - change in concentration of the reactants or products per unit of time (usually molarity per second, M/s)  By convention, rates are expressed as positive quantities, so include a (-) when calculating the rate of disappearance of a reactant  It is typical for rates to decrease as a reaction proceeds because the concentration of reactants decreases  Instantaneous rate – rate at a particular instant during the reaction (approximated using tangent lines, like in calculus)  To find a reaction rate that is balanced on stoichiometry, multiply the instantaneous rate by one over the coefficient  You do not need to do this to find the rate of a certain reactant or product 14.3  Rate law = k[A][B]  K is referred to as the “rate constant”  K must have the units that cancel out the individual reactants rates until the final units are M/s  Reaction order – the exponents on the reactants in the rate law  Overall reaction order – the sum of the exponents  Second order – increasing the concentration by two, quadruples the rate – exponent of 2  Zero order – takes the reactant out of the rate law, exponent of zero, reactant’s concentration has zero impact on the rate  Rate laws can only be determined experimentally, there is no correlation between exponents and reaction orders although it may seem as if there is  Most reaction orders are 0, 1, or 2, but instances occur where the reaction order is fractional 9  K of 10 or greater means a fast reaction, k of 10 or less indicates a slow reaction  Rate of a reaction depends on the concentration, however the rate constant does NOT 14.4  Relating a rate law to time involves integration –  First-order reaction: ln[A] t ln[A] =0-kt  Second-order reaction: 1/[A] = kt + 1/[A] 0  To find the order without rate data – plot ln[A] and 1/[A] vs time and see which is linear  Zero-order reaction: [A] = tkt + [A] 0  Half-life – (1/2 the time required for the concentration of a reactant to reach half of its initial value  To solve for half-life given a rate constant, simply plug in ½ the initial concentration for [A] t  In a first-order reaction, the concentration of the reactant decreases by ½ in each series of regularly spaced time intervals, each interval equal to t 1/2  In second-order and other reactions the half-life depends on the concentrations, and therefore changes as the reaction progresses 14.5  As temperature increases so does the reaction rate, this is due to the rate constant being dependent on temperature  Collision model – based on kinetic-molecular theory, this model accounts for the effects of both reactant concentrations and temperature on the rate of the reaction.  Collision model says as concentration increases, the likelihood of collisions increases, increasing reaction rates. Also, as temperature increases as does kinetic energy, increasing the force and frequency of collisions  Molecules must be oriented in a certain way during a collision for the reaction to take place  Activation energy – the minimum energy required to initiate a chemical reaction  Activated complex/transition state – the arrangement of atoms shown at the top of a reaction’s energy profile, at the highest activation energy  RATES ARE DEPENDENT ON ACTIVATION ENERGY  At higher temperatures, a larger fraction of molecules have sufficient kinetic energy to overcome the activation energy and cause a reaction  f = e^(-Ea/RT) where R is 8.314 J/mol-K  Arrhenius equation relates the temperature and activation energy to the reaction rate  Arrhenius: k = Ae^(-Ea/RT) (k is the rate constant, Ea is the activation energy, R is the gas constant previously mentioned, T is the absolute temperature, A is the frequency factor) 14.6  Reaction mechanism – steps by which a reaction occurs  Elementary reaction – reactions that occur in a single event of a single step  Molecularity – the number of molecules that participate as reactants in an elementary reaction  Unimolecular – molecularity of 1  Bimolecular and termolecular are self-explanatory using the above bullet  Termolecular reactions are very rare in comparison to uni- and bi-, four or more is so rare that they are never proposed as part of a reaction mechanism  Multistep mechanism – consists of a sequence of elementary reactions that result in the balanced chemical equation  Intermediate – a molecule that is created in one of the steps and consumed in the next, a reactant and product that is not in the final balanced chemical equation  Intermediates can be identified as the trough between two transition states in an energy profile of a reaction  Rate laws can only not be determined experimentally if the reaction is elementary, then it is based on the molecularity  Rate-determining step – step in the mechanism that is the slowest (has the highest activation energy)  Whenever a fast step precedes a slow one, we can solve for the concentration of an intermediate by