High School Chemistry/Gases & Gas Laws
Introduction[edit | edit source]
Background[edit | edit source]
- Atmospheric Pressure, or atm, is the force applied by gases present in the Earth's atmosphere.
- Pressure is the force per unit area from collisions of gas molecules with one another and with the walls in the container.
- Barometer is used to measure atm (in mmHg).
- Manometer is used to measure the pressure of any given gas.
- The air becomes thinner and thinner as the altitude increases, therefore a less exertion by gases on the Earth's surface is found: ATM ⇧ ALTITUDE ⇩
- STP, or standard temperature/pressure, is 1atm and 0 °C.
- Absolute zero is 0K = -273 °C. This is when the temperature is so low that the particles in motion come to a complete stop, therefore the gas has no average kinetic energy.
Properties of General Gases[edit | edit source]
- Gases are able to fill any container.
- Gases are easily compressed.
- Gases integrate into each other completely.
- Gases give off pressure on its surroundings.
Gases and KM Theory[edit | edit source]
This theory accurately models the behavior of gases, so as a result, an approximation will not always be true. This theory also tries to explain the properties present in an ideal gas.
Ideal Gas Properties[edit | edit source]
- There is no measurable volume, or the gas particles are very spread out, in these gases. [NOT ALWAYS TRUE]
- No intermolecular forces exist between the gas particles. [NOT ALWAYS TRUE]
- Gas particles have elastic collisions, or they do not lose kinetic energy upon a collision, and they are always in constant, random motion.
- Average kinetic energy increases as temperature increases (or, in other words, are directly proportional to each other).
Real Gas Properties[edit | edit source]
- Do not follow properties #1 and #2 above.
- KMT most likely is true for H2 and N2, which are held together by weaker intermolecular forces than for polar gases, such as ammonium, which have stronger intermolecular forces.
Gas Laws[edit | edit source]
These gas laws are used to compare alternating conditions for a given gas. These laws are simple, mathematical relationships between volume, pressure, temp and/or particles.
Boyle's Law[edit | edit source]
Pressure is indirectly proportional to the volume at STP. So: P ⇧ V ⇩
FORMULA: P1V1 = P2V2
Charles' Law[edit | edit source]
Volume is directly proportional to temperature. So: V ⇧ T ⇧
Temperature must be in Kelvin!
Gay-Lussac's Law[edit | edit source]
Pressure is directly proportional to temperature. So P ⇧ T ⇧.
The Combined Gas Law[edit | edit source]
This law adds all the laws together to accurately articulate the relationship between pressure, temperature and volume.
FORMULA: P1V1/T1 = P2V2/T2
Avagadro's Law[edit | edit source]
Volume and "n" are directly proportional so the volume increases while the moles increase. The pressure and temperature are constant.
Ideal Gas Law (relates P, V, T, and number of moles)[edit | edit source]
Relates the pressure, volume, temperature and the number of moles. The ideal gas law best models gases at low pressure and high temperature.
- P = Pressure in atm/kPa
- V = Volume in L
- N = Moles (m)
- R = Universal Gas Constant (0.0821 L.atm/mol.K)
- T = Temperature in K
Graham's Law of Effusion[edit | edit source]
Rates of effusion are dependent upon the molar mass of the gas molecules. Lighter molecules move faster than the heavier molecules at equivalent temperature. If given two gases, calculate the molar mass of both to find out which one effuses faster. Use Graham's Law to figure out how many times faster the lighter gas moves.
FORMULA: = √mm2/√mm1
Dalton's Law of Partial Pressure[edit | edit source]
The total pressure of a mixture of a gas = the partial pressure of its components. The main use of John Dalton's Law is with gases collected by H2O Displacement.