Name: chemical bonding and bond properties
Uploaded: 2020-09-03T13:38:28.000Z
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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

Types of Chemical Bonds

Les liaisons chimiques sont essentielles pour que les atomes forment une variété de composés, tels que le sel, l'eau ou les alliages. Comprendre la formation de liaisons chimiques, ou plus précisément l'attraction entre les atomes, aidera à comprendre et à prédire le comportement moléculaire. Le noyau d'un atome est constitué de neutrons et de protons chargés positivement, qui est entouré par les électrons chargés négativement.Lorsque deux atomes se rapprochent l'un de l'autre, les électrons d'un atome sont attirés vers le noyau de l'autre atome et vice versa. Simultanément, les noyaux des deux atomes se repoussent, tout comme les électrons de chaque atome. Lorsque ces interactions conduisent à une réduction nette de l'énergie potentielle, une liaison chimique est formée.Il existe trois types de liaisons chimiques. Un métal et un non-métal, comme le sodium et le chlore, forment une liaison ionique. Les métaux ont de faibles énergies d'ionisation permettant une élimination plus facile des électrons de l'atome.Ces électrons sont facilement acceptés par les non-métaux compte tenu de leurs affinités électroniques élevées et de leur désir d'obtenir une coquille de valence complète. Par conséquent, le métal transfère des électrons au non-métal formant un cation et un anion. Ces particules chargées s'attirent pour créer une liaison ionique.Le deuxième type est des liaisons covalentes formées entre deux non-métaux, comme un carbone et deux atomes d'oxygène formant du dioxyde de carbone. Les non-métaux ont des énergies d'ionisation élevées, ce qui rend difficile l'élimination et le transfert des électrons de valence d'un atome à un autre, d'où les électrons sont partagés entre les atomes. Ces électrons partagés interagissent avec les deux noyaux des atomes de liaison et réduisent l'énergie potentielle.Le dernier type est une liaison métallique formée entre deux atomes métalliques. Les atomes métalliques perdent facilement des électrons de valence, formant un pool d'électrons selon le modèle de mer d'électrons le plus simple. Les électrons de valence se délocalisent sur l'ensemble du métal et les atomes de métal chargés positivement sont attirés par ce pool d'électrons qui maintient les atomes ensemble.

types de liaisons chimiques

Le chimiste américain Gilbert N. Lewis a été le pionnier des théories de la liaison chimique. Il a conçu un modèle appelé le modèle de Lewis pour expliquer le type des différentes liaisons et leur formation. La liaison chimique est essentielle en chimie ; elle explique comment les atomes ou les ions se lient pour former des molécules. Elle explique pourquoi certaines liaisons sont fortes et d'autres faibles, ou pourquoi un carbone se lie avec deux oxygènes et non trois ; pourquoi l'eau est H2O et non H4O.  
 
 <h4>Liaisons ioniques </h4>
 
 Les ions sont des atomes ou des molécules portant une charge électrique. Un cation (un ion positif) se forme lorsqu'un atome neutre perd un ou plusieurs électrons de sa couche de valence, et un anion (un ion négatif) se forme lorsqu'un atome neutre gagne un ou plusieurs électrons sur sa couche de valence. Les composés constitués d'ions sont appelés des composés ioniques (ou sels), et leurs constituants ioniques sont maintenus ensemble par des liaisons ioniques : des forces d'attraction électrostatiques entre les cations et les anions de charges opposées. 
 
 Les propriétés des composés ioniques mettent en évidence la nature des liaisons ioniques. Les solides ioniques présentent une structure cristalline et ont tendance à être rigides et cassants ; ils ont également tendance à avoir des points de fusion et d'ébullition élevés, ce qui suggère que les liaisons ioniques sont très fortes. Les solides ioniques sont également de mauvais conducteurs d'électricité pour la même raison : la force des liaisons ioniques empêche les ions de se déplacer librement à l'état solide. Toutefois, la plupart des solides ioniques se dissolvent facilement dans l'eau. Une fois dissous ou fondus, les composés ioniques sont d'excellents conducteurs d'électricité et de chaleur car les ions peuvent se déplacer librement. 
 
 <h4>Liaisons covalentes </h4>
 
 Les atomes des non-métaux forment souvent des liaisons covalentes avec d'autres atomes de non-métaux. Les liaisons covalentes se forment lorsque les électrons sont mis en commun entre les atomes et qu'ils sont attirés par les noyaux des deux atomes. Si les atomes qui forment une liaison covalente sont identiques, comme dans H2, Cl2 et d'autres molécules diatomiques, alors les électrons de la liaison doivent être partagés de façon égale. C'est ce qu'on appelle une liaison covalente pure. Lorsque les atomes liés par une liaison covalente sont différents, les électrons liants sont partagés, mais plus de façon égale. Au lieu de cela, les électrons liants sont plus attirés par un atome que par l'autre, ce qui donne lieu à un déplacement de la densité électronique vers cet atome. Cette répartition inégale des électrons est connue sous le nom de liaison covalente polarisée. 
 