assuming that an equilibrium is established in the fast step  Manipulating the fast steps rate laws and solving for the intermediate allows one to create a rate law for the overall reaction based on the slow step that does not involve the intermediate 14.7  Catalyst – a substance that changes the speed of a chemical reaction without undergoing a permanent chemical change itself  Homogenous catalyst – a catalyst that is present in the same phase as the reactants in a reaction mixture  A catalyst is present at the start and end of a reaction, while an intermediate is created and consumed during the reaction  A catalyst lowers the overall activation energy of a reaction (possibly by offering a different mechanism for the reaction)  Heterogenous catalyst – a catalyst that exists in a phase different from the phase of reactant molecules, usually as a solid in contact with either gaseous reactants or with reactants in a liquid solution  Heterogenous catalysts are usually metals or metal oxides, and are given very large surface areas because the first step in catalysis is the adsorption of reactants to the surface of the catalyst  Enzyme – a biological catalyst  Most enzymes are large protein molecules that are very selective in the reactions they catalyze, sometimes an enzyme will only operate in only one reaction  Active site – where the enzyme catalysis takes place  Substrate – the substance that reacts at the active site  Lock-and-key model – the substrate acts as the key and the active site as the lock  Enzyme-substrate complex – the combination state of the substrate and the active site  Enzyme inhibitors – when anything other than the necessary substrate attaches to the active site it destroys the activity of the enzyme (nerve poison) Chapter 15: Chemical Equilibrium  Chemical equilibrium – occurs when opposing reactions proceed at equal rates, as a result concentrations cease to change and it appears as if the reaction has stopped 15.1  Equilibrium state/mixture – a mixture of reactants and products whose concentrations no longer change with time  Equilibrium mixtures occur because a reaction is reversible, symbolized by two half arrows pointing both in opposite directions  Equilibrium is DYNAMIC: the forward reaction is still occurring as is the reverse reaction  For equilibrium to occur, neither reactants nor products can leave the system  At equilibrium, a particular ratio of concentration terms equals a constant 15.2  Equilibrium condition is reached from either direction (you can start with all products or all reactants and the equilibrium condition reached will be the same)  Equilibrium constant – K ,cunitless  Equilibrium-constant expression: K = [c] [E] / [A] [B] a b  Subscript “c” means the concentrations in molarity are used to express the equilibrium  The equilibrium-constant expression depends only on the stoichiometry of the reaction, not on its mechanism  The value of K cnly depends on the particular reaction and on the temperature, not on initial concentrations  A similar equilibrium expression can be created based on pressures, K , p where in place of equilibrium concentrations there are equilibrium partial pressures  Kp= K (cT) Δn(where Δn = (moles of product) – (moles of reactant)) 15.3  If K >> 1 – equilibrium “lies to the right”, products predominate  If K << 1 – equilibrium “lies to the left”, reactants predominate  The equilibrium-constant expression for a reaction written in one direction is the reciprocal of the expression for the reaction written in the reverse direction  Because reactions can be balanced differently and be either forward/reverse, it is important to specify what the reaction is and at what temperature when listing an equilibrium constant  The equilibrium constant for a net reaction made up of two or more reactions is the product of the equilibrium constants for the individual reactions 15.4  Homogenous equilibria – equilibria where the substances are all in the same phase  Heterogenous equilibria – equilibria where the substances are in different phases (solid lead(II) chloride dissolving in water to form a saturated solution)  Whenever a pure solid or a pure liquid is onvolved in a heterogenous equilibrium, its concentration is not included in the equilibrium-constant expression 15.5  How to calculate the equilibrium constant with the concentration of at least one species (ICE TABLE) 1. Write the initial and equilibrium concentrations that appear in the equilibrium expression 2. For those species with initial concentrations, calculate the change in concentration that occurs to reach equilibrium 3. Use stoichiometry to calculate the change in concentration for the other species in the equilibrium-constant expression 4. Determine the value of the equilibrium constant 15.6  Reaction quotient (Q) – a number obtained by substituting reactant and product concentrations