Name: lattice energy patterns in ionic solids
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 54 s
Channel: JoVE

K12 Standards

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North Carolina K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.3.1

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Trends in Lattice Energy: Ion Size and Charge

La energía reticular está asociada con la formación o separación de una celosía iónica. Sin embargo, cuando se forman cloruro de sodio u óxido de magnesio, su energía reticular difiere significativamente. Por qué cada compuesto iónico tiene una energía reticular diferente y de qué factores depende?Un compuesto iónico consiste en una disposición ordenada de un gran número de iones cargados atraídos entre sí por interacciones electrostáticas. Según la ley de Coulomb, la energía potencial de dos iones es inversamente proporcional a la distancia entre los iones, que a su vez depende del radio iónico. En la tabla periódica, el radio iónico de los metales alcalinos y alcalinotérreos aumenta en la columna.A medida que aumenta el tamaño del ion metálico, la distancia entre los iones, o la longitud del enlace, se vuelve más grande también. Por ejemplo, la longitud del enlace del bromuro de litio y el bromuro de potasio difiere entre 217 picómetros y 282 picómetros, respectivamente. Debido al aumento de la distancia internuclear, la atracción entre los iones disminuye y resulta mucho más fácil separar los iones.Por lo tanto, la energía reticular para separar el bromuro de litio sólido es mayor que con el bromuro de potasio, siendo de 807 kilojulios/mol frente a 682 kilojulios/mol, respectivamente. Además del radio iónico, la magnitud de la energía de la red también depende de las cargas de los iones. Según la ley de Coulomb, la energía potencial de los iones es directamente proporcional al producto de sus cargas.Veamos dos compuestos iónicos, fluoruro de sodio y óxido de calcio. La distancia iónica en ambos compuestos es similar, pero la energía reticular del óxido de calcio es casi cuatro veces mayor que la del fluoruro de sodio. Tanto el sodio como el fluoruro son iones monovalentes;el producto de su carga es 1.Sin embargo, los iones de calcio y óxido son divalentes y el producto de su carga es 4, que es cuatro veces mayor que el fluoruro de sodio. Es por eso que se necesita casi cuatro veces más energía para separar el óxido de calcio en iones gaseosos en comparación con el fluoruro de sodio. Por lo tanto, la magnitud de la energía de la red es directamente proporcional al producto de las cargas de los iones e inversamente proporcional a la distancia entre los iones.

tendencias en energía de celosía: tamaño y carga iónica

Un compuesto iónico es estable debido a la atracción electrostática entre sus iones positivos y negativos. La energía reticular de un compuesto es una medida de la fuerza de esta atracción. La energía de celosía (ΔHreticular) de un compuesto iónico se define como la energía requerida para separar un mol de un sólido en los iones gaseosos que lo componen. Para el ion sólido cloruro de sodio, la energía reticular es el cambio de entalpía del proceso:
 
 <img alt="Eq1" src="https://www.jovecore.com/CDNSource/lessons/1058/1058_Equations-01.png" style="height: 40px;" />
 
 <h4>Convenciones</h4>
 
 Aquí, la convención se usa donde el sólido iónico se separa en iones, significando que las energías de celosía serán endotérmicas (valores positivos). Otra manera es usar una convención equivalente pero opuesta, en donde la energía reticular es exotérmica (valores negativos) y es descrita como la energía liberada cuando los iones se combinan para formar una red. Por lo tanto, asegúrese de confirmar qué definición se usa cuando busque energías reticulares en otra referencia. En ambos casos, una magnitud mayor para la energía de celosía indica un compuesto iónico más estable. Para el cloruro sódico, ΔHreticular = 769 kJ. Por lo tanto, se requieren 769 kJ para separar un mol de NaCl sólido en los iones gaseosos Na+ y Cl–. Cuando un mol de iones gaseosos Na+ y Cl– forman NaCl sólido, se liberan 769 kJ de calor.
 
