Name: lattice energy patterns in ionic solids
Uploaded: 2020-09-03T13:38:28.000Z
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Trends in Lattice Energy: Ion Size and Charge

L'énergie du réseau est associée à la formation ou à la séparation d'un réseau ionique. Pourtant, lorsque le chlorure de sodium ou l'oxyde de magnésium se forment, leur énergie de réseau diffère considérablement. Pourquoi chaque composé ionique a-t-il une énergie de réseau différente et de quels facteurs dépend-il?Un composé ionique consiste en un agencement ordonné d'un grand nombre d'ions chargés attirés les uns aux autres par des interactions électrostatiques. Selon la loi de Coulomb, l'énergie potentielle de deux ions est inversement proportionnelle à la distance entre les ions, qui dépend du rayon ionique. Dans le tableau périodique, le rayon ionique des métaux alcalins et alcalino-terreux augmente dans la colonne.Comme la taille de l'ion métallique augmente, la distance entre les ions, ou la longueur de liaison, augmente également. Par exemple, la longueur de liaison du bromure de lithium et du bromure de potassium diffère entre 217 pm et 282 pm, respectivement. En raison de la distance internucléaire accrue, l'attraction entre les ions diminue et il devient beaucoup plus facile de séparer les ions.Ainsi, l'énergie du réseau pour séparer le bromure de lithium solide est plus grande que pour le bromure de potassium étant de 807 kJ mol contre 682 kJ mol, respectivement. En plus du rayon ionique, l'amplitude de l'énergie du réseau dépend également des charges ioniques. Selon la loi de Coulomb, l'énergie potentielle des ions est directement proportionnelle au produit de leurs charges.Considérons deux composés ioniques fluorure de sodium et oxyde de calcium. La distance ionique des deux composés est similaire, mais l'énergie du réseau de l'oxyde de calcium est presque quatre fois supérieure à celle du fluorure de sodium. Le sodium et le fluorure sont des ions monovalents;le produit de leur charge est de 1.Cependant, les ions calcium et oxyde sont divalents et le produit de leur charge est de 4, ce qui est quatre fois plus grand que le fluorure de sodium. Voila pourquoi il faut près de quatre fois plus d'énergie pour séparer l'oxyde de calcium en ions gazeux par rapport au fluorure de sodium. Ainsi, l'amplitude de l'énergie du réseau est directement proportionnelle au produit des charges ioniques et inversement proportionnelle à la distance entre les ions.

tendances de l'énergie du réseau : taille et charge des ions

Un composé ionique est stable en raison de l'attraction électrostatique entre ses ions positifs et négatifs. L'énergie réticulaire d'un composé est une mesure de la force de cette attraction. L'énergie réticulaire (ΔHréticulaire) d'un composé ionique est définie comme l'énergie nécessaire pour séparer une mole du solide en ions gazeux le composant. Pour le chlorure de sodium solide ionique, l'énergie réticulaire est la variation d'enthalpie du processus : 
 
 <img alt="Eq1" src="https://www.jovecore.com/CDNSource/lessons/1058/1058_Equations-01.png" style="height: 40px;" />
 
 <h4>Conventions </h4>
 
 Ici, la convention de signe est utilisée lorsque le solide ionique est séparé en ions, ce qui signifie que les énergies réticulaires seront endothermiques (valeurs positives). Une autre façon est d'utiliser une convention équivalente mais de signe opposée, dans laquelle l'énergie réticulaire est exothermique (valeurs négatives) et décrite comme l'énergie libérée lorsque les ions se combinent pour former un réseau. Par conséquent, assurez-vous de confirmer la définition utilisée lorsque vous regardez des énergies réticulaires d'une autre référence. Dans les deux cas, une plus grande valeur algébrique pour l'énergie réticulaire indique un composé ionique plus stable. Pour le chlorure de sodium, ΔHréticulaire = 769 kJ. Ainsi, il faut 769 kJ pour séparer une mole de NaCl solide en ions Na+ et Cl– gazeux. Lorsqu'une mole de chaque ions Na+et Cl– forme du NaCl solide, 769 kJ de chaleur sont libérés. 
 
 <h4>Loi de Coulomb et énergie réticulaire </h4>
 
 L'énergie réticulaire ΔHréticulaire d’un cristal ionique peut être exprimée par l’équation suivante (déduite de la loi de Coulomb, qui régit les forces entre des charges électriques) :  
 
 &DeltaHréticulaire =  C(Z+)(Z−)/Ro 
 
 où C est une constante qui dépend du type de structure cristalline ; Z+ et Z– sont les charges sur les ions, et Ro est la distance interionique (la somme des rayons des ions positifs et négatifs). Ainsi, l'énergie réticulaire d'un cristal ionique augmente rapidement à mesure que les charges des ions augmentent et que les dimensions des ions diminuent. Lorsque tous les autres paramètres sont maintenus constants, doubler la charge à la fois du cation et de l'anion quadruplera l'énergie réticulaire.  
 