or partial pressures at any point during a reaction into an equilibrium-constant expression  Q is typically used to predict which direction the reaction will go given the temporary concentrations/pressures  If Q = K – the reaction is at equilibrium because the reactant quotient equals the equilbirum constant  If Q > K – the concentration/pressures of the products is too large, therefore the reaction proceeds from right to left (reverse direction)  If Q < K – the concentration/pressures of the reactants is too large, therefore the reaction proceeds from left to right (forward direction) 15.7  Le Chatelier’s Principle – If a system at equilibrium is disturbed by a change in temperature. Pressure, or a component concentration, the system will shift its equilibrium position so as to counteract the effect of the disturbance  Concentration (changed by adding/removing a reactant/product) – If a substance is added to a system at equilibrium, the system reacts to consume some of the substance. If a substance is removed from a system, the system reacts to produce more of a substance  Pressure (changed by changing volume) – At constant temperature, reducing the volume of a gaseous equilibrium mixture causes the system to shift in the direction that reduces the number of moles of gas (and vice versa)  Temperature – If the temperature of a system at equilibrium is increased, the system reacts as if we added a reactant to an endothermic reaction, or a product to an exothermic reaction. The equilibrium shifts in the direction that consumes the “excess reactant”, namely heat.  If the moles of gas are equal on each side of a reaction, a change in pressure will not affect the position of equilibrium  Changes in pressure or concentration DO NOT change the equilibrium- constant value (K /p )c only temperature changes affect the value.  Adding an inert gas, or a gas that does not participate in the reaction, will increase the total pressure but because it does not change the reactant’s/product’s partial pressures it does not cause a shift in equilibrium  In an endothermic reaction, heat is considered a reactant  In an exothermic reaction, heat is considered a product  Endothermic: increasing T results in a higher K value (and vice versa)  Exothermic: increasing T results in a lower K value (and vice versa)  A catalyst increases the rate at which equilibrium is achieved but does not change the composition of the equilibrium mixture Chapter 16: Acid Base Equilibria 16.1  Acids are sour and can cause certain dyes to change color (lemons)  Bases are bitter and feel slippery (soap)  Arrhenius acids/bases: Acid – a substance that, when dissolved in water, increases the concentration of H ions Base – a substance that, when dissolved in water, increases the - concentration of OH ions  Bronsted-Lowry acids and bases are based on the fact that acid-base + reactions involve the +ransfer of H ions from one substance to another  Hydronium ion – H O 3  Bronsted-Lowry acids/bases Acid – a substance that donates a proton to another substance Base – a substance that accepts a proton (In the creation of a hydronium ion, H O 2s the base and H is the acid)  Amphiprotic – a substance capable as acting as either an acid or a base  Amphiprotic substances act as a base when combined with something more strongly acidic than itself and acts as an acid when combined with something more strongly basic than itself  Conjugate base – the removal of a proton from an acid creates this conjugate base, which is considered a base due to its role in the reverse reaction  Conjugate acid – the addition of a proton to a base creates this conjugate acid, which is considered an acid due to its role in the reverse reaction  Conjugate acid-base pair – an acid and its conjugate base or a base and its - conjugate acid are considered one of these pairs (HX and X )  The stronger the acid the weaker its conjugate base, and the stronger the base the weaker its conjugate acid  Strong acid – completely transfers its protons to water, leaving no undissociated molecules in solution, its conjugate base shows negligible basicity  Weak acid – only partially dissociates in aqueous solution, creating a solution of both acid and conjugate base, the conjugate base of a weak acid is a weak base  Negligible acidity – contains hydrogen but does not demonstrate acidic behavior, the conjugate base of a substance with negligible activity is a strong base  In every acid-base reaction, equilibrium favors transfer of the proton from the stronger acid to the stronger base to form the weaker acid and the weaker base 16.3  Autoionization of water – water can act as a bronsted-lowry acid AND base, in fact they can donate protons to one another + -  Kc= [H O3][OH ] – this expression’s constant is replaced with “K ” w  K = 1.0 x 10 -14 w  By this convention, in an acid solution [H ] > 1.0 x 10 M, and