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More important at shallow depths is the higher temperature required to start bubble formation. Bubble formation deeper in the liquid requires a higher temperature due to the higher fluid pressure, because fluid pressure increases above the atmospheric pressure as the depth increases.
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With any incremental increase in that temperature, the vapor pressure becomes sufficient to overcome atmospheric pressure and lift the liquid to form vapor bubbles inside the bulk of the substance. The atmospheric pressure boiling point of a liquid (also known as the normal boiling point) is the temperature at which the vapor pressure equals the ambient atmospheric pressure. The vapor pressure of any substance increases non-linearly with temperature according to the Clausius–Clapeyron relation. As the kinetic energy of the molecules increases, the number of molecules transitioning into a vapor also increases, thereby increasing the vapor pressure. As the temperature of liquid increases, the kinetic energy of its molecules also increases.
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The pressure exhibited by vapor present above a liquid surface is known as vapor pressure. A substance with a high vapor pressure at normal temperatures is often referred to as volatile. It relates to the tendency of particles to escape from the liquid (or a solid). The equilibrium vapor pressure is an indication of a liquid's evaporation rate. Vapor pressure (or vapour pressure in English-speaking countries other than the US see spelling differences) or equilibrium vapor pressure is defined as the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phases (solid or liquid) at a given temperature in a closed system. By heating the alcohol, the vapors fill in the space, increasing the pressure in the tube to the point of the cork popping out. The tube contains alcohol and is closed with a piece of cork.
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