Name: shared electrons in covalent molecules
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 46 s
Channel: JoVE

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North Carolina K-12 Science Standard Course of Study

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Chemistry

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PS.Chm.3.1

Covalent Bonding and Lewis Structures

Covalente bindingen worden gevormd tussen niet-metalen door het delen van valentie-elektronen. Maar waarom delen niet-metalen elektronen liever dan ze over te dragen zoals bij ionische bindingen? Niet-metalen hebben hoge ionisatie-energieën, waardoor het moeilijk is om valentie-elektronen van het ene atoom naar het andere over te brengen.Denk aan een molecuul ammoniak. Het stikstofatoom heeft nog drie elektronen nodig om een octet te bereiken, en het waterstofatoom heeft een elektron nodig om een duet te bereiken. Daarom bindt het stikstofatoom zich met drie waterstofatomen zodat zowel stikstof als waterstof stabiele elektronenconfiguraties bereiken.Omdat elk van deze bindingen één paar elektronen deelt, wordt dit een enkele binding genoemd. Het gedeelde elektronenpaar in de covalente binding wordt een bindingspaar genoemd. De valentie-elektronen die niet deelnemen aan binding, worden het vrije paar of niet-bindende elektronen genoemd.Met 6 valentie-elektronen hebben zuurstofatomen nog twee elektronen nodig om een octet te bereiken. Daarom delen twee zuurstofatomen twee elektronenparen die een dubbele binding vormen. Stikstof, aan de andere kant, deelt de drie ongepaarde elektronen in zijn diatomische vorm, waardoor een drievoudige binding ontstaat.Enkele en meervoudige bindingen verschillen in termen van bindingslengte en sterkte. Drievoudige bindingen zijn korter dan dubbele bindingen, die korter zijn dan enkele bindingen. De bindingssterkte neemt toe met bindingsmultipliciteit.Dit is de reden waarom het moeilijk is om de drievoudige binding in stikstof te verbreken, waardoor het relatief niet-reactief wordt. Het Lewis-model helpt de structuur en stabiliteit van moleculen te voorspellen. Volgens het Lewis-model is H2O een stabiel molecuul;omdat zowel zuurstof als waterstof stabiele elektronenconfiguraties hebben bereikt.Als de zuurstof elektronen deelt met slechts één waterstofatoom, is het resulterende OH-molecuul niet stabiel omdat zuurstof slechts 7 valentie-elektronen heeft en het octet niet kan bereiken. Als er echter een extra elektron aan de zuurstof wordt toegevoegd om het octet te voltooien, wordt het resulterende hydroxide-ion stabiel met een negatieve lading. Covalente bindingen zijn directioneel omdat de gedeelde elektronen twee specifieke atoomparen verbinden.Ionische bindingen daarentegen zijn niet-directioneel en houden verschillende ionen in het rooster vast. In een covalente verbinding worden de individuele moleculen dus als fundamentele eenheden beschouwd.

