Name: valence electrons, lewis symbols, and bonding
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 36 s
Channel: JoVE

K12 Standards

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NC K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.3.1

Lewis Symbols and the Octet Rule

Chemische Bindung beinhaltet die gemeinsame Nutzung oder Übertragung von Valenzelektronen zwischen zwei oder mehr Atomen, die zur Senkung der potentielle Energien führen was somit zur Bindungsbildung beiträgt. Der Prozess der Bindungsbildung mittels Valenzelektronen wird durch das Lewis-Modell erklärt. Im Lewis-Modell wird jedes Valenzelektron eines Atoms aus dem s oder p Orbital als ein Punkt dargestellt um die Abkürzung des Elements, auch Lewis-Symbol genannt.Sauerstoff hat zum Beispiel die Elektronenkonfiguration von 1s2, 2s2, 2p4 und damit sechs Valenzelektronen. Die ersten vier Punkte befinden sich auf jeder Seite des Sauerstoffsymbols, ein Punkt in jede Richtung. Die verbleibenden zwei Valenzelektronen sind auf zwei anderen Seiten um das Atom angeordnet.Jede Seite ist gleichwertig und kann ein Maximum von zwei Elektronen aufnehmen. Die Anzahl der ungepaarten Punkte entspricht der Anzahl der Bindungen, die jedes Atom dieses Elements eingehen kann. Elemente der letzten Gruppe, wie Neon, sind Edelgase und beteiligen sich nicht bereitwillig an den chemischen Bindungen.Edelgase haben hohe Ionisierungsenergien und sind sehr stabil angesichts ihrer Elektronenkonfiguration mit der vollen äußeren Schale. Das Lewis-Symbol für Neon hat acht Punkte, zwei Punkte auf jeder Seite, die eine gefüllte Elektronenkonfiguration darstellen;mit anderen Worten, ein Oktett. Die Oktett-Regel besagt, dass ein Atom dazu neigt, Elektronen in Form von Bindungen zu verlieren, zu gewinnen oder zu teilen bis eine stabile Elektronenkonfiguration, ein Oktett, erreicht ist.Denken Sie an Kohlendioxid. Kohlenstoff hat vier ungepaarte Elektronen und Sauerstoff hat zwei. Die Anzahl der ungepaarten Elektronen entspricht der Anzahl der Elektronen, die benötigt werden, um ein Oktett zu erreichen.Daher teilt sich Kohlenstoff zwei Elektronen mit dem einem Sauerstoff, und die beiden anderen Elektronen mit dem anderen Sauerstoff, so dass jedes Atom ein Oktett erreichen kann. Wasserstoff und Helium sind eine Ausnahme zur Oktett-Regel, da sie nur einen s Orbital haben, welches nur 2 Elektronen enthalten kann;also sollten Wasserstoff oder Helium ein Duett erreichen.

