Chapter 11: Gases

Matter is paticulate and ts behavior can be understood in terms of particles

7.1 Describe measuring gas pressures using barometers and manometers. Convert between pressure units. (11.1 and 11.2) Explain and measure gas pressure. Relate pressure units. 

• The behavior of gases can be explained and predicted by a model called the kinetic molecular theory. The core of this model is that gases are composed of particles in constant motion.
• 1.) Observations on the properties of gases
• 2.) Observations led to scientific laws (that summarized the observed behavior)
• 3.) The observations and laws led to the development of the kinetic molecular theory-the model the nature of a gas. 
• Observations---> Scientific Laws ---> Kinetic Molecular Theory
• Gases:
• Have no definite shape and no definite volume for gases will expand to fill any container. 
• Characteristics of Gases:
• Compressibility: Change in volume of a sample resulting from a pressure change acting on the sample. 
• Solids and gases are not significantly compressible but gases can be compressed under sufficient pressure. 
• Gas molecules are widely spread and therefore have low densities. 
• Pressure: Particles in a gas collide with each and with the surfaces around them. Each collisions exerts only a small force, but when the forces of many particles are summed, they quickly add up. 
• The result of the constant collision between the atoms and molecules in a gas and the surfaces around them is pressure. 
• The pressure that a gas sample exerts is the force that results from the collisions of gas particles divides by the area of the surface with which they collide.
• Pressure= (Force/ Area) 
• P= (F/A)
• Pressure is the force (F) that acts on a given area (A)
• Atmospheric Pressure: The pressure that is exerted upon the earth's surface by the atmosphere. 
• Variation in pressure in Earth's atmosphere create wind, and changes in pressure help us to predict weather. 
• The number of gas particles in a given volume decreases with increasing altitude (See Below)*
• Pressure decreases with increasing altitude
• Pressure Observations:
• The total pressure exerted by a gas depends on several factors such as:
• Concentration: Number of gas particles in a given volume. The lower the concentration (fewer gas particles), the lower the pressure. Likewise, a high density of gas partiles results in high pressure. 
• Altitude *: Because the number of gas particles in a given volume decreases with increasing altitude, pressure decreases with increasing altitude. 
• Ex. In the air, airplane cabins have to be artificially pressurized due to lack of oxygen at high altitudes. 
• Ex. When you ascend in the plane or hike up a mountain, the external pressure (the pressure that surrounds you) drops, while the pressure within the ear cavity (internal pressure) stays the same. This imbalance, creates the ear popping effect. 
• Volume of the Container: A low density of gas particles results in low pressure.
• Average Speed of the Gas Particles: A high density of gas particles results in high pressure. 
• SI Units of  Pressure:
• Force: 
• (N) newton
• 1 N= 1 kg x m/s2
• Pressure:
• (Pa) Pascal
• 1 Pa=1 N/m2
• Non SI Units of Pressure that are Commonly Used:
• atmosphere (atm)
• millimeter of mercury (mm Hg)
• torr
• Note that 1 mmHg=1 torr

Manometer: The instrument used to measure the pressure of a gas sample in the laboratory. They are U shaped tubes partially filled with a liquid that are connected to the gas sample on one side and open to the air on the other.

7.2 Apply the ideal gas law to relate and calculate values for pressure, volume, temperature, and amount of a gas (11.3-11.5) Apply the ideal gas law to relate and calculate values for pressure, temperature, and amount of gas. 

Pressure/Gas Laws:

• Measurable properties of gases include: 
• Pressure (P) in atm
• Volume (V) in L
• Temperature (T) in Kelvin (C+273.15)
• Amount in moles (n)
• The properties are interrelated-when one changes, it affects the others. 
• The Gas Laws: mathematical relationships that describe the quantitative relationships and behavior of gases as they are mixed, or subjected to pressure or temperature changes. 

These simple gas laws describe the relationship between pairs of these properties.

• Boyle's Law: describes how volume (V) varies with pressure (P) at constant temperature and amount of gas. 
• Boyle and Hook used a J shaped tube to measure the volume of a gas sample of gas at different pressures. They trapped a sample of air in the J tube and added mercury to increase the pressure on the gas. 
• Boyle's law states that volume and pressure have an inverse relationships where an increase in one causes a decrease in the other. 
• If the volume is increased, the pressure will decrease
• We can use Boyle's ;aw to calculate the volume of a gas following a pressure change or the pressure of a gas following a volume change. 
• As long as the temperature change and the amount of gas remain constant. 

p1v1=v1v2; Where p1 and v1 are the initial pressure and volume of the gas and p2 and v1 are the final volume and pressure. 

