Dalton's Law of Partial Pressures

Examples of Dalton's Law of Partial Pressures in the following topics:

• Dalton's Law of Partial Pressure

  • Dalton's law of partial pressures states that the pressure of a mixture of gases is the sum of the pressures of the individual components.
  • Dalton's law states that the total pressure exerted by the mixture of inert (non-reactive) gases is equal to the sum of the partial pressures of individual gases in a volume of air.
  • Dalton's law also implies that the relative concentration of gasses (their partial pressures) does not change as the pressure and volume of the gas mixture changes, so that air inhaled into the lungs will have the same relative concentration of gasses as atmospheric air.
  • Dalton's Law is only truly applicable in every situation to ideal gasses.
  • Infer from Dalton's Law of Partial Pressure the sum of partial pressures in alveoli

• Dalton's Law of Partial Pressure

  • Dalton's Law of Partial Pressure states the total pressure exerted by a mixture of gases is equal to the sum of the partial pressure of each individual gas.
  • Dalton's Law (also called Dalton's Law of Partial Pressures) states that the total pressure exerted by the mixture of non-reactive gases is equal to the sum of the partial pressures of individual gases.
  • What is the partial pressure of He?
  • We know from Boyle's Law that the total pressure of the mixture depends solely on the number of moles of gas, regardless of the types and amounts of gases in the mixture; the Ideal Gas Law reveals that the pressure exerted by a mole of molecules does not depend on the identity of those particular molecules; Dalton's Law now allows us to calculate the total pressure in a system when we know each gas individual contribution.
  • Our measurements demonstrate that the partial pressure of N2 as part of the gas PN2 is 0.763 atm, and the partial pressure of O2 as part of the gas PO2, is 0.215 atm.

• Collecting Gases Over Water

  • The amount of gas present can be determined by collecting a gas over water and applying Dalton's Law.
  • The partial pressure of H2O is known as the vapor pressure of water and is dependent on the temperature.
  • The total pressure in the tube can be written using Dalton's Law of Partial Pressures:
  • When collecting oxygen gas and calculating its partial pressure by displacing water from an inverted bottle, the presence of water vapor in the collecting bottle must be accounted for; this is easily accomplished using Dalton's Law of Partial Pressures.
  • Apply Dalton's Law to determine the partial pressure of a gas collected over water.

• Atomic Theory of Matter

  • He based his study on two laws about chemical reactions that emerged (without referring to the notion of an atomic theory) in the late 18th century.
  • The second was the law of definite proportions, first proven by the French chemist Joseph Louis Proust.
  • For this reason, Dalton is considered the originator of modern atomic theory.
  • This effort led to the development of the kinetic theory of gases, where macroscopic properties of gases, such as pressure, temperature, and volume, are explained by considering their molecular composition and motion.
  • Various atoms and molecules as depicted in John Dalton's A New System of Chemical Philosophy (1808).

• Henry's Law

  • Henry's Law states that the amount of a gas that dissolves in liquid is directly proportional to the partial pressure of that gas.
  • Henry's Law states that at a constant temperature, the amount of a gas that dissolves in a liquid is directly proportional to the partial pressure of that gas in equilibrium with that liquid.
  • The practical description for the law is that the solubility (ie. equilibrium) of a gas in a liquid is directly proportional to the partial pressure of that gas.
  • Therefore, based on the properties of Henry's Law, both the partial pressure and solubility of the oxygen and carbon dioxide determine how they will behave during gas exchange.
  • Henry's Law states that when a gas is in contact with the surface of a liquid, the amount of the gas which will go into solution is proportional to the partial pressure of that gas.

• Expressing the Equilibrium Constant of a Gas in Terms of Pressure

  • For gas-phase reactions, the equilibrium constant can be expressed in terms of partial pressures, and is given the designation KP.
  • For gas-specific reactions, however, we can also express the equilibrium constant in terms of the partial pressures of the gases involved.
  • Our equilibrium constant in terms of partial pressures, designated KP, is given as:
  • The reason we are allowed to write a K expression in terms of partial pressures for gases can be found by looking at the ideal gas law.
  • Write the equilibrium expression, KP, in terms of the partial pressures of a gas-phase reaction

• Specific Heat for an Ideal Gas at Constant Pressure and Volume

  • The heat capacity at constant pressure of 1 J·K−1 ideal gas is:
  • It is easier to measure the heat capacity at constant pressure (allowing the material to expand or contract freely) and solve for the heat capacity at constant volume using mathematical relationships derived from the basic thermodynamic laws.
  • where the partial derivatives are taken at: constant volume and constant number of particles, and at constant pressure and constant number of particles, respectively.
  • For an ideal gas, evaluating the partial derivatives above according to the equation of state, where R is the gas constant for an ideal gas yields:
  • He is best known for his 1841 enunciation of one of the original statements of the conservation of energy (or what is now known as one of the first versions of the first law of thermodynamics): "Energy can be neither created nor destroyed. " In 1842, Mayer described the vital chemical process now referred to as oxidation as the primary source of energy for any living creature.

• Measurements: Gauge Pressure and the Barometer

  • In practice, pressure is most often measured in terms of gauge pressure.
  • Gauge pressure is the pressure of a system above atmospheric pressure.
  • However, it is important to determine whether it is necessary to use absolute (gauge plus atmospheric) pressure for calculations, as is often the case for most calculations, such as those involving the ideal gas law.
  • The height of the liquid within the glass column then gives a measure of the atmospheric pressure.
  • The aneroid cell is made from beryllium-copper alloy and is partially evacuated.

• The Rate Law

  • The rate law for a chemical reaction relates the reaction rate with the concentrations or partial pressures of the reactants.
  • The rate law for a chemical reaction is an equation that relates the reaction rate with the concentrations or partial pressures of the reactants.
  • The value of this coefficient k will vary with conditions that affect reaction rate, such as temperature, pressure, surface area, etc.
  • To reiterate, the exponents x and y are not derived from the balanced chemical equation, and the rate law of a reaction must be determined experimentally.
  • The rate of the reaction can be affected by the type of reaction as well as concentration, pressure, temperature and surface area.

• Solubility and Pressure

  • Increasing pressure will increase the solubility of a gas in a solvent.
  • In this equation, C is the concentration of the gas in solution, which is a measure of its solubility, k is a proportionality constant that has been experimentally determined, and Pgas is the partial pressure of the gas above the solution.
  • Henry's law does not apply to gases at extremely high pressures.
  • To prevent the bends, a diver must return to the surface slowly, so that the gases will adjust to the partial decrease in pressure and diffuse more slowly.
  • The discoverer of Henry's law, which states that the solubility of a gas in a solvent is directly proportional to the pressure of the gas.