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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Chemistry

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Trends in Lattice Energy: Ion Size and Charge

격자 에너지는 이온 격자의 형성 또는 분리와 관련이 있습니다. 그러나 염화나트륨이나 산화마그네슘이 형성될 때 그것의 격자 에너지는 크게 차이납니다. 왜 각각의 이온 화합물은 다른 격자 에너지를 가지고 있으며, 어떤 요인과 관계됩니까?

이온 화합물에는 정전기 상호작용으로 서로 끌어당기는 많은 전하를 띤 이온들이 질서정연하게 배열되어 있습니다. 쿨롱의 법칙에 따르면 두 이온의 위치 에너지는 이온 사이의 거리에 반비례하며, 결국 이온 반지름에 따라 달라집니다. 주기율표에서 알칼리 및 알칼리 토금속의 이온 반지름은 열 아래로 내려가면서 증가합니다.

금속 이온의 크기가 커질수록 이온 사이의 거리, 즉 결합 길이도 커집니다. 예를 들어 브롬화리튬과 브롬화칼륨의 결합 길이는 각각 217 피코미터와 282 피코미터로 차이가 있습니다. 원자핵 간의 거리가 증가하면 이온 사이의 인력은 감소하고 이온을 분리하는 것이 훨씬 쉬워집니다.

따라서 고체 브롬화리튬 분리하기 위한 격자 에너지는 브롬화 칼륨보다 크며, 각각 몰당807 킬로 줄, 682 킬로 줄 입니다. 이온 반지름 외에도 격자 에너지의 크기는 이온 전하에 따라 달라집니다. 쿨롱의 법칙에 따르면 이온의 위치 에너지는 그것의 전하를 곱한 값에 정비례합니다.

불화 나트륨과 산화칼슘 두 개의 이온 화합물을 고찰해보십시오. 두 화합물의 이온 거리는 비슷하지만 산화칼슘의 격자 에너지는 불화 나트륨보다 거의 4배나 큽니다. 나트륨과 불소는 모두 1가 이온들입니다;그것의 전하들을 곱한 값은 1입니다.

그러나 칼슘과 산화물 이온은 2가이며 그 전하들의 곱은 4로서 불소나트륨의 4배입니다. 이것은 산화칼슘을 기체 이온으로 분리하는 데 불소 나트륨에 비해 4배 되는 에너지가 필요한 이유입니다. 따라서 격자 에너지의 크기는 이온 전하들을 곱한 값에 정비례하며 이온 사이의 거리에 반비례합니다.

격자 에너지의 추세: 이온 크기 및 전하

이온 화합물은 양이온과 음이온 사이의 정전기적 인력으로 인해 안정적입니다. 화합물의 격자 에너지는 이러한 인력의 강도를 나타내는 척도입니다. 이온성 화합물의 격자 에너지(ΔH_lattice)는 고체 1몰을 성분 기체 이온으로 분리하는 데 필요한 에너지로 정의됩니다. 이온성 고체 염화나트륨의 경우 격자 에너지는 공정의 엔탈피 변화입니다.

<img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1058/1058_Equations-01.png" style="height: 40px;" />

규칙

여기서는 이온성 고체가 이온으로 분리되는 경우에 규칙이 사용됩니다. 이는 격자 에너지가 흡열(양수 값)이 됨을 의미합니다. 또 다른 방법은 동일하지만 반대되는 관례를 사용하는 것입니다. 여기서 격자 에너지는 발열(음수 값)이며 이온이 결합하여 격자를 형성할 때 방출되는 에너지로 설명됩니다. 따라서 다른 참고자료에서 격자에너지를 찾아볼 때 어떤 정의가 사용되는지 꼭 확인하시기 바랍니다. 두 경우 모두 격자 에너지의 크기가 클수록 이온 화합물이 더 안정하다는 것을 나타냅니다. 염화나트륨의 경우 ΔH_lattice = 769 kJ입니다. 따라서 고체 NaCl 1몰을 기체 Na^+ 및 Cl^- 이온으로 분리하려면 769 kJ가 필요합니다. 기체 Na^+ 및 Cl^- 이온 1몰이 각각 고체 NaCl을 형성할 때 769 kJ의 열이 방출됩니다.

