Name: finding empirical and molecular formulas
Uploaded: 2020-09-03T13:38:26.000Z
Duration: 2 min 37 s
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K12 Standards

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North Carolina K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.3.3

experimental-determination-of-chemical-formula

Experimental Determination of Chemical Formula

Chemische Verbindungen werden in der Regel mit einer empirischen oder molekularen Formel beschrieben. Diese Formeln geben Auskunft über die Menge der verschiedenen Atome der beteiligten Elemente. Aber wie werden diese Formeln aufgestellt?Die experimentelle Analyse wird, wie die Zersetzung von Verbindungen, verwendet, um die relativen Massen der beteiligten Elemente in der Verbindung abzuschätzen. Diese relativen Massen werden dann zur Berechnung der Molzahl jedes Elements verwendet, um die Formel einer chemischen Verbindung zu bestimmen. Zum Beispiel wird eine Probe einer Verbindung experimentell so bestimmt, dass sie 43, 64 Gramm Phosphor und 56, 36 Gramm Sauerstoff enthält.Unter Verwendung der Molmassen als Umrechnungsfaktoren werden die relativen Massen aus den experimentellen Daten in 1, 41 Mol für Phosphor und 3, 52 Mol für Sauerstoff umgerechnet. Diese Molwerte ergeben, wenn sie als vorläufige tiefgestellte Zahlen den Elementen zugeordnet werden, eine Pseudoformel der Verbindung. Wenn man die Molwerte durch den kleinsten Molwert teilt, erhält man die Molverhältnisse von etwa 2, 5 Mol Sauerstoff zu 1 Mol Phosphor, die sich direkt auf den tatsächlichen Anteil der Elemente in der Verbindung beziehen.Wenn einer der Quotienten noch ein Dezimalwert ist, werden alle Zahlen mit der kleinsten Zählzahl multipliziert, die das kleinste ganzzahlige Verhältnis der Subskriptionen ergibt, wodurch die empirische Formel P2O5 entsteht. Die Summenformel von Verbindungen kann aus ihrer Summenformel und entweder der Molmasse oder dem Molekulargewicht bestimmt werden. So wird beispielsweise die chemische Verbindung mit der Summenformel P2O5 experimentell so gemessen, dass sie eine Molmasse von 283, 89 Gramm pro Mol aufweist.Ihre Molekularformel ist ein ganzzahliges Vielfaches ihrer Summenformel, während ihre Molmasse ein ganzzahliges Vielfaches ihrer Masse nach der Summenformel ist. Das Verhältnis von Molmasse und Masse der Summenformel ergibt die Anzahl der Formeleinheiten. Die Multiplikation der Summenformel mit der Anzahl der Formeleinheiten ergibt die Summenformel.Ausgehend von der Molekülformel P4O10 wird die Verbindung daher als Tetraphosphordecaoxid oder, besser bekannt unter dem Namen der Summenformel, als Phosphorpentoxid bezeichnet.

