Reading...
Play button
Play button
Use LEFT and RIGHT arrow keys to navigate between flashcards;
Use UP and DOWN arrow keys to flip the card;
H to show hint;
A reads text to speech;
36 Cards in this Set
- Front
- Back
|
Standard temperature and pressure (STP)
|
Temp: 0 degrees Celsius = 273 K
Pressure: 1 atm (atmosphere) moles: 1 mole |
|
|
Mean free path
|
Distance traveled by a gas molecule between collisions
|
|
|
Ideal gas
|
Based on kinetic molecular theory
1. zero volume 2. no forces other than repulsive forces due to collisions 3. elastic collisions 4. kinetic energy is proportional to temperature of gas |
|
|
Kinetic molecular theory
|
Ideal gas lack certain real gas characteristics
|
|
|
Universal gas constant?
|
R = .08206 L*atm/K*mol
R = 8.314 J/K*mol |
|
|
Ideal gas law
|
PV = nRT
P: pressure V: volume n: number of moles R: universal gas constant = 8.314 J/Kmol T: temperature in kelvin |
|
|
Standard molar volume at STP
|
22.4 Liters
At STP, 1 mol of ANY ideal gas will occupy 22.4 L. |
|
|
Mole fraction (X)
|
number of mole of gas "a" divided by total number of moles of gas in sample
|
|
|
Partial pressure
|
Total pressure of gaseous mixture times mole fraction of particular gas
P(a) = (Xa) (Ptotal) P(a): partial pressure of gas "a" Xa: mole fraction of gas "a" (Mole fraction is # of moles of gas "a"/ total # of moles of gas in sample) Ptotal: total pressure of gaseous mixture |
|
|
Dalton's Law
|
Total pressure exerted by a gaseous mixture is the sum of the partial pressures of each of its gases
Ptotal = P1 + P2 + P3... |
|
|
Average translational kinetic energy
|
KE = (3/2) RT
KE: kinetic energy, found from RMS velocity This is the avg kinetic energy for one mol of gas molecules at a given moment |
|
|
Graham's Law
|
Ratio of RMS velocities of 2 gases in homogeneous mixture
V1/V2 = square root (m2)/square root (m1) V: RMS velocities m: mass of gas molecules Graham's law describes effusion and diffusion |
|
|
Effusion
|
Spreading of gas from high pressure to low pressure through a pinhole
effusion rate 1/effusion rate 2 = square root (M2)/square root (M1) M: molecular weights |
|
|
Diffusion
|
Spreading of one gas into another gas, or into empty space
diffusion rate 1/diffusion rate 2 = square root (M2)/square root (M1) |
|
|
How real gases deviate from ideal gases
|
Vreal > Videal
Intermolecular forces exist Preal < Pideal |
|
|
Collision model
|
In order for a reaction to occur, reacting molecules must collide
Rate of reaction is much lower than frequency of collisions Most collisions do not result in reaction |
|
|
Activation Energy (Ea)
|
Threshold energy required for collisions to create new molecules in a reaction
independent of temperature Rate of reaction (K) increases with increasing temperature Rate constant (kf) increases with increasing temperature??? |
|
|
Factors affecting rate of reaction
|
1. temperature
2. pressure (negligible effect) 3. concentration of reactants |
|
|
Elementary reaction's coefficients: can they be used to determine rate law?
|
Yes! Coefficient tells you how many molecules participate in a reaction producing collision
|
|
|
Intermediates
|
Species that are products of one reaction and reactants of a later reaction in a reaction chain
Concentration is very low because they are often unstable and react as quickly as they are formed |
|
|
Rate Law
|
Rate (forward) = Kf [A]^a [B]^b
Kf: rate constant (not rate of reaction) a: order of each respective reactant b: order of each respective reactant a + b: overall order of reaction Assumes no reverse reaction |
|
|
1st Order reaction
|
A --> product
rate = Kf[A] [A] decreases exponentially ln[A] vs. time graph: straight line slope: -Kf Constant half-life independent of concentration [A] No collisions take place in theory |
|
|
2nd Order Reaction (1 reactant)
|
2A --> products
rate = Kf[A]^2 1/[A] vs time graph: straight line slope: Kf Half-life dependent upon concentration [A] |
|
|
2nd Order Reaction (2 reactants)
|
A + B --> products
Rate = Kf[A][B] no predictable half-life |
|
|
3rd Order Reaction
|
3A --> products
Rate = Kf[A]^3 1/2[A]^2 vs. time graph: straight line slope: Kf |
|
|
Zero Order Reaction
|
[A] vs. time graph: straight line
slope: -Kf |
|
|
Rate determining step
|
Rate of slowest elementary step determines the rate of the overall reaction
|
|
|
Catalyst
|
Substance that increases the rate of reaction without being consumed or permanently altered
Enhances product selectivity Reduces energy consumption Lowers activation energy (Ea) Increases steric factor (p) Creates a new reaction pathway which typically includes an intermediate Cannot alter equilibrium constant (K) |
|
|
Heterogeneous catalyst
|
Different phase than reactants and products
Usually solids while reactants and products are liquids or gases Reaction rates can be enhanced by increasing surface area of catalyst |
|
|
Homogeneous catalyst
|
same phase as reactants and products, usually in gas or liquid phase
Aqueous acid or base solutions Autocatalysis = generate catalyst as product |
|
|
Chemical equilibrium
|
Rate of forward reaction equals rate of reverse reaction
No change in concentration of products or reactants Point of greatest entropy Reactions always move toward equilibrium, therefore Q will always change towards K |
|
|
Law of mass action
|
the relationship between the chemical equation and the equilibrium constant
K = 1 for solids or pure liquids like water |
|
|
Equilibrium constant (K)
|
Described by the law of mass action
Relationship between a chemical equation and the equilibrium constant aA + bB --> cC + dD K = ([C]^c [D]^d)/([A]^a [B]^b) = (products^coefficients)/(reactants^coefficients) K only depends upon temperature!!! Equilibrium constant for a series of reactions is equal to the product of the equilibrium constants for each of its elementary steps Pure solid or liquid is K=1, therefore not included in the law of mass action equation |
|
|
Reaction quotient (Q)
|
For reactions not at equilibrium. It is used to predict
Q = (products^coefficients)/(reactants^coefficients) Q is not constant, it can be any positive value Use Q to predict the direction in which a reaction will proceed Since reactions always move toward equilibrium, Q will always change towards K |
|
|
Comparison of Q (reaction quotient) & K (equilibrium constant)
|
Q = K: reaction at equilibrium
Q > K: leftward shift because there is more product than reactant Q < K: rightward shift because there is more reactant than product |
|
|
Le Chatelier's Principle
|
When a system at equilibrium is stressed, the system will shift in the direction that will reduce that stress
Types of stress: 1. addition or removal of a product or reactant 2. changing the pressure of the system 3. heating or cooling the system |
