Air pressure, the outside force we depend on
Air pressure is the force exerted by the weight of air molecules which obey the law of gravity. Atmospheric pressure is determined by the amount of air directly above a person or object. At sea level (mean sea level pressure (MSLP), the atmospheric pressure is 1013.25 mbar (101.325 kPa; 29.921 inHg; 760.00 mmHg) or PSI. At higher altitudes, the PSI decreases due to lower air pressure and density. Skin adjusts easily to changes in pressure, but the cavities within the body, such as the lungs, ears, and sinuses, do not adjust automatically. This is why many people experience a popping in their ears while taking off in an airplane or driving through mountains.
The highest adjusted-to-sea level barometric pressure ever recorded on Earth (above 750 meters) was 1084.8 hPa (32.03 inHg) measured in Tosontsengel, Mongolia on 19 December 2001. The highest adjusted-to-sea level barometric pressure ever recorded (below 750 meters) was at Agata in Evenk Autonomous Okrug, Russia (66°53’N, 93°28’E, elevation: 261 m, 856 ft) on 31 December 1968 of 1083.8 hPa (32.005 inHg). The Dead Sea, the lowest place on Earth at 430 meters (1,410 ft) below sea level, has a correspondingly high typical atmospheric pressure of 1065 hPa.
The lowest non-tornadic atmospheric pressure ever measured was 870 hPa (0.858 atm; 25.69 inHg), set on 12 October 1979, during Typhoon Tip in the western Pacific Ocean. The measurement was based on an instrumental observation made from a reconnaissance aircraft.
Pure water boils at 100 °C (212 °F) at earth’s standard atmospheric pressure. The boiling point is the temperature at which the vapor pressure is equal to the atmospheric pressure around the water. Because of this, the boiling point of water is lower at lower pressure and higher at higher pressure. Cooking at high elevations, therefore, requires adjustments to recipes. A rough approximation of elevation can be obtained by measuring the temperature at which water boils.
The pressure of air that is present outside the body is the same as that of the air present ‘inside’ the body. The air that is constantly present in the lungs, ears, and nose has the same atmospheric pressure as the air on the outside of the ears, nose, and chest. (Newton’s third law)
Just below the lungs is a muscle called the diaphragm. When a person breathes in, the lungs get air in it (or expands). The lungs on expansion move the diaphragm down. The diaphragm, which is a dome-shaped muscle becomes more “flattened”. When the lung volume increases, the pressure in the lungs decreases (Boyle’s law). Since air always moves from areas of high pressure to areas of lower pressure, air will now be drawn into the lungs because the air pressure outside the body is higher than the pressure in the lungs.
The opposite process happens when a person breathes out. When a person breathes out the diaphragm moves upwards and causes the volume of the lungs to decrease, the air inside lungs takes up the lesser volume or has now higher pressure. The pressure in the lungs will increase, and the air that was in the lungs is forced out towards the lower air pressure outside the body.
Human bodies are used to air pressure. The air pressure in the lungs, ears, and stomachs is the same as the air pressure outside of the body, which ensures that nobody gets crushed. The body is flexible enough to cope when the internal and external pressures aren’t exactly the same.
The effects of air pressure on the body can be illustrated by:
Blood pressure: Just as its name implies, blood moves through the body using a pressure system created by the heart. It makes sense that this pressure would be affected by the pressure in the air around. According to biometeorologist Jennifer Vanos, Ph.D., when the barometric pressure drops, so does the blood pressure. For some, this might mean a feeling of dizziness or even blurred vision.
Headaches: In an interview with the New York Times, Dr. Matthew Fink, a neurologist in chief at New York-Presbyterian Hospital/Weill Cornell Medical Center, explained that low barometric pressure can cause headaches or migraines by creating a pressure difference between the atmosphere and the air-filled sinuses. The problem is exacerbated when the sinuses are congested or blocked for any reason.
Joint pain: Researchers at Tufts-New England Medical Center in Boston surveyed 200 patients with knee osteoarthritis and found a link between changes in barometric pressure and ambient temperature and changes in knee pain severity. It’s not clear why a falling barometer would exacerbate joint pain and arthritis. It could be that barometric pressure affects the viscosity of the fluid that lines joint sacs, or it could be that it triggers the pain responses in the nerve endings of the joint. Either way, it’s what your grandma has been saying for years: Some people feel pain in their joints when a storm is approaching.
|Boyle’s law, also called Mariotte’s law, a relation concerning the compression and expansion of a gas at constant temperature. This empirical relation, formulated by the physicist Robert Boyle in 1662, states that the pressure (p) of a given quantity of gas varies inversely with its volume (v) at a constant temperature; i.e., in equation form, pv = k, a constant.
The British scientist James Clerk Maxwell and the Austrian physicist Ludwig Boltzmann, in the 19th century, led in establishing the theory, which became one of the most important concepts in modern science.
The simplest kinetic model is based on the assumptions that: (1) the gas is composed of a large number of identical molecules moving in random directions, separated by distances that are large compared with their size; (2) the molecules undergo perfectly elastic collisions (no energy loss) with each other and with the walls of the container, but otherwise do not interact; and (3) the transfer of kinetic energy between molecules is heat. These simplifying assumptions bring the characteristics of gases within the range of mathematical treatment.
Such a model describes a perfect gas (q.v.) and is a reasonable approximation to a real gas, particularly in the limit of extreme dilution and high temperature. Such a simplified description, however, is not sufficiently precise to account for the behavior of gases at high densities.
Based on the kinetic theory, pressure on the container walls can be quantitatively attributed to random collisions of molecules the average energy of which depends upon the gas temperature. The gas pressure can, therefore, be related directly to temperature and density. Many other gross properties of the gas can be derived, such as viscosity, thermal and electrical conduction, diffusion, heat capacity, and mobility. In order to explain observed deviations from perfect gas behavior, such as condensation, the assumptions must be appropriately modified. In doing so, considerable insight has been gained as to the nature of molecular dynamics and interactions.