in a basic + -7 solution [H ] < 1.0 x 10  Neutral solution – [H ] = [OH ] - 16.4 +  [H ] is usually very small, so it tends to be expressed in pH  pH = -log[H ]+  pH < 7 (acidic), pH = 7 (neutral), pH > 7 (basic)  pH decreases as [H ] increases  [H ] concentration is small, but can greatly impact the rates of reactions. Such that human blood has a pH range of 7.35 – 7.45, and any significant deviation from this range results in death -  pOH = -log[OH ]  pH + pOH = 14.00 16.5  Strong electrolytes – strong acids and bases, existing in aqueous solution entirely as ions  Strong acids – HCl, HBr, HI, HNO , HC3O , HCl3 , H SO 4 2 4  Reactions with strong acids are not in equilibrium, they go to completion. So for monoprotic acids (acids with a single proton), the concentration of the + acid equals [H ]  Strong bases – LiOH, KOH, NaOH, RbOH, CsOH, Ca(OH) , Ba(OH) , Sr2OH) 2 2  Strong bases act the same as strong acids, dissociating completely, thus the - - concentration of the base = [OH ] for bases with one OH group  Metal oxides (CaO for example) are also good at creating strongly basic solutions 16.6  Acid-dissociation constant – the equilibrium constant for an acidic solution, “Ka”  Ka= [H ][A ] / [HA]  The larger the value of K theastronger the acid  Always use an ICE table (previously mentioned) when working with weak acids/bases, as they do not dissociate completely and their concentrations can only be determined using an equilibrium-constant expression  Percent ionization = (concentration ionized) / (original concentration) X 100% (another measure of acid strength) +  Example: [H ] equilibriumHA] initial00%  Percent ionization decreases as the concentration increases for a weak acid  Polyprotic acids – acids with more than one ionizable H atom  It is always easier to remove the first proton from a polyprotic acid than to remove the second, K valaes become successively smaller as protons are removed 16.7  Base-dissociation constant – always refers to the equilibrium in which a base - reacts with H 2 to form the corresponding conjugate acid and OH , “K ” b  First category of weak bases – has a lone pair of electrons (for accepting a proton), most contain a nitrogen atom (ammonia)  “amines” – a subgroup of this first category, has a N – C bond and can form a N – H bond  Second category is comprised of anions of weak acids 16.8  The product of the acid-dissociation constant for an acid and the base- dissociation constant for its conjugate base equals the ion-product constant for water  Kax K b K w  This proves that the stronger the acid the weaker its conjugate base, because as Kaincreases (indicative of a stronger acid), K musb go down to keep the previous bullet true 16.9  Nearly all salts are strong electrolytes + -  Hydrolysis – reaction involving ions and water to generate either H or OH  A weak base in solution will form its conjugate acid, which will remove protons and create hydroxide ions, increasing the pH and making the solution more basic (somewhat counterintuitive)  For amphiprotic salts If a > K bhen the ion causes the solution to be acidic If b > K ahen the ion causes the solution to be basic  Many metal ions react with water to decrease the pH of an aqueous solution  Salts effects on pH: If the salt’s cation AND anion do not react with water, the pH will be neutral. This occurs if the anion is the conjugate base of a strong acid and the cation is from a strong base (NaCl, Ba(NO ) , RbClO ) 3 2 4 If the salt’s anion produces hydroxide ions and the cation does not react the pH will be basic. This occurs if the anion is the conjugate base of a weak acid, and the cation is from a strong base (NaClO, RbF, BaSO ) 3 If the salt’s cation produces hydronium ions and the anion does not react the pH will be acidic. This occurs when the cation is the conjugate acid of a weak base or a small cation with a charge >= 2+ (NH NO , Al4l ,3Fe(NO 3 ) 3 3 If both the cation and anion react, the pH is dependent on the relative ability to react of the ions (NH C4O, Al(CH CO3) , Cr3 ) 3 16.10  Strength of an acid is determined by the polarity of a bond, the strength of the bond, and the stability of the conjugate base  HCl for instance, H has a positive charge so it acts as a proton donor, in NaH however, H has a negative charge which makes it a proton acceptor, and in CH 4he bonds are nonpolar so it does not react to produce an acidic or basic solution  Binary acid – an HX acid where “X” represents members of the same group/period in the periodic table  For binary acids where “X” represents members of the same group, bond strength is the determining factor of acid strength. Typically, bonds become weaker as you go down the periodic table and elements become heavier. Bond strength decreases and strength of acid increases  For binary acids where “X” represents members of the same period, bond polarity is the determining factor of acid strength. Acidity increases as electronegativity of “X” increases, generally moving from left to right across a period  Acidity increases downward and to the right as a combination of the above two bullets  Oxyacids – acids in which OH groups and possibly additional oxygen atoms are bound to a central atom  For oxyacids, think of a strong base where the cation is a nonmetal, this requires the bond to be covalent and OH will not likely split off (general form YOH)  For oxyacids, as the electronegativity of Y increases, so does the acidity of the substance.  