Name: how bond polarity shapes molecular dipoles
Uploaded: 2020-09-03T13:38:28.000Z
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Bond Polarity, Dipole Moment, and Percent Ionic Character

Die Art der chemischen Bindungen, ob es sich um unpolare kovalente, polare kovalente oder ionische Bindungen handelt, wird weitgehend durch den Unterschied in der Elektronegativität zwischen den bindenden Atomen und ihrer Bindungslänge bestimmt. Bei Brom ist die Bindung zwischen den beiden Bromatomen unpolar, da der Unterschied in der Elektronegativität zwischen ihnen gleich Null ist. In Bromwasserstoff ist die Bindung polar kovalent, da das elektronegativere Brom die Elektronendichte von dem weniger elektronegativen Wasserstoffatom wegzieht.In Kaliumbromid ist die Bindung ionisch, da die Elektronegativitätsdifferenz größer ist, was zu einem vollständigen Elektronentransfer von Kalium zu Bromid führt. Wenn zwei gleich große, aber entgegengesetzt geladene Teilchen durch einen Abstand voneinander getrennt werden, entsteht ein Dipol. Das quantitative Maß für den Dipol wird als Dipolmoment bezeichnet, das durch den Buchstaben μ repräsentiert wird, der das Produkt aus der Größe der Ladungen Q, angegeben in Coulomb, und dem Abstand zwischen den Ladungen r, angegeben in Metern, ist.Das Dipolmoment wird in Debye-Einheiten angegeben, wobei ein Debye 3, 34 x 10^30 Coulomb*meter entspricht. Ob eine Verbindung polar kovalent oder ionisch ist, kann durch Berechnung des prozentualen Ionencharakters bestimmt werden, der das Verhältnis des gemessenen Dipolmoments einer Bindung zu dem Dipolmoment mit der Annahme eines vollständigen Elektronentransfers ist. Bindungen mit mehr als 50%ionischem Charakter gelten als ionisch.Betrachten Sie Fluorwasserstoff, bei dem die Wasserstoff-und Fluoratome durch einen Abstand von 92 picometern getrennt sind. Wenn die Bindung ionisch ist, wird das Dipolmoment unter der Annahme eines vollständigen Elektronentransfers berechnet. Die Ladung eines Elektrons wird mit dem Abstand zwischen den Atomen multipliziert, und der erhaltene Wert wird durch einen Debye geteilt, was ein Dipolmoment von 4, 41 Debye ergibt.Das experimentell gemessene Dipolmoment von Fluorwasserstoff beträgt jedoch 1, 18 Debye. Um den prozentualen ionischen Charakter von Fluorwasserstoff zu bestimmen, wird daher der gemessene Wert von 1, 18 Debye durch 4, 41 Debye geteilt, was einen prozentualen ionischen Charakter von 41%ergibt. Da der prozentuale ionische Charakter von Fluorwasserstoff weniger als 50%beträgt, handelt es sich um eine polare kovalente Bindung.

