Properties of Gases

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Properties of Gases

• Compiled by
• Mr. Walia

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Gas form is an essential form of the matter

• O2, O3, N2, CO2 are essential gases in the
  atmosphere.
• O2, and N2 provide the plants essential compounds
  for photosynthesis.
• O3 protects us from the harmful solar radiations.

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The proportion of these gases in the atmosphere
is measured using balloons by meteorologists
everyday.
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11 of the periodical table elements are in the
gas form.
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Many inorganic and organic compounds are gases
NO2, SO2, N2O, CH4 , C2H6 , C3H8 , CH3NH2
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Pressure of the Gases

• Pressure (P) is the force exerted on a surface
  divided by the area of the surface

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Properties of Pressure

• Pressure increases as more gas is added
• Conclusion - Pressure (P) is directly
  proportional to moles of gas (n)
• Pressure due to a gas is the same in all
  directions whereas pressure due to weight is
  directional
• Pressure unit in SI is Pascal
•  

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Theres another way to measure the gas pressure

Gas pressure can be measured by relating to the
atmospheric pressure.

• Barometer
• An apparatus used to measure pressure  derived
  from the Greek "baros" meaning "weight. 
• Created by Evangelista Torricelli in 1646.

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• Torricelli inverted a tube filled with mercury
  into a dish until the force of the Hg inside the
  tube balanced the force of the atmosphere on the
  surface of the liquid outside the tube.

• The hight of the mercury
• in the tube is a measure of
• the atmospheric pressure.

• At sea level and 0C
• this height  is 760 mmHg
• and the pressure supporting this height is
  called 1 atmosphere.

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Atmospheric Pressure is measured by different
units

• Atmospheric pressure is equal to 760 mmHg and is
  called 1 atm.
• 1 atm 760 mmHg
• 1 mmHg 1 torr so 1atm 760
  torr
• 1 atm 101.325 kPa
• 1 bar 10 5 Pa so 1 atm 1
  bar

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Change in average atmospheric pressure with
altitude
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Why Mercury?

• In theory, any liquid can be used in abarometer.
• Mercury is so dense that can form a usable height
  in the tube. A similar barometer made of water,
  in comparison, would have to be more than 34 feet
  (100 meters) high.
• Mercury also has a low vapor pressure, meaning it
  does not evaporate very easily. Water has a
  greater vapor pressure. Because of this, the
  pressure exerted by water vapor at the top of the
  barometer would affect the level of the mercury
  in the tube and the barometric reading.

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Aneroid barometer

• A major disadvantage of the mercury barometer is
  its bulkiness and fragility. The long glass tube
  can break easily, and mercury levels may be
  difficult to read under unsteady conditions, as
  on board a ship at sea.
• To resolve these difficulties, the French
  physicist Lucien Vidie invented the aneroid
  ("without liquid") barometer in 1843.

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Aneroid Barometer

• An aneroid barometer is a container that holds a
  sealed chamber from which some air has been
  removed, creating a partial vacuum. An elastic
  disk covering the chamber is connected to a
  needle or pointer on the surface of the container
  by a chain, lever, and springs. As atmospheric
  pressure increases or decreases, the elastic disk
  contracts or expands, causing the pointer to move
  accordingly.

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• This type of aneroid barometer has a pointer that
  moves from left to right in a semicircular motion
  over a dial, reflecting low or high pressure. The
  simple clock-like aneroid barometer hanging on
  the wall of many homes operates on this basis.

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Open- tube manometer

• The open-tube manometer is another device that
  can be used to measure pressure. The open-tube
  manometer is used to measure the pressure of a
  gas in a container.

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• Atmospheric pressure pushes on the mercury from
  one direction, and the gas in the container
  pushes from the other direction. In a manometer,
  since the gas in the bulb is pushing more than
  the atmospheric pressure, you add the atmospheric
  pressure to the height difference
• gas pressure atmospheric pressure h
• h is the difference in mercury levels.
• Gas pressure will be in units of torr or
  mmHg.

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Ideal Gas Laws

• There are some laws that explain
• (a) the relationship between the pressure and
  volume of the gas at a fixed temperature,
• (b) the relationship between the volume and
  temperature of it in a fixed pressure and,
• (c) the relationship between the pressure and
  temperature of a gas in a fixed volume.

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Boyle's Law (1662)

• The relationship between the pressure and the
  volume of a given sample of gas at fixed
  temperature.
• A sample of gas compresses if the external
  pressure applied to it increases and the product
  PV is constant.
• The Pressure (P) of a gas is inversely
  proportional to Volume (V) at constant
  Temperature (T) and moles of gas (n). 

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• Boyle's law, stated in mathematical terms for a
  gas whose pressure and volume is measured at two
  different pressure/volume states at a constant
  temperature is then,
• P1V1 P2V2

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Boyles Mathematical Law
What if we had a change in conditions
since PV k
P1V1 P2V2
Eg A gas has a volume of 3.0 L at 2 atm. What
is its volume at 4 atm?
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1. determine which variables you have

• P1 2 atm
• V1 3.0 L
• P2 4 atm
• V2 ?

