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

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Chemistry

Matter and its Interactions

PS.Chm.3.1

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Trends in Lattice Energy: Ion Size and Charge

Gitterenergie ist mit der Bildung oder Spaltung eines Ionengitters verbunden. Wenn sich jedoch Natriumchlorid oder Magnesiumoxid bilden, unterscheidet sich ihre Gitterenergie erheblich. Warum hat jede ionische Verbindung eine andere Gitterenergie, und von welchen Faktoren hängt sie ab?Eine ionische Verbindung besteht aus einer geordneten Anordnung einer großen Anzahl geladener Ionen, die durch elektrostatische Wechselwirkungen aneinander angezogen werden. Nach dem Coulomb'schen Gesetz ist die potentielle Energie zweier Ionen umgekehrt proportional zum Abstand zwischen den Ionen, der wiederum vom Ionenradius abhängt. Im Periodensystem nimmt der Ionenradius der Alkali-und Erdalkalimetalle spaltenabwärts zu.Mit zunehmender Größe des Metallions wird auch der Abstand zwischen den Ionen bzw. die Bindungslänge größer. Zum Beispiel unterscheidet sich die Bindungslänge von Lithiumbromid und Kaliumbromid mit 217 picometer und 282 picometer.Aufgrund des größeren Abstands zwischen den Kernen nimmt die Anziehung zwischen den Ionen ab und es wird viel einfacher, die Ionen zu trennen. Daher ist die Gitterenergie zur Trennung von festem Lithiumbromid größer als bei Kaliumbromid und beträgt 807 Kilojoule/mol gegenüber 682 Kilojoule/mol. Neben dem Ionenradius hängt die Größe der Gitterenergie auch von den Ionenladungen ab.Nach dem Coulomb'schen Gesetz ist die potentielle Energie der Ionen direkt proportional zum Produkt ihrer Ladungen. Betrachtet man zwei ionische Verbindungen, Natriumfluorid und Calciumoxid, ist der Ionenabstand bei beiden Verbindungen ähnlich, aber die Gitterenergie von Calciumoxid ist fast viermal so groß wie die von Natriumfluorid. Sowohl Natrium als auch Fluorid sind einwertige Ionen;das Produkt ihrer Ladung ist 1.Calcium-und Oxid-Ionen sind jedoch zweiwertig, und das Produkt ihrer Ladung ist 4, was viermal so hoch ist wie bei Natriumfluorid. Aus diesem Grund wird für die Spaltung des Calciumoxids in gasförmige Ionen fast viermal so viel Energie benötigt wie für Natriumfluorid. Die Größe der Gitterenergie ist also direkt proportional zum Produkt der Ionenladungen und umgekehrt proportional zum Abstand zwischen den Ionen.

