MCAT General Chemistry Lecture 2
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Created by:
campbellc24 Plus on May 29, 2012
Subjects:
Gases, Kinetics and Chemical Equilibrium
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186 terms
Terms | Definitions |
|---|---|
 Standard Temperature and Pressure: | for a gas, the temperature of 0 Celsius and pressure 1.00 atmosphere |
| Mean Free Path: | the average distance traveled by a molecule between collisions. Since the distances between molecules in a gas are so incredibly small, and the gas molecules move so fast, there are billions of collisions per second.. this is why some chemical reactions can appear to occur instantaneously. |
| Is a gas homogenous or heterogenous if it is a mixture of different compounds? | It would be homogenous regardless of polarity differences. |
| Homogenous: | Mixture that is the same throughout |
| Since polarity does not cause gas to separate, what does? | Colder temperatures cause denser gas to settle beneath less dense gas. |
| Idea Gas: | a gas with no intermolecular forces that obeys the ideal gas law (PV = nRT) at all temps, pressures and volumes. In other words, all gases, behaving ideally, will have the same volume, if they have the same temperature, pressure and number of molecules. |
| What 4 characteristics does an ideal gas have that real gases do not, within the Kinetic Molecular Theory? | 1) Gas molecules have zero volume in an ideal gas 2) Gas molecules exert no forces other than repulsive forces due to collisions within an ideal gas 3) Gas molecules make completely elastic collisions 4) The average kinetic energy of gas molecules is directly proportional to the temperature of the ideal gas. |
| Ideal Gas Law: | PV=nRT P= pressure in atm V= volume in litres n= number of moles of gas R= universal gas constant T= temperature in Kelvin |
| What do special cases of the ideal gas law include? | 1) Charle's Law: 2) Boyle's Law 3) Avogadro's Law |
| Charle's Law: | The volume of a gas is proportional to temperature at constant pressure |
| Boyle's Law: | The volume of a gas is inversely proportional to pressure at constant temperature |
| Avogadro's Law: | The volume of a gas is proportional to the number of moles at constant temperature and pressure. |
| Molar Mass: | MM=mRT/PV m= mass in grams #moles=grams/AMU |
| Gas Density: | ρ=PMM/RT |
| It is possible to cool a gas by ________ the volume, even the PV=nRT indicates that temperature is directly proportional to volume. | increasing; when air is let out of a tire, the gas rapidly expands and thus cools, this is why it feels so cold on your hands. |
| How does a gas do work? | As a gas expands in a container, it does work on its surroundings by pushing outwards on the walls of the container. The force on the walls times the distance that the walls expand is the work done by the gas. W=FxD This work represents a transfer of energy from the gas to the surroundings |
| Where does the energy for work done by gas come from? | The energy comes from the Kinetic Energy of the gas molecules. Both temperature and pressure are related to kinetic energy of the molecules. |
| Pressure is related to kinetic energy per ? | Volume |
| Temperature is related to kinetic energy per? | Mole |
| What happens to pressure and temperature as a gas expands? | As the gas expands, the pressure decreases due to both the loss of kinetic energy and the increase in volume. The temperature decreases due to the loss in kinetic energy via work done. |
| If the volume of gas were doubled and the Kinetic energy remained the same, what would happen to pressure? | Pressure would be reduced by a factor of 2, or halved. But if no heat was added, the kinetic energy would be reduced. |
| What would happen to pressure if the kinetic energy was reduced as the volume doubled? | The pressure would be reduced by more than a factor of 2, thus temperature would also decrease to preserve the equality PV=nRT. |
| What volume will one mole of any gas, behaving ideally, occupy at STP? | 22.4 Litres= The standard molar volume |
| Kinetic Molecular Theory: | the theory that all matter is composed of particles (atoms and molecules) moving constantly in random directions |
| How do individual gases in a mixture behave? | Each gas in a mixture essentially acts as if it were alone in a container. In a mixture of gases, each gas contributes to the pressure in the same proportion as it contributes to the number of molecules of the gas. The amount of pressure contributed by any gas in a gaseous mixture is called the partial pressure of that gas. |
| Partial Pressure: | the contribution each gas in a mixture of gases makes to the total pressure. The partial pressure of a particular gas is the total pressure of the gaseous mixture (all gases) times the mole fraction of the prticular gas. Pa=XaPtotal Xa= mole fraction of gas a= # moles of gas a/total # moles of gas in the sample. |
| What is the pressure of 2 moles of gas at 0°C, occupying 11.2 litres? | PV=nRT, STP= 0°c, 1 atm per 1 mole of ideal gas, therefore we doubled the number of moles and halved the volume, so we must multiply the pressure by a factor of 4.= 4 atm. |
| Dalton's Law: | States that the total pressure of a mixture of gasses is equal to the sum of the pressures of all the gases in the mixture. Ptotal= P₁+P₂+P₃...., This is a good way to understand an ideal gas. Each gas behaves as if it were in a container by itself so all the partial pressures added together equal the total pressure. |
| Average Kinetic Energy of a Gas: | K.E.avg= (3/2)RT, Can be calculated by the rms, root-mean square velocity. This equation is valid for any fluid system, including liquids. This KE, is the average kinetic energy for a mole of gas molecules and not the energy of every, or maybe even any, of the molecules. A gas molecule chosen at random may have almost ANY kinetic energy associated with it. |