bindungspolarität, dipolmoment und prozentualer ionencharakter

<h4>Bindungspolarität</h4> Der Absolutwert der Elektronegativitätsdifferenz (ΔEN) zweier gebundener Atome liefert ein grobes Maß für die zu erwartende Polarität der Bindung und damit für den Bindungstyp. Wenn der Unterschied sehr klein oder Null ist, ist die Bindung kovalent und unpolar. Wenn es groß ist, ist die Bindung polar kovalent oder ionisch. Die Absolutwerte der Elektronegativitätsunterschiede zwischen den Atomen in den Bindungen H–H, H–Cl und Na–Cl betragen 0 (unpolar), 0,9 (polar kovalent) bzw. 2,1 (ionisch). Der Grad, in dem Elektronen zwischen Atomen geteilt werden, variiert von völlig gleich (reine kovalente Bindung) bis gar nicht (ionische Bindung). <ul> <li>Beispielsweise haben die H- und F-Atome in HF einen Elektronegativitätsunterschied von 1,9 und die N- und H-Atome in NH3 einen Unterschied von 0,9, dennoch bilden beide Verbindungen Bindungen, die als polar kovalent gelten.</li> <li>Ebenso haben die Na- und Cl-Atome in NaCl einen Elektronegativitätsunterschied von 2,1 und die Mn- und I-Atome in MnI2 einen Unterschied von 1,0, beide Substanzen bilden jedoch ionische Verbindungen.</li> </ul> Der wichtigste Hinweis auf den kovalenten oder ionischen Charakter einer Bindung lässt sich in den beteiligten Atomtypen und ihren relativen Positionen im Periodensystem finden. <ul> <li>Bindungen zwischen zwei Nichtmetallen sind im Allgemeinen kovalent.</li> <li>Die Bindung zwischen einem Metall und einem Nichtmetall erfolgt oft ionisch.</li> </ul> Einige Verbindungen enthalten sowohl kovalente als auch ionische Bindungen. Die Atome in mehratomigen Ionen wie OH–, NO3−,   und NH4+ werden durch polare kovalente Bindungen zusammengehalten. Diese mehratomigen Ionen bilden jedoch ionische Verbindungen, indem sie sich mit Ionen der entgegengesetzten Ladung verbinden. Kaliumnitrat, KNO3, enthält beispielsweise das K+ Kation und das mehratomige NO3− Anion. Somit ist die Bindung in Kaliumnitrat ionisch und resultiert aus der elektrostatischen Anziehung zwischen den Ionen K+ und NO3− sowie einer kovalenten Bindung zwischen den Stickstoff- und Sauerstoffatomen in NO3−. <h4>Molekulare Polarität und Dipolmoment</h4> Wie bereits erwähnt, verbinden polare kovalente Bindungen zwei Atome mit unterschiedlicher Elektronegativität, sodass ein Atom eine teilweise positive Ladung (δ+) und das andere Atom eine teilweise negative Ladung (δ–) aufweist, da die Elektronen vom elektronegativeren Atom angezogen werden. Diese Ladungstrennung führt zu einem Bindungsdipolmoment. Die Größe eines Bindungsdipolmoments wird durch den griechischen Buchstaben mu (µ) dargestellt und durch die hier gezeigte Formel beschrieben, wobei Q die Größe der Teilladungen (bestimmt durch die Elektronegativitätsdifferenz) und r der Abstand zwischen den Ladungen ist: <img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1062/1062_Equations-01.png" style="height: 40px;"> Dieses Bindungsmoment kann als Vektor dargestellt werden, eine Größe, die sowohl Richtung als auch Größe hat. Dipolvektoren werden als Pfeile dargestellt, die entlang der Bindung vom weniger elektronegativen Atom zum elektronegativeren Atom zeigen. Am weniger elektronegativen Ende wird ein kleines Pluszeichen gezeichnet, um das teilweise positive Ende der Bindung anzuzeigen. Die Länge des Pfeils ist proportional zur Größe der Elektronegativitätsdifferenz zwischen den beiden Atomen. Abhängig von seiner Molekülstruktur und der Polarität jeder seiner Bindungen kann es auch bei einem ganzen Molekül zu einer Ladungstrennung kommen. Liegt eine solche Ladungstrennung vor, spricht man von einem polaren Molekül (oder Dipol); andernfalls spricht man von unpolaren Molekülen. Das Dipolmoment misst das Ausmaß der Nettoladungstrennung im gesamten Molekül. Das Dipolmoment wird durch Addition der Bindungsmomente im dreidimensionalen Raum unter Berücksichtigung der Molekülstruktur bestimmt. Bei zweiatomigen Molekülen gibt es nur eine Bindung, daher bestimmt das Bindungsdipolmoment die Molekülpolarität. Homonukleare zweiatomige Moleküle wie Br2 und N2 haben keinen Unterschied in der Elektronegativität, daher ist ihr Dipolmoment gleich Null. Bei heteronuklearen Molekülen wie CO gibt es ein kleines Dipolmoment. Bei HF gibt es ein größeres Dipolmoment, da der Unterschied in der Elektronegativität größer ist. Wenn ein Molekül mehr als eine Bindung enthält, muss seine Geometrie berücksichtigt werden. Wenn die Bindungen in einem Molekül so angeordnet sind, dass sich ihre Bindungsmomente aufheben (Vektorsumme gleich Null), dann ist das Molekül unpolar. Dies ist die Situation bei CO2. Jede der Bindungen ist polar, aber das Molekül als Ganzes ist unpolar. Aus der Lewis-Struktur und unter Verwendung der VSEPR-Theorie kann hergeleitet werden, dass das CO2-Molekül linear und die polaren C=O-Bindungen auf den gegenüberliegenden Seiten des Kohlenstoffatoms liegen. Die Bindungsmomente heben sich auf, da sie in entgegengesetzte Richtungen zeigen. Im Fall des Wassermoleküls zeigt die Lewis-Struktur wiederum, dass es zwei Bindungen zu einem Zentralatom gibt, und die Elektronegativitätsdifferenz zeigt wiederum, dass jede dieser Bindungen ein Bindungsmoment ungleich Null aufweist. In diesem Fall ist die Molekülstruktur jedoch aufgrund der freien Elektronenpaare am O gebogen und die beiden Bindungsmomente heben sich nicht auf. Daher hat Wasser ein Nettodipolmoment und ist ein polares Molekül (Dipol). Dieser Text wurde adaptiert von <a href="https://openstax.org/books/chemistry-2e/pages/7-2-covalent-bonding">Openstax, Chemistry 2e, Chapter 7.2: Covalent Bonding and  Openstax, Chemistry 2e, Chapter 7.6 Molecular Structure and Polarity.</a>