 <h4>Exemples </h4>
 
 <ol>
 <li>L'énergie réticulaire de LiF (Z+ et Z– = 1) est de 1023 kJ/mol, alors que celle de MgO (Z+ et Z– = 2) est de 3900 kJ/mol (Ro est presque le même, soit environ 200 pm pour les deux composés). </li>
 <li>Différentes distances interatomiques produisent des énergies réticulaires différentes. Par exemple, comparez l'énergie réticulaire de MgF2 (2957 kJ/mol) à celle de MgI2 (2327 kJ/mol) pour observer l'effet sur l'énergie réticulaire de la taille ionique plus petite de F– par rapport à I–. </li>
 <li>La gemme précieuse rubis est de l'oxyde d'aluminium Al2O3, contenant des traces de Cr3+. Le composé Al2Se3 est utilisé dans la fabrication de certains semi-conducteurs. Dans ces deux composés ioniques, les charges Z+ et Z– sont les mêmes, de sorte que la différence d'énergie réticulaire dépend de Ro. Comme l'ion O2– est plus petit que l'ion Se2–, Al2Se3 a une distance interionique plus courte que Al2O3 et a donc une énergie réticulaire plus grande. </li>
 <li>Un autre exemple est l'oxyde de zinc, ZnO, comparé à NaCl. ZnO a une énergie réticulaire plus grande parce que les valeurs de Z du cation et de l'anion dans ZnO sont plus grandes et que la distance interionique de ZnO est plus petite que celle de NaCl. </li>
 </ol>
 
 Ce texte est adapté de <a href="https://openstax.org/books/chemistry-2e/pages/7-5-strengths-of-ionic-and-covalent-bonds">Openstax, Chimie 2e, Section 7.5 : Forces des liaisons ioniques et covalentes.</a>

Tendances de l'énergie du réseau : taille et charge des ions

Un composé ionique est stable en raison de l’attraction électrostatique entre ses ions positifs et négatifs. L’énergie du réseau d’un composé est une mesure ...

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lattice energy patterns in ionic solids

Lattice energy shows how strongly ions hold an ionic solid together. It is the energy required to separate one mole of a solid into its gaseous ions. For sodium chloride, this process requires 769 kJ per mole. The opposite convention may also be used, where lattice energy is the energy released when gaseous ions form the solid. In either case, a larger magnitude means a more stable ionic compound.

For ionic crystals, lattice energy follows a pattern based on charge and ion size. The relationship can be written as ΔH lattice = C(Z+)(Z−)/Ro. In this equation, C depends on the crystal structure, Z+ and Z− are the ion charges, and Ro is the distance between the ions. Because of this, lattice energy rises quickly when ion charges increase and when ion sizes decrease. If both ion charges double, the lattice energy becomes four times larger.

Examples show these trends clearly. LiF, with ion charges of 1 and 1, has a lattice energy of 1023 kJ/mol, while MgO, with charges of 2 and 2, has a lattice energy of 3900 kJ/mol. Their interionic distances are nearly the same, at about 200 pm. Ion size also matters. MgF2 has a lattice energy of 2957 kJ/mol, while MgI2 has a lower value of 2327 kJ/mol because F− is smaller than I−.

The same idea explains differences in other compounds. Ruby is aluminum oxide, Al2O3, and it contains traces of Cr3+. Al2Se3 is used in some semiconductor devices. These compounds have the same ion charges, so the difference in lattice energy depends on Ro. O2− is smaller than Se2−, so Al2O3 has the shorter interionic distance and the larger lattice energy. ZnO also has a larger lattice energy than NaCl because the ion charges are greater and the interionic distance is smaller.

Lattice Energy Patterns in Ionic Solids

types of chemical bonds

Chemical Bonding and Bond Properties

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Ionic Compounds and Crystal Lattices

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trends in lattice energy: ion size and charge

Lattice energy shows how strongly ions hold an ionic solid together. It is the energy required to separate one mole of a solid into its gaseous ions. For sodium chloride, this process requires 769 kJ per mole. The opposite convention may also be used, where lattice energy is the energy released when gaseous ions form the solid. In either case, a larger magnitude means a more stable ionic compound. For ionic crystals, lattice energy follows a pattern based on charge and ion size. The...

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lewis structures of molecular compounds and polyatomic ions

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Drawing Lewis Structures with Valence Electron Counts

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Resonance Forms in Lewis Structures

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Lewis Structures with Formal Charge

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Electron sea model explains how metallic bonding holds metals together. In this model, most metal atoms do not have enough valence electrons to form ionic or covalent bonds in the usual way. Their valence electrons are held loosely because metals have low electronegativity, and their ionization energy is low, so electrons are removed easily. When a metal atom loses these outer electrons, it becomes a positively charged metal ion. The freed electrons form a delocalized electron cloud around...

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Learn how ion charge and ion size affect lattice energy in ionic solids, with examples like NaCl, LiF, MgO, MgF2, MgI2, Al2O3, and ZnO.

Lattice Energy Patterns in Ionic Solids | JoVE K12

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Lattice Energy Patterns in Ionic Solids

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