| What dictates the average kinetic energy of the molecules in a gas and what does this mean? | Temperature dictates the average kinetic energy of the molecules in a gas, thus the gas molecules of each gas in a mixture must have the same average kinetic energy if at the same temperature. SInce different gases will have different molecular masses, they will have different rms velocities. Therefore, by setting their velocities equal to each other, we can derive a relationship between their rms velocities= called Graham's Law. |
| Graham's Law: | Graham's law states that the rate of effusion of a gas is inversely proportional to the square root of its molecular weight. Thus, if the molecular weight of one gas is four times that of another, it would diffuse through a porous plug or escape through a small pinhole in a vessel at half the rate of the other. Graham's law is most accurate for molecular effusion which involves the movement of one gas at a time through a hole. It is only approximate for diffusion of one gas in another or in air, as these processes involve the movement of more than one gas. |
| How do you write Graham's Law? | ν₁/ν₂=√m₂/√m₁ |
| Graham's law tells us that the average speed of the molecules of a pure gas is inversely proportional to the _____ ____ of the ____ of the gas molecules. | square root of the mass of the gas molecules. Meaning, those molecules with the greater molecular mass will travel at speed less than those with small molecular masses. |
| What are the 2 types of gaseous spreading that Graham's law gives information for? | effusions and diffusion |
| Effusion: | is the spreading of a gas from high pressure to low pressure through a pinhole. A pinhole is defined as an opening much smaller than the average distance between the gas molecules. |
| Explain why molecules with higher rms velocities find the pinhole more frequently. | Molecules with higher rms velocities, thus lower molecular masses, travel much more quickly within mixture thus will experience more collisions with the walls of the container, therefore the rate at which molecules from such a gas find the pinhole will be much greater. |
| Ratio of the Rates of Effusion for 2 gases: | is equal to the inverse of the ratio of the square roots of their molecular weights and equal to the ratio of their rms velocities. Effusion rate₁/Effusion rate₂=√m₂/√m₁ |
| Diffusion: | is the spreading of one gas into another gas or into an empty space. The ratio of the diffusion rates of two gases is approximated by Graham's Law. The |
| Why would the diffusion rate be much slower than the rms velocity of the molecules? | The diffusion rate is slower than the rms velocity of the molecules because gas molecules collide with each other and with molecules of other gases as they diffuse, making the rate of diffusion a slow process, having the gases move slower than their actual velocities. |
| If the ratio for an effusion or diffusion rate is greater than 1, what does that mean? | If the ratio, X is greater than 1, the numerator, or top value has a rate of diffusion or effusion of X times greater than the denominator or bottom value. It could also mean that X travels X times further than the denominator. |
| When and why do real gases deviate from ideal behavior? | Real gases deviate from ideal behavior at low temperature and high pressure, when their molecules are close together, thus at lower volume. When molecules are close together, their volume become significant compared to the volume around the molecules. Also, when molecules are closer together there are greater electrostatic forces between them and they too become significant. High pressure pushes the molecules together and low temperatures cause gas molecules to settle close together, each resulting in deviations from ideal behavior. Gases deviate from ideal behavior at 10atm and around their BP. |
| Gases only behave ideally... | at high temperature and low pressure, when the molecules are far apart. In other words, gases behave ideally at high volumes |
| How can ideal behavior deviations occur with either an increase in pressure of a decrease in temperature? | To create deviations in ideal behavior, we need only move the molecules closer together, thus decrease the volume. According to PV=nRT, if we decrease the volume, the Pressure must increase and the temperature must decrease. |
| What are 2 ways in which we can decrease the volume? | 1) Pushing molecules closer together with greater pressure, or 2) Lower the temperature and let molecules settle closer together. |
| Why do gases deviate from ideal behavior as pressure increases? | As pressure increases, the ratio of PV/RT should stay at 1, however for real gases, as pressure increases, the attractive intermolecular forces between molecules lead to lower pressure than predicted by ideal gas law. Also, real gas molecules take up space, whereas the ideal gas molecules are assumed to have zero volume. As a result, real gases occupy volumes greater than predicted. |
| When is the ratio of PV/RT less than 1? | As atmospheric pressure begins to increase, attractive intermolecular forces between gas molecules increase- due to shorter mean free path between gas molecules. Therefore, the early negative deviation in the ratio is due to the attractive intermolecular forces that are prevalent, such that the impact of a given molecule against the wall of the container will be reduced by this attraction to nearby molecules. This produces a partial pressure of the gas that is less than predicted from the ideal gas law. PV/RT < 1 due to attractive intermolecular forces |
| When is the ratio of PV/RT greater than 1? | As pressure continues to increase, the molecular volume effect predominates and creates the late positive deviation from ideal gas behavior. The combined volume of the gas molecules becomes a larger fraction of the total space available. Therefore, at high external atmospheric pressure, the volume of the gas molecules no longer remains negligible relative to the total volume occupied by the gas. This produces a volume larger than predicted by the ideal gas law, thus PV/RT is >1 due to molecular volume effect. |