 Les composés qui contiennent des liaisons covalentes présentent des propriétés physiques différentes de celles des composés ioniques. Comme l'attraction entre les molécules, qui sont électriquement neutres, est plus faible que celle qui existe entre les ions chargés électriquement, les composés covalents ont généralement des points de fusion et d'ébullition beaucoup plus faibles que les composés ioniques. En outre, alors que les composés ioniques sont de bons conducteurs d'électricité lorsqu'ils sont dissous dans l'eau, la plupart des composés covalents sont insolubles dans l'eau ; comme ils sont électriquement neutres, ils sont de mauvais conducteurs d'électricité quel que soit l'état. 
 
 <h4>Liaisons métalliques </h4>
 
 Des liaisons métalliques se forment entre les atomes de deux métaux. Paul Drüde a conçu un modèle simplifié pour décrire la liaison métallique, appelé le &#39 ; modèle de la mer d'électrons &#39 ;. Basé sur les faibles énergies d'ionisation des métaux, le modèle indique que les atomes d'un métal perdent facilement leurs électrons de valence et deviennent des cations. Ces électrons de valence créent une mer d'électrons délocalisés entourant les cations sur l'ensemble du métal.  
 
 Les solides métalliques, tels que les cristaux de cuivre, d'aluminium et de fer, sont formés par les atomes des métaux et présentent tous une conductivité thermique et électrique élevée, un éclat métallique et une malléabilité. Beaucoup sont très durs et assez solides. En raison de leur malléabilité (la capacité de se déformer sous une pression ou un martelage), ils ne se brisent pas et, par conséquent, font des matériaux de construction utiles. Les points de fusion des métaux varient considérablement. Le mercure est un liquide à température ambiante et les métaux alcalins fondent en dessous de 200 °C. Plusieurs métaux de post-transition ont également des points de fusion bas, alors que les métaux de transition fondent à des températures supérieures à 1000 °C. Ces différences reflètent les différences au niveau de la force de la liaison métallique parmi les métaux.  
 
 Ce texte est adapté de <a href="https://openstax.org/books/chemistry-2e/pages/7-1-ionic-bonding">Openstax, Chimie 2e, Section 7.1 : Liaison ionique</a>, <a href="https://openstax.org/books/chemistry-2e/pages/7-2-covalent-bonding">Openstax, Chimie 2e, Section 7.2 : Liaison covalente</a>, et <big><a href="https://openstax.org/books/chemistry-2e/pages/10-5-the-solid-state-of-matter">Openstax, Chimie 2e, Section 10.5 : L'état solide de la matière</a>. </big>

Types de liaisons chimiques

Les théories des liaisons chimiques ont été mises au point par le chimiste américain Gilbert N. Lewis. Il a développé un modèle appelé le modèle de Lewis pour expliquer le type et ...

Education

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Chemical Bonding: Basic Concepts

Liaisons chimiques : concepts de base

KEY TERMS AND CONCEPTS

TERMES ET CONCEPTS CLÉS

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 shell. An anion is a negative ion formed when a neutral atom gains one or more electrons in its valence shell. Ionic compounds, also called salts, are held together by electrostatic attraction between oppositely charged cations and anions.

The properties of ionic compounds show why ionic bonds are strong. Ionic solids usually have a crystalline structure and are rigid and brittle. They also tend to have high melting and boiling points. In the solid state, they are poor conductors of electricity because the ions cannot move freely. Many ionic solids dissolve readily in water. When they are dissolved or melted, they conduct electricity and heat well because the ions can move about freely.

Covalent bonding is common between nonmetal atoms. It forms when atoms share electrons, and the shared electrons are attracted by the nuclei of both atoms. When identical atoms bond, as in H2 and Cl2, the electrons are shared equally. This is called a pure covalent bond. When different atoms bond, the electrons are still shared, but not equally. This creates a polar covalent bond, where electron density shifts toward one atom.

Covalent compounds have different properties from ionic compounds. Because the attraction between neutral molecules is weaker than the attraction between charged ions, covalent compounds usually have much lower melting and boiling points. Most are insoluble in water. They are also poor conductors of electricity in any state because they are electrically neutral.

Metallic bonding forms between metal atoms. Paul Drude's electron sea model describes this bond by saying that metal atoms lose valence electrons easily and become cations. The valence electrons form a pool of delocalized electrons around the cations throughout the metal. Metallic solids, such as copper, aluminum, and iron, conduct heat and electricity well. They also have metallic luster and are malleable, so they can deform under pressure or hammering without shattering.

Metals can have very different melting points. Mercury is liquid at room temperature, and alkali metals melt below 200 degrees Celsius. Several post-transition metals also have low melting points. Transition metals, however, melt above 1000 degrees Celsius. These differences reflect differences in the strength of metallic bonding among metals.

Chemical Bonding and Bond Properties

types of chemical bonds

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

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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

Explore ionic, covalent, and metallic bonding, plus the properties of each. Learn how electrons, ions, and metals explain structure, conductivity, and melting points.

Chemical Bonding and Bond Properties | JoVE K12

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Bonding Types: From Nonpolar Covalent to Ionic

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PS.Chm.1.3

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