 <h4>Ley de Coulomb y Energía Reticular</h4>
 
 La energía de celosía ΔHreticular de un cristal iónico puede expresarse mediante la siguiente ecuación (derivada de la ley de Coulomb, que rige las fuerzas entre cargas eléctricas): 
 
 ΔHreticular = C(Z+)(Z−)/Ro 
 
 En la que C es una constante que depende del tipo de estructura cristalina; Z+ y Z– son las cargas sobre los iones, y Ro es la distancia interiónica (la suma de los radios de los iones positivos y negativos). Así, la energía de celosía de un cristal iónico aumenta rápidamente a medida que las cargas de los iones aumentan y los tamaños de los iones disminuyen. Cuando todos los demás parámetros se mantienen constantes, al doblar la carga tanto del catión como del anión se cuadruplica la energía reticular. 
 
 <h4>Ejemplos</h4>
 
 <ol>
 <li>La energía de celosía de CsF (Z+ y Z– = 1) es de 1023 kJ/mol, mientras que la de MgO (Z+ y Z– = 2) es de 3900 kJ/mol (Ro es casi igual, aproximadamente 200 pm para ambos compuestos).</li>
 <li>Las diferentes distancias interatómicas producen diferentes energías reticulares. Por ejemplo, compare la energía de celosía de MgF2 (2957 kJ/mol) con la de MgI2 (2327 kJ/mol) para observar el efecto sobre la energía reticular del tamaño iónico más pequeño del F– en comparación con I–.</li>
 <li>El gema preciosa rubí es el óxido de aluminio, Al2O3, que contiene trazas de Cr3+. El compuesto Al2Se3 se utiliza en la fabricación de algunos dispositivos semiconductores. En estos dos compuestos iónicos, las cargas Z+ y Z– son las mismas, por lo que la diferencia en la energía reticular depende de Ro. Dado que el ion O<sub2– es más pequeño que el ion Se2–, el ion Al2O3 tiene una distancia interiónica más corta que el Al2Se3 y, por lo tanto, tiene mayor energía reticular.</li>
 <li>Otro ejemplo es el óxido de zinc, ZnO, comparado con el NaCl. ZnO tiene una mayor energía de celosía porque los valores Z tanto del catión como del anión en ZnO son mayores, y la distancia interiónica del ZnO es menor que la del NaCl.</li>
 </ol>
 
 Este texto es adaptado de<a href="https://openstax.org/books/chemistry-2e/pages/7-5-strengths-of-ionic-and-covalent-bonds">Openstax, Química 2e, Sección 7.5: Fortalezas de los Enlaces Iónicos y Covalentes.</a>

Tendencias en Energía de Celosía: Tamaño y Carga Iónica

Un compuesto iónico es estable debido a la atracción electrostática entre sus iones positivos y negativos. La energía reticular de un compuesto es una medida ...

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Chemical Bonding: Basic Concepts

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KEY TERMS AND CONCEPTS

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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

Chemical bonding explains how atoms and ions join to form molecules and compounds. American chemist Gilbert N. Lewis pioneered chemical bonding theory with the Lewis model. This model helps explain why bonds form, why some are stronger than others, and why compounds have fixed formulas like H2O instead of H4O. Ionic bonding involves ions, which are atoms or molecules with an electrical charge. A cation is a positive ion formed when a neutral atom loses one or more electrons from its valence...

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Predicting Ionic Stability with Lattice Energy

Lattice energy helps explain why ionic compounds are stable. It measures the strength of attraction between positive and negative ions in a crystal. In one common definition, lattice energy is the energy required to separate one mole of an ionic solid into gaseous ions, so it has a positive value. Another definition uses the opposite sign and describes the energy released when gaseous ions form a lattice. When checking data, it is important to confirm which convention is being used. In either...