| All gases deviate from ideal behavior as ______ increases and _______ decreases. | pressure increases; temperature decreases |
| At moderate pressure, a real gas exerts ____ pressure than predicted by the ideal gas law. | less pressure due to attractive intermolecular forces |
| At high pressure, molecules of a real gas have a volume ______ than predicted by the ideal gas law. | greater volume due to the volume of the gas molecules becoming significant. |
| As temperature _______, a gas will approach ideal behavior because why? | As temperature increases, a gas will approach ideal behavior because the attractive intermolecular forces between molecules are minimized due to overwhelming distance between molecules, as well as the volume per molecule becomes negligible relative to the overall volume of the gas. |
| PV/nRT=1 | For an ideal gas at high temp and low pressure No IMF or volume of molecule |
| PV/nRT<1= negative deviation for pressure | Real gas at low temp and moderate pressure There are IMF as the predominating factor |
| PV/nRT>1= positive deviation for volume | Real gas at low temp and high pressure There are molecular volumes associated with each molecule of gas as the predominating factor |
| Van der Waals equation: | approximated the real pressure and real volume of a gas. [P+a(n/V)²](V-nb)=nRT] b= measure of actual volume occupied by a mole of gas. a= reflects the strength of the attractive intermolecular forces V- volume of real gas The values of a and b generally increase with the molecular mass and molecular complexity of gas. |
| What are the 2 ideal gas deviations for real gases? | 1) Vreal>Videal Molecules of a real gas do have volume Videal can be calculated from PV=nRT 2) Preal<Pideal Molecules in a real gas do exert forces on each other, and those forces are attractive when the molecules are far apart. In a gas, repulsive forces are only significant during or near collisions. Since the predominant intermolecular forces in a gas are attractive, gas molecules are pulled inward toward the centre of the gas, and slow before colliding with container walls... therefore, having been slightly slowed, they strike the container walls with slightly less force than predicted by ideal gas law, thus the pressure exerted on the container wall is less for a real gas. |
| What do we expect PV/RT to equal for an IDEAL gas? | For an Ideal gas, we expect PV/RT to equal 1, for 1 mole of an ideal gas at any temp or pressure. Though volume deviates positively from ideal behavior, and pressure deviates negatively from ideal behavior. Thus, if PV/RT>1 for one mole of gas, then the deviation due to molecular volume must be greater than the deviation due to intermolecular forces. If PV/RT<1 for 1 mole of gas, then the deviation due to intermolecular forces must be greater than the deviation due to molecular volume. |
| Density of a gas: | ρ=m/V... since ρ=nxMM/V--> ρ=PMM/RT |
| If a FORCE is applied to a container of gas, reducing its volume by half, what happens to the temp and why? | Since there is work being done on the container, there is also work being done on the gas, thus the energy from the force applied will be transferred to the gas, increasing its internal energy. Since the internal energy of the gas increases and the number of moles remains the same, the temperature, which is average kinetic energy per mole also increases. |
| What is the square root of 2? | 1.4 |
| What is the square root of 3? | 1.7 |
| Chemical Kinetics: | the area of chemistry that is concerned with reaction rates and reaction mechanisms. Kinetics deals with the rate of reaction as it moves towards equilibrium, while thermodynamics deals with the balance of reactants and products after they have achieved equilibrium. Kinetics tells us how fast equilibrium is achieved, while thermodynamics tells us what equilibrium looks like. |
| Collision Model: | In order for a reaction to occur, the reacting molecules must collide, but since the rate of a given reaction is found to be much lower than the frequency of collisions, most collisions do not result in a reaction. |
| The 2 requirements for a given collision to create new molecules in a reaction are: | 1) The relative kinetic energies of the colliding molecules must reach a threshold energy called the Activation Energy, due to relative velocity only. Therefore, velocity in a direction away from another molecule decreases the relative kinetic energy of collision. 2) The colliding molecules must have the proper spatial orientation. |
| Thermodynamics: | direction of reaction |
| Kinetics: | speed or rate of reaction |
| Does the rate of an exothermic reaction increase with temp? | Yes, temp can be thought of as a reactant, so if more heat, there will be an increased rate. Increasing the rate simply means that the equilibrium is reached more quickly. |
| Arrhenius Equation: | the mathematical equation k=Ae^-Ea/RT, which expresses the dependence of the rate constant on temperature. A formula for the temperature dependence of the reaction rate constant, and therefore, rate of a chemical reaction. In short, the Arrhenius equation gives "the dependence of the rate constant k of chemical reactions on the temperature T (in absolute temperature kelvins) and activation energy[3] Ea", |
| What does the value of the Rate Constant depend on? | Pressure, catalysts and temperature. Pressure dependence is typically negligible. Temp dependence is seen in the Arrhenius equation. The rate constant generally doubles to triples with every increase in temp by 10°C. |
| Why does the rate of a reaction increase with temperature? | The fraction of collisions that have at least the activation energy (e^(-Ea/RT)), increases with temperature. This in turn indicates that the rate constant K increases with increasing temp for nearly all reactions. The rate constant is directly proportional to the rate of a reaction. The rate of a reaction increases with temp mainly because more collisions with sufficient relative kinetic energy occur each second. |
