Gases & Atmospheric Chemistry IntroductIon to Gases & Atmospheric Chemistry

Introduction Gas behaviour is well understood and can be explained in terms of simple mathematical models  A study of gas behaviour will allow us to comprehend the nature of gaseous elements and compounds. FYI: a gas is a substance that is normally in the gaseous state at ordinary pressures and temperature. However, a vapour is the gaseous form of a substance that is normally a liquid or a solid at ordinary pressure and temperature (eg. Oxygen gas versus H2O vapour).  Our gaseous atmosphere provides a means of transferring energy and material throughout the globe, and it is a source of live-giving chemical, hence a study of the nature of the gas phase will also further enhance our understanding of our atmosphere.

Review of Inter/Intra-molecular forces (IM) The Particle Theory of Matter states that the strength of forces depends on the: type of the force temperature of the system   Types of attractive forces: Intramolecular forces between atoms: are very strong forces; ex. Covalent and ionic bonds Intermolecular forces between molecules; Dipole-dipole: between polar molecules London dispersion: very weak, short-lived forces between non-polar molecules. Become significant in larger molecules. Temperature: the higher the temperature, the more kinetic energy (more motion), thus the more attractive forces are overcome.

Motion of Particles All substances contain particles that are in constant, random Motion  3 types of motion:   translational (straight line motion) rotational (spinning) vibrational (back & forth motion of atoms in a molecule)

States of Matter Volume & Shape -Fixed volume & Shape   SOLID LIQUID GAS Volume & Shape -Fixed volume & Shape - Virtually incompressible -does not flow readily -Fixed volume, variable shape -virtually incompressible -flows readily -Variable volume & shape -highly compressible Densit y High density Med-high density Very low density Interactions between particles -particles are held tightly together - many strong attractive forces between particles -particles are held more loosely together - many weak attractive forces between particles -particles are highly spaced apart -virtually no attractive forces between molecules -forces increase with molecular size Motion Vibrational motion only - molecules held tightly Rotational & vibrational, loosely held All three types, very loosely held

Characteristics & Descriptions of Liquids Volatile – when two liquids are compares, the one that evaporates more readily is said to be volatile.   Boiling point – when the vapour pressure of the liquid equals the atmospheric pressure the liquid boils. Only the fastest moving molecules are able to overcome the attractive forces of their neighbours and leave the surface of the liquid. The slower molecules are left behind, thus the average kinetic energy (i.e. temperature) decreases. Vapour pressure - the pressure exerted by a gas above a liquid in a closed container. Increases with increasing temperature, but decreases with increasing IM forces. Ask yourself, how fast does the liquid evaporate (water versus rubbing alcohol).

Kinetic Molecular Theory – (Clausius 1857) describes the behaviour of gases at the molecular (or atomic) level Any law is an empirical generalization which describes the results of many experiments; it does not explain why the results have been obtained. A theory is a description which explains the results of an experiment. In this unit we will be looking at several laws that describe the behaviour of gases. The Kinetic Molecular Theory of ideal gases explains these observed behaviours described by different gas laws (Boyle’s law, Charles’ Law, Ideal Gas Law, etc.)

The Kinetic Molecular Theory can be stated in five (5) postulates:   A gas consists of molecules in constant random (and rapid) trans-locational motion (motion in a straight line). The volume actually occupied by the molecules of a gas is negligibly small; the vast majority of the volume of the gas is “empty” space. The attractive and repulsive forces between gas molecules are negligible. All collisions between gas molecules are perfectly elastic. This means that when gas molecules collide, there is no loss of kinetic energy. The average kinetic energy of gas molecules is directly related to the temperature. The > the temperature, the > the average kinetic motion, the > their average kinetic energy.

Kinetic Molecular Theory Therefore: All molecules, regardless of their size, have the same average kinetic energy at the same temperature Each gas molecule “behaves” as if it were alone in the container (due to #2 & #3 above)   Pressure According to KMT, pressure is created when particles of gas hit the sides of the container Increase Temp, increase KE, increase # collisions, increase Pressure.

“Ideal” versus “Real” Gases Experiments into gas behaviour demonstrate that under normal temperature and pressures, nearly all gases behave in similar and predictable ways. The gas laws we will be studying are used to predict gas behaviour are “ideal gas laws” The Kinetic Molecular Theory describes this hypothetical gas called an ideal gas. Ideal gases take up no space and do not attract each other. An ideal gas is considered as a collection of rebounding particles, separated by great distances, so that the gas is mostly empty space, and there are no forces of attraction between the molecules.

