Name: how bond polarity shapes molecular dipoles
Uploaded: 2020-09-03T13:38:28.000Z
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Chemistry

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Bond Polarity, Dipole Moment, and Percent Ionic Character

Rodzaj wiązań chemicznych, niezależnie od tego, czy są to niepolarne kowalencyjne, polarne kowalencyjne czy jonowe, zależy w dużej mierze od różnicy elektroujemności między atomami wiążącymi i długością ich wiązania. 

W bromie wiązanie między dwoma atomami bromu jest niepolarne, ponieważ różnica elektroujemności między nimi wynosi zero. 

W bromku wodoru wiązanie jest kowalencyjne polarnie, ponieważ bardziej elektroujemny brom przyciąga gęstość elektronów z mniej elektroujemnego atomu wodoru. 

W bromku potasu wiązanie jest wiązaniem jonowym ze względu na większą różnicę elektroujemności, co powoduje całkowite przeniesienie elektronów z potasu do bromku.

Gdy dwie równo, ale przeciwnie naładowane cząstki są oddalone od siebie o odległość, powstaje dipol. Ilościowa miara dipola nazywana jest momentem dipolowym reprezentowanym przez literę μ, która jest iloczynem wielkości ładunków, Q, wyrażonej w kulombach, i odległości między ładunkami, r, podanej w metrach. 

Moment dipolowy jest podawany w jednostkach debye'a, gdzie jeden debye równa się 3,34 × 10−30 C·m. 

To, czy związek jest polarny, kowalencyjny czy jonowy, można określić, obliczając procentowy charakter jonowy, który jest stosunkiem zmierzonego momentu dipolowego wiązania do momentu dipolowego przy założeniu całkowitego transferu elektronów. Wiązania o charakterze jonowym większym niż 50% są uważane za jonowe.

Weźmy pod uwagę fluorowodór, w którym atomy wodoru i fluoru są oddalone od siebie o 92 pm.

Jeśli wiązanie jest jonowe, moment dipolowy jest obliczany przy założeniu całkowitego przeniesienia elektronu. Ładunek elektronu mnoży się przez odległość między atomami, a uzyskaną wartość dzieli się przez jeden debye, co daje moment dipolowy 4,41 D.

Jednak eksperymentalnie zmierzony moment dipolowy fluorowodoru wynosi 1,18 D. Dlatego, aby określić procentowy charakter jonowy fluorowodoru, zmierzoną wartość 1,18 D dzieli się przez 4,41 D, co daje procent charakteru jonowego 41%. 

Ponieważ procentowy charakter jonowy fluorowodoru jest mniejszy niż 50%, jest to polarne wiązanie kowalencyjne.

polaryzacja wiązania, moment dipolowy i procentowy charakter jonowy

<h4>Polaryzacja wiązania</h4>

Wartość bezwzględna różnicy elektroujemności (ΔEN) dwóch związanych atomów stanowi przybliżoną miarę polarności, jaką powinny mieć wiązania w tym samym rodzaju wiązania. Gdy różnica jest bardzo mała lub równa zeru, wiązanie jest kowalencyjne i niepolarne. Gdy jest duże, wiązanie jest polarne, kowalencyjne lub jonowe. Bezwzględne wartości różnic elektroujemności między atomami w wiązaniach H – H, H – Cl i Na – Cl wynoszą odpowiednio 0 (niepolarne), 0,9 (polarne kowalencyjne) i 2,1 (jonowe).

Stopień współdzielenia elektronów między atomami waha się od całkowicie równego (czyste wiązanie kowalencyjne) do całkowitego braku (wiązanie jonowe).

Na przykład atomy H i F w HF mają różnicę elektroujemności wynoszącą 1,9, a atomy N i H w NH3 różnicę 0,9, mimo to oba te związki tworzą wiązania uważane za kowalencyjne polarne.

Podobnie atomy Na i Cl w NaCl mają różnicę elektroujemności wynoszącą 2,1, a atomy Mn i I w MnI2 mają różnicę 1,0, mimo że obie te substancje tworzą związki jonowe.

Najlepszym przewodnikiem po kowalencyjnym lub jonowym charakterze wiązania jest wzięcie pod uwagę rodzajów poszczególnych atomów i ich względnych pozycji w układzie okresowym.

Wiązania między dwoma niemetalami są na ogół kowalencyjne.

Wiązanie pomiędzy metalem i niemetalem jest często jonowe.

