Chapter 5 - Gases

5.1 - Substances That Exist as Gases

5.2 - Pressure of a Gas

• Any place exposed to Earth's atmosphere feels a force equal to the weight of the air column above it.
  • The pressure exerted by the Earth's atmosphere is known as atmospheric pressure.
• The barometer is arguably the most well-known tool for determining atmospheric pressure.
• At 0°C at sea level, standard atmospheric pressure (1 atm) is equal to the pressure that sustains a 760 mm (or 76 cm) tall column of mercury.
• A manometer is a device that is used to determine the pressure of gases other than air.

5.3 - The Gas Laws

• The volume occupied by the gas reduces when the pressure is increased. If the applied pressure is reduced, the volume occupied by the gas increases.

• Boyle's law is the name given to this relationship.

  • Boyle's law states that under constant temperature, the pressure of a fixed amount of gas is inversely proportional to the volume of the gas.

• Lord Kelvin recognized the significance of this occurrence in 1848.

  • Absolute zero, or 273.15°C, is the lowest temperature that may theoretically be reached.

• After that, he established an absolute temperature scale, currently known as the Kelvin temperature scale, with absolute zero as the beginning point.

  • One kelvin (K) is the same as one degree Celsius on the Kelvin scale.

• The sole difference between the absolute and Celsius temperature scales is that the zero position is changed in the former.

  • Important points on both scales are in agreement.

• The volume of a fixed amount of gas maintained at constant pressure is exactly proportional to the absolute temperature of the gas, according to Charles' law.

  • Avogadro's law asserts that the volume of a gas is directly proportional to the number of moles present under constant pressure and temperature.

5.4 - The Ideal Gas Equation

• The proportionality constant, R, is referred to as the gas constant.
• The relationship between the four variables P, V, T, and n is described by the ideal gas equation.
  • A hypothetical gas whose pressure-volume-temperature behavior can be perfectly explained by the ideal gas equation is referred to as an ideal gas.

5.5 - Gas Stoichiometry

• We can also use the relationships between quantities (moles, n) and volume (V) to answer analogous problems when the reactants and/or products are gases.

5.6 - Dalton’s Law of Partial Pressures

• The total gas pressure is related to partial pressures in all circumstances involving gas mixtures.
  • Individual gas component pressures in the mixture.
• Dalton established the law of partial pressures in 1801, which states that the total pressure of a mixture of gases is equal to the sum of the pressures that each gas would exert if it were present alone.

5.7 - The Kinetic Molecular Theory of Gases

• Gas Compressibility is the ability of a gas to be compressed.
  • Because gas molecules are separated by huge distances, gases may readily be compressed to take up less space.
• Boyle's Law is a set of rules that governs how things are done.
• The impact of a gas's molecules on the container's walls causes pressure to be exerted by it.
• The number of molecule collisions with the walls every second, or collision rate, is proportional to the gas's number density.
  • When the volume of a gas is reduced, the number density rises, and the collision rate rises as well.
  • As a result, a gas's pressure is inversely proportional to the volume it takes up; as volume drops, pressure rises, and vice versa.
• The Law of Charles.
  • The average kinetic energy of gas molecules is proportional to the absolute temperature of the sample, hence increasing the temperature increases the average kinetic energy.
  • As a result, if the gas is heated, molecules will collide with the container's walls more frequently and with more force, increasing the pressure.
  • The volume of gas expands until the pressure of the gas is balanced by the constant external pressure.
• The Law of Avogadro is related to both its density and its temperature.
  • Density is represented as n/V because the mass of the gas is directly proportional to the number of moles (n) in the gas.
• The Law of Partial Pressures by Dalton, The pressure generated by one type of molecule is unaffected by the presence of another gas if molecules do not attract or repel one another.
• Diffusion, the slow mixing of molecules of one gas with molecules of another due to their kinetic properties, provides direct evidence of gaseous random motion.
• Thomas Graham, a Scottish scientist, discovered in 1832 that rates of diffusion for gases are inversely proportional to the square roots of their molar weights under the same conditions of temperature and pressure.
  • Diffusion is the process of a gas progressively mixing with another, while effusion is the process of a gas under pressure escaping from one compartment of a container to another through a small opening.

5.8 - Deviation from Ideal Behavior

• The gas laws and kinetic molecular theory assume that molecules in the gaseous state have no attraction or repulsive force on one another.
• The other assumption is that the molecules' volume is insignificant in comparison to the container's. Ideal behavior is defined as a gas that meets both of these criteria.