lewis-symbole und die oktettregel

Chemische Bindungen sind komplexe Wechselwirkungen zwischen zwei oder mehr Atomen oder Ionen, die die potentielle Energie des Moleküls verringern. Gilbert N. Lewis entwickelte ein Modell namens Lewis-Modell, das die Darstellung der Bildung chemischer Bindungen vereinfachte und simple Erklärungsmodelle für die chemischen Bindungen lieferte, die in den meisten Verbindungen vorkommen. <h4>Lewis-Modell</h4> Das Lewis-Modell stellt die Bildung chemischer Bindungen als gemeinsame Nutzung oder Übertragung von Valenzelektronen dar, die somit zu einer stabile Elektronenkonfiguration beider Atome führen. Eine ionische Bindung entsteht, wenn Elektronen zwischen einem Metall und einem Nichtmetall übertragen werden, wohingegen eine kovalente Bindung entsteht, wenn Elektronen zwischen zwei Nichtmetallen geteilt werden. Das Lewis-Modell wird ausschließlich zur Beschreibung der Bildung von Bindungen verwendet, ohne die Energieänderungen zu berücksichtigen, die mit der Anziehung und Abstoßung zwischen Elektronen und Kernen benachbarter Atome einhergeht. Während diese Wechselwirkungen für die chemische Bindung von zentraler Bedeutung sind, ist die genaue Bestimmung der Werte kompliziert. Anstelle sich darauf zu fokusieren entwarf Lewis spezielle Zeichnungen zur Darstellung chemischer Bindungen mithilfe von Valenzelektronen, durch sogenannte Lewis-Symbole. <h4>Lewis-Symbole</h4> Lewis-Symbole beschreiben die Elektronenkonfigurationen der Valenzschale von Atomen und einatomigen Ionen. Ein Lewis-Symbol besteht aus einem Elementsymbol, das für jedes seiner Valenzelektronen von einem Punkt umgeben ist. Natrium hat beispielsweise ein Valenzelektron; also wird ein Punkt um das Symbol Na gezeichnet. <img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1055/1055_Image 1.png" style="height: 20px;"> Bei Hauptgruppenelementen wird die Anzahl der Valenzelektronen durch eine mit Buchstaben versehene Gruppennummer im Periodensystem angegeben. Lithium (Li) gehört beispielsweise zur Gruppe IA und hat ein Elektron; Beryllium (Be) ist ein Element der Gruppe IIA und hat zwei Valenzelektronen. Es gibt Ausnahmen vom Lewis-Modell. Bei Helium ist die Zahl der Valenzelektronen nicht gleich der Gruppenzahl. Die Übergangsmetalle, Lanthaniden und Aktiniden haben unvollständig gefüllte Innenschalen; daher können sie nicht in einfachen Lewis-Punkt-Symbolen geschrieben werden. <h4>Die Oktettregel</h4> Die Halogenmoleküle (F2, Br2, I2 und At2) bilden Bindungen wie das Chlormolekül: eine Einfachbindung zwischen Atomen und drei freie Elektronenpaare pro Atom. Dadurch kann jedes Halogenatom eine Edelgaselektronenkonfiguration haben. Die Tendenz von s- oder p-Blockatomen, genügend Bindungen zu bilden, um acht Valenzelektronen zu erhalten, wird als Oktettregel bezeichnet. Die Oktettregel sagt die Kombinationen von Atomen voraus, die eine niedrigere potentielle Energie haben wenn sie sich verbinden. Die Anzahl der Bindungen, die ein Atom bilden kann, lässt sich oft anhand der Anzahl der Elektronen vorhersagen, die zum Erreichen eines Oktetts (acht Valenzelektronen) erforderlich sind. Dies gilt insbesondere für die Nichtmetalle der zweiten Periode des Periodensystems (C, N, O und F). <ul> <li>Elemente der Gruppe 14 haben vier Elektronen in ihrer äußersten Schale und benötigen daher vier weitere Elektronen, um ein Oktett zu erreichen. Diese vier Elektronen können durch die Bildung von vier kovalenten Bindungen gewonnen werden, wie Kohlenstoff in CH4 (Methan) und Silizium in SiH4 (Silan).</li> <li>Elemente der Gruppe 15 wie Stickstoff haben fünf Valenzelektronen im atomaren Lewis-Symbol: ein freies Elektronenpaar und drei ungepaarte Elektronen. Diese Atome bilden drei kovalente Bindungen, wie in NH3 (Ammoniak).</li> <li>Elemente der Gruppe 16 wie Sauerstoff und andere Atome erhalten ein Oktett, indem sie zwei kovalente Bindungen bilden – wie die Bindung mit zwei Wasserstoffatomen in H2O (Wasser).</li> </ul> Es gibt Ausnahmen von der Oktettregel. Da Wasserstoff nur zwei Elektronen benötigt, um seine Valenzschale zu füllen, stellt er eine Ausnahme von der Oktettregel dar. In diesem Fall möchte Wasserstoff den Zustand eines Duetts anstelle des Oktetts erreichen. Auch die Übergangselemente und inneren Übergangselemente folgen nicht der Oktettregel. Dieser Text wurde adaptiert von <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Chemistry 2e, Chapter 7.3: Lewis Symbols and Structures</a>.

Lewis-Symbole und die Oktettregel

Chemische Bindungen sind komplexe Wechselwirkungen zwischen zwei oder mehr Atomen oder Ionen, die die potentielle Energie des Moleküls reduzieren. Gilbert N. Lewis entwickelte eine ...

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JoVE Core

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Chemical Bonding: Basic Concepts

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KEY TERMS AND CONCEPTS

KERNBEGRIFFE UND KONZEPTE

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 electrons. The model does not include the detailed energy changes from attractions and repulsions between electrons and nuclei on nearby atoms.

Lewis symbols are special drawings that show valence electrons for atoms and monatomic ions. A Lewis symbol uses the element symbol with one dot for each valence electron. For example, sodium is written with one dot because it has one valence electron. For many main group elements, the group number on the periodic table matches the number of valence electrons, such as lithium in group IA and beryllium in group IIA.

There are exceptions to this pattern. Helium does not match its group number in a simple Lewis symbol. Transition metals, lanthanides, and actinides also do not fit the Lewis symbol pattern well because they have incompletely filled inner shells. They cannot be shown with simple Lewis dot symbols.

The octet rule helps predict how many bonds many s- and p-block atoms will form. These atoms often form enough bonds to reach eight valence electrons, which gives them a noble gas electron configuration. This idea is especially useful for second-period nonmetals such as carbon, nitrogen, oxygen, and fluorine. It can help predict which combinations of atoms are likely to have lower potential energy when they bond.

Group 14 elements have four valence electrons and need four more to reach an octet. They can do this by forming four covalent bonds, as carbon does in methane, CH4, and silicon does in silane, SiH4. Group 15 elements such as nitrogen have five valence electrons, shown as one lone pair and three unpaired electrons, so they usually form three covalent bonds, as in ammonia, NH3. Group 16 elements such as oxygen form two covalent bonds to complete an octet, as in water, H2O.

Not every atom follows the octet rule. Hydrogen is an exception because it only needs two electrons in its valence shell. When hydrogen reaches this state, it has a duet, not an octet. Transition elements and inner transition elements also do not follow the octet rule.

Valence Electrons, Lewis Symbols, and Bonding

valence-electrons-lewis-symbols-and-bonding

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

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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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...

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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...

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...

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

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...

bonding in metals

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 valence electrons, Lewis symbols, and the octet rule help predict ionic and covalent bonding in common high school chemistry examples.

Valence Electrons, Lewis Symbols, and Bonding | JoVE K12

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Valence Electrons, Lewis Symbols, and Bonding

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