trends in der gitterenergie: ionengröße und ladung

Eine ionische Verbindung ist aufgrund der elektrostatischen Anziehung zwischen ihren positiven und negativen Ionen stabil. Die Gitterenergie einer Verbindung ist ein Maß für die Stärke dieser Anziehung. Die Gitterenergie (ΔHGitter) einer ionischen Verbindung ist definiert als die Energie, die erforderlich ist, um ein Mol des Feststoffs in seine gasförmigen Ionen zu zerlegen. Für den ionischen Feststoff Natriumchlorid ist die Gitterenergie die Enthalpieänderung des Prozesses: <img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1058/1058_Equations-01.png" style="height: 40px;"> <h4>Konventionen</h4> Hier wird die Konvention verwendet, das der ionische Feststoff in Ionen aufgetrennt wird und die Gitterenergien somit endotherm sind (positive Werte). Eine andere Möglichkeit besteht darin, eine äquivalente, aber entgegengesetzte Betrachtungsweise zu wählen, bei der die Gitterenergie exotherm ist (negative Werte) und als die Energie beschrieben wird, die freigesetzt wird, wenn sich Ionen zu einem Gitter verbinden. Achten Sie daher darauf, sich zu versichern, welche Definition verwendet wird, wenn Sie in einer anderen Referenz nach Gitterenergien suchen. In beiden Fällen weist ein größerer Betrag der Gitterenergie auf eine stabilere ionische Verbindung hin. Für Natriumchlorid beträgt ΔHGitter = 769 kJ. Somit sind 769 kJ erforderlich, um ein Mol festes NaCl in gasförmige Na+ und Cl–Ionen zu trennen. Wenn jeweils ein Mol gasförmiges Na+ und Cl–Ionen festes NaCl bilden, werden 769 kJ Wärme freigesetzt. <h4>Coulombsches Gesetz und Gitterenergie</h4> Die Gitterenergie ΔHGitter eines Ionenkristalls kann durch die folgende Gleichung ausgedrückt werden (abgeleitet vom Coulombschen Gesetz, welches die Kräfte zwischen elektrischen Ladungen beschreibt): ΔHGitter = C(Z +)(Z−)/Ro  wobei C eine Konstante ist, die von der Art der Kristallstruktur abhängt; Z+ und Z− sind die Ladungen der Ionen und Ro ist der interionische Abstand (die Summe der Radien der positiven und negativen Ionen). Somit nimmt die Gitterenergie eines Ionenkristalls stark zu, wenn die Ladung der Ionen zunimmt oder die Größe der Ionen abnimmt. Wenn alle anderen Parameter konstant gehalten werden, führt eine Verdoppelung der Ladung sowohl des Kations als auch des Anions zu einer Vervierfachung der Gitterenergie. <h4>Beispiele</h4> <ol> <li>Die Gitterenergie von LiF (Z+ und Z− = 1) beträgt 1023 kJ/mol, während die von MgO (Z+ und Z− = 2) 3900 kJ/mol beträgt (Ro ist nahezu gleich – etwa 200 pm für beide Verbindungen).</li> <li>Unterschiedliche interatomare Abstände erzeugen unterschiedliche Gitterenergien. Die unterschiedlichen Gitterenergien von MgF2 (2957 kJ/mol) und die von MgI2 (2327 kJ/mol) zeigen die Auswirkung der kleineren Ionengröße von F− im Vergleich zu I− und ihren Einfluss auf die Stärke der Bindung.</li> <li>Der Edelstein Rubin ist Aluminiumoxid, Al2O3, das Spuren von Cr3+ enthält. Die Verbindung Al2Se3 wird bei der Herstellung einiger Halbleiterbauelemente verwendet. In diesen beiden ionischen Verbindungen sind die Ladungen Z+ und Z− gleich, sodass der Unterschied in der Gitterenergie von Ro abhängt. Da das O2− Ion kleiner ist als das Se2− Ion, hat Al2O3 einen kürzeren interionischen Abstand als Al2Se3 und daher eine größere Gitterenergie.</li> <li>Ein weiteres Beispiel ist Zinkoxid, ZnO, im Vergleich zu NaCl. ZnO hat eine größere Gitterenergie, da die Z-Werte sowohl des Kations als auch des Anions in ZnO größer sind und der interionische Abstand von ZnO kleiner ist als der von NaCl.</li> </ol> Dieser Text wurde adaptiert von <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>

Trends in der Gitterenergie: Ionengröße und Ladung

Eine ionische Verbindung ist aufgrund der elektrostatischen Anziehungskraft zwischen ihren positiven und negativen Ionen stabil. Die Gitterenergie einer Verbindung ist ein Maß für ...

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Chemical Bonding: Basic 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...

the born-haber cycle

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...

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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...

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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

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...

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,...

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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...

key terms and concepts

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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Bonding Types: From Nonpolar Covalent to Ionic

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Lattice Energy&#xA0;

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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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Analyze the structure of atoms and isotopes.

Lattice Energy Patterns in Ionic Solids

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