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*What is a gas?*

Gases come into view as materials of very low density which must be enclosed to keep them altogether. Dissimilar solids, gases encompass no definite shape. Unlike liquids, gases have no definite volume however they entirely fill the container. The volume of the container is the volume of the gas in it. A gas applies a pressure on all sides of the container which holds it.

Gas can be compressed via pressures greater than the pressure of the gas on its container. The terms fume, vapor and air as well illustrate a gas. Air illustrates the common mixture of gases in the atmosphere. The terms vapor and fume recommend that the gas comes from a specific liquid.

*Gaseous State of Matter:*

Gaseous state is one of the physical states of matter. Close to absolute zero of temperature, a substance exists as a solid. As heat is added to this substance, it melts to a liquid at its melting point, boils to a gas at its boiling point. A pure gas might be made up of individual atoms (example: a noble gas such as neon), elemental molecules made up from one kind of atoms (example: oxygen) or compound molecules made up from a variety of atoms (example: carbon dioxide). A gas mixture would have a variety of pure gases much similar to the air.

** Physical characteristics of Gases**:

In a gas, the particles are the farthest apart, encompass the least force of attraction among them and encompass the highest freedom of motion to the particles as compared to that in solids and liquids. In a gas, the particles have adequate kinetic energy in such a way that the effect of the intermolecular force is small (or zero for the ideal gas) and the typical distance between neighboring particles is much bigger than the size of the particles themselves. A gas consists of no definite shape or volume however occupies the full container in which it is limited. A gas can be liquefied through compression devoid of cooling or a liquid can be transformed to a gas via decreasing the pressure at constant temperature.

** The Ideal Gas Law**:

The ideal gas law is the equation of state of the hypothetical ideal gas. This is a good estimation to the behavior of numerous gases under numerous conditions; however it consists of lots of limitations.

The ideal gas law ignores both molecular size and intermolecular attractions therefore the law is most precise for monoatomic gases at high temperatures and low pressures. The ignorance of molecular size becomes less significant for bigger volumes that is, for lower pressures. The relative significance of intermolecular attractions reduces with increasing thermal kinetic energy that is, by increasing temperature. More complicated equations of states like Van-der-Waals equation, let deviations from ideality caused by the molecular size and intermolecular forces to be taken into account.

The ideal gas law is mainly based on the given physical estimations:

i) The gas molecules are moving in the arbitrary directions.

ii) Molecules of a gas contain negligible size as compared to the average distance between them.

iii) The interaction between gas molecules can be ignored as the average distance between the molecules is extremely large.

iv) Whenever molecules collide by one other or the container, the collisions are elastic that is, no kinetic energy is lost.

v) The average energy of gas molecules is proportional to the temperature of the gas.

A gas might be fully illustrated by its makeup - pressure, volume and temperature. The ideal gas law represents an equation PV = n RT, here 'P' is the pressure, 'V' is the volume, 'T' is the absolute temperature, 'n' is the number of moles of the gas and 'R' is the Universal Gas Constant.

** Different Gas Laws**:

The early gas laws were mainly developed at the end of the 18th century, when scientists start to realize that the relationships between the pressure, volume and temperature of a gas could be obtained that would hold for all gases. The former gas laws are now taken as special cases of the ideal gas law having one or more of the variables kept constant.

Boyle's Law:

Boyle's law is one of the some gas laws and a special case of the ideal gas law. It illustrates the inversely proportional relationship between the absolute pressure and volume of a gas at constant temperature. The Boyle's law was introduced in the year 1662. It is stated as:

Temperature remaining constant, the volume of a particular mass of an ideal gas is inversely proportional to the pressure applied by the gas.

Mathematically, Boyle's law is represented as follows:

P α 1/V or PV = K (constant)

Charles' Law:

Charles' law is the other gas law that presents a special case of the ideal gas law. It illustrates the direct relationship between the absolute temperature and volume of a gas at constant pressure. The law was first introduced in the year 1787 and then confirmed in the year 1802. Charles' law is represented as follows:

Pressure remaining constant, the volume of a particular mass of gas increases or decreases by 1/273 times its volume at 0^{o}C for every one degree increase or fall in temperature correspondingly.

The modern version of Charles law is represented as:

At constant pressure, the volume of a given mass of the ideal gas is directly proportional to the absolute temperature.

Mathematically, Charles' law is represented as:

V α T or V/T = K (constant)

** Absolute Scale of Temperature**:

Charles' law appears to mean that as we decrease the temperature, volume of a gas goes on reducing. Ultimately, at some phase (at - 273 ^{o}C), the volume of gas would become zero. However this is not possible. Gas molecules themselves encompass some volume in such a way that they will occupy some volume. This signifies the volume of a gas will never be zero.

The temperature at which the volume of a gas would theoretically be decreased to zero is termed as absolute zero of temperature.