“Ideal” versus “Real” Gases However, we find that no gas is “truly” an “ideal gas”, that there are attractive forces that cause deviations from the ideal gas laws, and finally lead to the formation of liquids. Real gases deviate from the Kinetic Theory at various temperatures and pressures The greatest deviation from ideal gas behaviour occurs at: High pressure Higher density of gas molecules (molecules are closer together, so volume and attractive forces play a larger role) The properties and behaviours of “real gases” can be generalized into a theory of an “ideal gas” – thus allowing us to make accurate inferences

Graham’s Law of Diffusion Diffusion is the process by whereby a substance spreads from a region of high concentration to one of lower concentration. This is the intermingling of molecules as a result of random molecular motion.  The rate at which a gas diffuses depends directly on the molecular speed – molecules with high speeds diffuse faster.  Graham’s Law: the rates of diffusion of two gases are inversely proportional to the square room of their molar masses.  Therefore the molecules with smaller mass have a higher speed, and hence travel faster than a heavier gas, and therefore travel the furthest.

Will a gas diffuse faster or slower in a vacuum? Gas particles travel in a straight line until they hit another particle. Since a vacuum is devoid of molecules, particles of gas in a vacuum will not hit anything until the hit the side of the container, thus travelling faster in a vacuum

Introduction to Gas Laws In a closed system, the FOUR Gas Law variables are: Volume, V ( in L) Temperature, T (in Kelvins) Quantity of gas, n (in moles) Pressure, P (in kPa)   A closed system is one that is not open to the atmosphere (i.e. there is no net loss or gain of amount of substance).

1) Volume All gases must be enclosed in a leak-proof container Symbol for variable: V, Units: L or mL, dm3 or cm3 When discussing such things as Boyle’s law and other laws were volume is a variable, it is important to remember that the container changes sizes with no loss of gas. (imagine a container with a sealed movable wall.

2) Temperature Usually measured in degrees celsius (symbol = 0C) Symbol of variable: T, Units: uses the Kelvin scale, K NOTE: all gas law problems will use Kelvin scale Conversions: 1 Kelvin = Celsius + 273.15 For example: 250C = 298 K The Kelvin scale is directly proportional to the amount of kinetic energy: double the Kelvin temperature, you double the kinetic energy.

3) Amount of Gas   the amount of gas present is measured in moles (mol) or in grams (g) Typically, if grams are used, you will need to convert to moles at some point.

4) Pressure Gas molecules are in constant motion, colliding with each other and with the walls of their container  pressure – the force exerted on an object per unit of surface area and Who would exert a greater pressure on a wooden floor, a woman with a mass of 50.0kg wearing high heeled shoes, or a woman also with a mass of 50.0 kg wearing running shoes?

4) Pressure Who would exert a greater pressure on a wooden floor, a woman with a mass of 50.0kg wearing high heeled shoes, or a woman also with a mass of 50.0 kg wearing running shoes? Since, P = F/A, a decrease in surface area can dramatically increase the pressure. UNITS OF PRESSURE One commonly used SI unit for pressure is the pascal (Pa), which is equal to 1 N/m2. In most cases, pressure is reported in kilopascals, (kPa). 1000 Pa is equal to 1 kPa. 1 kPa=1000Pa=1kN/m2

4) Pressure atmospheric pressure – Force per unit area exerted by air on an object. Standard Pressure = 101.325 kPa = 1 atm Standard Ambient Pressure = 100kPa STP= 0ºC and 101.325 kPa SATP=25ºC and 100kPa Know these!

*** Know how to convert these units 4) Pressure UNIT NAME UNIT SYMBOL DEFINITION/CONVERSION Pascal Pa 1 Pa = 1N/m2 Atmosphere atm 1 atm = 101.325 kPa mm of mercury mm Hg 760mm of Hg=1 atm=101.325kPa Torr torr 1 torr = 1mm of Hg bar barr 1 bar = 100 kPa Thus, there are FOUR (4) common units of pressure used in chemistry: *** Know how to convert these units

Sample Calculation Sample Calculation:   Convert 100 kPa to a) atm and b ) mm of Hg a) 100 kPa x __1 atm = 0.987 atm 101.325 kPa b) 100kPa x 760 mm of Hg = 750. mm of Hg 101.325kPa

How does a Gas exert Pressure? a gas cannot exert a measurable pressure in the same way that a solid or a liquid does. rather the pressure of a gas is determined by the kinetic motion of its component particles. Force = mass x velocity of particle The pressure is proportional to:  Number of molecules (per volume) Mass of the molecules velocity of molecules (Temperature)   unlike solids and liquids, particles in the gaseous state are able to move independently of one another, and thus have a high degree of disorder  exhibit translational motion, vibrational motion, and rotational motion. Density!