Name: how bond energy predicts reaction heat
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 49 s
Channel: JoVE

K12 Standards

north-carolina-k-12-science-standard-course-of-study

NC K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.3.1

Bond Energies and Bond Lengths

Jede chemische Reaktion ist mit einer Enthalpieänderung verbunden, die mitbestimmt, ob bei der Reaktion Energie freigesetzt oder benötigt wird. Diese Enthalpieänderung kann anhand der durchschnittlichen Bindungsenergien abgeschätzt werden. Die Energie, die erforderlich ist, um ein Mol einer bestimmten Bindung in eine gasförmige chemische Verbindung umzuwandeln, wird als Bindungsenergie bezeichnet und in kJ/mol ausgedrückt.Diese Energie hängt von der Art der gebundenen Atome und der Anzahl der gemeinsamen Elektronenpaare ab. Die Bindungsenergie wird normalerweise als Durchschnitt der Bindungsenergien der gleichen Bindung in verschiedenen Verbindungen ausgedrückt. Die durchschnittliche Bindungsenergie kann verwendet werden, um zu bestimmen, ob eine chemische Reaktion exotherm oder endotherm ist.Man betrachte Ethen und Brom, die mit einander reagieren um 1, 2-Dibromethan zu bilden. Zuerst brechen die Kohlenstoff-Doppelbindung und die Brom-Einfachbindung ein Prozess, der Energiezufuhr benötigt, wodurch die potentielle Energie der Atome erhöht wird. Daher ist das Bindungsbrechen ein endothermer Prozess mit einer positiven Änderung der Enthalpie.Anschließend werden neue Bindungen zwischen Kohlenstoff-und Bromatomen gebildet, die das Produkt ergeben. Die Bindungsbildung erhöht die Stabilität des Moleküls, indem die potentielle Energie verringert wird. Daher ist die Bindungsbildung ein exothermer Prozess, der eine negative Änderung der Enthalpie bewirkt.Nach dem Hess'schen Gesetz ist die Summe der Enthalpieänderungen von Reaktanten und Produkten gleich der Änderung der Gesamtenthalpie einer Reaktion. Die Enthalpieänderung der Reaktanten ist die Summe der Enthalpie der aufgebrochenen Bindungen, während die Enthalpieänderung der Produkte die Summe der Enthalpie der neu gebildeten Bindungen ist. Daher ist die Bildung von 1, 2-Dibromethan mit einer Enthalpie von 255 kJ/mol eine endotherme Reaktion.Neben den Bindungsenthalpien beeinflusst auch die Bindungsart und die Bindungslänge der verbundenen Atome, die den durchschnittlichen Abstand zwischen den Kernen zweier verbundener Atome angibt. Betrachten Sie die unterschiedlichen Bindungen zwischen zwei Stickstoffatomen und zwei Kohlenstoffatomen. Atome mit Mehrfachbindungen, wie z.B.Dreifachbindungen, werden fester zusammengehalten, was zu kürzeren und stärkeren Bindungen führt. Folglich ist das Molekül stabiler und benötigt zur Dissoziation eine höhere Energie. Im Allgemeinen ist die Bindungsstärke bis auf einige Ausnahmen indirekt proportional zur Bindungslänge.