Charles's Law: The volume of a fixed quantity of gas maintained at constant pressure is directly proportional to its absolute temperature * (T in Kelvin) assuming that the pressure and the number of moles are constant. 

• AJ Charles was interested in gases and was among the first people to ascend in a hydrogen filled balloon. 
• When the temperature of a gas sample increases, the gas particles move faster; collisions with the walls are more frequent 
• The volume of a gas increases with increasing temperature. 
• Volume and temperature are linearly related. If two variables are linearly related, plotting one against the other produces a straight line. 
• Absolute zero * refers to the temperature at 0 K-colder temperatures do no exist. Anything less would have a negative volume which is physically impossible. 
• If a balloon is moved from an ice water bath to a boiling water bath, its volume expands as the gas particles within the balloon move faster (due to the increased temperature) and collectively occupy more space. 
• Ice water (low kinetic energy)----> Boiling water (high kinetic energy)

(v1/t1)=v2/t2 at constant pressure (P) and number of moles (n)

The pressure exerted on a sample of a fixed amount of gas is doubled at constant temperature, and then the temperature of the gas in kelvins is doubled at constant pressure. What is the final volume of the gas?

(a) The final volume of the gas is twice the initial volume.

(b) The final volume of the gas is four times the initial volume.

(c) The final volume of the gas is one-half the initial volume.

(d) The final volume of the gas is one-fourth the initial volume.

(e) The final volume of the gas is the same as the initial volume.

Avogradro's Law: The volume (v) of a gas maintained at constant temperature and pressure is directly proportional to the number of moles of the gas (n)

• When the amount of gas in a sample increases at constant temperature and pressure, its volume increases in direct proportion because the greater number of gas particles fills more space. 
• More gas molecules=larger volume
• We can use Avogrado's law to calculate the volume of a gas following a change in the amount of the gas as long as the pressure and temperature are constant. 
• The volume of a gas sample increases linearly with the number of moles in the gas sample. 
• When the amount of gas in a sample increases at constant temperature and pressure, its volume increases in direct proportion because the greater number of gas particles fill more space. 
• Equal volumes of gases contain equal numbers of molecules (gas doesn't matter)

• R (gas constant) is the proportionality constant in the ideal gas equation
• R relates pressure, volume, temperature, and the number of moles of gas in the ideal gas equation.
• Molar Volume: Volume of one mole of a gas determines at STP (standard temperature and pressure)
• An ideal gas is a gas whose volume, pressure, and temperature may be described by the ideal gas equation

7.3 Apply Dalton's Law of Partial Pressure to calculate the pressure of combined gases and to calculate the partial pressures of gases in mixtures. (11.6)

7.4 Describe gases in terms of the Kinetic Molecular Theory (11.7, 11.8)

• Kinetic Model Theory: The simplest model for the behavior of gases. In this theory, a gas is modeled as a collection of particles (either molecules or atoms, depending on the gas) in constant motion. The basic postulates (or assumption) of kinetic molecular theory are listed below:
• 1.) The size of a particle is negligibly small. KMT assumes that the particles themselves occupy no volume, even though they have mass. 
• Smaller molecules move faster then larger molecules. 
• 2.) The average kinetic energy of a particle is proportional to the temperature in kelvins. The motions of atoms ad molecules in a gas is due to thermal energy, which distributes itself among the particles in the gas.  
• The average kinetic energy of the gas particles is directly proportional to the Kelvin temperature. Nothing else. 
• The temperature of the gas increases, the average speed of the particle increases. 
• However, not all the gas particles are moving at the same speed. 

• 3.) The collision of one particle with another (or with the walls of its container) is completely elastic. This means that when two particles collide, they may exchange energy, but there is no overall loss of energy. 

7.5 Differentiate Effusion and Diffusion (11.9)

Diffusion: The process by which gas molecules spread out in response to a concentration gradient, and even though the particles undergo many collisions, the root square velocity still influence the rate of diffusion.

Effusion: The process by which a gas escapes from a container into a vacuum through a small hole.

7.6 Distinguish between Ideal and Real Gases (11.11)

• Ideal gases are hypothetical-there is no gas that will exactly follow the ideal gas behavior. 
• However, may gases will behave very closely to ideal gases under certain conditions. 
• According to KMT, ideal gas laws assume:
• No attractions between gas molecules
• gas molecules do no take up space
• At low temperatures and high pressures these assumptions are not valid. 
• Real gases will behave more like ideal gases under conditions of:
• Low pressure (at higher pressures, the volume of the gas molecules become significant
• High temperatures (at lower temperatures, the inter molecular attractions can lower the expected pressures)