The temperature scale having absolute zero as the beginning point is termed as the absolute scale of temperature.

This scale was recommended by Thomson and Kelvin in the year 1848. It is termed as Kelvin scale or absolute scale. All the temperatures on Kelvin scale are positive. Kelvin temperature can be acquired by adding 273 to the Celsius temperature. Kelvin temperature is represented in 'K' or 'A' however no degree symbol is linked to it.

Combined Gas Law Equation:

The combined gas law illustrates the interdependence of variables such as pressure, volume and temperature. The statement can be represented as:

The ratio of pressure-volume product and the temperature of a system remain constant for a particular amount of gas.

Mathematically, it is represented as PV/T = K (constant)

For comparing the similar substance under different sets of conditions, the law can be represented as:

P_{1}V_{1}/T_{1 }= P_{2}V_{2}/T_{2}

Gay Lussac's Law of Combining Volumes:

In the year 1808, Gay Lussac, while studying the reactions of gases made an examination that gases react in simple proportion of their volumes. The Gay Lussac's law of combining volumes is represented as:

At constant pressure and temperature, the volumes of the gaseous reactants and products can be represented as ratios of whole numbers.

For illustration: hydrogen and oxygen react in volume ratio 2/1 to form a volume of water vapor equivalent to the initial volume of hydrogen or two times the initial volume of oxygen. In other illustration: nitrogen and oxygen can react in volume ratio 2/1 or 1/1 based on whether the product is dinitrogen oxide N_{2}O or NO.

Such reactions can be represented as:

2 H_{2} (g) + O_{2} (g) → 2 H_{2}O (g)

2 N_{2} (g) + O_{2} (g) → 2 N_{2}O (g)

N_{2} (g) + O_{2} (g) → 2 NO (g)

Avogadro's Law:

In the year 1811, Avogadro established a relationship between the volume and number of moles of various gases under identical conditions of pressure and temperature. Avogadro's law is as well known as Avogadro's principle or Avogadro's hypothesis. Avogadro's law is represented as:

Equivalent volumes of different ideal or perfect gases, at similar pressure and temperature, contain similar number of molecules.

Therefore, the number of molecules in a particular volume of a gas is independent of the size or mass of gas molecules at a particular pressure and temperature.

Mathematically, it is represented as V/n = K (constant)

Here, 'V' is the volume and 'n' is the number of moles of the gas.

Dalton's Law of Partial Pressure:

One of the significant predictions made via Avogadro is that the identity of a gas is insignificant in finding out P-V-T properties of the gas. This behavior signifies that a gas mixture acts in exactly the similar way as a pure gas. In the year 1801, Dalton introduced the law of partial pressures. Dalton's law of partial pressures is represented as:

The net or total pressure of a mixture of non-reacting gases is equivalent to the sum of the pressures that each and every gas would apply, if it was present separately.

In this case, the pressure applied by each and every gas in the mixture is known as the partial pressure of that gas.

Mathematically, it is represented as P total = P_{1} + P_{2} + ------- P_{n}

Here, P_{1}, P_{2}, P_{n} and so on are the partial pressures of gas 1, 2 --- n.

Dalton's law of partial pressure states us that each and every gas behaves independently of the other gases in the mixture that is; each and every gas applies its own pressure (that is, partial pressure).

Graham's Law of Diffusion:

Gases encompass a tendency to spontaneously intermix and form a homogeneous mixture devoid of the help of external agency. This is termed as the diffusion of gases. Diffusion is a physical phenomenon. The gases diffuse very rapidly due to the presence of large empty spaces between the gas molecules which makes their movement rapid to one other. Gases move from an area of higher concentration to a area of lower concentration till the mixture acquires uniform concentration.

Graham studied the rate of diffusion of different gases and provides a quantitative relation between the density and rate of diffusion of gases in the year 1837. Graham's law of diffusion can be stated as:

The rate of diffusion of a gas is inversely proportional to the square root of its density.

Mathematically, it is represented as r α 1/(d)^{1/2} or r = K/(d)^{1/2} (K = constant)

It follows then that the comparative rates of diffusion of two gases are inversely proportional to the square root of their densities. Under identical conditions of pressure and temperature, the relation between the comparative rates of diffusion and densities is represented as:

r_{1}/r_{2 }= (d_{2}/d_{1})^{1/2}

As density is proportional to the molecular mass that is, d α M, we can substitute density via molecular mass and represent the equation as:

r_{1}/r_{2} = (M_{2}/M_{1})^{1/2 }

The rate of diffusion is represented by the formula:

Rate of diffusion = Volume of the gas diffused/Time taken for diffusion

Whenever the gas contained in the vessel is allowed to escape via a small aperture or pin hole, the phenomenon is termed as effusion. Graham's law is as well applicable to the effusion of gases.

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