쿨롱의 법칙과 격자 에너지

이온 결정의 격자 에너지 ΔH격자는 다음 방정식으로 표현될 수 있습니다(전하 사이의 힘을 지배하는 쿨롱의 법칙에서 파생됨).

ΔH_lattice = C(Z^+)(Z^−)/R_o

C는 결정 구조의 유형에 따라 달라지는 상수입니다. Z^+ 및 Z^–는 이온의 전하이고 R_o는 이온간 거리(양이온과 음이온의 반경의 합)입니다. 따라서 이온 결정의 격자 에너지는 이온 전하가 증가하고 이온 크기가 감소함에 따라 급격히 증가합니다. 다른 모든 매개변수가 일정하게 유지되면 양이온과 음이온의 전하가 두 배로 늘어나 격자 에너지가 네 배로 늘어납니다.

예

<ol>
	<li>LiF(Z^+ 및 Z^– = 1)의 격자 에너지는 1023 kJ/mol인 반면 MgO(Z^+ 및 Z^– = 2)의 격자 에너지는 3900 kJ/mol입니다(Ro는 거의 동일합니다. 두 화합물 모두 약 200pm입니다). .</li>
	<li>서로 다른 원자간 거리는 서로 다른 격자 에너지를 생성합니다. 예를 들어, MgF_2(2957 kJ/mol)의 격자 에너지를 MgI_2(2327 kJ/mol)의 격자 에너지와 비교하여 I^-에 비해 F^-의 더 작은 이온 크기가 격자 에너지에 미치는 영향을 관찰합니다.</li>
	<li>귀중한 보석 루비는 Cr^3+의 흔적을 함유한 산화알루미늄, Al_2O_3입니다. Al_2Se_3 화합물은 일부 반도체 장치 제조에 사용됩니다. 이 두 이온 화합물에서는 Z^+와 Z^-의 전하가 동일하므로 격자 에너지의 차이는 R_o에 따라 달라집니다. O_2^- 이온은 Se_2^- 이온보다 작기 때문에 Al_2O_3는 Al_2Se_3보다 이온 간 거리가 짧고 따라서 격자 에너지가 더 큽니다.</li>
	<li>또 다른 예는 NaCl과 비교되는 산화아연(ZnO)입니다. ZnO는 양이온과 음이온의 Z 값이 더 크고 ZnO의 이온 간 거리가 NaCl보다 작기 때문에 격자 에너지가 더 큽니다.</li>
</ol>

이 문서는 에서 발췌되었습니다 <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>

이온 화합물은 양이온과 음이온 사이의 정전기 인력으로 인해 안정적입니다. 화합물의 격자 에너지는 측정 ...

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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...

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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...

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Electronegativity and Polar Bonds

Electronegativity helps explain why some covalent bonds are nonpolar and others are polar covalent. It is a measure of how strongly an atom attracts shared electrons toward itself. That pull controls how electron density is distributed in a bond. Linus Pauling proposed the first electronegativity values by studying bond energies in heteronuclear molecules such as hydrogen fluoride. He expected the bond energy of HF to be about the average of the bond energies of H2 and F2, or 296 kJ/mol. The...

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...

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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...

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lewis structures of molecular compounds and polyatomic ions

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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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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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Construct an explanation to infer the atomic size, reactivity, electronegativity, and ionization energy of an element based on its position in the Periodic Table.

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The Born-Haber Cycle

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An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a compound is a measure ...

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Bond Polarity, Dipole Moment, and Percent Ionic Character

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:


	...

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.

Lattice Energy Patterns in Ionic Solids

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