bindungsenergien und bindungslängen

Stabile Moleküle existieren, weil kovalente Bindungen die Atome zusammenhalten. Die Stärke einer kovalenten Bindung wird an der Energie gemessen, die erforderlich ist, um sie zu brechen, also an der Energie, die erforderlich ist, um die gebundenen Atome zu trennen. Die Trennung eines Paares gebundener Atome erfordert Energie – je stärker eine Bindung, desto größer ist die Energie, die zum Aufbrechen erforderlich ist. Die Energie, die erforderlich ist, um eine bestimmte kovalente Bindung in einem Mol gasförmiger Moleküle aufzubrechen, wird Bindungsenergie oder Bindungsdissoziationsenergie genannt. Die Bindungsenergie für ein zweiatomiges Molekül, DX–Y, ist definiert als die Standard-Enthalpieänderung für die endotherme Reaktion: <img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1067/1067_Equations-01.png" style="height: 40px;"> Beispielsweise beträgt die Bindungsenergie der reinen kovalenten H–H-Bindung, DH–H, 436 kJ/mol gebrochener H–H-Bindungen: <img alt="Eq2" src="https://www.jove.com/files/core/CDNSource/lessons/1067/1067_Equations-02.png" style="height: 40px;"> Moleküle mit drei oder mehr Atomen haben zwei oder mehr Bindungen. Die Summe aller Bindungsenergien in einem solchen Molekül entspricht der Standard-Enthalpieänderung für die endotherme Reaktion, die alle Bindungen im Molekül aufbricht. Beispielsweise entspricht die Summe der vier C–H-Bindungsenergien in CH4, 1660 kJ, der Standardenthalpieänderung der Reaktion: <img alt="Figure1" src="https://www.jove.com/files/core/CDNSource/lessons/1067/1067Image_1.svg" style="height: 60px;"> Die durchschnittliche C–H-Bindungsenergie, DC–H, beträgt 1660/4 = 415 kJ/mol, da pro Mol der Reaktion vier Mol C–H-Bindungen gebrochen werden. Obwohl die vier C–H-Bindungen im ursprünglichen Molekül äquivalent sind, benötigen sie nicht jeweils die gleiche Energie, um zu brechen; Sobald die erste Bindung gebrochen ist (was 439 kJ/mol erfordert), sind die verbleibenden Bindungen leichter zu brechen. Der Wert von 415 kJ/mol ist der Durchschnitt und nicht der genaue Wert, der zum Aufbrechen einer Bindung erforderlich ist. Die Stärke einer Bindung zwischen zwei Atomen nimmt mit zunehmender Anzahl der Elektronenpaare in der Bindung zu. Im Allgemeinen nimmt die Bindungslänge mit zunehmender Bindungsstärke ab. Somit sind Dreifachbindungen stärker und kürzer als Doppelbindungen zwischen denselben beiden Atomen; Ebenso sind Doppelbindungen stärker und kürzer als Einfachbindungen zwischen denselben beiden Atomen. Wenn ein Atom an mehrere Atome in einer Gruppe bindet, nimmt die Bindungsstärke normalerweise ab, je weiter wir uns in der Gruppe bewegen. Beispielsweise beträgt C–F 439 kJ/mol, C–Cl 330 kJ/mol und C–Br 275 kJ/mol. Bindungsenergien können verwendet werden, um ungefähre Enthalpieänderungen für Reaktionen, auch Bindungsenthalpien genannt, zu berechnen, bei denen keine Bildungsenthalpien verfügbar sind. Berechnungen dieser Art geben auch Aufschluss darüber, ob eine Reaktion exotherm oder endotherm ist. Eine exotherme Reaktion (ΔH negativ, erzeugte Wärme) entsteht, wenn die Bindungen in den Produkten stärker sind als die Bindungen in den Reaktanten. Eine endotherme Reaktion (ΔH positiv, Wärme absorbiert) entsteht, wenn die Bindungen in den Produkten schwächer sind als die in den Reaktanten. Die Enthalpieänderung ΔH für eine chemische Reaktion entspricht ungefähr der Summe der Energie, die zum Aufbrechen aller Bindungen in den Reaktanten erforderlich ist (Energie „in“, positives Vorzeichen) plus der Energie, die freigesetzt wird, wenn alle Bindungen in den Produkten gebildet werden (Energie). „out“, negatives Vorzeichen). Mathematisch lässt sich dies wie folgt ausdrücken: <img alt="Eq3" src="https://www.jove.com/files/core/CDNSource/lessons/1067/1067_Equations-03.png" style="height: 40px;"> In diesem Ausdruck ist das Symbol ? bedeutet „die Summe von“ und D stellt die Bindungsenergie in kJ/mol dar, die immer eine positive Zahl ist. Die Bindungsenergie hängt davon ab, ob es sich bei der jeweiligen Bindung um eine Einfach-, Doppel- oder Dreifachbindung handelt. Daher ist es bei der Berechnung der Enthalpien auf diese Weise wichtig, dass die Bindung in allen Reaktanten und Produkten berücksichtigt wird. Da es sich bei den D-Werten typischerweise um Durchschnittswerte für einen Bindungstyp in vielen verschiedenen Molekülen handelt, liefert diese Berechnung eine grobe Schätzung und keinen exakten Wert für die Reaktionsenthalpie. Betrachten Sie die folgende Reaktion: <img alt="Eq4" src="https://www.jove.com/files/core/CDNSource/lessons/1067/1067_Equations-04.png" style="height: 40px;"> oder <img alt="Eq5" src="https://www.jove.com/files/core/CDNSource/lessons/1067/1067_Equations-05.png" style="height: 40px;"> Um zwei Mol HCl zu bilden, müssen ein Mol H–H-Bindungen und ein Mol Cl–Cl-Bindungen aufgebrochen werden. Die zum Aufbrechen dieser Bindungen erforderliche Energie ist die Summe der Bindungsenergie der H–H-Bindung (436 kJ/mol) und der Cl–Cl-Bindung (243 kJ/mol). Während der Reaktion werden zwei Mol H–Cl-Bindungen gebildet (Bindungsenergie = 432 kJ/mol), wodurch 2 × 432 kJ freigesetzt werden; oder 864 kJ. Da die Bindungen in den Produkten stärker sind als die in den Reaktanten, wird bei der Reaktion mehr Energie freigesetzt als verbraucht: <img alt="Eq6" src="https://www.jove.com/files/core/CDNSource/lessons/1067/1067_Equations-06.png" style="height: 40px;"> <img alt="Eq7" src="https://www.jove.com/files/core/CDNSource/lessons/1067/1067_Equations-07.png" style="height: 40px;"> Diese überschüssige Energie wird als Wärme freigesetzt, die Reaktion ist also exotherm. Dieser Text wurde angepasst von <a href="https://openstax.org/books/chemistry-2e/pages/7-5-strengths-of-ionic-and-covalent-bonds">Openstax, Chemistry 2e, Section 7.5: Bond Strength: Covalent Bonds.</a>

