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IT’S A GAS…. IT’S A GAS…. IT’S A GAS…. The Nature of Gases. Gases have some interesting characteristics that have fascinated scientists for 300 years.
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The Nature of Gases • Gases have some interesting characteristics that have fascinated scientists for 300 years. • The first gas to be studied was air & it was a long time before it was discovered that air was actually a mixture of particles rather than a single gas. • But this realization did not make the study of gas behavior more difficult. • Although air is a mixture of several different gases, it behaves much the same as any single gas.
The Nature of Gases • Regardless of their chemical identity, gases tend to exhibit similar physical behaviors • Gas particles can be monatomic (Ne), diatomic (N2), or polyatomic (CH4) – but they all have some common characteristics: • Gases have mass. • Gases are compressible. • Gases fill their containers. • Gases diffuse. • Gases exert pressure. • Pressure is related to Temperature
Kinetic Molecular Theory • There is a theory that modern day chemist’s use to explain the behaviors and characteristics of ideal gases - the Kinetic Molecular Theory of Matter. • The theory states that the tiny particles in all forms of matter are in continuous motion. • There are 3 basic assumptions of the KMT as it applies to ideal gases. • Ideal gases are “perfect” gases that are used as a model to describe characteristics of real gases.
KMT Assumption #1 • A gas is composed ofsmall hardparticles. • The particles have an insignificant volume and are relatively far apart from one another. • There isempty spacebetween particles. • No attractive or repulsive forces between particles.
KMT Assumption #2 • The particles in a gas move inconstant random motion. • Particles move in straight paths and are completely independent of each of other • Particles path is only changed bycolliding with another particle or the sides of its container. Molecular motion
elastic collisions inelastic collisions KMT Assumption #3 • All collisions a gas particle undergoes areperfectlyelastic. • They exert a pressure but don’t lose any energy during the collisions.
Gases have mass. • Gases seem weightless, but they are classified as matter, therefore, they must have mass. • The density of a gas is much less than the density of a liquid or solid. • It’s this very low density that allows us to be able to walk through the room without concerning ourselves with air resistance. • The mass is really only noticeable if we have a collection of gas in a container.
Gases “R” squeezable • Gas particles have a high velocity, relative to their low masses. • Which means they have a great deal of kinetic energy • This high velocity and low mass means the particles are spread far apart, with empty space in between particles. • Therefore, if you squeeze on a gas, its volume can be reduced considerably
Gases “R” squeezable • The gas particles empty space can be compressed by added pressure giving the gas particles less room to bounce around thus decreasing the overall volume. • Compression allows us to package a lot of gas in a relatively small volume. • And compression generates a spring like character to develop to the gas collection
Gases “R” squeezable • There are a huge number of applications • Storm door closers • Pneumatic tube delivery devices • Tires • Air tanks • Etc…
Gases fill their containers • Gases expand until they take up as much room as they possibly can. • Gases spread out to fill containers until the concentration of gases is uniform throughout the entire space. • This is why that nowhere around you is there an absence of air.
Gases fill their containers • Since the particles are in constant random motion, according to the KMT, then the gases move in a straight line until they collide with other particles or the sides of the container, which causes them to change directions until they collide with something else. • This random bouncing motion, allows for the mixing up and spreading of the particles until they are uniform throughout the entire container.
Gases diffuse • Gases can move through each other rapidly. • The movement of one substance through another is called diffusion. • Because of all of the empty space between gas molecules, gas molecules can pass between each other until each gas is mixed evenly throughout the entire container. • Since gases are in constant random motion, they are moving and colliding with everything around them, and there is so much empty space, the gases mix uniformly.
Gases diffuse • This doesn’t happen at the same speeds for all gases. • Some gases diffuse more rapidly then other gases based on their size and their energy. • KE=1/2mv2 • Diffusion explains why gases are able to spread out to fill their containers. • It’s why we can all breath oxygen anywhere in the room.
Gases exert pressure • Gas particles exert pressure by colliding with objects in their path. • The sum of all of the collisions makes up the pressure the gas exerts. • The KMT says since they are in constant random motion, the particles will colliding with anything in their path. • The definition of pressure is the force per unit area – so the total of all of the tiny collisions makes up the pressure exer- ted by the gas
Gases exert pressure • It’s the pressure exerted by the gases that hold the walls of a container out • The pressure of gases is what keeps our tires inflated, makes our basketballs bounce, makes hairspray come out of the can, helps our lungs inflate, allow vacuum cleaners to work, etc.
Pressure depends on Temp • The average kinetic energy of the particles that make up an object is defined as temperature. • Therefore, the higher the temperature the more energy the gas particle has. • So the collisions are more often and with a higher force. • And since pressure is a function of force, the pressure increases inside the container. • Think about the pressure of a set of tires on a car.