| Does the energy of activation change with increasing temp? | No, Ea is independent of temp. |
| Reaction Rate: | the decrease in reactant concentration or increase in product concentration per unit of time as a reaction proceeds. Essentially, how quickly the concentration of a reactant or product is changing. The reaction rate (rate of reaction) or speed of reaction for a reactant or product in a particular reaction is intuitively defined as how fast or slow a reaction takes place. |
| What are the factors that affect the rate of reaction? | temp, pressure, and concentration of certain substances in the reacting system, though pressure effects are usually small enough to be ignored. |
| What does the rate of a reaction tell us? | it tells is how quickly the concentration of a reactant or product is changing. |
| What does it mean when they say that a reaction is "elementary"? | An elementary reaction is a reaction that occurs in a single step.The stoichiometric coefficients in an elementary reaction give the molecularity of the reaction, it tells you how many molecules participate in a reaction producing collisions. |
| Molecularity: | the number of molecules colliding at one time to make a reaction. There are 3 possible molecularities: unimolecular, bimolecular or termolecular. |
| How do you determine the molecularity of the following elementary reaction? aA+bB--> cC+dD | a+b = molecularity |
| How do you distinguish an elementary reaction from a complex or composite reaction? | There is no way to distinguish an elementary reaction from a complex reaction by inspection of the chemical equation, on the MCAT, the only way to know if a reaction is elementary is if you are told so. |
| Elementary Reaction: | a reaction in which reactants are converted to products in a SINGLE STEP. An elementary reaction is a chemical reaction in which one or more of the chemical species react directly to form products in a single reaction step and with a single transition state. |
| Unimolecular Elementary Reaction: | A reaction involving one molecular entity is called unimolecular. |
| Bimolecular Elementary Reaction: | A reaction involving two molecular entities |
| Average Reaction Rate for any brief time interval t during the reaction: aA+bB--> cC+dD | rate= -(1/a)(Δ[A]/t)= -(1/b)(Δ[B]/t)= +(1/c)(Δ[C]/t)=+(1/d)(Δ[D]/t) the negative sign indicates that the reactant concentrations are decreasing as the reaction moves forward. The Positive sign indicates that the concentration of the products is increasing as the reaction moves forward. This rate equation is strictly for elementary reactions, though can approximate complex reactions is concentration of intermediates is kept low. |
| Intermediates: | Species that are products of one reaction and reactants of a later reaction in a reaction chain. |
| Why is the concentration of intermediates often very low? | Because intermediates are often very unstable and react as quickly as they are formed. |
| What is the Rate Law for the following reaction? aA+bB--> cC+dD | Rate of forward Rxn= Kf[A]^a[B]^b |
| Given the following rate law: Rate of forward Rxn= Kf[A]^a[B]^b, What is the overall order of the reaction and what are the order of the individual reactants? | Order of reactant A=a, B=b. Overall order of Reaction= a+b |
| In order to determine the rate law via experiment, what must be done? | Since both the order of the reactants and the value of the rate constant must be determined through experiment. 1) Find order of each reactant by by comparing the rates between two trials in which only the concentration of one of the reactants is changed. 2) Derive the rate law from the newly derived reactant orders and concentrations. 3) After deriving the rate law, plug in the values and derive the rate constant, kf. |
| What does the graph look like when plotting [A] with respect to time, t, for a zeroth order reaction? | It results in a straight line, with a slope of -kf. The Line line goes from top left to bottom right as time passes, indicating that the reactant [A] decreases linearly with time. |
| Is the half-life for an irreversible second order reaction with a single reactant, dependent on the [A]? | Yes, each consecutive half-life is twice as long as the last. For instance, the time required to reduce the concentration of A from 100% to 50% is half as long as the time required to reduce the concentration from 50% to 25%. |
| How are the complications of reverse reactions often avoided? | By employing the technique of initial rates. In the initials moments of a reaction starting with all reactants and no products, the rate of the reverse reaction is zero. Usually, the rate law is determined with initial rates. |
| Any complex reaction can be separated into ________ steps. | elementary |
| Rate Determining Step: | Slowest step in a reaction mechanism determining the overall rate of reaction. The rate of the slowest elementary reaction determines the rate of the overall reaction. |
| What determines the rate of the overall reaction in a complex reaction? | The RDS |
| If the RDS is the first step, how can the rate law of the overall reaction be determined? | By the RDS and no other step. |
| If the slow step is any other than the first step, how can the rate law be determined? | The slow step is still the determining step but steps prior to the slow step will contribute to the rate law. Steps after the slow step will make no contribution to the overall rate law. |
| If the slow elementary step of a complex reaction is the first step what is the RDS? | The RDS would be the first, slow step. |