1. determine which law is being represented

P and V Boyles Law
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3) Rearrange the equation for the variable you
dont know
4) Plug in the variables and chug it on a
calculator
V2 1.5L
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Example 1

• At 0 C and 5.00 atm, a given sample of a gas
  occupies 75.0 L. The gas is compressed to a final
  volume of 30.0 L at 0 C. What is the final
  pressure?
• Answer 12.5 atm

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Charles' Law (1787)

• (It was developed by Guy Lussac in 1802)
• The volume of any gas increases directly with
  increasing temperature at constant pressure.
• If we plot a graph of the volume of a sample of
  gas versus the temperature at constant pressure
  we get something that looks like the following

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Absolute Temperature

• Experimental data show that 1C increase in the
  temperature of an ideal gas would increase its
  volume as 1/273 of the volume at 0C .
• So if the volume in 0C was 273 mL, it would
  increase 1/273 x 273 mL 1 mL at 1C and the
  total volume will be 274 mL.
• A 10C increase will increase the volume
• 10 x 1/273 x 273 mL 10 mL, the total volume
  283 mL.
• A 273C increase will increase the volume
• 273 x 1/273 x 273 mL 273 mL, the total
  volume 546 mL.

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• Although the the volume increases in a regular
  manner with increase in temperature, it is not
  directly proportional to the Celsius temperature.
  An increase in temperature from 1C to 10 C ,
  eg., does not increase the volume 10- fold, but
  only from 274 mL to 283 mL.
• An absolute temperature scale, kelvin
  temperature, is defined in such a way the volume
  is directly proportional to kelvin temperature.
• A 2-fold increase in the absolute temperature
  would increase the volume the same an increase
  from 273 K (0C) to 546 K (273 C) increases the
  volume from 273 mL to 546 mL.
• A Kelvin reading is (T) is obtained by adding 273
  to the Celsius temperature (t)
• T t 273

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• From the extrapolated line, we can determine the
  temperature at which an ideal gas would have a
  zero volume. Since ideal gases have infinitely
  small atoms the only contribution to the volume
  of a gas is the pressure exerted by the moving
  atoms bumping against the walls of the container.
  If no volume then there must be no kinetic energy
  left. Thus, absolute zero is the temperature at
  which all kinetic energy (motion) has been
  removed. NOTE This does not mean all energy has
  been removed, merely all kinetic energy.

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To avoid the need to know k, we use ratios. The
ratio of V to T of an ideal gas at constant
pressure is constant over all temperatures. Or...

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• A balloon racer uses the Charles law. When the
  air in the balloon gets warmer it expands and
  will become less dense and balloon floats in the
  air. When the air gets colder it will become more
  dense and it will come down in the air.

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Example 2

• A sample of a gas has a volume of 79.5 mL at 45
  C. What volume will the sample occupy at 0 C
  when the pressure is held constant?
• Answer 68.2 mL

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Amonton's Law (1703)

• The Pressure of a gas is directly proportional
  to the Temperature (Kelvin) at constant V and n.
• Example 3
• A 10.0 L container is filled with a gas to a
  pressure of 2.00 atm at 0 C. At what temperature
  will the pressure inside the container be 2.50
  atm?
• Answer 341 K 68 C

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Molar Volume

• Molar volume (Vm) is the volume that every mole
  of the material occupies (L. mol-1)
• Vm V. occupied by the material/ No. of moles
  of it
• Vm V / n
• Physical data show that molar volumes of gases
  are equal at the same pressure and temperatures.
  Vm for some gases at 0 C and 1 atm
• Argon
  22.09
• Carbon dioxide
  22.26
• Nitrogen
  22.40
• Oxygen
  22.40
• Hydrogen
  22.43

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Avogadro's Law(1811)

• The Volume of a gas is directly proportional to
  the moles of the gas, n at constant P and T. 
• According to Avogadro, equal volumes of different
  (ideal) gases at the same temperature and
  pressure contain equal numbers of molecules
  (moles) of the different gases.

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Ideal Gas Law

• If we take the three of the gas laws we've
  studied so far, we can combine them into a single
  law called the Ideal Gas law. This law covers the
  relationship between temperature, pressure,
  volume and number of moles of an Ideal gas.
• Avogadro's Law V k1n T,P
• Boyle's Law V k2/P T,n
• Charles' Law V k3T n,P

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• After some consideration and algebra, we arrive
  at
• V koverall nT/P
• where koverall turns out to be the Ideal gas
  constant
• (or universal gas constant) 
• We're more familiar with the equation written as
• PV nRT
• This is the Ideal Gas Law or the equation of
  state for an ideal gas. At ordinary conditions of
  temperature and pressure, most gases conform well
  to the behavior described by this equation.
  Deviations occur, however, under extreme
  conditions (low temperature and high pressure).

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• The molar volume of an ideal gas at STP (0 C and
  1 atm) 22.4136 L
• So we can calculate the ideal gas constant
• R PV / nT
• (1 atm) (22.4136 L) / (1 mol) (273.15 K)
• 0.082056 8.2056 x 10 2 L.atm / K.mol
• Another form of the equation of state for an
  ideal gas
• Since n g / MW
• So PV (g/MW) RT

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• Example 4
• The volume of a sample of gas is 462 mL at 35 C
  and 1.15 atm. Calculate the volume of the sample
  at STP.
• Example 5
• What is the density of NH3 (g) at 100 C and 1.15
  atm?