CHEM 1411 - General Chemistry I

Course (Lecture and Lab) Syllabus

Summer 2007

 

 

Instructor: Dr. Shawn Amorde

e-mail address: samorde@austincc.edu

web page: http://www.austincc.edu/samorde

Office Hours: TThF, 7-8am

 

 

 

Gases and the Kinetic-Molecular Theory

 

Text Chapters 12

Recommended Problems (2, 6, 10, 14, 18, 20, 26, 28, 34, 40, 44, 46, 48, 50, 54)

 

            Gases

 

1.      Composition of the Atmosphere and Properties of Gases

2.      Pressure

3.      Boyle’s Law

4.      Charles’s Law

5.      Standard Temperature and Pressure

6.      The Combines Gas Law Equation

7.      Avogadro’s Law

8.      Ideal Gas Equation

9.      Molecular Weights and Molecular Formulas

10.  Dalton’s Law

11.  Mass – Volume Relationships

12.  The Kinetic-Molecular Theory

13.  Diffusion and Effusion of Gases

14.  Deviations from the Ideal Gas Equation

 

 

Gases

 

Let’s review some of the general physical properties of gases.  Gases and liquids are known as fluids because they flow freely and are significantly less dense than solids.  Gases are the least dense of all the states of matter and fill up a container entirely.  These properties tell us atoms or molecules in a gaseous state are far apart and interactions between them are weak.  A large number of liquids are converted to gas states upon heating and vice versa, the temperature at which a liquid becomes a gas is the liquids boiling point.  A low boiling point reflects the liquids ability to vaporize readily.

 

Compostion of the Atmosphere and Some Common Properties of Gases

 

 

Abundance of Atmospheric Gases

 

 

 

 

 

We are basing our laws on the work of four major scientists, Toricelli (1643), Boyle (1660), Charles (1787), Graham (1831). 

 

Here are some of the major properties of gases;

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Pressure

 

The official definition of pressure is force per unit area, i.e. pounds per square inch, or psi.  A device used for measuring atmospheric pressure is called a barometer.  The most common barometers are mercury barometers.  Let’s draw one.

 

 

 

 

 

 

 

 

 

 

 

Pressures are often expressed in mm Hg, millimeters of mercury.  One torr is equal to 1 mm Hg.

 

 

A manometer is used to compare the pressure exerted by an unknown gas compared to atmospheric pressure.  Let’s draw one.

 

 

 

 

 

 

 

 

 

 

 

We need to keep in mind that atmospheric pressure is location dependant!!  Remember so is weight!!  At sea level, the atmospheric pressure raises a column of mercury 760 mm, or 760 mm Hg which is 760 torr.

 

 

Just like our conversion factors, we have an SI unit for pressure which is a Pascal.  Here is the definition of a Pascal.

 

 

 

 

 

 

 

 

Boyle’s Law

 

Using a manometer Robert Boyle (1627-1691), studied the pressures exerted by several gases.  He then added more Hg to the tube and measured the changes in height.  What he discovered was that pressure and volume are inversely proportional and the product of pressure and volume is always the same number.

 

 

 

 

 

 

 

This relationship is Boyle’s Law.  It is important to remember that k is dependant on the number of moles of gas and the temperature.  These are left out of the law, assuming standard temperatures and comparing the same amounts.

Most gases follow Boyle’s Law at normal temperatures and pressures.

In fact we can compare a fixed amount of gas at a constant temperature.

 

 

 

 

 

 

 

Let’s do a problem.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Charles’s Law

 

During his studies Boyle noted that heating the sample changed the volume, but he did not follow up on this observation.  Jacques Charles (1746-1823) and Joseph Gay-Lussac (1778-1850), balloonists, studied the relationship between the expansion of a gas upon heating.  This is the basis for the absolute temperature scale we have been using.  Lord Kelvin noticed by plotting temperature versus volume of different gases leads back to a common intercept, -273.15 K.  The volume-temperature relationship is known as Charles’s Law.

 

 

 

 

 

 

 

 

 

Charles’s Law is defined as; at constant pressure, the volume occupied by a definite mass of a gas is directly proportional to its absolute temperature.

 

Let’s do a problem.

 

 

 

 

 

 

 

 

 

 

 

 

Boyle’s Law relates pressure and volume and Charles’s Law relates volume to temperature and we can combine these equations to form the Combined Gas Law Equation.  It is important to remember that this is for a constant amount of gas.

 

 

 

 

 

 

 

Let’s work a problem.

 

 

 

 

 

 

 

 

 

 

Back to Avogadro, in 1811, he postulated that at the same temperature and pressure, equal volumes of all gases contain the same number of molecules!!

 

This is Avogadro’s Law

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

The volume occupied by a mole of gas at STP, standard temperature and pressure, is called the standard molar volume.  This is constant for nearly all gases.

 

 

 

 

 

 

 

Now let’s put all three together.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

This is called the Ideal Gas Law.

 

 

 

 

 

Let’s work a problem.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

We can use the Ideal Gas Law to calculate the molecular weights of gaseous molecules. 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

We can also use the Ideal Gas Law to calculate the molecular formula of a gaseous compound.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

But, what if we have a mixture of gases?  How are all these laws applied?

 

Dalton’s Law of Partial Pressures

 

First we need to calculate the total number of moles in the gaseous mixture.

 

 

 

 

Then we can apply this to the Ideal Gas Law.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Let’s work a problem.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

We can now apply these Gas Laws to chemical reactions.  Let’s consider the gas forming reactions we just talked about.  We can calculate the mass of the gas produced by calculating the volume and vice versa.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

We already know that reactions proceed with conservation of matter and therefore we produce specific amounts of products, the same is true with gaseous reactions.  Instead of moles to grams we use moles to volume (liters). 

This is summed up in the Gay-Lussac’s Law of Combining Volumes.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Let’s work a problem.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

In 1738, Daniel Bernoulli (1700-1782) envisioned gaseous molecules colliding against the walls of their container and exerting pressure.  Later, 1857, Clausius, attempted to explain everything known about gases so far…..

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

These summations are known as the kinetic-molecular theory.

 

 

Let’s look at the kinetic energy of gaseous molecules in relation to each Gas Law.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Diffusion and Effusion of Gases

 

Because gas molecules are in constant motion they will fill any container over time, this is called diffusion.  For example, if I open a vial of a foul smelling gas, you will all smell it eventually.

 

If the container has porous walls a gaseous substance will pass through the tiny openings, this is called effusion. 

 

 

 

 

Non-Ideal Gas Behavior

 

What are non-ideal conditions?

What about gases at very low temperature or very high pressure?

 

 

 

 

 

 

 

The van der Waal’s equation.