| If we derive the rate law from an elementary equation, and in the elementary equation 2 of the same molecules collide as reactants, we use that as the ? | order of the reactant for that equation. NO₂+NO₂-->NO₃+NO, The rate law is: rate=k₁[NO]², automatically using the coefficient of the balanced equation for the exponent in the rate law works only if the equation is elementary. |
| What happens when the first step of a reaction series is the fast step? | The rate of the overall reaction is still equal to the RDS, however, now one of the products of the fast step is a reactant in the slow step. Such a species is called an intermediate. The concentration of the intermediate is tricky to predict because it is unstable and reacts quickly. If we assume that the fast reaction reaches equilibrium very quickly, then the concentration of the intermediate remains at its equilibrium concentration. Thus, we can use the equilibrium concentration of the intermediate to predict the slow step. |
| Equilibrium Approximation: | Assumes that all steps prior to the RLS or the RDS are in equilibrium. This requires that the slow step be significantly slower than the fast step. |
| Steady State Approximation: | Frequently, the kinetic equations include intermediates, even if they remain undetected. The Steady-State Approximation is which the concentration of the intermediate is assumed to be small and essentially unchanging during much of the reaction. |
| If a reaction is in equilibrium, what can we assume about the forward and reverse reactions? | They are equal to each other. |
| Catalyst: | A substance that initiates or accelerates a chemical reaction without itself being affected. A substance that increases the rate of a reaction without being consumed or permanently altered. They are capable of enhancing product selectivities and reducing energy consumption. A catalyst does NOT lower the activation energy, it finds a different reaction path with a lower activation energy. A catalyst may also increase the steric factor. Usually the catalyst participates in the slowest step if the reaction |
| Catalysts work by providing an alternative reaction mechanism that ______ with the uncatalyzed mechanism. | competes |
| Does a catalyst alter the equilibrium constant of a reaction? | No, a catalyst cannot alter the equilibrium constant of a reaction, so it must increase the rate of both the forward and reverse reactions. |
| A _______ will increase the rate of both the forward and reverse reaction. | catalyst |
| For the MCAT, catalysts ______ change the equilibrium composition! | DO NOT |
| A heterogenous catalyst: | is in a different phase than the reactants and products. They are usually solids while the reactants and products are liquids or gases. A reactant may physically adsorb (VDW forces) or, more often, chemically adsorb (via covalent bonds) to the surface of a solid catalyst. |
| Adsorption vs. absorption | Adsorption: is the binding of molecules to a surface. Absorption: is the uptake of molecules into an interior. |
| How does binding of molecules to a heterogenous catalyst happen? | Molecules bind to a metal surface because, unlike metal atoms in the interior, metal atoms at the surface have unfilled valence requirements. The binding is almost always exothermic and the rate of catalysis depends upon the strength of the bond between the reactant and the catalyst. |
| What does the rate of catalysis depend on with a heterogenous catalyst? | The rate of catalysis depends upon the strength of the bond between the reactant and the catalyst. If bonds are too weak, not enough adsorption occurs; if bonds are too strong, too much energy is required to remove the reactant. |
| How can the rate of catalysis with a heterogenous catalyst be increased and why? | The rate of catalysis with a heterogenous catalyst can be enhanced by increasing the surface area of a catalyst because once adsorbed, molecules migrate from one adsorption site to the next, the more binding that occurs, the greater the reaction rate. |
| A homogenous catalyst: | is in the same phase as the reactants and products, usually in the gas or liquid phase. Aqueous acid or base solutions often act as homogenous catalysts. |
| Autocatalysis: | catalysis in which the catalyst is one of the products of the reaction |
| If the concentration of a catalyst is small compared to the concentration of the reactants and products, how would you increase the rate of the reaction? | By increasing the concentration of the catalyst |
| If the concentration of the catalyst is large compared to that of the reactants and products, will the rate of the reaction be affected by increasing [catalyst]? | The rate will change very little if at all by increasing [catalyst] when it is already high compared to the concentrations of the reactants and products. |
| Since catalysts alter the reaction mechanism, reactions with catalysts require separate ____ _____. | rate constants. This is because the catalyst does not prevent the original reaction from proceeding, so the total rate is given by the sum of the rates for both reactions. Though, typically the rate of the original reaction is negligible compared to the catalyzed rate. |
| A catalyst increases reaction rate by lowering activation energy for a different path for the reaction to proceed, but a catalyst does not change the? | does not change the equilibrium and it is not used up or consumed. The catalyst creates a new reaction pathway, that typically includes an intermediate. |
| Turnover Number: | The number of subtrate molecules converted to product per enzyme molecule in one second. |
| The total reaction rate with a catalyst is the sum of the ? | Rates of the original uncatalyzed reaction and the catalyzed reaction. |
| Do molecules in liquid or molecules in gas make more collisions per second? | Liquid molecules make more collisions per second than do gas molecules, roughly 100 times more because the liquid molecules are much closer together. However, most of the collisions within a liquid are with the solvent resulting in no collision. |