The strength of an oxyacid also increases as additional electronegative atoms bond to the central atom, Y Hypochlorous (HClO) < Chlorous (HClO ) < Chl2ric (HClO ) < Perch3oric (HClO 4 (in terms of strength)  Carboxylic acids – acids containing a carboxyl group (COOH), the additional O – C bond draws electron density from the O – H bond, making it more polar and increasing the stability of the conjugate base  Conjugate bases of carboxylic acids (referred to as “carboxylate anions”) can exhibit resonance 16.11  Lewis acid – electron pair acceptor  Lewis base – electron pair donor Chapter 17: Additional Aspects of Aqueous Equilibria 17.1  Common-ion effect – whenever a weak electrolyte ad a strong electrolyte containing a common ion are together in solution, the weak electrolyte ionizes less than it would if it were alone in a situation  Example: CH COO3 and CH COONa ha3e a common ion – CH COO 3 - + - CH C3ONa ionizes completely and produces Na and CH COO 3 CH C3OH is a weak electrolyte so it only ionizes partially into H and CH COO 3 - According to Le Chatelier, if CH COO3a were added to a solution containing CH C3OH, the ionization of the weak acid would decrease – because the salt ionizes completely adding CH COO ,3a product, to the solution and pushing the equilibrium to the left, reforming the weak acid  Common-ion effect also affects base-dissociation 17.2  Buffered solutions or “buffers” – solutions that contain a weak conjugate acid- base pair, these solutions resist drastic changes to pH when strong acid/base is added  Buffers work because the weak acid and base pair coexist, so if strong acid is added the weak base can neutralize, and vice versa  Because the acid and base are conjugates of each other they can coexist in a solution with neutralizing one another  Henderson-Hasselbalch equation – pH = pK + log( [base] / [acid] )  Buffer capacity – the amount of acid or base the buffer can neutralize before the pH begins to change to an appreciable degree  Buffer capacity is greater the greater the concentrations of the buffer pair, even if the ratio and pH are the same as one of lesser concentrations  pH range – pH range over which a buffer acts effectively  Buffers usually have a usable range within +/- 1 pH unit of pK a  Reactions between strong acids and weak bases proceed essentially to completion, as do those between strong bases and weak acids (we can assume the strong acid/base is completely consumed by the reaction with the buffer) 17.3  Equivalence point – the point at which stoichiometrically equal quantities of acid and base have been brought together  pH titration curve – a graph of pH as a function of volume of titrant added, the data is collected using a pH meter  A titration curve can be used to find the equivalence point as well as determine the K oa the weak acid or K of tbe weak base being titrated  Strong acid – strong base titrations The initial pH is low, due to the acid being strong initially The pH at equivalence point is very close if not 7.00  Weak acid – strong base titrations pH is slightly higher than that of a strong acid due to the lower ionization of a weak acid The pH equivalence point will be higher than 7.00, because the strong base will dissociate and neutralize the weak acid leaving the weak conjugate base, leaving the solution slightly basic  The process is the mirror opposite for the previous two titrations if the strong/weak base is titrated with strong acid  A titration of a polyprotic acid will have multiple equivalence point as each form of the acid is neutralized (resembles rounded stairs)  When choosing an indicator, choose one where the equivalence point of the titration falls within the indicator’s color-change interval 17.4  Solubility-product constant (or “solubility product”) – equilibrium constant responsible for referring to how soluble the solid is in water, “K ” sp  The solubility product K of a compound equals the product of the sp concentration of the ions involved in the equilibrium, each raised to the power of its coefficient in the equilibrium equation  Molar solubility – number of moles of solute that dissolve in forming 1L of saturated solution of the solute (mol/L)  Many factors can affect the value of solubility including pH of the solvent and the presence of other ions in