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38 Cards in this Set
- Front
- Back
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Standard temperature and pressure (STP)
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Temp: 0 degrees Celsius = 273 K
Pressure: 1 atm (atmosphere) moles: 1 mole |
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Mean free path
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Distance traveled by a gas molecule between collisions
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Ideal gas
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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 |
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Kinetic molecular theory
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Ideal gas lack certain real gas characteristics
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Ideal gas law
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PV = nRT
P: pressure V: volume n: number of moles R: universal gas constant = 8.314 J/Kmol T: temperature in kelvin |
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Charles' Law
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Volume of gas is proportional to temperature at constant pressure
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Boyle's Law
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Volume of gas is inversely proportional to pressure a constant temperature
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Avogadro's Law
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Volume of gas is proportional to number of moles at constant temperature and pressure
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Standard molar volume at STP
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22.4 Liters
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Partial pressure
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Total pressure of gaseous mixture times mole fraction of particular gas
Pa = (Xa) (Ptotal) Pa: partial pressure of gas "a" Xa: mole fraction of gas "a" Ptotal: total pressure of gaseous mixture |
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Mole fraction (X)
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number of mole of gas "a" divided by total number of moles of gas in sample
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Dalton's Law
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Total pressure exerted by a gaseous mixture is the sum of the partial pressures of each of its gases
Ptotal = P1 + P2 + P3... |
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Average translational energy of ideal gas
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KE = (3/2) RT
KE: kinetic energy, found from RMS velocity R: universal gas constant T: temperature |
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Graham's Law
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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 Describes effusion and diffusion |
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Effusion
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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 |
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Pinhole
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opening much smaller than average distance between gas molecules
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Diffusion
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Spreading of one gas into another gas or empty space
diffusion rate 1/diffusion rate 2 = square root (M2)/square root (M1) |
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How real gases deviate from ideal gases
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Vreal > Videal
Intermolecular forces exist Preal < Pideal |
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Collision model
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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 |
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Activation Energy
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Threshold energy required for collisions to create new molecules in a reaction (Ea) - independent of temperature
Rate of reaction (K) increases with temperature |
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Factors affecting rate of reaction
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1. temperature
2. pressure (negligible effect) 3. concentration of reactants |
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Elementary reaction
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Coefficient tells you how many molecules participate in a reaction producing collision
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Intermediates
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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 |
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Rate Law
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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 Assume no reverse reaction |
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1st Order reaction
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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 |
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2nd Order Reaction (1 reactant)
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2A --> products
rate = Kf[A]^2 1/[A] vs time graph: straight line slope: Kf Half-life dependent upon concentration [A] |
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2nd Order Reaction (2 reactants)
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A + B --> products
Rate = Kf[A][B] no predictable half-life |
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3rd Order Reaction
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3A --> products
Rate = Kf[A]^3 1/2[A]^2 vs. time graph: straight line slope: Kf |
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Zero Order Reaction
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[A] vs. time graph: straight line
slope: -Kf |
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Rate determining step
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Rate of slowest elementary step determines the rate of the overall reaction
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Catalyst
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Substance that increases the rate of reaction without being consumed or permanently altered
Enhance product selectivity Reduce energy consumption Lower activation energy (Ea) Increase steric factor (p) Creates a new reaction pathway which typically includes an intermediate Cannot alter equilibrium constant (K) |
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Heterogeneous catalyst
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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 |
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Homogeneous catalyst
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same phase as reactants and products, usually in gas or liquid phase
Aqueous acid or base solutions Autocatalysis = generate catalyst as product |
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Chemical equilibrium
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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 |
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Equilibrium constant (K)
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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) Only depends upon temperature Equilibrium constant for a series of reaction is equal to the product of the equilibrium constants for each of its elementary steps Pure solid or liquid is K=1, therefore no included in the law of mass action equation |
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Reaction quotient (Q)
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For reactions not at equilibrium
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 |
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Comparison of Q & K
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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 |
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Le Chatelier's Principle
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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 |