| What is the rate constant in a liquid a function of? | The rate constant in a liquid is a function of the solvent as well as the temperature. The reactant in a liquid is solvated. is the process of attraction and association of molecules of a solvent with molecules or ions of a solute. As ions dissolve in a solvent they spread out and become surrounded by solvent molecules. These solvent-reactant bonds must be broken before a reaction can take place. These bonds may also stabilize the intermediate. Therefore, the degree of solvation affects k. |
| What are the 3 ways that a solvent can affect k, the rate constant of a reaction? | 1) Degree of solvation 2) Dielectric of the solvent 3) Solvent viscosity can affect k via the "cage effect". |
| How does the degree of solvation affect k, the rate constant of a reaction? | The solvent makes intermolecular bonds with the reactants, thus stabilizing the intermediates. |
| How can the dielectric of a solvent affect k? | The solvent can electrically insulate reactants reducing the electrostatic forces between them. |
| Cage Effect: | The cage effect in chemistry describes how properties of a molecule are affected by its surroundings. In a solvent a molecule is often more accurately described existing in a cage of solvent molecules, the so-called solvent cage. Reactions occur when a molecule occasionally "jumps out" and meets another molecule. Reactants in a liquid can be trapped in a cage of solvent molecules. They rattle around in the cage at such tremendous rates making hundreds of collisions before squeezing between solvent molecules and into a new solvent cage. If they are trapped in the solvent cage with another reactant, many of their collisions are with the other reactant and a reaction is likely to occur, but if there is not another reactant in the solvent cage, they cannot react until they escape the cage. The net result is that the reactants in a liquid make an approximately equal number of collisions with other reactants as they would in a gas with equal concentrations; collisions in a liquid occur at about the same rate as in gas. |
| What change to a reaction will always increases the rate of the reaction? | increase the temperature |
| Catalysts: | - are not used up in the reaction - lower the activation energy - increase the rates of the forward and reverse reactions - do not alter the equilibrium |
| When determining the Rate Law by experiment, given a table with numerous trials, what info can you derive when you double a reactant concentration and the rate also doubles? | If the concentration of a single reactant causes the doubling of a reaction rate, that reactant is thus directly proportional to the concentration of that reactant, thus that reactant gets an exponent of 1. |
| When determining the Rate Law by experiment, given a table with numerous trials, what info can you derive when you double a reactant concentration and the rate of the reaction quadruples? | The rate of that reaction is proportional to the square of the concentration of that reactant, thus that reactant will be given an exponent of 2. |
| When determining the Rate Law by experiment, given a table with numerous trials, what info can you derive when you double a reactant concentration and the rate of the reaction remains the same? | The rate of the reaction is independent of the concentration of that reaction, thus the reactant is given an exponent of 0. |
| With an increase in temperature in an exothermic reaction: | 1) Reaction rate increases 2) Rate constant increases 3) rms molecular velocity increases, ... activation energy is NOT changed with temperature. |
| A+B+Heat-->C+D | Endothermic |
| A+B-->C+Heat | Exothermic |
| A first order reaction has a ______ half life. | constant. |
| If in a first order reaction the concentration of the reactant is doubled, what happens to the reaction rate? | A-->products, this is a first order reaction. The reaction rate is thus doubled since, Reaction Rate=k[A]... The concentration of the reactants is directly proportional to the reaction rate. |
| If a radioactive isotope undergoes nuclear decay, and the concentration of the isotope decreases exponentially with a constant half life, what order is the reaction? |  first order. |
| What does the graph look like for a zeroth order reaction? |  |
| What is the net reaction rate at equilibrium? | The net reaction rate at equilibrium is zero because there is a forward and a reverse reaction rate at equilibrium that cancel out. Equilibrium is a dynamic process. |
| Entropy, ΔS: | Entropy is a thermodynamic property that can be used to determine the energy not available for work in a thermodynamic process, such as in energy conversion devices, engines, or machines. Such devices can only be driven by convertible energy, and have a theoretical maximum efficiency when converting energy to work. During this work, entropy accumulates in the system, which then dissipates in the form of waste heat. Entropy of an isolated system always increases or remains constant. Thus, entropy is also a measure of the tendency of a process, such as a chemical reaction, to be entropically favored, or to proceed in a particular direction. It determines that thermal energy always flows spontaneously from regions of higher temperature to regions of lower temperature, in the form of heat. These processes reduce the state of order of the initial systems, and therefore entropy is an expression of disorder or randomness. |
| Why is Entropy an expression of disorder and randomness? | It is an expression of disorder and randomness because it states that thermal energy must always flow from a region with high thermal energy to a region with low thermal energy, thus destabilizing the two regions. In this sense, the entropy of an isolated system always increases or remains constant. |
| What happens to the rates of a chemical reaction as the reaction proceeds? | The rate of the forward reaction begins to slow as reactants decrease and the rate of the reverse reaction quickens as the concentration of the products increases. The rates of a reaction are related to concentrations. |
