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3.9:

Experimentele Bepaling van Scheikundige Formules

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Chemistry
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JoVE Core Chemistry
Experimental Determination of Chemical Formula

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Chemische verbindingen worden meestal beschreven met behulp van een empirische of moleculaire formule. Deze formules geven informatie over het aantal verschillende atomen van betrokken elementen. Maar hoe komen deze formules tot stand?Experimentele analyse, zoals ontleding van verbindingen, wordt gebruikt om de relatieve massa van samenstellende elementen in de verbinding te schatten. Deze relatieve massa’s worden vervolgens gebruikt om het aantal mol van elk element te berekenen om de formule van een chemische verbinding te bepalen. Zo wordt experimenteel bepaald dat een monster van een verbinding 43, 64 gram fosfor en 56, 36 gram zuurstof bevat.Met behulp van molaire massa’s als conversiefactoren worden de relatieve massa’s uit de experimentele gegevens omgezet in 1, 41 mol voor fosfor en 3, 52 mol voor zuurstof. Deze molwaarden, wanneer ze worden toegekend als voorlopige subscripts aan de elementen, leveren een pseudo-formule van de verbinding op. Door de molwaarden te duiken met de kleinste molwaarde, worden de molverhoudingen van ongeveer 2, 5 mol zuurstof op 1 mol fosfor verkregen, die rechtstreeks verband houden met het werkelijke aandeel van de elementen in de verbinding.Als een van de quotiënten nog steeds een decimaal getal is, worden alle getallen vermenigvuldigd met het kleinste natuurlijke getal dat de kleinste gehele getalverhouding van subscripts geeft, waardoor de empirische formule van P2O5 wordt gegenereerd. De molecuulformule van verbindingen kan worden bepaald uit hun empirische formule en de molaire massa of het molecuulgewicht. De chemische verbinding met de empirische formule van P2O5 wordt bijvoorbeeld experimenteel gemeten om een molaire massa van 283, 89 gram per mol te hebben.Zijn molecuulformule is een veelvoud van zijn empirische formule, terwijl zijn molaire massa een veelvoud is van zijn empirische formule. De verhouding tussen molaire massa en empirische formulemassa levert het aantal formule-eenheden op. Door de empirische formule te vermenigvuldigen met het aantal formule-eenheden, wordt de molecuulformule verkregen.Daarom wordt de verbinding volgens de molecuulformule P4O10 geïdentificeerd als tetrafosfordecaoxide, of beter bekend onder de empirische formulenaam als fosforpentoxide.

3.9:

Experimentele Bepaling van Scheikundige Formules

The elemental makeup of a compound defines its chemical identity, and chemical formulas are the most concise way of representing this elemental makeup. When a compound’s formula is unknown, measuring the mass of its constituent elements is often the first step in determining the formula experimentally.

Determination of Empirical Formulas

The most common approach to determining a compound’s chemical formula is first to measure the masses of its constituent elements. However, chemical formulas represent the relative numbers and not masses of atoms in the substance. Therefore, any experimentally derived data involving mass must be used to obtain the corresponding numbers of atoms in the compound. This is accomplished using molar masses to convert the mass of each element to its number of moles. These molar amounts are used to compute whole-number ratios that can be used to derive the empirical formula of the substance.

Consider a sample of a compound determined to contain 1.71 grams of carbon and 0.287 grams of hydrogen. The corresponding numbers of atoms are 0.142 moles of carbon and 0.284 moles of hydrogen. Thus, this compound may be represented by the formula C0.142H0.284. Per convention, formulas contain whole-number subscripts, which can be achieved by dividing each subscript by the smallest subscript (0.142). The empirical formula for this compound is thus CH2. Subscripts of “1” are not written but rather assumed if no other number is present. This may or not be the compound’s molecular formula; however, additional information is needed to make that determination.

As a second example, a sample of a compound is determined to contain 5.31 grams of chlorine and 8.40 grams of oxygen. The same approach yields a tentative empirical formula of ClO3.5. In this case, dividing by the smallest subscript still leaves a decimal in the empirical formula. To convert this into a whole number, multiply each of the subscripts by two, retaining the same atom ratio and yielding Cl2O7 as the final empirical formula.

Deriving Empirical Formulas from Percent Composition

In instances where the percent composition of a compound is available, it is used to calculate the masses of elements present in the compound. Since the scale for percentages is 100, it is convenient to calculate the mass of elements present in a sample weighing 100 grams. The masses obtained are used to derive the empirical formula.

For example, suppose a gaseous compound contains 27.29% C and 72.71% O. The mass percentages, therefore, are expressed as fractions:

Eq1

Eq2

The mass of carbon, 27.29 g, corresponds to 2.272 moles of carbon, and the mass of oxygen, 72.71 g, corresponds to 4.544 moles of oxygen. The representative formula is, therefore, C2.272O4.544. Dividing each subscript with 2.272 provides the empirical formula: CO2

Derivation of Molecular Formulas

Determining the absolute numbers of atoms that compose a single molecule of a covalent compound requires knowledge of both its empirical formula and its molecular mass or molar mass. These quantities may be determined experimentally by various measurement techniques. Molecular mass, for example, is often derived from the mass spectrum of the compound.

Molecular formulas are derived by comparing the compound’s molar mass or molecular mass to its empirical formula mass. As the name suggests, an empirical formula mass is the sum of the average atomic masses of all the atoms represented in an empirical formula. If the known molar mass of a substance is divided by the empirical formula mass, it yields the number of empirical formula units per molecule (n). 

Eq3

The molecular formula is then obtained by multiplying each subscript in the empirical formula by n, as shown by the generic empirical formula AxBy:

Eq4

For example, the empirical formula of a covalent compound is determined to be CH2O, and its empirical formula mass is approximately 30 amu. If the compound’s molecular mass is determined to be 180 amu, this indicates that molecules of this compound contain six times the number of atoms represented in the empirical formula. 

Eq5

Molecules of this compound are then represented by a molecular formula with subscripts that are six times greater than those in the empirical formula: (CH2O)6 = C6H12O6.

This text is adapted from Openstax, Chemistry 2e, Section 3.2: Determining Empirical and Molecular Formulas.