G R O U P 4

Mole-Volume Stoichiometry
• You are given the moles of one component and needed to find the volume of another gaseous component. The temperature and pressure must be stated in a problem such as this.The following steps are applied:
1) Convert moles of given to moles of needed using the coefficients of the balanced chemical equation
2) Convert moles of needed to volume of needed using the
Ideal Gas Law Equation

• If you are given the volume of a gaseous component and asked to find the number of moles involved of some other component, then you would simply reverse the two steps above.(First use the formula PV = nRT, to convert volume of the gaseous component to moles of gaseous component and then convert moles of gaseous component to moles of other component by using the mole to mole ratio from the balanced equation.)

Mass-Volume Stoichiometry Problem
• Here you are given the mass of one of the components and asked to find the Volume of a gaseous component at a stated temperature and pressure. Here are the steps that one would take:
1) Convert mass of given to moles of given by dividing by the formula mass of the given
2) Convert moles of given to moles of needed using the balanced equation
3) Convert moles of needed to liters of requested using PV = nRT

• The volume-volume problems are the easiest since according to the Law of Combining Gas Volumes, gases combine at the same temperature and pressure in simple whole number of volumes. What this means is that we can use the coefficients in the balanced equation to form volume relationships just as we did in the earlier
Stoichiometry problems when we used the coefficients to form mole relationships.

G R O U P 4

• assumes that a model can be used to explain why gases behave the way they do.
• The experimental observations about the behaviour of gases discussed so far can be explained with a simple theoretical model known as the kinetic molecular theory. This theory is based on the following postulates, or assumptions.

• Gases are made of very small molecules, which are separated by a very great distance between them. The dimension of the molecules is very much smaller than the distance between them.
• Because of this very great distance between them, the force of attraction between the molecules is negligible. The molecules are independent of each other.
• The molecules are in constant motion, moving in randomly in all directions.
• Due to the great number of molecules and their random motion, it is unavoidable that the molecules will collide with each other and with the walls of the container.
• During these collisions, there is no change in the momentum of the molecules
• The average kinetic energy of the molecules is determined only by the absolute temperature of the gas

• The average kinetic energy of a collection of gas particles depends on the temperature of the gas and nothing else.
• Any object in motion has a kinetic energy that is defined as onehalf of the product of its mass times its velocity squared.
KE = 1 /
2 mv 2

Kinetic Molecular Theory

a. Gases are composed of molecules. The distances from molecule to molecule are far greater than the molecules dimensions. These molecules can be considered as spherical bodies’ which possess negligible mass and volume.

b. Gas molecules are always in constant random motion and they frequently collide with each other and with the walls of the container. Collisions among molecules are perfectly elastic, that is, energy may transfer from molecule to molecule as the result of collision, but the total energy of the molecules in the system remains the same or constant.

c. There is neither attractive nor repulsive force between or among gas molecules.
d. Movement of gas molecules is affected by temperature. The average kinetic of the molecules is directly related to the temperature of gas.

• Kinetic explanation of Boyle’s Law: Boyle's law is easily explained by the kinetic molecular theory. The pressure of a gas depends on the number of times per second that the molecules strike the surface of the container. If we compress the gas to a smaller volume, the same number of molecules are now acting against a smaller surface area, so the number striking per unit of area, and thus the pressure, is now greater.
• Kinetic explanation of Charle’s Law: Kinetic molecular theory states that an increase in temperature raises the average kinetic energy of the molecules. If the molecules are moving more rapidly but the pressure remains the same, then the molecules must stay farther apart, so that the increase in the rate at which molecules collide with the surface of the container is compensated for by a corresponding increase in the area of this surface as the gas expands.

• Kinetic explanation of Avogadro's law: If we increase the number of gas molecules in a closed container, more of them will collide with the walls per unit time. If the pressure is to remain constant, the volume must increase in proportion, so that the molecules strike the walls less frequently, and over a larger surface area.
• Kinetic explanation of Dalton’s law: "Every gas is a vacuum to every other gas". This is the way Dalton stated what we now know as his law of partial pressures. It simply means that each gas present in a mixture of gases acts independently of the others. This makes sense because of first fundamental tenet of KMT theory is that gas molecules have negligible volumes. So Gas A in mixture of A and B acts as if Gas B were not there at all. Each contributes its own pressure to the total pressure within the container, in proportion to the fraction of the molecules it represents