| What is the point of greatest entropy? | Equilibrium |
| We only ever use the stoichiometric coefficients as the exponents in the rate law when? | We only ever use the stoichiometric coefficients as the exponents in the rate law when we are told that the reaction is elementary. |
| What does it say about the Forward rate constant compared to the reverse rate constant when the forward rate of a reaction changes faster than the reverse rate of the reaction? | This indicates that the forward rate constant is greater then the reverse rate constant. |
| The forward rate constant is ______ than the reverse rate constant is the forward rate changes faster than the reverse rate. | greater |
| If there is only one reactant, the rate of the forward reaction will be directly proportional to the [A], thus what would happen if the rate of the forward reaction was reduced by more than half? A-->B | The [A] would also be reduced by more than half. FOr every molecule of A lost, a molecule of B is created, so at equilibrium [B] must be greater than [A]. Setting the rates equal, Kf[A]=Kr[B], we see that Kf must be greater than Kr if [B] is greater than [A] at equilibrium. |
| What is the rate of the forward reaction? A-->B | Rf=Kf[A] |
| What is the rate of the reverse reaction? A-->B | Rr=Kr[B] |
| If the rate constant is much greater for the forward reaction than it is for the reverse reaction, what can be said about the state of equilibrium? | For all practical purposes the reaction proceeds to completion |
| If a product is continually removed from a reaction, in the form of gas leaving an aqueous solution, the reaction can? | run to completion, because there will be no reverse reaction |
| At equilibrium, the forward rate of reaction ___ the reverse rate of reaction? A-->B | equal to. Thus, Rf=Rr... Kf[A]=Kr[B] |
| At equilibrium, if the [A] is less than [B], Kf ___ Kr? | At equilibrium, if [A] is less than [B], Kf > Kr |
| In a ________ mixture, with all species in the same phase, there will always be some of each species present at equilibrium. | homogenous |
| What are the forward and reverse rate laws for the following elementary reaction? 1A+2B-->3C+4D | Rate forward=Kf[A]¹[B]² Rate reverse=Kr[C]³[D]⁴ - we can only use the stoichiometric coefficients as exponents here because we were told that the reaction was elementary. However, for the MCAT, unless told it is elementary, we never use the coefficients as exponents. |
| Since equilibrium occurs when the forward rate law is equal to the reverse rate law, how do we derive the equilibrium constant, K? | Since, Rate forward=Kf[A]¹[B]² Rate reverse=Kr[C]³[D]⁴. At equilibrium, Rf=Rr, thus: Kf=Kr=Kf[A]¹[B]²=Kr[C]³[D]⁴... Therefore solve for Kf/Kr=[C]³[D]⁴/[A]¹[B]² kf/kr=K k=[C]³[D]⁴/[A]¹[B]² This simple relationship between K equilibrium and K rate is only true for elementary equations. |
| Law of Mass Action: | Kc=[C][D]/[A][B]=Products/Reactants -where * is the coefficient -The value of K has no dimensions because the concentrations are actually approximations for a dimensionless quantity called an activity. -The law of mass action is good for all chemical equations, including non-elementary equations. In other words, for equilibrium constants, K, use the chemical equation coefficients as the exponents of the concentrations regardless of molecularity. - The eq, constant is a capital K, the rate constant is a lower case k. |
| The rate constant is represented by a _____ case ___, the equilibrium constant is represented by an ____ case ___. | The rate constant is represented by a lower case k, the equilibrium constant is represented by an upper case K. |
| Regardless of whether or not the reaction is elementary, the equilibrium constant for the reverse reaction is the _______ of the equilibrium constant for the forward reaction. | reciprocal |
| The equilibrium constant for a series of reactions is _____ to the product of the equilibrium constants for each of its elementary steps. | equal to |
| What do the rate constants and equilibrium constants depend on? | temperature. The eq constant K, depends upon temp only. |
| When writing an equilibrium expression, do we use the concentration of a pure solid or pure liquid? | No, the concentrations of such species are given a value of 1 for the equilibrium expression, thus can essentially be ignored in the expression. - Pure solids can still participate in the equilibrium and when they do, they must be present in order for equilibrium to exist. Therefore, we do NOT use pure solid or liquids such as water in the law of mass action. |
| The equilibrium constant Kc is ________ dependent. | temperature dependent, it is not influenced by changes in pressure, volume or the presence of a catalyst. |
| For gases, the equilibrium constant is often expressed in... | Partial pressures(atmospheres) instead of concentrations. Kp is then used; Kp=PCPD/PAPB |
| Principle for detailed balance: | The principle of detailed balance is formulated for kinetic systems which are decomposed into elementary processes (collisions, or steps, or elementary reactions): At equilibrium, each elementary process should be equilibrated by its reverse process. Thus, at equilibrium, the forward and reverse reaction rates for each step must be equal, and any two or more single reactions or series of reactions resulting in the same products from identical reactants must have the same equilibrium constant for a given temperature. The equilibrium constant does not depend on whether or not other substances are present. |
| How are the partial pressure equilibrium constant Kp, and the concentration equilibrium constant Kc, related? | Kp=Kc(RT)* - where * is equal to the sum of the coefficients of the products minus the sum of the coefficients of the reactants. |