solutions (common ions), however, K sponly has one value for a given solute at any specific temperature 17.5 (STUDY THIS)  The solubility of a slightly soluble salt is decreased by the presence of a second solute that furnishes a common ion (think of the common-ion effect previously mentioned)  The pH of a solution affects the solubility of any substance whose anion is basic  The solubility of a compound containing a basic anion (that is, the anion of a weak acid) increases as the solution becomes more acidic  The solubility of slightly soluble salts containing basic anions increases as [H ] increases (as pH is lowered)  Complex ion – the assembly of a metal ion and the Lewis bases bonded to it  Formation constant “K ” –fequilibrium expression for the formation of a complex ion  When forming a complex ion, the addition of the lewis base consumes the metal ion, and according to Le Chatelier this drives the dissolution of the salt to the right, as the removal of product tends to do  Amphoteric oxides/hydroxides – metal oxides and hydroxides that are relatively insoluble in water but dissolve in strongly acidic and strongly basic solutions (such as Al(OH) fo3ming Al(OH) (a com4lex ion) in a basic solution, 3+ or Al in an acidic solution) 17.6  If Q > K spprecipitation occurs, reducing ion concentrations until Q = K sp  If Q = K spequilibrium exists, the solution is saturated  If Q < K spsolid dissolves, increasing ion concentrations until Q = K sp  Selective precipitation – separation of ions in an aqueous solution by using a reagent that forms a precipitate with one or more (but not all) of the ions 17.7  Quantitative analysis – determines how much of a given substance is present  Qualitative analysis – determines the presence or absence of a particular metal ion  THIS ENTIRE SECTION IS: the solubility flowchart on methods used to determine the content of a solution using precipitations and logic Chapter 19: Chemical Thermodynamics  Chemical thermodynamics – the area of chemistry that deals with energy relationships 19.1  First law of thermodynamics – energy is conserved, it may be transferred from one system to another but it cannot be destroyed nor created  ΔE = q + w (ΔE is the change in the internal energy of a system, q is the heat absorbed or released by the system from/to the surroundings, and w is the work done on the system by the surroundings or on the surroundings by the system)  If q > 0 then the system absorbing heat from its surroundings  If w > 0 then the surroundings are doing work on the system  Spontaneous process – one that proceeds on its own without any outside assistance  Spontaneous processes can only occur in one direction, the reverse process is nonspontaneous  Experimental conditions (temperature, pressure, etc) can be important in determining whether or not a process is spontaneous  Thermodynamics tells us the direction and extent of a reaction but nothing about the speed  State function – a property that define a state but do not depend on how the state is reached (temperature, internal energy, and enthalpy)  q and w are NOT state functions  ideal engine – operates under an ideal set of conditions in which all processes are reversible  A reversible change produces the maximum amount of work that can be done by a system on its surroundings  Isothermal process – a process carried out under a constant temperature  When two objects of different temperatures are in contact heat flows spontaneously from the hotter object to the colder one  Reversible processes are those that reverse direction whenever an infinitesimal change is made in some property if the system  Any spontaneous process is irreversible 19.2  Entropy – “S” – thermodynamic quantity associated either with the extent of randomness in a system or with the extent to which energy is distributed among the various motions of the molecules of the system  Entropy is a state function  ΔS = q rev T (for an isothermal process) (q revs the heat transferred if the process were reversible, and T is the absolute temperature)  Second law of thermodynamics – any irreversible process results in an increase in total entropy, whereas any reversible process results in no overall change in entropy  Reversible process: ΔS univ ΔS sys ΔS surr 0  Irreversible process: ΔS univ ΔS sys+ ΔS surr 0  Entropy of the universe increases in any spontaneous process 19.3  Statistical thermodynamics – a field of study which uses the tools of statistics and probability to link the microscopic and macroscopic worlds  Microstate – a single possible arrangement of the positions and kinetic energies of the gas molecules when the gas is in a specific thermodynamic state  “W” is used to describe the number of microstates  Boltzmann’s Equation – S = k ln(W) (where k is Boltzmann’s constant (1.38 x -23 10 J/K)  Entropy is a measure of how many microstate are associated