lewis symbols and the octet rule

Valence Electrons, Lewis Symbols, and Bonding

Chemical bonds form when atoms or ions interact in ways that lower the potential energy of a molecule. Gilbert N. Lewis created a simple model to show bond formation using valence electrons, the electrons in the outer shell. In this model, electrons are shared or transferred to help atoms reach a stable electron configuration. The Lewis model helps describe two common bond types. An ionic bond forms when electrons move from a metal to a nonmetal. A covalent bond forms when two nonmetals share...

electronegativity

Electronegativity and Polar Bonds

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ionic bonding and electron transfer

Ionic Compounds and Crystal Lattices

Ionic compounds form when positive and negative ions attract each other. A cation is a positive ion, and it forms when an atom loses one or more valence electrons. An anion is a negative ion, and it forms when an atom gains one or more valence electrons. Compounds made of these ions are called ionic compounds, or salts, and the force between them is an ionic bond. The properties of ionic solids help show how strong this attraction is. They usually form rigid, brittle crystals and have high...

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...

covalent bonding and lewis structures

Shared Electrons in Covalent Molecules

Covalent molecules form when atoms share electrons instead of transferring them. This bonding usually happens between nonmetal atoms with similar tendencies to attract electrons. Because the atoms pull on the electrons in a shared way, the bond is held together by mutual attraction. Covalent compounds often have physical properties that differ from ionic compounds. The attraction between neutral molecules is weaker than the attraction between charged ions. As a result, many covalent compounds...

bond polarity, dipole moment, and percent ionic character

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How Bond Polarity Shapes Molecular Dipoles

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lewis structures of molecular compounds and polyatomic ions

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Drawing Lewis Structures with Valence Electron Counts

Drawing Lewis structures starts with counting valence electrons and building the atom arrangement step by step. Valence electrons are the outer-shell electrons that help show how atoms bond. For ions, a positive charge means one electron is removed for each plus sign, and a negative charge means one electron is added for each minus sign. The next step is to sketch the skeleton structure. The least electronegative atom usually goes in the center, while hydrogen is almost never central. Fluorine...

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Resonance Forms in Lewis Structures

Resonance forms in Lewis structures help explain molecules and ions that cannot be shown correctly with just one drawing. In nitrite ion, NO2−, the Lewis structure can be drawn in two ways. The N-O and N=O bonds switch positions in the two drawings, but the atoms stay in the same arrangement. If nitrite really had one single bond and one double bond, the two bond lengths should be different. A double bond is shorter and stronger than a single bond between the same atoms. Experiments show that...

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Lewis Structures with Formal Charge

Lewis structures can sometimes show more than one reasonable arrangement for a molecule or polyatomic ion. Formal charge helps chemists choose the best Lewis structure when that happens. It is a useful way to compare possible structures in chemistry. Formal charge is the hypothetical charge an atom would have if bonding electrons were shared evenly between the atoms. It can also be found by starting with the valence electrons of the neutral atom, subtracting the nonbonding electrons, and then...

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Octet Rule Exceptions in Lewis Structures

Some covalent molecules do not follow the octet rule in their Lewis structures. Their central atoms can fall into three groups: odd-electron molecules, electron-deficient molecules, and hypervalent molecules. These cases show why a Lewis structure may not always give every atom eight valence electrons. Odd-electron molecules have an odd number of valence electrons, so one electron remains unpaired. Molecules with an odd number of electrons are called radicals. Nitric oxide, NO, is one example,...

bond energies and bond lengths

How Bond Energy Predicts Reaction Heat

Bond energy helps explain why molecules stay together and how much heat a reaction may release or absorb. A covalent bond is stronger when more energy is needed to separate the atoms. The energy needed to break one specific covalent bond in one mole of gaseous molecules is called bond energy, or bond dissociation energy. For a diatomic molecule, bond energy is the standard enthalpy change for an endothermic reaction that breaks that bond. The H-H bond is a common example. Its bond energy is...

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Electron Sea Model in Metals

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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Structure And Properties Of Matter

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Analyze the structure of atoms and isotopes.

Lattice Energy Patterns in Ionic Solids

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