Name: predicting ionic stability with lattice energy
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 44 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

The Born-Haber Cycle

A formação de ligações iónicas requer uma transferência de elétrons de um átomo metálico para um não-metálico um processo que é frequentemente endotérmico. No entanto, quando o sódio elementar e o cloro reagem para formar um cristal sólido de cloreto de sódio, é um processo altamente exotérmico. Mas de onde vem a energia, de acordo com a lei de Coulomb, os catíons e os âníons são atraídos uns pelos outros por fortes forças eletrostáticas numa matriz sólida, ou grelha.A estrutura da malha resultante é estabilizada através da diminuição da energia potencial, que é libertada como calor, uma reação exotérmica. A energia total associada à formação ou à avaria de uma malha cristalina nos seus constituintes gasosos chama-se energia da malha. Num composto iónico sólido, um grande número de partículas carregadas interagem umas com as outras tornando difícil determinar o valor exato da malha de energia a título experimental.No entanto, pode ser calculada utilizando a lei de Hess numa série hipotética de passos chamada de ciclo de Born-Haber, que representa a formação de um composto iónico dos seus elementos constituintes. Por exemplo, o ciclo de Born-Haber para a formação de cloreto de sódio considera duas vias alternativas, uma direta e outra indireta. A rota direta representa a entalpia padrão da formação de NaCl a partir do sódio elementar e cloro.A via indireta tem cinco etapas. No primeiro passo, o sódio em estado sólido é convertido para a sua forma gasosa. A seguir, as moléculas de cloro diatómicas dissociam-se em átomos de cloro gasoso.O terceiro e quarto passos são responsáveis pela transferência do elétron para formar íons. Um elétron é removido do sódio gasoso para formar um cátion de sódio. O elétron é então absorvido pelo cloro gasoso para formar um ânion de cloreto.Na etapa final, a atração eletrostática entre os íons gasosos leva à formação da estrutura da malha. A lei de Hess afirma que a alteração na entalpia geral de um processo por etapas é a soma das alterações da entalpia de cada passo. Significa que o valor da entalpia da rota direta é igual à soma das entalpias dos cinco passos.Resolvendo a equação da energia da malha é determinado um valor grande e negativo, o que significa uma reação exotérmica.

o ciclo de born-haber

<h4>Energia Reticular </h4>

Um composto iónico é estável devido à atração eletrostática entre os seus iões positivos e negativos. A energia reticular de um composto é uma medida da força desta atração. A energia reticular (ΔHreticular) de um composto iónico é definida como a energia necessária para separar um mole do sólido nos seus iões gasosos constituintes. 

Aqui, é usada a convenção em que o sólido iónico é separado em iões, o que significa que as energias reticulares serão endotérmicas (valores positivos). Outra forma é usar uma convenção equivalente, mas oposta, na qual a energia reticular é exotérmica (valores negativos) e descrita como a energia libertada quando iões se combinam para formar uma rede. Assim, certifique-se de confirmar que definição é usada ao procurar energias reticulares em outra referência. 

Em ambos os casos, uma magnitude maior para a energia reticular indica um composto iónico mais estável. Para o cloreto de sódio, ΔHreticular = 769 kJ. Assim, requer 769 kJ para separar um mole de NaCl sólido nos iões gasosos Na+ e Cl–. Quando um mole de cada um dos iões gasosos Na+ e Cl– formam NaCl sólido, são libertados 769 kJ de calor.

<h4>Determinação da Energia Reticular de um Composto Iónico</h4>

Não é possível medir energias reticulares diretamente. No entanto, a energia reticular pode ser calculada utilizando um ciclo termoquímico. O ciclo Born-Haber é uma aplicação da lei de Hess que divide a formação de um sólido iónico em uma série de passos individuais:

<table style="margin-left: auto; margin-right: auto;">
<tbody style="border: 1px solid black;">
<tr>
<td style="color: black; border: 1px solid black; text-align: center;">Entalpia da sublimação de Cs (s)</td>
<td colspan="2" style="color: black; border: 1px solid black; text-align: center;">Cs (s) → Cs (g)</td>
<td style="color: black; border: 1px solid black; text-align: center;">ΔH = ΔHs° = 76,5 kJ/mol</td>
</tr>
<tr>
<td style="color: black; border: 1px solid black; text-align: center;">Metade da energia de ligação de F2</td>
<td colspan="2" style="color: black; border: 1px solid black; text-align: center;">½ F2 (g) → F (g)</td>
<td style="color: black; border: 1px solid black; text-align: center;">ΔH = ½ D = 79,4 kJ/mol</td>
</tr>
<tr>
<td style="color: black; border: 1px solid black; text-align: center;">Energia de ionização de Cs (g)</td>
<td colspan="2" style="color: black; border: 1px solid black; text-align: center;">Cs (g)  → Cs+ (g) + e−</td>
<td style="color: black; border: 1px solid black; text-align: center;">ΔH = IE = 375,7 kJ/mol</td>
</tr>
<tr>
<td style="color: black; border: 1px solid black; text-align: center;">Afinidade eletrónica de F</td>
<td colspan="2" style="color: black; border: 1px solid black; text-align: center;">F (g) + e− → F− (g)</td>
<td style="color: black; border: 1px solid black; text-align: center;">ΔH = EA = −328,2 kJ/mol</td>
</tr>
<tr>
<td style="color: black; border: 1px solid black; text-align: center;">Negativo da energia reticular de CsF (s)</td>
<td colspan="2" style="color: black; border: 1px solid black; text-align: center;">Cs+ (g) + F− (g) → CsF (s)</td>
<td style="color: black; border: 1px solid black; text-align: center;">ΔH = −ΔHreticular = ?</td>
</tr>
<tr>
<td style="color: black; border: 1px solid black; text-align: center;">Entalpia da formação de CsF (s), adicionar passos 1–5</td>
<td colspan="2" style="color: black; border: 1px solid black; text-align: center;">ΔH = ΔHf° = ΔHs°+ ½ D + IE + (EA) + (−ΔHreticular) 
Cs (s) + ½ F2 (g) → CsF (s)</td>
<td style="color: black; border: 1px solid black; text-align: center;">ΔH = −553,5 kJ/mol</td>
</tr>
</tbody>
</table>

<ol>
<li>Considere os elementos nos seus estados mais comuns, Cs (s) e F2 (g). </li>
<li>A ΔHs° representa a conversão do césio sólido em gás (sublimação), e depois a energia de ionização converte os átomos gasosos de césio em catiões. </li>
<li>No passo seguinte, a energia necessária para quebrar a ligação F–F para produzir átomos de flúor precisa ser contabilizada. </li>
<li>Converter um mole de átomos de flúor em iões fluoreto é um processo exotérmico, pelo que este passo produz energia (a afinidade eletrónica) </li>
<li>Agora, um mole de catiões de Cs e um mole de aniões de F está presente. Estes iões combinam-se para produzir fluoreto de césio sólido. A alteração da entalpia neste passo é o negativo da energia reticular, por isso é também uma quantidade exotérmica. </li>
<li>A energia total envolvida nesta conversão é igual à entalpia da formação experimentalmente determinada, ΔHf°, do composto a partir dos seus elementos. Neste caso, a alteração geral é exotérmica.</li>
</ol>

As energias reticulares calculadas para compostos iónicos são tipicamente muito mais altas do que as energias de dissociação de ligação medidas para ligações covalentes. Enquanto que as energias reticulares caem normalmente na faixa de 600 – 4000 kJ/mol (algumas ainda mais altas), as energias de dissociação de ligações covalentes estão tipicamente entre 150 – 400 kJ/mol para ligações simples. Tenha em mente, no entanto, que estes não são valores diretamente comparáveis. Para compostos iónicos, as energias reticulares estão associadas a muitas interações, à medida que catiões e aniões se juntam em um retículo estendido. Para ligações covalentes, a energia de dissociação de ligações está relacionada com a interação de apenas dois átomos.

<h4>Energia Reticular como uma Função do Raio e Carga Iónicos </h4>

A energia reticular de um cristal iónico aumenta rapidamente à medida que as cargas dos iões aumentam e os tamanhos dos iões diminuem. Quando todos os outros parâmetros são mantidos constantes, a duplicação da carga tanto do catião como do anião quadruplica a energia reticular. Por exemplo, a energia reticular de LiF (Z+ e Z– = 1) é de 1023 kJ/mol, enquanto que a de MgO (Z+ e Z– = 2) é de 3900 kJ/mol (Ro = a distância interiónica definida como a soma dos raios dos iões positivos e negativos, é quase a mesma — aproximadamente 200 pm para ambos os compostos).

Diferentes distâncias interatómicas produzem diferentes energias reticulares. Por exemplo, compare a energia reticular de MgF2 (2957 kJ/mol) com a de MgI2 (2327 kJ/mol), que demonstra o efeito na energia reticular do tamanho iónico menor de F– em comparação com I–.

<h4>Outras Aplicações do Ciclo Born-Haber</h4>

O ciclo Born-Haber também pode ser usado para calcular qualquer uma das outras quantidades na equação para a energia reticular, desde que o restante seja conhecido. Por exemplo, se a entalpia da sublimação relevante  ΔHs°,  a energia de ionização (IE), a entalpia de dissociação de ligações (D), a energia reticular ΔHreticular, e a entalpia de formação padrão  ΔHf°  forem conhecidas, o ciclo Born-Haber pode ser usado para determinar a afinidade eletrónica de um átomo.

