In scuba-diving, while ascending towards the water surface, there is a danger of bursting the lungs. It is because of

1. Kinetic Theory and the Concept of Gas Pressure (basic)

To understand the atmosphere or the behavior of gases in any environment, we must first look at matter through a microscope. According to the Kinetic Theory of Gases, matter is not a continuous block but is made of tiny particles. In a gaseous state, these particles possess high kinetic energy, allowing them to overcome attractive forces and move freely in all directions at high speeds Science, Class VIII, Particulate Nature of Matter, p.112. Because they are constantly zooming around, they inevitably collide with one another and, more importantly, with the walls of whatever container they are in.

This brings us to the fundamental definition of Gas Pressure. In physics, pressure is defined as force per unit area (P = F/A) Science, Class VIII, Pressure, Winds, Storms, and Cyclones, p.94. When billions of gas particles strike a surface, their collective impacts exert a steady force. The SI unit for this measurement is the Pascal (Pa). It is important to remember that gases exert pressure in all directions equally—whether it is the air in a balloon or the atmospheric pressure pressing against our bodies at sea level.

State of Matter	Particle Movement	Pressure Characteristics
Solid	Vibrate in fixed positions	Exerts pressure primarily downwards due to gravity.
Liquid	Slide over each other	Exerts pressure on the bottom and sides of the container.
Gas	Rapid, random motion	Exerts high pressure in all directions on container walls.

One of the most critical relationships to master is Boyle’s Law. It states that at a constant temperature, the pressure of a gas is inversely proportional to its volume (P ∝ 1/V). Think of it this way: if you squeeze a gas into a smaller space (decreasing volume), the particles have less room to move and will hit the walls much more frequently. This increase in the frequency of collisions results in higher pressure. Conversely, if you expand the volume, the particles hit the walls less often, and the pressure drops. This principle is why warm air rising in our atmosphere expands as it encounters lower environmental pressure Science, Class VIII, Pressure, Winds, Storms, and Cyclones, p.94.

Sources: Science, Class VIII, Particulate Nature of Matter, p.112; Science, Class VIII, Pressure, Winds, Storms, and Cyclones, p.94

2. Hydrostatic Pressure: The Weight of the Ocean (basic)

To understand why the ocean exerts such immense force, we must start with the fundamental definition of pressure: it is the force exerted per unit area (Science, Class VIII NCERT, Pressure, Winds, Storms, and Cyclones, p.94). When you are standing on land, you are at the bottom of an "ocean of air." This air has weight, and it pushes down on you with a standard sea-level pressure of approximately 1,013.25 millibars (mb) (Physical Geography by PMF IAS, Pressure Systems and Wind System, p.305).

However, when you dive into the ocean, the situation changes dramatically. Hydrostatic pressure is the pressure exerted by a fluid at rest due to the force of gravity. Because water is much denser than air, it is significantly heavier. As you descend, the weight of the water column above you increases. Consequently, pressure increases with increasing depth (NCERT Class XI Geography, The Origin and Evolution of the Earth, p.19). While it takes miles of atmosphere to double the pressure we feel on land, it only takes 10 meters (about 33 feet) of sea water to add another full atmosphere of pressure.

A crucial characteristic of fluid pressure is that it is exerted in all directions. Whether it is a gas or a liquid, the substance pushes against every surface it touches—including the walls of a container or the tissues of a diver's body (Science, Class VIII NCERT, Pressure, Winds, Storms, and Cyclones, p.94). This means a submerged object isn't just being pushed "down" by the weight above; it is being squeezed from every side simultaneously. This mounting external pressure has profound effects on any gas-filled spaces, such as the lungs or a diving bell, because as the external pressure rises, the volume of the gas inside must respond.

Depth	Approximate Pressure	Comparison
0 m (Sea Level)	1 Atmosphere (1 ATM)	Standard air pressure at the surface.
10 m (33 ft)	2 Atmospheres (2 ATM)	Pressure is doubled compared to the surface.
20 m (66 ft)	3 Atmospheres (3 ATM)	Triple the surface pressure.