experimentelle bestimmung der chemischen formel

Die elementare Zusammensetzung einer Verbindung definiert ihre chemische Identität, und chemische Formeln sind die prägnanteste Möglichkeit, diese elementare Zusammensetzung darzustellen. Wenn die Formel einer Verbindung unbekannt ist, ist die Messung der Masse ihrer Bestandteile oft der erste Schritt zur experimentellen Bestimmung der Formel. <h4>Bestimmung empirischer Formeln</h4> Der gebräuchlichste Ansatz zur Bestimmung der chemischen Formel einer Verbindung besteht darin, zunächst die Massen ihrer Bestandteile zu messen. Chemische Formeln stellen jedoch die relativen Zahlen und nicht die Massen der Atome in der Substanz dar. Daher müssen alle experimentell abgeleiteten Massendaten verwendet werden, um die entsprechende Anzahl von Atomen in der Verbindung zu erhalten. Dies wird erreicht, indem Molmassen verwendet werden, um die Masse jedes Elements in seine Molzahl umzurechnen. Aus diesen Molmengen werden ganzzahlige Verhältnisse berechnet, aus denen sich die Summenformel des Stoffes ableiten lässt. Betrachten Sie eine Probe einer Verbindung, bei der festgestellt wurde, dass sie 1,71 Gramm Kohlenstoff und 0,287 Gramm Wasserstoff enthält. Die entsprechenden Atomzahlen betragen 0,142 Mol Kohlenstoff und 0,284 Mol Wasserstoff. Somit kann diese Verbindung durch die Formel C0.142H0.284 dargestellt werden. Gemäß der Konvention enthalten Formeln ganzzahlige Indizes, was erreicht werden kann, indem jeder Index durch den kleinsten Index (0,142) dividiert wird. Die Summenformel für diese Verbindung lautet somit CH2. Indizes von „1“ werden nicht geschrieben, sondern angenommen, wenn keine andere Zahl vorhanden ist. Dies kann die Summenformel der Verbindung sein oder auch nicht; Für diese Entscheidung sind jedoch zusätzliche Informationen erforderlich. Als zweites Beispiel wird festgestellt, dass eine Probe einer Verbindung 5,31 Gramm Chlor und 8,40 Gramm Sauerstoff enthält. Der gleiche Ansatz ergibt eine vorläufige empirische Formel von ClO3.5. In diesem Fall verbleibt bei der Division durch den kleinsten Index immer noch eine Dezimalstelle in der empirischen Formel. Um dies in eine ganze Zahl umzuwandeln, multiplizieren Sie jeden Index mit zwei, behalten Sie dabei das gleiche Atomverhältnis bei und erhalten Sie Cl2O7 als endgültige empirische Formel. <h4>Ableitung empirischer Formeln aus der prozentualen Zusammensetzung</h4> In Fällen, in denen die prozentuale Zusammensetzung einer Verbindung verfügbar ist, wird sie zur Berechnung der Massen der in der Verbindung vorhandenen Elemente verwendet. Da die Prozentskala 100 beträgt, ist es praktisch, die Masse der in einer Probe mit einem Gewicht von 100 Gramm vorhandenen Elemente zu berechnen. Aus den ermittelten Massen wird die Summenformel abgeleitet. Angenommen, eine gasförmige Verbindung enthält 27,29 % C und 72,71 % O. Die Massenprozente werden daher als Bruchteile ausgedrückt: <img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1018/1018_Equations-01v2.png" style="height: 35px;"> <img alt="Eq2" src="https://www.jove.com/files/core/CDNSource/lessons/1018/1018_Equations-02v2.png" style="height: 35px;"> Die Kohlenstoffmasse von 27,29 g entspricht 2,272 Mol Kohlenstoff und die Sauerstoffmasse von 72,71 g entspricht 4,544 Mol Sauerstoff. Die repräsentative Formel lautet daher C2.272O4.544. Die Division jedes Indexes durch 2,272 ergibt die empirische Formel: CO2. <h4>Ableitung molekularer Formeln</h4> Die Bestimmung der absoluten Anzahl von Atomen, aus denen ein einzelnes Molekül einer kovalenten Verbindung besteht, erfordert die Kenntnis sowohl ihrer empirischen Formel als auch ihrer Molekülmasse oder Molmasse. Diese Größen können experimentell durch verschiedene Messtechniken bestimmt werden. Beispielsweise wird die Molekülmasse oft aus dem Massenspektrum der Verbindung abgeleitet. Molekularformeln werden durch den Vergleich der Molmasse oder Molekülmasse der Verbindung mit ihrer empirischen Formelmasse abgeleitet. Wie der Name schon sagt, ist eine empirische Formelmasse die Summe der durchschnittlichen Atommassen aller in einer empirischen Formel dargestellten Atome. Dividiert man die bekannte Molmasse eines Stoffes durch die Summenformelmasse, so erhält man die Anzahl der Summenformeleinheiten pro Molekül (n). <img alt="Eq3" src="https://www.jove.com/files/core/CDNSource/lessons/1018/1018_Equations-03v2.png" style="height: 40px;"> Die Summenformel erhält man dann, indem man jeden Index in der empirischen Formel mit n multipliziert, wie die allgemeine empirische Formel AxBy zeigt: <img alt="Eq4" src="https://www.jove.com/files/core/CDNSource/lessons/1018/1018_Equations-04v2.png" style="height: 17px;"> Die empirische Formel einer kovalenten Verbindung lautet beispielsweise CH2O und ihre empirische Formelmasse beträgt etwa 30 amu. Wenn die Molekülmasse der Verbindung mit 180 amu bestimmt wird, bedeutet dies, dass Moleküle dieser Verbindung sechsmal so viele Atome enthalten wie in der empirischen Formel dargestellt. <img alt="Eq5" src="https://www.jove.com/files/core/CDNSource/lessons/1018/1018_Equations-05v2.png" style="height: 35px;"> Moleküle dieser Verbindung werden dann durch eine Summenformel dargestellt, deren Indizes sechsmal größer sind als die in der empirischen Formel: (CH2O)6 = C6H12O6. Dieser Text wurde angepasst von <a href="https://openstax.org/books/chemistry-2e/pages/3-2-determining-empirical-and-molecular-formulas">Openstax, Chemistry 2e, Section 3.2: Determining Empirical and Molecular Formulas.</a>

Experimentelle Bestimmung der chemischen Formel

Die elementare Zusammensetzung einer Verbindung definiert ihre chemische Identität, und chemische Formeln sind die prägnanteste Art, diese elementare Zusammensetzung darzustellen. ...

