Gas Properties (Hydrocarbon Exploration & Development)

Continued from Physical Properties of Hydrocarbon Part 2 (Hydrocarbon Exploration & Development Notes #4)

Gas properties

  • a state of matter in which the molecules move freely, thereby permitting the matter to expand indefinitely
  • virtually have no fixed volume as they always occupy the whole of the container
  • gas molecules in constant motion (Brownian Movement)

1. Mole Concept

  • a mole is defined as the number of grams of a compound such that, the number of grams is the same as the molecular weight
  • Example, 1 mole of H2O is equal to 18 grams. A glass is filled with 360 grams of H2O, therefore, the number of mole it would contain is 360/18 = 20 moles of water (gm-mole)

2. Mole Fraction

  • Mole fraction = (No. of Moles of a component) / (Total No. of moles in the mixture)
  • or, relative mass / molecular weight

3. Atmospheric Pressure

  • when measured at sea level, it exerts a pressure of 14.69 psi, 760 mm Hg or 101.325 kPa.
  • different result at different alleviation
  • therefore, atmospheric pressure is measured at sea level to give a constant standard
  • gauge readings (denoted by letter,g) is based on pressure gauge which read ‘0’ at atmospheric pressure. So, if the pressure gauge reading is 10 psig, meaning, it’s 10 psi above atmospheric pressure.
  • to get absolute pressure, psia, add 14.69 psi to psig.

4. Boyle’s Law

  • Boyle’s law is concerned with the relationship between the volume occupied by the gas and the pressure of the gas.
  • state that, “that for a given amount of gas, a change in volume is inversely proportional to a change in its absolute pressure, providing the temperature is constant”
  • P = 1 / C or PV=T, where T is a constant (temperature)
  • Isothermal process = a process which is carried out a constant temperature.
  • important point, there is a definite relationship between the change in volume and the change in pressure, that is, volume and pressure can be calculated
  • P1V1=T & P2V2=T, therefore, P1V1=P2V2
  • if volume decrease , pressure increases, and vice versa

5. Charles’ Law

  • Charles’ law is concerned with the relationship between the volume of a gas and its absolute temperature.
  • Charles’ law states, “That for a given amount of gas, a change in it’s absolute temperature is directly proportional to a change in its volume provided the pressure remains constant.”
  • P=V / T, P (Pressure is a constant)
  • V1/T1 = P = V2/T2

6. Combined Gas Laws

  • the separate relations of boyle’s and charles’ laws can be combining to give the combined gas law (below)
  • PV/T -> a constant or P1V1 = (P2V2T1) / T2

7. Gas measurement and conversion

  • SI Units, T= 288K or 15 DegC , P = 101.325 kPa
  • Imperial Units, T 520 DegR or 60 DegF, P = 14.69 psia
  • to offset the distortion due to the compression of gas, a compressibility factor or correction factor is included with the Combined Gas Laws. This is called the supercompressibility factor or Z-factor, such that :
    • P1V1 / T1Z1 = P2V2 / T2Z2
  • The need for a compressibility factor – due to difference between ideal and real gases.

8. Ideal and Real Gases

  • Avogadro Law, “under the same conditions of temperature and pressure, equal volume of all ideal gases contained the same number of molecules”
  • PV = nRT,
    • n: number of pound moles
    • P: absolute pressure
    • V: volume in cubit feet
    • R: the universal gas constant (10.732 psia.ft3/mole.R)
    • T: absolute temperature (Deg R)
  • The equation above only truly applicable for simple gases at pressure close to atmospheric. It consistent with the kinetic theory of gases which assumes:
    • there are no inter-molecular forces
    • volume of the molecules is negligible compared with their separation
  • ideal gas or perfect is a hypothetical gas that obeys the gas law perfectly. An ideal gas would consist of molecules that occupy negligible space and negligible forces between them
  • real gases don’t have the properties assigned to ideal gases but the molecules have a finite size and forces exist between them.
  • under compression gases behave differently to the gas laws

9. Bubble point pressure and dew point temperature

Bubble point:

  • mixture of hydrocarbon liquids will, as pressure is reduced, evolved a vapor phase
  • for a given temperature, there is a point where pressure is reduced sufficiently for bubbles of the most volatile hydrocarbon fractions to form.
  • given that pressure remains constant, there’s also a point where temperature will cause the most volatile fractions to vaporize and form bubbles in the liquid. This is the bubble point  pressure of the liquid phase.
  • main problem with bubble point is a sudden loss of pressure in a vessel or pump causing the light fractions to boil at the new pressure-temperature conditions (changing state), and becoming gas.

Dew point:

  • the temperature where a vapor, cooled at constant pressure, at which the 1st drop of condensing liquid is formed.
  • can be defined as the temperature at which vapor condenses out of a gas at 14.7 psia.

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