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Lecture 8sf

# Lecture 8sf - Laws of Definite Proportions and Multiple...

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Laws of Definite Proportions and Multiple Proportions One of Dalton’s postulates, written in 1808, was that atoms combine in simple whole-number ratios to form compounds. This ratios are represented in our familiar formulas today, such as H 2 O, CO 2 , NH 3 , C 12 H 22 O 11 . This was suggested to Dalton by the Law of Definite Proportions, which states that all compounds have a definite percent by mass of each element making it up. For example, 100 g of carbon dioxide always contains 27.3 g of carbon (27.3% carbon by mass), regardless of the source of the carbon dioxide. In the same way, 100 g of water always contains 88.9 g of oxygen (88.9% oxygen by mass). This suggested to Dalton that each compound had a definite ratio, in terms of simple whole numbers, of the atoms making it up—what we call a formula. Dalton also had information about compounds which contain the same elements but with different mass percentages. Say we had compound X (known today as carbon dioxide) which had 27.3% by mass carbon, meaning that in 100 g total, there are 27.3 g carbon and 72.7 g oxygen. Another compound Y also contains carbon and oxygen, and 100 g of that compound contains 42.9 g carbon and 57.1 g oxygen. This means that compound X has 3 . 27 7 . 72 or 2.66 times more O than C (by mass), and compound Y has 9 . 42 1 . 57 = 1.33 times more O than C (by mass). Summarizing these result in a table: Compound Mass C Mass O Mass C Mass O X 27.3 72.7 1 3 . 27 7 . 72 = 2.66 Y 42.9 57.1 1 9 . 42 1 . 57 = 1.33 This table shows that for a fixed 1 unit of mass of carbon, compound X has twice as much O than compound Y. This means that if compound X is CO 2 , compound Y must be CO. Or if compound Y is CO 2 , compound X is CO 4 . Dalton didn’t have a table of atomic masses in 1808—these were developed gradually over the 19 th century, but data such as this indicates that there are definite relationships between two related compounds, and further suggested to Dalton and others, the idea that compounds have definite formulas with atoms in simple whole number ratios. The data illustrated in the CO, CO 2 example is an illustration of the Law of Multiple Proportions. If two elements combine to form more than compound, the masses of one of these elements that combine with a fixed mass of the other, is in a simple whole number ratio.

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Molar Mass, Percent Composition, Empirical Formula Molar mass is the mass of a mole of substance. We have already discussed molar masses of atoms. The molar mass of a molecule is simply the sum of the molar mass of its atoms. To calculate the molar mass of methane, CH 4 : C has a molar mass of 12.0 g, and each H has a molar mass of 1.01 g. The molar mass of the molecule CH 4 is the sum of the mass of carbon + 4 hydrogens or 16.04 g. We use the unit mol g for molar mass, and usually 3 significant figures is sufficient. Molar masses ( mol g ) CH 4 16.0 CO 2 44.0 H 2 O 18.0 C 12 H 22 O 11 342 NH 3 17.0 NaCl 58.5 What is the mass of 7.50 x 10 20 molecules of methane, CH 4 ?
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