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

Os tipos de ligações químicas, quer se trate de não polar covalente, polar covalente, ou iónico, são em grande parte determinadas pela diferença na eletronegatividade entre a colagem átomos e o seu comprimento de ligação. No bromo, a ligação entre os dois átomos de bromo é não polar já que a diferença na eletronegatividade entre eles é zero. No brometo de hidrogênio, a ligação é polar covalente uma vez que o bromo, mais eletronegativo atrai a densidade do elétron para longe do átomo de hidrogênio menos eletronegativo.No brometo de potássio, a ligação é uma ligação iónica devido à maior diferença na eletronegatividade, resultando em transferência electrónica completa de potássio para brometo. Quando iguais, mas com carga oposta as partículas são separadas por uma distância em que se forma um dipolo. A medida quantitativa do dipolo é chamado o momento dipolo representado pela carga mu, que é o produto da magnitude das cargas, Q, relatado em coulombs, e a distância entre as cargas, r, relatado em metros.O momento dipolo é dado em unidades Debye, onde, um debye equivale a 3, 34 x 10^30 statcoulomb-centímetros. Se um composto é polar covalente ou iónico pode ser determinado através do cálculo da percentagem de carácter iónico, que é a proporção de um dipolo medido de uma ligação momento para o momento dipolo assumindo a transferência de um elétron completo. As ações com mais de 50%de caráter iónico são consideradas iónicas.Considere o fluoreto de hidrogênio, onde o hidrogênio e o flúor átomos são separados por uma distância de 92 picometros. Se a ligação é iónica, o momento dipolo é calculado assumindo uma transferência eletrónica completa. A carga de um elétron é multiplicada com a distância entre os átomos e o valor obtido é dividido por um debye, que rende o momento dipolo de 4, 41 Debye.No entanto, a medida experimental momento dipolo de fluoreto de hidrogênio é 1, 18 Debye. Portanto, para determinar a percentagem carácter iónico do fluoreto de hidrogênio, o valor medido de 1, 18 Debye é dividido por 4, 41 Debye, o que dá uma percentagem de carácter iónico de 41%Desde a percentagem de carácter iónico do fluoreto de hidrogênio é inferior a 50%é uma ligação polar covalente.

polaridade de ligações, momento dipolar, e carácter iónico percentual

<h4>Polaridade de Ligações</h4>

O valor absoluto da diferença na eletronegatividade (ΔEN) de dois átomos ligados fornece uma medida aproximada da polaridade a ser esperada na ligação e, assim, o tipo de ligação. Quando a diferença é muito pequena ou zero, a ligação é covalente e não polar. Quando é grande, a ligação é covalente polar ou iónica. Os valores absolutos das diferenças da eletronegatividade entre os átomos nas ligações H–H, H–Cl, e Na–Cl são 0 (não polar), 0,9 (covalente polar), e 2,1 (iónica), respectivamente. 

O grau em que os eletrões são partilhados entre átomos varia entre ser completamente igual (ligação covalente pura) a não o ser de todo (ligação iónica). 

<ul>
<li>Por exemplo, os átomos de H e F em HF têm uma diferença de eletronegatividade de 1,9, e os átomos de N e H em NH3 uma diferença de 0,9, mas ambos formam ligações que são consideradas covalentes polares. </li>
<li>Da mesma forma, os átomos de Na e Cl em NaCl têm uma diferença de eletronegatividade de 2,1, e os átomos de Mn e I em MnI2 têm uma diferença de 1,0, mas ambas as substâncias formam compostos iónicos.</li>
</ul>

O melhor guia para o carácter covalente ou iónico de uma ligação é considerar os tipos de átomos envolvidos e as suas posições relativas na tabela periódica. 