Sources: Science, Class VIII NCERT, Pressure, Winds, Storms, and Cyclones, p.94; Physical Geography by PMF IAS, Pressure Systems and Wind System, p.305; NCERT Class XI Geography, The Origin and Evolution of the Earth, p.19

3. Henry’s Law: Gas Solubility and 'The Bends' (intermediate)

To understand why divers face risks underwater, we must first look at how gases behave when they meet liquids. Henry's Law states that the amount of gas dissolved in a liquid is directly proportional to the partial pressure of that gas above the liquid. In simpler terms: the more pressure you apply, the more gas you can 'force' to dissolve into the liquid. This forms a uniform mixture or a solution Science, Class VIII NCERT, Chapter 9, p.139. You see this every time you open a carbonated drink; the 'hiss' is the sound of gas escaping because the pressure inside the bottle has suddenly dropped, reducing the gas's solubility.

Under the sea, the weight of the water creates immense pressure. As a scuba diver descends, the increased atmospheric and water pressure forces more nitrogen from their air tank to dissolve into their blood and tissues. While this isn't a problem at depth, it becomes a major health risk during the ascent. If a diver rises to the surface too quickly, the external pressure drops rapidly. Following Henry's Law, the nitrogen can no longer stay dissolved in the blood and begins to form bubbles, much like the bubbles in a freshly opened soda. These bubbles can block blood flow or damage tissues, a painful and dangerous condition known as 'The Bends' (Decompression Sickness).

It is also important to distinguish how pressure and temperature affect solubility differently. While higher pressure increases the solubility of a gas, higher temperatures generally decrease it Science, Class VIII NCERT, Chapter 9, p.139. This is why aquatic life, which depends on dissolved oxygen, often thrives better in cold water than in very warm water Science, Class VIII NCERT, Chapter 9, p.149.

Factor	Effect on Gas Solubility	Real-world Example
Increased Pressure	Increases Solubility	Nitrogen dissolving in a diver's blood at depth.
Decreased Pressure	Decreases Solubility	CO₂ escaping a soda bottle when opened.
Increased Temperature	Decreases Solubility	Warm water holding less dissolved oxygen for fish.

Sources: Science, Class VIII NCERT, Chapter 9 — The Amazing World of Solutes, Solvents, and Solutions, p.139; Science, Class VIII NCERT, Chapter 9 — The Amazing World of Solutes, Solvents, and Solutions, p.149

4. Dalton’s Law of Partial Pressures (intermediate)

To understand Dalton’s Law of Partial Pressures, we must first look at what air actually is. As we have learned, air is not a single substance but a uniform mixture of various gases—primarily Nitrogen (78%), Oxygen (21%), and traces of Argon and Carbon Dioxide Science, Class VIII NCERT, Nature of Matter: Elements, Compounds, and Mixtures, p.118. In thermal physics, we know that pressure is the result of gas molecules colliding with the walls of their container Science, Class VIII NCERT, Pressure, Winds, Storms, and Cyclones, p.81.

John Dalton formulated a fundamental principle: in a mixture of non-reacting gases, each gas acts independently of the others. Dalton’s Law states that the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of each individual gas in the mixture. A "partial pressure" is simply the pressure that a specific gas would exert if it were the only occupant of that entire volume at the same temperature.

The mathematical expression for this is straightforward:

P_total = P₁ + P₂ + P₃ + ... + Pₙ

For example, if you have a container of dry air at sea level where the total atmospheric pressure is 101.3 kPa, that pressure is the combined result of the partial pressure of Nitrogen (≈79 kPa), Oxygen (≈21 kPa), and other trace gases Physical Geography by PMF IAS, Earth's Atmosphere, p.270. This concept is vital because it explains how gases like water vapour exert their own independent pressure, known as vapour pressure, which influences processes like evaporation and storm formation Physical Geography by PMF IAS, Tropical Cyclones, p.358.

