Dr Bikash Subedi

Moderator: Prof. Dr Baburaja Shrestha

The Gas Laws & Clinical

Applications

Elements that exist as gases at 250C and 1 atmosphere

Physical Characteristics of Gases

• Assume the volume and shape of their containers.

• most compressible state of matter

• mix evenly and completely when confined to the same container

• much lower densities than liquids and solids.

Ideal gas

• Theoretical

• Negligible intermolecular forces

• collisions between atoms or molecules are perfectly elastic

• Obeys universal gas law PV= nRT at all temp & pressures

• Where P = Pressure V = Volume n = Numbers of moles R = Universal gas constant = 8.3145 J/mol K T = Temperature

Real gas

• Real gases H2, N2 , O2

• exhibit properties that cannot be explained entirely using the ideal gas law

• Behave like ideal gas at STP

• Air at atmospheric pressure is a nearly ideal

• STP = standard temperature and pressure

Standard Temperature & Pressure(STP)

• Variable

• IUPAC has, since 1982

• standard reference conditions as being 0 °C and 100 kPa (1 bar), in contrast to its old standard of 0 °C and 101.325 kPa (1 atm)

Gas Laws

Boyle’s Law

• Robert boyle,1662

• At constant temperature, V 1/P

• PV = K (constant)

• P1V1 = P2V2

P

V

Calculation of Amount of gas in a cylinder

1300 Litres @ 1 atm pressure

10x130= V x 1 bar

10 litre O2 cylinder @ 130 bar

Charles law

• Jacques Charles, 1678

• At constant pressure volume of a given mass of gas varies directly with temperature , that is

V T ( in kelvin)

or V/T = Constant (k2)

V

T

• Gases expand when heated, become less dense , thus hot air rises >> convection

• LMA inflatable cuff expands in an autoclave

3rd gas law /Gay-Lussac’s law

• At constant volume the absolute pressure of a given mass of gas varies directly with the absolute temperature

P T

or P/T = Constant

P

T

• Hydrogen thermometer or constant volume gas thermometer

• Internal combustion engines

Combined Gas Law

• Boyle’s + Charle’s + Gay Lussac’s law

• P1V1 / T1 = P2V2 / T2

• useful for converting gas volumes collected under one set of conditions to a new volume for a different set of conditions

Spirometry

• BTPS 37 0 C and H2O pressure 47 mm of Hg• Standard room temp & H2O pressure

250C• spirometer records volume under room air, not

body conditions• Thus, conversion factor or 1.07

• BTPS – body temp & pressure

Avogadro’s Hypothesis

• For a given mass of an ideal gasvolume amount (moles) of the gasif temperature and pressure are constant

Amedeo Avogadro, 1811

1 MOLE OF A SUBSTANCE

• Quantity of a substance containing the same number of particles as there are atoms in 0.012kg of carbon12

• There are 6.022 x 1023 atoms in 12 g of carbon 12. This is called Avogadro’s Number

• equal volumes of gases at the same temperature and pressure contain equal numbers of molecules

• One mole of any gas at STP occupies 22.4litres !

• 2 g of Hydrogen or 32 g of Oxygen or 44 g of Carbon dioxide occupy 22.4 litres at STP

Calculating the volume of nitrous oxide in a cylinder :

• A nitrous oxide cylinder contains 3.4 kg of nitrous oxide .

• The molecular weight of nitrous oxide is 44

• One mole is 44 g

• At STP , 44 g occupies 22.4 Litres . Therefore 3,400 g occupies 22.4 x 3,400/44 = 1730 litres.

Ideal Gas Equation

Charles’ law: V a T (at constant n and P)

Avogadro’s law: V a n (at constant P and T)

Boyle’s law: V a (at constant n and T)1

P

V anT

P

V = constant x = RnT

P

nT

PR is the gas constant

PV = nRT

Practical application ; use of pressure gauges to assess the contents of acylinder

Dalton’s Law of Partial Pressures

V and T are

constant

P1 P2 Ptotal = P1 + P2

John Dalton , 1801

In a mixture of gases , pressure exerted by each gas is the same as that which it would exert if it alone occupied the container .

Dalton’s Law

• The total pressure of a mixture of gases equals the sum of the partial pressures of the individual gases.

Ptotal = P1 + P2 + ...

AIR

O2 – 104CO2 – 40H2O – 47N2 - 569

Pulmonary capillary

vein arteryO2- 40CO2- 46

O2- 100CO2- 40

Alveolus at 760 mm Hg

At everest

• Atm pressure almost one third at sea level

• Thus, alveolar O2 pressure about 35 mm Hg

• Supplemental Oxygen

Henry’s Law

• William Henry in 1803

• At constant temperatureSolubility of gas Partial Pressure of gas

Solubility of gas :

• Depends on type of gas and liquid

• Decreases with increase in temperature

• Caisson’s disease/ decompression sickness

Adiabatic changes of state in a gas

• If the state of a gas is altered without a change in heat energy , it is said to undergo adiabatic change

• Heat energy neither received nor given to surrounding

• If a gas is rapidly compressed ; its temperature rises (the Joule – Kelvin principle).

• Conversely , If a compressed gas expands rapidly, cooling occurs (cryoprobe)

Application

• Compression of air rapidly in compressor >> ↑ temp >> need of coolant

• Cylinder connected to an anesthetic machine rapidly turned on >> ↑↑ temperature in gauges & pipelines >> fire or explosion

Cryoprobe

• Rapidly expanding gas through a capillary tube causes cooling

• N2O, He, Argon, N2

• Cooling causes degeneration, necrosis

• Wart, mole removal. Nerve degeneration for pain

Critical temperature

• Temperature above which a gas cannot be liquefied

• No matter how much pressure!

• For N2O 36.5 0C, - 119 0C for O2

• for CO2 = 31.1oC

Critical Pressure

• Minimum pressure that causes liquefaction

of a gas at its critical temperature

(for CO2 pc = 73 atmospheres)

• So CO2 liquefies ↓ 73 atm at 31.1 0C

Pseudocritical temperature

• Deals with gas mixture

• Temperature at which gas mixture may separate out into constituents

• Entonox 50% O2 50% N2O