Este texto é adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/7-5-strengths-of-ionic-and-covalent-bonds">Openstax, Chemistry 2e, Section 7.5: Strengths of Ionic and Covalent Bonds.</a>

O Ciclo de Born-Haber

Energia da Estrutura

Um composto iônico é estável devido à atração eletrostática entre seus íons positivos e negativos. A energia da rede de um ...

Education

JoVE Core

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Chemical Bonding: Basic Concepts

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KEY TERMS AND CONCEPTS

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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 case, a larger lattice energy magnitude means a more stable ionic compound.

Sodium chloride gives a simple example. Its lattice energy is 769 kJ/mol, which means 769 kJ is needed to pull one mole of solid NaCl apart into Na+ and Cl- ions in the gas phase. The same amount of heat is released when those ions come together to form solid NaCl. This value shows how strongly the ions attract each other in the crystal.

Lattice energy cannot be measured directly. Instead, it is found with a thermochemical cycle. The Born-Haber cycle uses Hess’s law to break the formation of an ionic solid into separate steps. For cesium fluoride, these steps include sublimation of Cs(s) to Cs(g), half of the F2 bond energy to make F(g), ionization of Cs(g) to Cs+(g) plus an electron, electron affinity of F(g) to make F-(g), and formation of solid CsF from the gaseous ions. The enthalpy change for forming the lattice is the negative of the lattice energy.

For CsF, the overall reaction Cs(s) + 1/2 F2(g) -&gt; CsF(s) has a standard enthalpy of formation of -553.5 kJ/mol. The cycle combines the sublimation enthalpy of cesium, 76.5 kJ/mol, one-half of the F2 bond energy, 79.4 kJ/mol, the ionization energy of cesium, 375.7 kJ/mol, and the electron affinity of fluorine, -328.2 kJ/mol, with the lattice term. Because the total is exothermic, energy is released overall when the solid forms from the elements in their most common states, Cs(s) and F2(g).

Lattice energies for ionic compounds are usually much larger than bond dissociation energies for covalent bonds. Ionic lattice energies often fall between 600 and 4000 kJ/mol, and some are even higher. Single covalent bond dissociation energies are usually about 150 to 400 kJ/mol. These values are not directly comparable because lattice energy reflects many attractions in an extended crystal, while a covalent bond energy describes the interaction of just two atoms.

Ion charge and ion size strongly affect lattice energy. When all other factors stay the same, larger charges give much larger lattice energies, and smaller ions also increase lattice energy. Doubling the charge on both ions can quadruple the lattice energy. For example, LiF has a lattice energy of 1023 kJ/mol, while MgO is 3900 kJ/mol, even though their interionic distances are nearly the same at about 200 pm. Size also matters in compounds such as MgF2 and MgI2, where the smaller F- ion leads to a higher lattice energy than the larger I- ion.

The Born-Haber cycle can also solve for other unknowns in the lattice energy equation. If the other terms are known, the cycle can be used to find a missing quantity such as electron affinity. This makes the method useful for linking measured enthalpy values to the energy changes inside ionic compounds.

Predicting Ionic Stability with Lattice Energy

predicting-ionic-stability-with-lattice-energy

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...

the born-haber cycle

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...

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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...

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lewis structures of molecular compounds and polyatomic ions

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Learn how lattice energy, ion charge, ion size, and the Born-Haber cycle help explain ionic stability and calculate values for compounds like CsF.

Predicting Ionic Stability with Lattice Energy | JoVE K12

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Bond Polarity

The absolute value of the difference in electronegativity (&Delta;EN) of two bonded atoms provides a rough measure of the polarity to be ...

Lewis Structures of Molecular Compounds and Polyatomic Ions

To draw Lewis structures for complicated molecules and molecular ions, it is helpful to follow a step-by-step procedure as outlined:


	Determine the ...

Resonance

The Lewis structure of a nitrite anion (NO2&minus;) may actually be drawn in two different ways, distinguished by the locations of the N-O and N=O ...

Formal Charges

In some cases, there are seemingly more than one valid Lewis structures for molecules and polyatomic ions. The concept of formal charges can be used to ...

Exceptions to the Octet Rule

Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:


	...

Bond Energies and Bond Lengths

Stable molecules exist because covalent bonds hold the atoms together. The strength of a covalent bond is measured by the energy required to break it, ...

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.

Predicting Ionic Stability with Lattice Energy

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