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G A S E O U S ST A T E1.INTRODUCTION
1.1 States of Matter
Matter can be defined as anything that occupies volume and has mass. Matter can be classified
into three states – solid, liquid and gas. Plasma is regarded as the fourth state of matter, which
exists only at very high temperatures (at interiors of stars, 107 K). At very high temperatures, all
gases become ionized, which results in the formation of the fourth state of matter, the so-called
plasma state.
A solid state has definite shape and volume at a given temperature and pressure. A liquid has
definite volume but no definite shape, whereas a gas has neither definite volume nor definite
shape.
A substance can exist in either of the three states depending on the temperature and pressure
under which it exists,
e.g. at ordinary temperature and pressure, water exists as liquid and can be passed into gaseous
state at 100°C. A substance can also exist in all the three states simultaneously, e.g. water has all
the three phases in equilibrium at 4.58 mm Hg pressure and 0.0098°C, which is known as triple
point, i.e. the point at which three phases of a component exist together.
Ice Water Vapor
Thus, an increase in the forces of attraction (by increasing pressure) and a decrease in the kinetic
energy (by lowering temperature) may result in the conversion of the gaseous state into the liquid
state and then into the solid state. Different states of matter are thus, associated with definite
energy contents and are interconvertible.
1.2 The Gaseous State
Measurable Properties of Gases
(a) Mass: The amount of a gas is expressed in terms of its number of moles. For a gas with a
molar mass M, the mass in gram (w) is related to the number of moles (n) as n = w/M
(b) Volume: The volume of a gas is the space occupied by its molecules under a given set of
conditions. Volume of the container in which a gas is enclosed is expresses as 1 m3 = 103 L =
103 dm3 = 106 cm3
(c) Temperature: The extent of hotness or coldness of a body is known as temperature. The
measurement of temperature is based on the principle that substances expand on heating.
The units used for the measurement of temperature are as follows:
1
, (i) Centigrade or Celsius scale (named after Anders Celsius)
(ii) Fahrenheit scale (named after Daniel Fahrenheit, a German instrument maker)
(iii) Kelvin scale (name after Lord Kelvin). Also, K = oC + 273.15
F'– 32 C
The celsius and fahrenheit scales are related by the following: =
9 5
0°C = 32°F and 37°C = 98.6°F (human body temperature)
(d) Pressure: The force experiended by the walls of a container due to the bombardment of gas
molecules. This force per unit area of the walls is known as gas pressure.
The pressure of pure gas is measured by manometer while that of mixture of gas is mesaurd
using barometer.
A standard or normal atmospheric pressure is the pressure exerted by a mercury
column of exactly 76 cm at 0°C, which is the pressure exerted by the atmosphere at the sea
level.
The smaller unit commonly used for expressing the pressure of a gas is
mm or torr (after the name of Torricelli, who invented the barometer).
Thus,
1 atm = 76 cm = 760 mm or 760 torr
The unit of pressure commonly used
is ‘bar.’ 1 atm = 1.01325 bar or 1 bar
= 0.987 atm
The SI unit of pressure is pascal (Pa). Pa is defined as the pressure exerted by a force of 1
newton on an area of 1 m2.
1 Pa = 1 Nm–2 = 1 kg m–1 s–2
1.3 Atmosphere and Atmospheric Pressure
A thick blanket of air that surrounds the earth is called atmosphere. Molecules of various gases
that are present in the atmosphere are under constant pull of the gravitational force of the earth.
As a result of this, the atmosphere is dense near the surface of the earth than that at high
altitudes. Force experienced by molecules in any area of the earth exposed to the atmosphere is
equal to the weight of the column of the air above it. This force per unit area of the earth is known
as atmospheric pressure.
1 atm = 76.0 cm of mercury = 760 mm of mercury = 760 torr = 1.01325 × 105 Pa
CONCEPTS
The high density of mercury (13.6 g/mL) leads to shorter length of glass tube. The closed-end
manometer should not contain water droplets adhered inside its long arm. If it is so, the observed
pressure would be lower than the real pressure exerted by the gas.
The figure below shows a manometer and a barometer. A barometer is used to measure
atmospheric pressure. The basic concept used in all pressure-measuring instruments is given
below:
PA = PB + gh, where h is the height difference between the points A and B
2
, CONCEPTS
B
Pgas
Gas
h
Pgas Atmospheric
h cm
Hg A
Pressure
Pgas= Patm + hdg
Figure 3.1: An open arm manometer Figure 3.2: Barometer
Name Symbol Value
Things to Remember:
Pascal 1 Pa 1 Nm–2, 1 kg m–1s–2
Bar 1 bar 105 Pa
Atmosphere 1 atm 101.325 kPa
Torr 1 torr 101 325/760 Pa = 133.32 Pa
Millimeters of mercury 1 mmHg 133.322 Pa
Pound per square inch 1 psi 6.894 757 kPa
The pressure is independent of the shape and cross-sectional area of the column. The mass of the colum
This difficulty can be solved by carrying oxygen cylinders.
STP conditions: 0ºC or 273.15 K temperature and 1 atm (=1.01325 bar) pressure Standard Ambient Tem
SATP conditions: 298.15 K (25°C) and 1 bar (105 Pa) pressure
The molar volume of an ideal gas at SATP conditions is 24.789 L mol–1.
2.GAS LAWS
Among the three common states of matter, the gaseous state is the simplest. The laws of gaseous
behavior are more uniform and better understood. The well-known laws of gaseous behavior are
the Boyle’s law, Charles’s law, Graham’s law and Avogadro’s law.
2.1 Boyle’s Law
Boyle’s law states that at constant temperature, the volume of a given mass of a gas is inversely
proportional to pressure.
V
1 (T and mass of gas constant); PV = constant
P
Log P + log V = constant
3
, Constant T
Constant T
V V PV log P
Constant T Constant P
1/P P P log V
Figure 3.3: Various plots of pressure (P) vs
Volume (V)
2.2 Charles’s
Law
Charles’s law states that at constant pressure, the volume of a given mass of gas is directly
proportional to its
absolute temperature, i.e. V T (P and m constant) V = = constant = K
V
KT
T
V V Constant P V/P
Constant P
Constant P
T 1/T T
Figure 3.4: Various plot of volume (V) vs Temperature (T)
2.3 Gay-Lussac’s Law or Pressure-Temperature Law
Gay-Lussac’s law states that at constant volume, the pressure of a given mass of gas is directly
proportional to its absolute temperature, i.e. P T ( V and mass of gas constant).
P1
P2
T1 =
Constant P P/T ConstantV T
T2 ConstantV
T
Figure 3.5: Plots of pressure
(P) vs Temperature
(T)
2.4 Avogadro’s Law
Avogadro’s law states that equal volumes of all gases under similar conditions of temperature and
pressure will contain equal number of molecules or vice versa, i.e. V n (at constant T and P)
4
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