<ul>
<li>As ligações entre dois não-metais são geralmente covalentes.</li>
<li>As ligações entre um metal e um não-metal é frequentemente iónica.</li>
</ul>

Alguns compostos contêm ligações covalentes e iónicas. Os átomos em iões poliatómicos, tais como OH–, NO3−,  e NH4+, são mantidos juntos por ligações polares covalentes. No entanto, estes iões poliatómicos formam compostos iónicos combinando-se com iões de carga oposta. Por exemplo, o nitrato de potássio, KNO3, contém o catião K+ e o anião poliatómico NO3−. Assim, as ligações no nitrato de potássio são iónicas, resultantes da atração eletrostática entre os iões K+ e NO3−, bem como covalentes entre os átomos de nitrogénio e oxigénio em NO3−.

<h4>Polaridade Molecular e Momento Dipolar</h4>

Como discutido anteriormente, as ligações covalentes polares ligam dois átomos com diferentes eletronegatividades, deixando um átomo com carga positiva parcial (δ+) e o outro átomo com carga negativa parcial (δ–), à medida que os eletrões são puxados para o átomo mais eletronegativo. Esta separação de carga dá origem a um momento dipolar de ligação. A magnitude de um momento dipolar de ligação é representada pela letra grega mu (µ) e é dada pela fórmula aqui apresentada, onde Q é a magnitude das cargas parciais (determinada pela diferença de eletronegatividade) e r é a distância entre as cargas:

<img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1062/1062_Equations-01.png" style="height: 40px;" />

Esse momento de ligação pode ser representado como um vector, uma quantidade que tem tanto direção como magnitude. Os vectores dipolares são mostrados como setas a apontar ao longo da ligação do átomo menos eletronegativo para o átomo mais eletronegativo. Um pequeno sinal de mais é desenhado na extremidade menos elétrica para indicar a extremidade parcialmente positiva da ligação. O comprimento da seta é proporcional à magnitude da diferença da eletronegatividade entre os dois átomos.

Uma molécula inteira pode também ter uma separação de carga, dependendo da sua estrutura molecular e da polaridade de cada ligação. Se tal separação de carga existir, diz-se que a molécula é uma molécula polar (ou dipolar); caso contrário, diz-se que a molécula é não-polar. O momento dipolar mede a extensão da separação de carga líquida na molécula como um todo. O momento dipolar é determinado pela adição dos momentos de ligação no espaço tridimensional, levando em conta a estrutura molecular.

Para moléculas diatómicas, há apenas uma ligação, portanto, o momento dipolar de ligação determina a polaridade molecular. Moléculas diatómicas homonucleares, como Br2 e N2, não têm diferença na eletronegatividade, portanto o seu momento dipolar é zero. Para moléculas heteronucleares como CO, há um pequeno momento dipolar. Para HF, há um maior momento dipolar porque há uma maior diferença na eletronegatividade.

Quando uma molécula contém mais do que uma ligação, a geometria tem de ser tida em conta. Se as ligações em uma molécula estão dispostas de tal forma que os seus momentos de ligação se cancelam (soma de vector igual a zero), então a molécula é não-polar. Esta é a situação do CO2. Cada ligação é polar, mas a molécula como um todo é não-polar. A partir da estrutura de Lewis, e usando a teoria VSEPR, a molécula de CO2 é determinada como linear com ligações C polares em lados opostos do átomo de carbono. Os momentos de ligação cancelam porque estão apontados em direções opostas. No caso da molécula de água, a estrutura de Lewis mostra novamente que existem duas ligações para um átomo central, e a diferença da eletronegatividade mostra novamente que cada ligação tem um momento de ligação diferente de zero. Neste caso, no entanto, a estrutura molecular está dobrada por causa dos pares solitários em O, e os dois momentos de ligação não cancelam. Portanto, a água tem um momento dipolar líquido e é uma molécula polar (dipolar).

Este texto é adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/7-2-covalent-bonding">Openstax, Chemistry 2e, Chapter 7.2: Covalent Bonding e  Openstax, Chemistry 2e, Chapter 7.6 Molecular Structure and Polarity.</a>

Polaridade de Ligações, Momento Dipolar, e Carácter Iónico Percentual

Polaridade de ligação

O valor absoluto da diferença de eletronegatividade (ΔEN) de dois átomos ligados fornece uma medida aproximada da polaridade a ser ...

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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.

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How Bond Polarity Shapes Molecular Dipoles

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