covalente binding en lewis-structuren

Vergeleken met ionische bindingen, die het resultaat zijn van de overdracht van elektronen tussen metallische en niet-metallische atomen, zijn covalente bindingen het gevolg van de wederzijdse aantrekking van atomen voor een ‘gedeeld’ paar elektronen. Covalente bindingen worden gevormd tussen twee atomen wanneer beide vergelijkbare neigingen hebben om elektronen naar zich toe te trekken (dat wil zeggen wanneer beide atomen identieke of redelijk vergelijkbare ionisatie-energieën en elektronenaffiniteiten hebben). <h4>Fysische eigenschappen van covalente verbindingen</h4> Verbindingen die covalente bindingen bevatten, vertonen andere fysische eigenschappen dan ionische verbindingen. Omdat de aantrekkingskracht tussen moleculen, die elektrisch neutraal zijn, zwakker is dan die tussen elektrisch geladen ionen, hebben covalente verbindingen over het algemeen veel lagere smelt- en kookpunten dan ionische verbindingen. In feite zijn veel covalente verbindingen bij kamertemperatuur vloeistoffen of gassen, en in hun vaste toestand zijn ze doorgaans veel zachter dan ionische vaste stoffen. Terwijl ionische verbindingen goede elektriciteitsgeleiders zijn wanneer ze in water worden opgelost, zijn de meeste covalente verbindingen bovendien onoplosbaar in water; omdat ze elektrisch neutraal zijn, zijn ze in welke staat dan ook slechte elektriciteitsgeleiders. <h4>Vorming van covalente obligaties</h4> Niet-metaalatomen vormen vaak covalente bindingen met andere niet-metaalatomen. Het waterstofmolecuul, H2, bevat bijvoorbeeld een covalente binding tussen de twee waterstofatomen. Twee afzonderlijke waterstofatomen met een bepaalde potentiële energie naderen elkaar, hun valentie-orbitalen (1s) beginnen elkaar te overlappen. De afzonderlijke elektronen op elk waterstofatoom interageren vervolgens met beide atoomkernen en bezetten de ruimte rond beide atomen. De sterke aantrekkingskracht van elk gedeeld elektron op beide kernen stabiliseert het systeem, en de potentiële energie neemt af naarmate de bindingsafstand kleiner wordt. Als de atomen elkaar blijven naderen, beginnen de positieve ladingen in de twee kernen elkaar af te stoten en neemt de potentiële energie toe. De bindingslengte wordt bepaald door de afstand waarop de laagste potentiële energie wordt bereikt. Het is essentieel om te onthouden dat energie moet worden toegevoegd om chemische bindingen te verbreken (een endotherm proces), terwijl bij het vormen van chemische bindingen energie vrijkomt (een exotherm proces). In het geval van H2 is de covalente binding erg sterk; er moet een grote hoeveelheid energie, 436 kJ, worden toegevoegd om de bindingen in één mol waterstofmoleculen te verbreken en de atomen te laten scheiden: <img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1059/1059_Equations-01.png" style="height: 40px;"> Omgekeerd komt dezelfde hoeveelheid energie vrij wanneer één mol H2-moleculen wordt gevormd uit twee mol H-atomen: <img alt="Eq2" src="https://www.jove.com/files/core/CDNSource/lessons/1059/1059_Equations-02.png" style="height: 40px;"> <h4>Lewis-structuren</h4> Lewis-symbolen kunnen worden gebruikt om de vorming van covalente bindingen aan te duiden, die worden weergegeven in Lewis-structuren, tekeningen die de binding in moleculen en polyatomaire ionen beschrijven. Wanneer twee chlooratomen bijvoorbeeld een chloormolecuul vormen, delen ze één paar elektronen: <img alt="Figure1" src="https://www.jove.com/files/core/CDNSource/lessons/1059/1059Image_1.svg" style="height: 90px;"> De Lewis-structuur geeft aan dat elk Cl-atoom drie elektronenparen heeft die niet worden gebruikt bij de binding (de zogenaamde alleenstaande of eenzame paren) en één gedeeld elektronenpaar (geschreven tussen de atomen). Soms wordt een streepje (of lijn) gebruikt om een gedeeld elektronenpaar aan te duiden: Cl-Cl. <ul> <li>Een enkel gedeeld elektronenpaar wordt een enkele binding genoemd. Elk Cl-atoom interageert met acht valentie-elektronen: de zes in de alleenstaande paren en de twee in de enkele binding.</li> <li>Het kan echter nodig zijn dat een paar atomen meer dan één paar elektronen delen om het vereiste octet te bereiken. Een dubbele binding ontstaat wanneer twee elektronenparen worden gedeeld tussen een paar atomen, zoals tussen de koolstof- en zuurstofatomen in CH2O (formaldehyde) en tussen de twee koolstofatomen in C2H4 (ethyleen). <img alt="Figure2" src="https://www.jove.com/files/core/CDNSource/lessons/1059/1059Image_2.svg" style="height: 120px;"></li> <li>Er ontstaat een drievoudige binding wanneer drie elektronenparen worden gedeeld door een paar atomen, zoals bij koolmonoxide (CO) en het cyanide-ion (CN–). <img alt="Figure3" src="https://www.jove.com/files/core/CDNSource/lessons/1059/1059Image_3.svg" style="height: 90px;"></li> </ul> Het periodiek systeem kan worden gebruikt om het aantal valentie-elektronen in een atoom te voorspellen en het aantal bindingen dat zal worden gevormd om een octet te bereiken. Groep 18-elementen zoals argon en helium hebben een gevulde elektronenconfiguratie en nemen dus zelden deel aan chemische bindingen. Atomen uit groep 17, zoals broom of jodium, hebben echter slechts één elektron nodig om het octet te bereiken. Daarom kunnen atomen die tot groep 17 behoren een enkele covalente binding vormen. De atomen van groep 16 hebben 2 elektronen nodig om een octet te bereiken; daarom kunnen ze twee covalente bindingen vormen. Op dezelfde manier heeft koolstof, die tot groep 14 behoort, nog 4 elektronen nodig om een octet te bereiken; aldus kan koolstof vier covalente bindingen vormen.

Covalente binding en Lewis-structuren

Compared to ionic bonds, which results from the transfer of electrons between metallic and nonmetallic atoms, covalent bonds result from the mutual ...

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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 have lower melting and boiling points, and some are liquids or gases at room temperature. When they are solids, they are usually softer than ionic solids. Most covalent compounds are also insoluble in water and are poor conductors of electricity in any state.

Hydrogen gas shows how a covalent bond forms. In H2, two hydrogen atoms approach each other and their 1s valence orbitals overlap. Each electron interacts with both nuclei, which lowers the potential energy and stabilizes the molecule. If the atoms move too close, the positive nuclei repel each other, and the potential energy rises. The bond length is the distance where the potential energy is lowest. Forming a bond releases energy, so it is an exothermic process. Breaking a bond requires energy, so it is endothermic. For H2, 436 kJ must be added to break one mole of molecules, and the same amount is released when one mole forms from hydrogen atoms.

Lewis structures show covalent bonding with Lewis symbols. In chlorine gas, Cl2, the two chlorine atoms share one pair of electrons. Each chlorine atom also has three lone pairs, which are pairs not used in bonding. A shared pair can be shown with a dash, so Cl—Cl represents a single bond. Each chlorine atom then has eight valence electrons around it, including lone pairs and the bonding pair.

Some atoms need more than one shared pair to reach an octet, which is a full outer shell of eight electrons. A double bond forms when two pairs of electrons are shared, as in CH2O and C2H4. A triple bond forms when three pairs are shared, as in CO and CN−. The periodic table helps predict bonding patterns from valence electrons. Group 18 elements such as argon and helium already have filled electron configurations and rarely bond. Group 17 atoms, such as bromine and iodine, usually need one electron and can form one covalent bond. Group 16 atoms need two electrons and can form two covalent bonds. Carbon in group 14 needs four more electrons, so it can form four covalent bonds.

Shared Electrons in Covalent Molecules

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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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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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

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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lewis structures of molecular compounds and polyatomic ions

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key terms and concepts

Explore how nonmetal atoms share electrons to form covalent bonds, from H2 and Cl2 to double and triple bonds, Lewis structures, and bonding patterns.

Shared Electrons in Covalent Molecules | JoVE K12

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Shared Electrons in Covalent Molecules

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