The study of Gas Laws begins by exploring the fundamental relationship between the volume, pressure, and temperature of a gas. It starts with Boyle’s Law, which states that for a fixed mass of gas at a constant temperature, the volume of the gas is inversely proportional to its pressure. This means if you squeeze a gas into a smaller space (increasing pressure), its volume decreases, provided the temperature doesn’t change. Following this is Charles’ Law, which explains how gases expand when heated. It establishes that for a fixed mass of gas at constant pressure, the volume is directly proportional to its absolute temperature (measured in Kelvin). So, as you heat a gas, its volume increases, and as you cool it, the volume contracts. These two laws lay the groundwork for understanding how gases behave under different physical conditions.
Building upon these individual laws, the chapter introduces the Combined Gas Law, which merges the principles of Boyle’s and Charles’ Laws. This powerful relationship shows how pressure, volume, and temperature are all interconnected for a given mass of any gas. It logically leads to the concept of an “ideal gas” and culminates in the Ideal Gas Equation (PV = nRT), a cornerstone formula in chemistry. This equation allows for calculations involving the number of moles of a gas. Finally, the chapter often covers Avogadro’s Law, which states that equal volumes of all gases, under the same conditions of temperature and pressure, contain an equal number of molecules. This principle is crucial as it provides the link between the volume of a gas and the number of particles it contains, solidifying our understanding of gaseous behavior at the molecular level.
Understanding these laws is not just theoretical; they have significant practical applications in daily life and technology. For instance, the principles explain why a balloon inflates when you blow air into it (increasing the number of moles and thus the volume), why a pressurized deodorant can warns against incineration (heating increases pressure, risking explosion as per Gay-Lussac’s Law), and how scuba divers must manage pressure changes to avoid conditions like “the bends.” In essence, the Gas Laws provide a predictable and mathematical framework for describing the behavior of gases, which is essential for fields ranging from meteorology and medicine to engineering.
Exercise 7 (A)
Question 1.
What do you understand about gas?
Ans:
Gas is one of the fundamental states of matter, distinct from solids and liquids. Its core trait is that its particles are spaced far apart and move freely and rapidly in all directions.
This results in a few key characteristics:
- No Fixed Shape or Volume: A gas will completely fill any sealed container it’s placed in, expanding to fit the entire space.
- Compressibility: Because of the empty space between particles, gases can be squeezed into a much smaller volume under pressure.
- Low Density: Gases are much lighter for their size compared to solids and liquids.
- Diffusion: Gases naturally mix and spread out through each other over time. You can smell perfume across a room because its gaseous particles diffuse into the air.
Question 2.
Give the assumptions of the kinetic molecular theory.
Ans:
- Negligible Volume: Gas particles are so small that their total volume is insignificant compared to the total volume of their container.
- Constant, Random Motion: Particles are in continuous, rapid, and straight-line motion.
- Elastic Collisions: Particles colliding with each other and the container walls do so without any loss of kinetic energy.
- No Forces: There are no attractive or repulsive forces between the particles.
- Temperature Dependence: The average kinetic energy of the particles is directly proportional to the absolute temperature of the gas.
Question 3.
During the practical session in the lab when hydrogen sulphide gas having offensive odour is prepared for some test, we can smell the gas even 50 metres away. Explain the phenomenon.
Ans:
The gas spreads so far due to the natural process of diffusion, where particles move from a high concentration (the lab) to a low concentration (the surrounding area).
This is particularly noticeable with hydrogen sulphide because our sense of smell is extremely sensitive to it. Our noses can detect incredibly tiny amounts of the gas in the air, meaning we can smell it long before it reaches a high or dangerous concentration 50 metres away.