with a particular macroscopic state  Entropy increases with the number of microstates of the system (therefore entropy increases with increasing temperatures and increasing volumes)  Translational motion – an entire molecule moving in one direction  Vibrational motion – atoms in the molecule move periodically toward and away from one another  Rotational motion – the molecule spins about an axis  All three types are collected under the umbrella of “motional energy” of a molecule  The number of microstates possible for a system increases with an increase in volume, an increase in temperature, or an increase in the number of molecules because ant of these changes increases the possible positions and kinetic energies of the molecules making up the system  Entropy will increase if: Gases form from either solids or liquids Liquids or solutions form from solids The number of has molecules increases during a chemical reaction  Third law of thermodynamics – the entropy of a pure crystalline substance at absolute zero is zero 19.4  Standard molar entropies – S – molar entropies for substances in their standard states  Standard molar entropies generally increase with more degrees of motion, increasing molar mass, and increasing number of atoms in a formula of a substance  ΔS = Σ nS (products) – Σ mS (reactants) (this is the general form for all standard measures)  ΔSsurr= - ΔH sys T  When ΔS > 0 the process is spontaneous, as entropy of the universe increases for spontaneous processes 19.5  A spontaneous process that is endothermic must be accompanied by an increase in entropy of the system  Spontaneous processes that result in a decrease in the system’s entropy are always exothermic o o o  ΔG = ΔH - TΔS  If ΔG < 0, the reaction is spontaneous in the forward direction  If ΔG = 0, the reaction is at equilibrium  If ΔG > 0, the reaction in the forward direction is nonspontaneous (work must be done to make it occur) but the reverse reaction is spontaneous  Gibbs free energy – G – a state function used in thermodynamics  In any spontaneous process carried out at a constant temperature and pressure, the free energy always decreases  Again: ΔG = Σ nΔG (prodfcts) – Σ mΔG (reactanfs) 19.6  The sign of ΔG, and thus the determination of spontaneity, is based on the signs and magnitudes of ΔS and ΔH  Recap: +ΔH – endothermic -ΔH – exothermic +ΔS – increased entropy (increased randomness) -ΔS – decreased entropy (decreased randomness) 19.7 o  Using standard conditions (ΔG ) we can calculate the free energy under non- standard conditions o  ΔG = ΔG + RT ln(Q) (where R is the ideal gas constant 8.314 J/mol-K, T is the absolute temperature, and Q is the reaction quotient of the mixture)  Under standard conditions Q = 1, so ln(Q) = 0, so ΔG = ΔG o  At equilibrium ΔG = 0 and Q = K, therefore: ΔG = -RT ln(K)  Be careful of units, the ideal gas constant R is in joules (J), while free energy is measures in kilojoules (kJ) Chapter 20: Electrochemistry  Electrochemistry – the study of the relationships between electricity and chemical reactions  Oxidation – loss of electrons  Reduction – gain of electrons 20.1  Oxidation numbers (“oxidation states”) – used to identify whether a reduction-oxidation reaction is occurring (“redox”)  In a redox reaction both oxidation and reduction must occur  Oxidizing agent (“oxidant”) – the substance that makes it possible for another substance to be oxidized  Reduction agent (“reductant”) – the substance that makes it possible for another substance to be reduced  The reducing agent is oxidized, the oxidizing agent is reduced 20.2  2 rules must be observed when balancing redox reactions 1. Conservation of mass 2. The gains and losses of electrons must be balanced  Half-reactions – equations that show either oxidation or reduction alone, method used to balance redox reactions  Balancing half-reactions – (1) balance all elements other than H and O, (2) balance O atoms with H O a2 needed, (3) balance H atoms with H as + - needed, (4) balance charge with e as needed  Then multiply the half reaction by integers so the electrons lost in the oxidation half-reaction match the electrons gained in the reduction half- reaction  Finally, combine the half-reactions, canceling out any species that appear on both sides of the reaction  If the redox reaction is occurring in a BASIC solution, the only difference in - + process is balancing with OH and H O, r2ther than H and H O (when 2t says “in a basic solution”, add the amount of OH required to match the amount of added H to both sides of the equation. This will make the H into H O, and 2 leave the other OH on the other side) 20.3  Voltaic or galvanic cell – device in which the transfer of electrons takes place through an external pathway rather than directly between reactants present in the same reaction vessel  Cathode – the cell that is reduced, determined by having