| The equilibrium constant only describes _______ conditions, where as the _______ ______ describes information about the direction of the reaction. | The equilibrium constant only describes equilibrium conditions and concentrations, where as the reaction quotient describes information about the direction of the reaction. |
| We can use the Reaction Quotient to predict? | To predict the direction in which the reaction will proceed. |
| What is the Reaction Quotient formula? | Q=Products/Reactants, - where * are the associated coefficients - Q can have any positive value and it is not constant |
| If Q=K, then | the reaction is at equilibrium |
| If Q<K, then | then Q must increase as the system moves toward equilibrium. Since the products are in the numerator of the Q expression, the reaction must be 'product-deficient' and must proceed forward/to the right to reach equilibrium. There are too many reactants relative to products. The forward reaction is favored |
| If Q>K, then | the Q must decrease as the system moves toward equilibrium. There is too much product relative to reactant. The reaction is 'reactant-deficient' and will proceed to the left/reverse in order to reach equilibrium. The reverse reaction is favored. |
| Le Chatelier's Principle: | When a system at equilibrium is stressed, the system will shift in a direction that will reduce that stress. |
| What are the 3 types of stress that usually obey Le Chatelier's Principle? | 1) Temperature 2) Pressure or Volume 3) Change in concentration of reactants or products 4) Catalysts allow a reaction to reach equilibrium more quickly, but do not alter the equilibrium position. No change in the value of K either. |
| For Le Chatelier's purposes we can think of heat as what? | A reactant for an endothermic reaction and a product for an exothermic reaction |
| Consider the following reaction: N₂+3H₂-->2NH₃+Heat, If we add N₂ or H₂ to the reaction at equilibrium, what happens? | Forward Reaction is Favored There will be a shift to the right, therefore there will be more NH₃+Heat created, in this exothermic reaction. |
| Consider the following reaction: N₂+3H₂-->2NH₃+Heat What happens to the reaction if we add heat or NH₃ at equilibrium? | Reverse Reaction is Favored There will be a shift to the left, therefore creating more N₂+H₂ |
| Consider the following reaction: N₂+3H₂-->2NH₃+Heat If this reaction is taking place in an enclosed container, what happens to the reaction if the size of the container is reduced at constant temperature? | Total pressure increases while volume decreases. Since there are 4 gas molecules on the left and only 2 on the right, the reaction will proceed to the right producing more heat and NH₃. Thus, the equilibrium shifts to the side of the reaction with less moles of gas. This also occurs in a reaction wherein the solution is concentrated or diluted, the equilibrium will shift to the side with fewer moles when the solution is concentrated. |
| Describe a few situations when Le Chatelier's Principle will not predict the correct shift in reaction equilibrium. | 1) Solvation reactions: The solubility of salts generally increases with increasing temperature, even when reaction is exothermic. This is due to the significant entropy increase that occurs with dissolution. The entropy factor becomes more important as the temperature increases. 2) Pressure increases due to the addition of a nonreactive gas: If He, an inert gas, is added to the reaction: N₂+H₂-->2NH₃+Heat the equilibrium position is not affected, even though total pressure has increased. This is because adding He to a rigid container does not change the partial pressures of the other gases, so the equilibrium does not shift. |
| If volume decreases, the total pressure would increase and equilibrium would be shifted to? | the side of the reaction with fewer moles of gas in order to decrease pressure There is no change in the value of K with pressure changes. |
| When has the Gibbs free energy of a reaction reached its minimum? | at equilibrium |
| Gibbs Free Energy: | A thermodynamic potential that measures the "useful" or process-initiating work obtainable from a thermodynamic system at a constant temperature and pressure (isothermal, isobaric). Just as in mechanics, where potential energy is defined as capacity to do work, similarly different potentials have different meanings. The Gibbs free energy is the maximum amount of non-expansion work that can be extracted from a closed system; this maximum can be attained only in a completely reversible process. When a system changes from a well-defined initial state to a well-defined final state, the Gibbs free energy ΔG equals the work exchanged by the system with its surroundings, minus the work of the pressure forces, during a reversible transformation of the system from the same initial state to the same final state. |
| Enthalpy: | Enthalpy is a measure of the total energy of a thermodynamic system. It includes the internal energy, which is the energy required to create a system, and the amount of energy required to make room for it by displacing its environment and establishing its volume and pressure. |
| For a reverse reaction, the equilibrium constant Kc, is the _____ of the equilibrium constant Kc, for the forward reaction. | inverse, the products are now the reactants and the reactants are now the products. |
| If the coefficients of a balanced reaction are multiplied by a given factor, the equilibrium constant is taken to the _______ of that same factor. | power Ex) If reaction 1, Kc1, is multiplied by 3, the resulting Kc2=(Kc1)³ |
| If the coefficients of a balanced reaction are divided by a given factor, the equilibrium constant is taken to the _____ of that same factor. | Root, Ex) If reaction 1, Kc1, is divided by 4, the resulting Kc2=(Kc1)^1/4=⁴√Kc1 |
| When multiple reactions are combined, the equilibrium constants of the reaction are? | Multiplied together. Ex) Reaction 1= Kc1 Reaction 2=Kc2 Therefore, Overall Reaction=Kc1 x Kc2=Kc |
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