Pressure depends on Temp Today’s temp: 35°F Pressure Gauge
Pressure depends on Temp Today’s temp: 85°F Pressure Gauge
Measuring Gases • The conditions under which a gas is studied is very important to its behavior. • Experimental work in chemistry requires the measurement of such quantities as volume, temperature, pressure, and the amount of sample. • These quantities are called variables and if they are not accounted for then the results of the experiment might be jeopardized.
Amount (n) • The quantity of gas in a given sample is generally given in terms of moles of gas. • This of course is in terms of 6.02 x 1023 molecules per mole of the gas. • Don’t forget to convert mass to moles you just divide by the molar mass of the gas. • So amount of a gas, refers to how many gas particles are in the sample.
Volume (V) • The volume of a gas is simply the volume of the container it is contained in. • The metric unit of volume, liter (L), is often used. • There might also be problems that use cubic meters as the unit for volume. • 1000 L = 1 m3
Temperature (T) • The temperature of a gas is generally measured with a thermometer in Celsius. • All calculations involving gases should be made after converting the Celsius to Kelvin temperature. Kelvin = C° + 273 Celsius = K - 273
Pressure (P) • The pressure of a gas is the force exerted on the wall of the container, in which a gas is trapped. • There are several units for pressure depending on the instrument used to measure it including: 1) atmospheres (atm) 2) Millimeters of Mercury (mmHg) 3) Kilopascal (kPa)
S T P • The behavior of a gas depends very strongly on the temperature and the pressure at which the gas is held. • To make it easier to discuss the behavior of a gas, it is convenient to designate a set of standard conditions, called STP. • Standard Temp and Standard Pressure • Standard Temperature = 0°C or 273K • Standard Pressure = 1atm or 760mmHg or 101.3kPa (depending on the method of measure)
Atmospheric Pressure • The gases in the air are exerting a pressure called atmospheric pressure • Atmospheric pressure is a result of the fact that air has mass is and is colliding with everything under the sun with a force. • Atmospheric pressure is measured with a barometer.
Atmospheric Pressure • Atmospheric pressure varies with altitude • The lower the altitude, the longer and heavier is the column of air above an area of the earth. • Look on the back of a box of cake mix for the difference in baking times based on the atmospheric pressure in your region.
Atmospheric Pressure • Knowing this atmospheric pressure and predicting changes in the atmospheric pressure is how forecasters predict the weather. • Low pressure or dropping pressure indicates a change of weather from fair to rain. • High pressure is an ind- ication of clear skies & sun.
Gas Laws • Studies of the behavior of gases played a major role in the development of physical sciences in the 7th and 8th centuries. • The Kinetic Molecular theory marked a significant achievement in understanding the behavior of gases. • Observations have become mathematical laws which we can use to predict quantitative outcomes.
Boyle’s Law • Robert Boyle was among the first to note the inverse relationship between pressure and volume of a gas. • As the pressure on a gas increased the volume of the gas will decrease. • He measured the volume of air at different pressures, and observed a pattern of behavior. • During his experimentsTemperatureandamount of gasweren’t allowed to change 1621-1691
Volume Pressure How does Pressure and Volume of gases relate graphically? PV = k Temperature, # of particles remain constant
Boyle’s Mathematical Law: If we have a given amount of a gas at a starting pressure and volume, what would happen to the pressure if we changed the volume? Or to the volume if we changed the pressure? since PV equals a constant P1V1 = P2V2 Ex: A gas has a volume of 3.0 L at 2 atm. What will its volume be at 4 atm?
Boyle’s Mathematical Law: • List the variables or clues given: • P1 = 2 atm • V1 = 3.0 L • P2 = 4 atm • V2 = ? • determine which law is being represented: P1V1 = V2 P2 3) Plug in the variables & calculate: (2 atm) (3.0 L) = (4 atm) (V2) 1.5 L
Charles’s Law • Jacques Charles studied the direct mathematical relationship between temp-erature and volume of a gas. • As temperature increases the volume of the gas increases • Charles measured the volume of air at different temperatures, and recorded the results. • During his experimentspressureof the system andamount of gaswere held constant. 1746-1823
Volume Temp How does Temperature and Volume of gases relate graphically? V/T = k Pressure, # of particles remain constant
V1 V2 = T1 T2 Charles’s Mathematical Law: If we have a given amount of a gas at a starting volume and temperature, what would happen to the volume if we changed the temperature? Or to the temperature if we changed the volume? since V/T = k Ex: A gas has a volume of 3.0 L at 400K. What is its volume at 500K?