Bindungspolarität, Dipolmoment und prozentualer Ionencharakter

Polarität der Bindung

Der Absolutwert der Differenz in der Elektronegativität (ΔEN) zweier gebundener Atome liefert ein grobes Maß für die Polarität, die ...

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how bond polarity shapes molecular dipoles

Bond polarity depends on electronegativity, the ability of an atom to attract shared electrons. The absolute value of the electronegativity difference, or ΔEN, gives a rough guide to bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is larger, the bond may be polar covalent or ionic.

Examples show how this works in real compounds. H–H has a ΔEN of 0, so it is nonpolar. H–Cl has a ΔEN of 0.9, so it is polar covalent. Na–Cl has a ΔEN of 2.1, so it is ionic. Even so, ΔEN is only a guide. HF has a ΔEN of 1.9 and NH3 has a ΔEN of 0.9, yet both have polar covalent bonds. NaCl has a ΔEN of 2.1 and MnI2 has a ΔEN of 1.0, yet both are ionic compounds.

The type of atoms in the periodic table is often the best clue. Bonds between two nonmetals are generally covalent. Bonds between a metal and a nonmetal are often ionic. Some compounds include both kinds of bonding. Polyatomic ions such as OH−, NO3−, and NH4+ have polar covalent bonds inside the ion. They form ionic compounds when they combine with ions of the opposite charge. Potassium nitrate, KNO3, is one example. It contains K+ and NO3−, so the attraction between those ions is ionic, while the N–O bonds inside nitrate are covalent.

Polar covalent bonds create partial charges, written as δ+ and δ−. The more electronegative atom pulls the electrons closer, so one end of the bond becomes partially negative and the other becomes partially positive. This charge separation creates a bond dipole moment, symbolized by μ. A dipole is shown as a vector, or arrow, that points from the less electronegative atom toward the more electronegative atom. A small plus sign marks the positive end. A longer arrow means a larger electronegativity difference.

A whole molecule can also have a net dipole moment. That depends on both bond polarity and molecular shape. The dipole moment of the molecule is found by adding the bond moments in three-dimensional space. For diatomic molecules, there is only one bond, so the bond dipole sets the molecular polarity. Homonuclear diatomic molecules such as Br2 and N2 have no electronegativity difference, so their dipole moment is zero. Heteronuclear molecules such as CO have a small dipole moment, and HF has a larger one because its electronegativity difference is larger.

Molecular geometry can make polar bonds cancel. CO2 is linear, so its two polar C=O bond moments point in opposite directions and cancel. The molecule is nonpolar even though its bonds are polar. Water is different. Its Lewis structure shows two bonds to oxygen, and each bond has a nonzero dipole moment. Because the molecule is bent, the bond moments do not cancel. Water therefore has a net dipole moment and is a polar molecule.

How Bond Polarity Shapes Molecular Dipoles

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bond polarity, dipole moment, and percent ionic character

Bond polarity depends on electronegativity, the ability of an atom to attract shared electrons. The absolute value of the electronegativity difference, or ΔEN, gives a rough guide to bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is larger, the bond may be polar covalent or ionic. Examples show how this works in real compounds. H–H has a ΔEN of 0, so it is nonpolar. H–Cl has a ΔEN of 0.9, so it is polar covalent. Na–Cl has a ΔEN of 2.1, so it...

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How Bond Polarity Shapes Molecular Dipoles

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