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Dalton's law

For the law of stoichiometry, see Law of multiple proportions.

In chemistry and physics, Dalton's law (also called Dalton's law of partial pressures) states that the total pressure exerted by a gaseous mixture is equal to the sum of the partial pressures of each individual component in a gas mixture. This empirical law was observed by John Dalton in 1801 and is related to the ideal gas laws.  Mathematically, the pressure of a mixture of gases can be defined as the summation

or :P_{total} = \sum_{i=1} ^ n {p_i}      or      P_{total} = p_1 +p_2 + \cdots + p_n where p_{1},\ p_{2},\ p_{n} represent the partial pressure of each component.  It is assumed that the gases do not react with each other.

\ P_{i} =P_{total}m_i

where m_i\ = the mole fraction of the i-th component in the total mixture of m components.

The relationship below provides a way to determine the volume based concentration of any individual gaseous component.

P_i =\frac{P_{total}C_i}{1,000,000}

where: C_i\ = is the concentration of the ith component expressed in ppm.

Dalton's law is not exactly followed by real gases. Those deviations are considerably large at high pressures. In such conditions, the volume occupied by the molecules can become significant compared to the free space between them. Moreover, the short average distances between molecules raises the intensity of intermolecular forces between gas molecules enough to substantially change the pressure exerted by them. Neither of those effects are considered by the ideal gas model.

stoichiometry - difinition
Stoichiometry rests upon the law of conservation of mass, the law of definite proportions (i.e., the law of constant composition) and the law of multiple proportions. In general, chemical reactions combine in definite ratios of chemicals. Since chemical reactions can neither create nor destroy matter, nor transmute one element into another, the amount of each element must be the same throughout the overall reaction. For example, the amount of element X on the reactant side must equal the amount of element X on the product side.

Stoichiometry is often used to balance chemical equations. For example, the two diatomic gases, hydrogen and oxygen, can combine to form a liquid, water, in an exothermic reaction, as described by the following equation:

2H_2 + O_2 \rightarrow 2H_2O\,  [[Media:2H_2 + O_2 \rightarrow 2H_2O\,]]

The term stoichiometry is also often used for the molar proportions of elements in stoichiometric compounds. For example, the stoichiometry of hydrogen and oxygen in H2O is 2:1. In stoichiometric compounds, the molar proportions are whole numbers (that is what the law of definite proportions is about).

Compounds for which the molar proportions are not whole numbers are called non-stoichiometric compounds.

Stoichiometry is not only used to balance chemical equations but also used in conversions, i.e., converting from grams to moles, or from grams to milliliters. For example, to find the number of moles in 2.00 g of NaCl, one would do the following:

\frac{2.00 \mbox{ g NaCl}}{58.44 \mbox{ g NaCl mol}^{-1}} = 0.034 \ mol

In the above example, when written out in fraction form, the units of grams form a multiplicative identity, which is equivalent to one (g/g=1), with the resulting amount of moles (the unit that was needed), is shown in the following equation,

\left(\frac{2.00 \mbox{ g NaCl}}{1}\right)\left(\frac{1 \mbox{ mol NaCl}}{58.44 \mbox{ g NaCl}}\right) = 0.034\ mol

Stoichiometry is also used to find the right amount of reactants to use in a chemical reaction. An example is shown below using the thermite reaction,

Fe_2O_3 + 2Al \rightarrow Al_2O_3 + 2Fe

So, to completely react with 85.0 grams of iron (III) oxide, 28.7 grams of aluminum are needed.

m Al = \left(\frac{85.0 \mbox{ g }Fe_2O_3}{1}\right)\left(\frac{1 mol\mbox{ mol }Fe_2 O_3}{159.7 \mbox{ g }Fe_2 O_3}\right)\left(\frac{2 \mbox{ mol }Al}{1 \mbox{ mol }Fe_2 O_3}\right)\left(\frac{27.0 \mbox{ g }Al}{1 \mbox{ mol }Al}\right) = 28.7 \mbox{ g }Al <