the higher potential of the two electrodes  Anode – the cell that is oxidized, determined by having the lower potential of the two electrodes  Salt bridge – Anions migrate toward the anode Cations migrate toward the cathode  Electrical current – created by electrons flowing through a wire and ions moving in solution  Electrodes – two solid metals connected by the external circuit  Half-cell – each compartment of a voltaic cell  Salt bridge – a porous glass disc separating the two half-cells allowing the ions to migrate and maintain the electrical neutrality  The electrolyte’s ions that are in the salt bridge will not react with the ions in solution 20.4  Electrons flow spontaneously toward the electrode with more positive electrical potential  Potential difference – energy per electrical charge between two electrodes, measured in volts  Volts (V) – potential difference required to impart 1 joule (J) of energy to a charge of 1 coulomb (C) 1 V = 1 J/C  Cell potential (E cell– the potential difference between the two electrodes of a voltaic cell  Cell potential is the driving force behind the flow of electrons, also called the “electromotive force” or “emf”, also commonly called the “voltage” o o o o  E cell E redcathode) – E redanode) (E cell 0 in a voltaic cell)  Standard hydrogen electrode (SHE) – an electrode designed to have a standard reduction potential of exactly 0 V.  Whenever we assign an electrical potential to a half-reaction, we write the reaction as a reduction  Changing the stoichiometric coefficient in a half reaction does not affect the value of the standard reduction potential o  The more positive the value of E redthe greater the tendency for reduction under standard conditions. In addition, the greater the tendency to reduce, the better the oxidizing agent  The half-reaction with the most negative reduction potential is the one most easily reversed and run as an oxidation 20.5  Simplifying to include energy not just from voltaic cells E = E oredreduction process) – E oredoxidation process)  A positive value of E indicates a spontaneous process, a negative value of E indicates a nonspontaneous process  “E” is used to represent emf under nonstandard conditions, “E ” is used to indicate the standard emf  ΔG = -nFE (n is the number of moles of electrons, F is Faraday’s constant 96,485 C/mol (J/V-mol), and E is the emf of the redox reaction)  A positive value of E and a negative value of ΔG both indicate a spontaneous reaction  Relating ΔG to E and K E = ((RT)/(nF))*ln(K) 20.6 o  Nernst equation: E = E – (RT/nF) ln(Q)  Concentration cell – a cell constructed using the same species in both half- cells creating an emf based solely on a difference in concentration (use the Nernst equation) 20.7  Battery – a portable, self-contained electrochemical power source that consists of one or more voltaic cells  When cells are connected in series (cathode of one connected to anode of another) the battery produces a voltage that is the sum of the voltages of the individual cell  Primary cell – a non-rechargeable battery that must be discarded or recycled after the voltage drops to zero  Secondary cell – a battery that can be recharged from an external power source after its voltage has dropped  Lead-Acid battery – used in automobiles, consists of six voltaic ells in series, each producing 2 V (total 12 V), both electrodes are immersed in sulfuric acid, can be recharged  Alkaline battery – most common primary cell  Nickel-Cadmium, Nickel-Metal-Hydride, and Lithium-Ion batteries – created to fulfill the need of portable devices, lightweight rechargeable batteries  Nicad batteries are the most common; each cell produces 1.30 V, usually 3+ cells in succession. However, they are toxic and dense  NiMH batteries last up to 8 years, they are being used in gas-electric hybrid vehicles  Li-ion batteries are very light so they achieve a greater energy density  Fuel cells – not batteries as they are not self-contained. These cells directly crease electricity from fuels such as CH o4 H 2 20.8  Corrosion – spontaneous redox reactions in which a metal is attacked by some substance in its environment and converted to an unwanted compound  Rusting – the corrosion of iron, requires oxygen and water  Cathodic protection – protecting a metal from corrosion by making it the cathode in an electrochemical cell  Sacrificial anode – the metal that is oxidized while protecting the cathode 20.9  Electrolysis reactions – take place in electrolytic cells and use electrical energy to drive the reactions in the nonspontaneous direction  Electroplating – uses electrolysis to deposit a thin layer of one metal on another metal to improve beauty or resistance to corrosion  Coulombs – the measure of the quantity of charge passing through an electrical circuit (coulombs = amperes x seconds)


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