General Chemistry Latest Exam |Questions
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Ideal Gas (Kinetic Molecular Theory) - ✔️✔️An ideal gas has the following 4 characteristics that
are different from a real gas:
1) Gas molecules have zero volume.
2) Gas molecules exert no forces other than repulsive forces due to collisions.
3) Gas molecules make completely elastic colisions.
4) The average KE of gas molecules is directly proportional to the temperature of the gas.
Ideal Gas Law - ✔️✔️*PV = nRT*
P is in atmospheres (atm)
V is in liters (L)
n is the number of moles of gas
T is in Kelvin (K)
R is the universal gas constant 0.08206 L*atm/K*mol or 8.314 J/K*mol (for differences in units,
recall: pressure = energy/volume --> P = nRT/V so nRT = energy!)
Can also think of temperature as (kinetic) energy/mol.
Partial Pressure - ✔️✔️The total pressure of the gaseous mixture times the mole fraction of the
particular gas (# of moles of gas 'a' divided by the total # of moles of gas in the sample).
--> The pressure of a gas as if it were alone in a container.
P(a) = X(a)P(total)
Dalton's Law - ✔️✔️The total pressure exerted by a gaseous mixture is the sum of the partial
pressures of each of its gases.
,P(total) = P(1) + P(2) + P(3) + ....
Average Kinetic Energy of a Gas - ✔️✔️KE(avg) = (3/2)RT
Where the average translational KE is found fromt he root-mean-square (rms) velocity.
-> rms velocity is the square root of the average squares of the molecular velocities; slightly
greater than average speed.
-> For one mole of gas (not individual molecules).
-> Valid for any fluid system (gases and liquids).
Graham's Law - ✔️✔️In a sample of gas, the KE of the molecules will vary but there will be an
average KE of the molcules that is proportional to the temperature and independent of the
type of gas. Two different molecules in a gaseous mixture will have the same KE but have
different masses so they must have different rms velocities (KE = 1/2mv^2). By setting their KEs
equal to each other, we can derive a relationship between their rms velocities.
v1/v2 = √m2/√m1
Effusion - ✔️✔️The spreading of a gas from high pressure to very low pressure through a
'pinhole'.
-> Molecules with higher rms velocity (lower molecular weight, per p = mv) will experience
more collisions with the walls of a container so the rate at which molecules find the pinhole and
go through is likely to be greater.
(effusion rate 1)/(effusion rate 2) = √M2/√M1 [= v1/v2]
--> From Graham's Law.
,Diffusion - ✔️✔️The spreading of one gas into another gas or into empty space.
(diffusion rate 1)/(diffusion rate 2) = √M2/√M1 [= v1/v2]
--> From Graham's Law.
The diffusion rate is much slower than the rms velocity of the molecules bcause gas molecules
collide with each other and molecules of other gases as they diffuse.
Real Gases - ✔️✔️Deviate from ideal behavior when molecules are close together - low V caused
by high P or low T.
Since molecules of a real gas do have volume, their volume must be added to the ideal volume.
V(real) > V(ideal), where V(ideal) is calculated from PV=nRT.
Since molecules in a real gas do exhibit forces on each other (predominantly attractive), gas
molecules are pulled inward toward the center of the gas and slow before colliding with
container walls.
P(real) < P(ideal), where P(ideal) is calculated from PV=nRT.
--> From PV=nRT, we expect PV/RT to equal one for one mole of gas at any T or P. Since V
deviates positively from ideal behavior and P deviates negatively:
-If PV/RT > 1 for one mole of gas, then the deviation due to molecular V must be greater than
the deviation due to the intermolecular forces.
-If PV/RT < 1 for one mole of gas, then the deviation due to intermolecular forces must be
greater than the deviation due to molecular V.
Chemical Kinetics - ✔️✔️Deals with the rate of the reaction as it moves towards equilibrium; tells
us how fast equilibrium is achieved.
, Chemical Thermodynamics - ✔️✔️Deals with the balance of the reactants and products after
they have achieved equilibrium; tells us what equilibrium looks like. Based on probabilities and
are valid mostly for systems with a large number of molecules (macroscopic).
-> Divide the universe into systems and surroundings.
-> Uses two different types of properties to describe *state of a system; extensive and
intensive.
Collision Model - ✔️✔️Most molecular collisions do not result in a reaction. A collision must have
2 requirements to create new molecules in a reaction:
1) The relative KEs (due to relative velocity only) of the colliding molecules must reach the
threshold energy -> "activation energy".
Note: velocity in a diretion away from another molecule decreases the relative KE of a collision.
2) The colliding molecules must have the proper spatial orientation.
-> Consider activation energy to be independent of temperature. However, at higher T, more
collisions will reach the activation energy.
Reaction Rate - ✔️✔️Tells us how quickly the concentration of a reactant or product is changing.
Typically given in molarity per second (mol*L/s). Factors affecting the rate of a reaction are:
-Temperature
-Pressure (usually small enough to be ignored)
-Concentration of certain substances in the reacting system
Effects of Temperature on Reaction Rate - ✔️✔️The rate of a reaction increases with
temperature, mainly because more collisions with sufficient relative KE occur each second.
*Do not confuse kinetics with thermodynamics! Increasing the rate is not necessarily a
statement about the equilibrium. It simply means that the equilibrium is achieved more quickly.