Sources: Science, Class VIII NCERT, Nature of Matter: Elements, Compounds, and Mixtures, p.118; Science, Class VIII NCERT, Pressure, Winds, Storms, and Cyclones, p.81; Physical Geography by PMF IAS, Earth's Atmosphere, p.270; Physical Geography by PMF IAS, Tropical Cyclones, p.358

5. Charles’s Law and Gay-Lussac’s Law: Temperature Effects (intermediate)

When we study the behavior of gases under the influence of heat, we are essentially looking at how kinetic energy affects the arrangement of particles. According to the particulate nature of matter, as the temperature of a substance increases, its particles move faster and tend to spread out Science, Class VIII, NCERT (Revised ed 2025), The Amazing World of Solutes, Solvents, and Solutions, p.147. How the gas reacts to this movement depends on whether the container is flexible or rigid, leading us to two fundamental gas laws.

Charles’s Law describes what happens when pressure is kept constant. In this scenario, as the temperature (T) increases, the volume (V) of the gas increases proportionally (V ∝ T). This expansion occurs because the faster-moving particles push outward to maintain a steady pressure. A classic example is a hot air balloon or a rising air parcel in our atmosphere. When an air parcel is heated, its volume increases, making it less dense than the surrounding cooler air, which causes it to rise Physical Geography by PMF IAS, Vertical Distribution of Temperature, p.297.

Gay-Lussac’s Law, on the other hand, applies when the volume is kept constant (such as in a rigid metal tank). If you heat a gas in a fixed volume, the particles cannot spread out to increase volume; instead, they collide with the container walls more frequently and with greater force. This results in an increase in pressure (P) proportional to the temperature (P ∝ T). This explains why pressure cookers work or why sealed containers are hazardous when exposed to high heat.

Law	Constant Factor	Relationship	Physical Result of Heating
Charles’s Law	Pressure	V ∝ T	Gas expands (Volume increases)
Gay-Lussac’s Law	Volume	P ∝ T	Gas exerts more force (Pressure increases)

The link between these laws and density is crucial for UPSC topics like meteorology. Since Density = Mass / Volume, heating a gas in a flexible environment increases its volume while the mass remains the same. Consequently, the density decreases Science, Class VIII, NCERT (Revised ed 2025), The Amazing World of Solutes, Solvents, and Solutions, p.147. This lower density is the driving force behind atmospheric convection and the vertical distribution of temperature.

Sources: Science, Class VIII, NCERT (Revised ed 2025), The Amazing World of Solutes, Solvents, and Solutions, p.147; Physical Geography by PMF IAS, Vertical Distribution of Temperature, p.297

6. Boyle’s Law: The Inverse Pressure-Volume Relationship (exam-level)

Boyle’s Law is a fundamental principle of fluid mechanics and thermodynamics that describes the inverse relationship between the pressure and volume of a gas. To put it simply: if you keep the temperature constant, as the pressure on a gas increases, its volume decreases, and vice versa. This happens because gases are highly compressible; the space between gas particles is vast, and applying pressure forces those particles closer together Science Class VIII NCERT, Particulate Nature of Matter, p. 105.

The mathematical expression of this law is P ∝ 1/V (where P is pressure and V is volume), or more commonly used for problem-solving: P₁V₁ = P₂V₂. This equation tells us that the product of pressure and volume remains a constant value for a fixed amount of gas at a steady temperature. In practical terms, if you double the pressure (P), the volume (V) must be cut in half to keep the equation balanced. This concept is vital for understanding how a parcel of rising air expands as it moves up into the atmosphere where the ambient pressure is lower Physical Geography by PMF IAS, Chapter 22, p. 297.

A classic application of Boyle’s Law, often cited in safety and environmental science, is scuba diving. When a diver is 33 feet (10 meters) underwater, they experience 2 atmospheres of pressure (2 ATA)—one from the air above and one from the weight of the water. If the diver breathes in air at this depth and then ascends to the surface (1 ATA) without exhaling, the surrounding pressure drops by half. According to Boyle’s Law, the volume of that air in their lungs must double. Because the human lungs have a limited capacity, this expansion can cause the lung tissue to rupture, a condition known as pulmonary barotrauma. This is why divers are trained to never hold their breath while ascending.