Education

Lehre

JoVE Core

Chemie

Eins

Molecules, Compounds, and Chemical Equations

Moleküle, Verbindungen und chemische Gleichungen

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KEY TERMS AND CONCEPTS

KERNBEGRIFFE UND KONZEPTE

finding empirical and molecular formulas

Chemical formulas show the elemental makeup of a compound. When a formula is unknown, scientists often start by measuring the masses of the elements in the sample. Those mass data are then used to find the simplest whole-number ratio of atoms, which gives the empirical formula.

To get from mass to atom ratio, the masses of the elements are first converted to moles using molar masses. Chemical formulas describe numbers of atoms, not masses, so this step is necessary. In one example, 1.71 grams of carbon and 0.287 grams of hydrogen become 0.142 moles of carbon and 0.284 moles of hydrogen. Dividing by the smallest value gives the empirical formula CH2, with the 1 omitted by convention.

Sometimes the numbers do not become whole right away. A sample containing 5.31 grams of chlorine and 8.40 grams of oxygen gives a tentative ratio of ClO3.5. Multiplying both subscripts by two keeps the same atom ratio and gives the whole-number empirical formula Cl2O7.

Percent composition can also be used to find an empirical formula. Because percentages are based on 100, it is convenient to imagine a 100-gram sample. For a gas that is 27.29% carbon and 72.71% oxygen, those masses correspond to 2.272 moles of carbon and 4.544 moles of oxygen. Dividing by the smaller amount gives the empirical formula CO2.

A molecular formula gives the actual number of atoms in one molecule of a covalent compound. To find it, both the empirical formula and the molecular mass or molar mass are needed. The empirical formula mass is the sum of the average atomic masses in the empirical formula, and dividing the molar mass by that value gives the number of empirical units per molecule, n. If the empirical formula is CH2O and the molecular mass is 180 amu, then n equals 6, so the molecular formula is C6H12O6.

Finding Empirical and Molecular Formulas

molecules and compounds

Atomic Elements, Molecules, and Bonds

Matter is made of elements, and elements are built from atoms. An atom is the smallest unit of an element that still keeps that element’s properties. A silver coin, for example, is made of silver atoms. Each atom keeps the unique properties of silver. Elements are pure substances made of identical atoms. They cannot be broken down into simpler substances by chemical changes. Most elements are atomic elements, which use single atoms as their basic unit. Carbon and sodium are examples of atomic...

molecular models

Molecular Model Types in Chemistry

Molecular models help chemists and students visualize the structures and shapes of atoms and molecules. They represent the molecular architecture of chemical compounds in a way that is easier to study than formulas alone. Skeletal models are simple two-dimensional drawings of compounds. They show the molecular framework, or the bonds, without drawing every atom. In this model, many carbon and hydrogen atoms are not shown directly. Their positions are implied by the ends and junctions of the...

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Organic molecules are carbon-based compounds found in living things and in many everyday products. Most organic compounds have carbon as the main element, bonded to hydrogen and other carbon atoms. Some carbon-containing substances, such as carbonates, cyanides, and simple oxides like CO and CO2, are not classified as organic. Carbon can form up to four strong bonds, which helps explain why organic compounds can make chains and rings of many sizes and shapes. These molecules appear in...

formula mass and mole concepts of compounds

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Chemical Formula Mass and Percent Composition

Chemical formula mass and percent composition help describe the makeup of compounds. A chemical formula shows which atoms are present and how many of each atom appear in a substance. For covalent compounds, the formula shows the atoms in one molecule. That is why formula mass can also be called molecular mass. Chloroform, CHCl3, is one example. Its formula shows one carbon atom, one hydrogen atom, and three chlorine atoms. The molecular mass is the sum of their average atomic masses. Ionic...

experimental determination of chemical formula

Chemical formulas show the elemental makeup of a compound. When a formula is unknown, scientists often start by measuring the masses of the elements in the sample. Those mass data are then used to find the simplest whole-number ratio of atoms, which gives the empirical formula. To get from mass to atom ratio, the masses of the elements are first converted to moles using molar masses. Chemical formulas describe numbers of atoms, not masses, so this step is necessary. In one example, 1.71 grams...

chemical equations

Reading Chemical Equations and Balancing Them

Chemical equations show the identities and relative amounts of substances in a chemical reaction. The substances that react are called reactants, and their formulas are written on the left side of the equation. The substances formed are called products, and their formulas are written on the right side. A plus sign (+) separates formulas on the same side, and an arrow (→) shows the change from reactants to products. The numbers in a chemical equation are shown with coefficients. A coefficient...

key terms and concepts

Learn how to use mass, percent composition, and molar mass to determine empirical formulas and molecular formulas from experimental data.

Finding Empirical and Molecular Formulas | JoVE K12

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Ohio’s Learning Standards and Model Curriculum

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CH.3.a

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New York State Science Learning Standards

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Predicting Chemical Formulas

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PS.Chm.3.1

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PS.Chm.3.2

Analyze the structure of atoms and isotopes.

Finding Empirical and Molecular Formulas

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