Bindungsenergien und Bindungslängen

Stabile Moleküle existieren, weil kovalente Bindungen die Atome zusammenhalten. Die Stärke einer kovalenten Bindung wird durch die Energie gemessen, die erforderlich ist, um sie zu brechen, ...

Education

Lehre

JoVE Core

Chemie

Eins

Chemical Bonding: Basic Concepts

Chemische Bindungen: Grundkonzepte

KEY TERMS AND CONCEPTS

KERNBEGRIFFE UND KONZEPTE

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 436 kJ/mol. Molecules with three or more atoms have several bonds, so the total bond energy is the sum of all the bonds that must be broken.

Methane shows why average bond energy is useful. Breaking all four C-H bonds in CH4 requires 1660 kJ, so the average C-H bond energy is 415 kJ/mol. That is an average value, not the exact energy for every single C-H bond. After the first bond breaks, which requires 439 kJ/mol, the remaining bonds are easier to break.

Bond strength usually increases as the number of electron pairs in the bond increases. As bond strength increases, bond length usually decreases. Triple bonds are therefore stronger and shorter than double bonds between the same atoms, and double bonds are stronger and shorter than single bonds. Bond strength can also decrease down a group. For example, C-F is 439 kJ/mol, C-Cl is 330 kJ/mol, and C-Br is 275 kJ/mol.

Bond energies can also estimate reaction enthalpy, especially when enthalpies of formation are not available. These calculations can show whether a reaction is exothermic or endothermic. An exothermic reaction has a negative ΔH and releases heat. An endothermic reaction has a positive ΔH and absorbs heat.

To estimate ΔH, add the energy needed to break all bonds in the reactants and subtract the energy released when bonds form in the products. In equation form, this is the sum of bond energies broken plus the sum of bond energies formed, with formed bonds counted as negative. Because bond energy values are usually averages for a bond type, the result is an estimate, not an exact value.

The formation of hydrogen chloride is a useful example. To make two moles of HCl, one mole of H-H bonds and one mole of Cl-Cl bonds must be broken. That requires 436 kJ/mol plus 243 kJ/mol. At the same time, two moles of H-Cl bonds form and release 864 kJ because each H-Cl bond has a bond energy of 432 kJ/mol. Since the product bonds are stronger than the reactant bonds, the reaction releases more energy than it uses and is exothermic.

How Bond Energy Predicts Reaction Heat

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...

electronegativity

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

trends-in-lattice-energy-ion-size-and-charge

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...

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

bond-polarity-dipole-moment-and-percent-ionic-character

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

lewis-structures-of-molecular-compounds-and-polyatomic-ions

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...

resonance

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...

formal charges

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...

exceptions to the octet rule

Octet Rule Exceptions in Lewis Structures

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bond energies and bond lengths

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...

bonding in metals

Electron Sea Model in Metals

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key terms and concepts

Learn how bond energy, bond length, and bond dissociation energy help predict whether reactions absorb or release heat, with examples like H-H, CH4, and HCl.

How Bond Energy Predicts Reaction Heat | JoVE K12

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Pennsylvania State Academic Standards

Physical Science

Chemical Reactions

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High School

Matter And Its Interactions

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Study Of Matter

Bonding and compounds

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HIgh School

CH.3

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PH.9

PH.9.e

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New Jersey Student Learning Standards for Science

Matter and Its Interactions

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Motion and Stability: Forces and Interactions

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IPC.7

IPC.7.A

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South Carolina College and Career Ready Science Standards

C-PS1-4

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New York State Science Learning Standards

HS. Chemical Reactions

HS-PS1-2.

HS-PS1-4.

HS-PS1-12.

Georgia Standards of Excellence

SC2.

Energy Changes in Chemical Reactions

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Radioactive Decay and Radiometric Dating

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Nuclear Power

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PS.Chm.1.2

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The Energies of Atomic Orbitals

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The Aufbau Principle and Hund's Rule

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Electron Configuration of Multielectron Atoms

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PS.Chm.1.3

Periodic Classification of the Elements

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The Periodic Table

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Analyze and interpret data to explain the mechanisms and properties of the two main types of intramolecular (ionic and covalent) bonds.

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How Bond Energy Predicts Reaction Heat

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