Sources: Science Class VIII NCERT, Particulate Nature of Matter, p.105; Physical Geography by PMF IAS, Vertical Distribution of Temperature, p.297

7. Pulmonary Barotrauma and Diving Physiology (exam-level)

To understand why diving can be hazardous to the human respiratory system, we must apply Boyle's Law from thermal physics. This law states that for a fixed amount of gas at a constant temperature, the volume (V) of the gas is inversely proportional to its pressure (P). In mathematical terms, P₁V₁ = P₂V₂. When a scuba diver is underwater, they are subjected to the weight of the water above them. For every 10 meters (33 feet) of depth, the pressure increases by approximately 1 atmosphere (ATA). Therefore, at 10 meters, a diver experiences 2 ATA of pressure (1 from the atmosphere and 1 from the water).

The danger of Pulmonary Barotrauma (lung overexpansion injury) occurs specifically during the ascent. As the diver moves toward the surface, the ambient water pressure decreases. According to Boyle's Law, as pressure (P) drops, the volume (V) of the air trapped in the lungs must increase. If a diver holds their breath while ascending, the air inside the alveoli—the tiny, balloon-like structures responsible for gas exchange—expands beyond the lungs' physical capacity Science, class X (NCERT 2025 ed.), Life Processes, p.90. This expansion can cause the lung tissue to rupture, leading to air escaping into the chest cavity or the bloodstream.

The physiological impact of pressure changes is much more volatile than typical environmental lung stressors. While conditions like Pneumoconiosis (Black Lung) result from the long-term inhalation of coal dust Environment, Shankar IAS Academy (ed 10th), Environment Issues and Health Effects, p.416, or nitrogen oxides cause chronic inflammation Environment and Ecology, Majid Hussain (3rd ed.), Environmental Degradation and Management, p.40, pulmonary barotrauma is an acute mechanical failure. Just ascending from 10 meters to the surface causes the air volume in the lungs to double. This is why controlled ascent and continuous breathing are the most critical safety protocols in diving physiology.

Sources: Science, class X (NCERT 2025 ed.), Life Processes, p.90; Environment, Shankar IAS Academy (ed 10th), Environment Issues and Health Effects, p.416; Environment and Ecology, Majid Hussain (3rd ed.), Environmental Degradation and Management, p.40

8. Solving the Original PYQ (exam-level)

Now that you have mastered the fundamental gas laws, this question tests your ability to apply the Pressure-Volume relationship to a physiological scenario. As you learned in Physical Geography by PMF IAS, any parcel of rising air expands as it encounters lower ambient pressure. In the context of scuba diving, your lungs act as the container for that air. The core concept here is understanding how a change in the environment (water depth) translates into a physical change in the gas trapped inside the body.

To arrive at the correct answer, you must apply Boyle's Law, which states that at a constant temperature, the volume of a gas is inversely proportional to its pressure. As a diver ascends, the weight of the water column above them decreases, leading to a significant drop in ambient pressure. According to the formula P1V1 = P2V2, when pressure (P) decreases, the volume (V) must increase. If a diver holds their breath, the air trapped in the lungs expands rapidly; for example, ascending from just 33 feet to the surface causes the air volume to double. This expansion can exceed the lungs' physical capacity, leading to "bursting" or pulmonary barotrauma. Thus, (B) Boyle's Law is the only principle that explains this mechanical expansion.

UPSC frequently uses "distractor" laws to test whether you can distinguish between related physical phenomena. Archimedes' principle is a common trap because it relates to buoyancy (why a diver floats), but it does not govern gas expansion. Gay-Lussac’s law of combining volumes refers to the ratios of gases in chemical reactions, which is irrelevant to pressure changes. Similarly, Graham’s law of diffusion explains how gases mix and move (relevant to nitrogen absorption in the blood), but it does not account for the physical volume change that causes lung injury. By focusing on the inverse relationship between pressure and volume, you can confidently eliminate these decoys.