Niektóre związki zawierają zarówno wiązania kowalencyjne, jak i jonowe. Atomy jonów wieloatomowych, takich jak OH–, NO3− i NH4+, są połączone polarnymi wiązaniami kowalencyjnymi. Jednakże te jony wieloatomowe tworzą związki jonowe, łącząc się z jonami o przeciwnym ładunku. Na przykład azotan potasu, KNO3, zawiera kation K+ i wieloatomowy anion NO3-. Zatem wiązanie w azotanie potasu ma charakter jonowy, wynikający z przyciągania elektrostatycznego pomiędzy jonami K+ i NO3-, a także kowalencyjnego pomiędzy atomami azotu i tlenu w NO3-.

<h4>Polaryzacja molekularna i moment dipolowy</h4>

Jak omówiono wcześniej, polarne wiązania kowalencyjne łączą dwa atomy o różnej elektroujemności, pozostawiając jeden atom z częściowym ładunkiem dodatnim (δ+), a drugi z częściowym ładunkiem ujemnym (δ–). Elektrony są przyciągane w stronę atomu bardziej elektroujemnego . To rozdzielenie ładunku powoduje powstanie momentu dipolowego wiązania. Wielkość momentu dipolowego wiązania jest reprezentowana przez grecką literę mu (µ) i jest wyrażona wzorem pokazanym tutaj, gdzie Q jest wielkością ładunków cząstkowych (określoną przez różnicę elektroujemności), a r jest odległością między ładunkami :

<img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1062/1062_Equations-01.png" style="height: 40px;">

Ten moment wiązania można przedstawić jako wektor, czyli wielkość mającą zarówno kierunek, jak i długość. Wektory dipolowe pokazano jako strzałki wskazujące wzdłuż wiązania od atomu mniej elektroujemnego w stronę atomu bardziej elektroujemnego. Na mniej elektroujemnym końcu narysowany jest mały znak plus, aby wskazać częściowo dodatni koniec wiązania. Długość strzałki jest proporcjonalna do wielkości różnicy elektroujemności między dwoma atomami.

Cała cząsteczka może również mieć rozdzielony ładunek, w zależności od jej struktury molekularnej i polarności każdego z jej wiązań. Jeśli istnieje takie rozdzielenie ładunków, cząsteczkę nazywa się cząsteczką polarną (lub dipolem). W przeciwnym razie mówi się, że cząsteczka jest niepolarna. Moment dipolowy mierzy stopień separacji ładunków netto w cząsteczce jako całości. Moment dipolowy wyznacza się poprzez dodanie momentów wiązań w przestrzeni trójwymiarowej, biorąc pod uwagę strukturę molekularną.

W przypadku cząsteczek dwuatomowych istnieje tylko jedno wiązanie, więc moment dipolowy wiązania określa polarność molekularną. Homonuklearne cząsteczki dwuatomowe, takie jak Br2 i N2, nie różnią się elektroujemnością, więc ich moment dipolowy wynosi zero. W przypadku cząsteczek heterojądrowych, takich jak CO, występuje mały moment dipolowy. W przypadku HF występuje większy moment dipolowy, ponieważ istnieje większa różnica elektroujemności.

Jeżeli cząsteczka zawiera więcej niż jedno wiązanie, należy wziąć pod uwagę geometrię. Jeśli wiązania w cząsteczce są ułożone w taki sposób, że momenty ich wiązań znoszą się (suma wektorów wynosi zero), to cząsteczka jest niepolarna. Tak właśnie jest w przypadku CO2. Każde z wiązań jest polarne, ale cząsteczka jako całość jest niepolarna. Ze struktury Lewisa i stosując teorię VSEPR, cząsteczka CO2 jest określana jako liniowa z polarnymi wiązaniami C=O po przeciwnych stronach atomu węgla. Momenty wiązania znoszą się, ponieważ są skierowane w przeciwne strony. W przypadku cząsteczki wody struktura Lewisa ponownie pokazuje, że istnieją dwa wiązania z atomem centralnym, a różnica elektroujemności ponownie pokazuje, że każde z tych wiązań ma niezerowy moment wiązania. Jednak w tym przypadku struktura molekularna jest wygięta z powodu samotnych par na O, a momenty dwóch wiązań nie znoszą się. Z tego powodu woda ma moment dipolowy netto i jest cząsteczką polarną (dipol).

Polaryzacja wiązania, moment dipolowy i procentowy charakter jonowy

Polaryzacja wiązania

Bezwzględna wartość różnicy elektroujemności (ΔEN) dwóch związanych atomów stanowi przybliżoną miarę polarności, która ma być ...

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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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How Bond Polarity Shapes Molecular Dipoles | JoVE K12

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Analyze and interpret data to explain the mechanisms and properties of the two main types of intramolecular (ionic and covalent) bonds.

Analyze the structure of atoms and isotopes.

How Bond Polarity Shapes Molecular Dipoles

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