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Ideal Gas (Kinetic Molecular Theory) - ✔️✔️An ideal gas has the following 4 characteristics that
are different from a real gas:
1) Gas molecules have zero volume.
2) Gas molecules exert no forces other than repulsive forces due to collisions.
3) Gas molecules make completely elastic colisions.
4) The average KE of gas molecules is directly proportional to the temperature of the gas.
Ideal Gas Law - ✔️✔️*PV = nRT*
P is in atmospheres (atm)
V is in liters (L)
n is the number of moles of gas
T is in Kelvin (K)
R is the universal gas constant 0.08206 L*atm/K*mol or 8.314 J/K*mol (for differences in units,
recall: pressure = energy/volume --> P = nRT/V so nRT = energy!)
Can also think of temperature as (kinetic) energy/mol.
Partial Pressure - ✔️✔️The total pressure of the gaseous mixture times the mole fraction of the
particular gas (# of moles of gas 'a' divided by the total # of moles of gas in the sample).
--> The pressure of a gas as if it were alone in a container.
P(a) = X(a)P(total)
Dalton's Law - ✔️✔️The total pressure exerted by a gaseous mixture is the sum of the partial
pressures of each of its gases.
,P(total) = P(1) + P(2) + P(3) + ....
Average Kinetic Energy of a Gas - ✔️✔️KE(avg) = (3/2)RT
Where the average translational KE is found fromt he root-mean-square (rms) velocity.
-> rms velocity is the square root of the average squares of the molecular velocities; slightly
greater than average speed.
-> For one mole of gas (not individual molecules).
-> Valid for any fluid system (gases and liquids).
Graham's Law - ✔️✔️In a sample of gas, the KE of the molecules will vary but there will be an
average KE of the molcules that is proportional to the temperature and independent of the
type of gas. Two different molecules in a gaseous mixture will have the same KE but have
different masses so they must have different rms velocities (KE = 1/2mv^2). By setting their KEs
equal to each other, we can derive a relationship between their rms velocities.
v1/v2 = √m2/√m1
Effusion - ✔️✔️The spreading of a gas from high pressure to very low pressure through a
'pinhole'.
-> Molecules with higher rms velocity (lower molecular weight, per p = mv) will experience
more collisions with the walls of a container so the rate at which molecules find the pinhole and
go through is likely to be greater.
(effusion rate 1)/(effusion rate 2) = √M2/√M1 [= v1/v2]
--> From Graham's Law.
,Diffusion - ✔️✔️The spreading of one gas into another gas or into empty space.
(diffusion rate 1)/(diffusion rate 2) = √M2/√M1 [= v1/v2]
--> From Graham's Law.
The diffusion rate is much slower than the rms velocity of the molecules bcause gas molecules
collide with each other and molecules of other gases as they diffuse.
Real Gases - ✔️✔️Deviate from ideal behavior when molecules are close together - low V caused
by high P or low T.
Since molecules of a real gas do have volume, their volume must be added to the ideal volume.
V(real) > V(ideal), where V(ideal) is calculated from PV=nRT.
Since molecules in a real gas do exhibit forces on each other (predominantly attractive), gas
molecules are pulled inward toward the center of the gas and slow before colliding with
container walls.
P(real) < P(ideal), where P(ideal) is calculated from PV=nRT.
--> From PV=nRT, we expect PV/RT to equal one for one mole of gas at any T or P. Since V
deviates positively from ideal behavior and P deviates negatively:
-If PV/RT > 1 for one mole of gas, then the deviation due to molecular V must be greater than
the deviation due to the intermolecular forces.
-If PV/RT < 1 for one mole of gas, then the deviation due to intermolecular forces must be
greater than the deviation due to molecular V.
Chemical Kinetics - ✔️✔️Deals with the rate of the reaction as it moves towards equilibrium; tells
us how fast equilibrium is achieved.
, Chemical Thermodynamics - ✔️✔️Deals with the balance of the reactants and products after
they have achieved equilibrium; tells us what equilibrium looks like. Based on probabilities and
are valid mostly for systems with a large number of molecules (macroscopic).
-> Divide the universe into systems and surroundings.
-> Uses two different types of properties to describe *state of a system; extensive and
intensive.
Collision Model - ✔️✔️Most molecular collisions do not result in a reaction. A collision must have
2 requirements to create new molecules in a reaction:
1) The relative KEs (due to relative velocity only) of the colliding molecules must reach the
threshold energy -> "activation energy".
Note: velocity in a diretion away from another molecule decreases the relative KE of a collision.
2) The colliding molecules must have the proper spatial orientation.
-> Consider activation energy to be independent of temperature. However, at higher T, more
collisions will reach the activation energy.
Reaction Rate - ✔️✔️Tells us how quickly the concentration of a reactant or product is changing.
Typically given in molarity per second (mol*L/s). Factors affecting the rate of a reaction are:
-Temperature
-Pressure (usually small enough to be ignored)
-Concentration of certain substances in the reacting system
Effects of Temperature on Reaction Rate - ✔️✔️The rate of a reaction increases with
temperature, mainly because more collisions with sufficient relative KE occur each second.
*Do not confuse kinetics with thermodynamics! Increasing the rate is not necessarily a
statement about the equilibrium. It simply means that the equilibrium is achieved more quickly.
