Name: chemical bonding and bond properties
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 37 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

Types of Chemical Bonds

Los enlaces químicos son esenciales para que los átomos formen una variedad de compuestos, como sal, agua o aleaciones. Comprender la formación de enlaces químicos, o más específicamente la atracción entre átomos, ayudará a comprender y predecir el comportamiento molecular. El núcleo de un átomo está formado por neutrones y protones cargados positivamente, que está rodeado por los electrones cargados negativamente.Cuando dos átomos se acercan, los electrones de un átomo son atraídos al núcleo del otro átomo y viceversa. Simultáneamente, los núcleos de ambos átomos se repelen, al igual que los electrones de cada átomo. Cuando estas interacciones conducen a una reducción neta de la energía potencial, se forma un enlace químico.Hay tres tipos de enlaces químicos. Un metal y un no metal, como el sodio y el cloro, forman un enlace iónico. Los metales tienen energías de ionización bajas, lo que permite una eliminación más fácil de los electrones del átomo.Estos electrones son fácilmente aceptados por los no metales debido a sus altas afinidades electrónicas y su afán por lograr una capa de valencia completa. Por lo tanto, el metal transfiere electrones al no metal formando un catión y un anión. Estas partículas cargadas se atraen entre sí para crear un enlace iónico.El segundo tipo son los enlaces covalentes formados entre dos no metales, como un átomo de carbono y dos átomos de oxígeno que forman dióxido de carbono. Los no metales tienen altas energías de ionización, lo que dificulta la eliminación y transferencia de los electrones de valencia de un átomo a otro, por lo que los electrones se comparten entre los átomos. Estos electrones compartidos interactúan con ambos núcleos de los átomos de enlace y reducen la energía potencial.El último tipo es el enlace metálico formado entre dos átomos de metal. Los átomos de metal pierden electrones de valencia fácilmente, formando un grupo de electrones según el modelo del conjunto de electrones más simple. Los electrones de valencia se deslocalizan en todo el metal y los átomos de metal cargados positivamente son atraídos por este grupo de electrones que mantiene unidos a los átomos.

tipos de enlaces químicos

Las teorías del enlace químico fueron promovidas por el químico americano Gilbert N. Lewis. Él desarrolló un modelo llamado el modelo de Lewis para explicar el tipo y la formación de los diferentes enlaces. El enlace químico es fundamental para la química; explica cómo los átomos o iones se unen para formar moléculas. Explica por qué algunos enlaces son fuertes y otros son débiles, o por qué un enlace de carbono se une a dos oxígenos y no a tres; por qué el agua es H2O y no H4O.  
 
 <h4>Enlaces Iónicos</h4>
 
 Los iones son átomos o moléculas que llevan una carga eléctrica. Un catión (un ion positivo) se forma cuando un átomo neutro pierde uno o más electrones de su capa de valencia, y un anión (un ion negativo) se forma cuando un átomo neutro gana uno o más electrones en su capa de valencia. Los compuestos conformados por iones se llaman compuestos iónicos (o sales), y sus iones constituyentes se mantienen juntos por enlaces iónicos: Fuerzas electrostáticas de atracción entre cationes y aniones con cargas opuestas.
 
 Las propiedades de los compuestos iónicos arrojan algo de luz sobre la naturaleza de los enlaces iónicos. Los sólidos iónicos presentan una estructura cristalina y tienden a ser rígidos y quebradizos; también tienden a tener altos puntos de fusión y ebullición, lo que sugiere que los enlaces iónicos son muy fuertes. Los sólidos iónicos también son malos conductores de la electricidad por la misma razón: La fuerza de los enlaces iónicos impide que los iones se muevan libremente en estado sólido. Sin embargo, la mayoría de los sólidos iónicos se disuelven fácilmente en agua. Una vez disueltos o fundidos, los compuestos iónicos son excelentes conductores de electricidad y calor, ya que los iones pueden moverse libremente.
 
 <h4>Enlaces covalentes</h4>
 
 Los átomos no metálicos frecuentemente forman enlaces covalentes con otros átomos no metálicos. Los enlaces covalentes se forman cuando los electrones son compartidos entre átomos y son atraídos por los núcleos de ambos átomos. Si los átomos que forman un enlace covalente son idénticos, como en H2, Cl2, y otras moléculas diatómicas, entonces los electrones en el enlace deben ser compartidos por igual. Esto se conoce como enlace covalente puro. Cuando los átomos unidos por un enlace covalente son diferentes, los electrones del enlace son compartidos, pero ya no por igual. En cambio, los electrones del enlace son más atraídos a un átomo que por el otro, dando lugar a un cambio de densidad electrónica hacia ese átomo. Esta distribución desigual de los electrones se conoce como enlace covalente polar.
 
 Los compuestos que contienen enlaces covalentes presentan propiedades físicas diferentes a las de los compuestos iónicos. Debido a que la atracción entre moléculas, que son eléctricamente neutras, es más débil que la que existe entre los iones cargados eléctricamente, los compuestos covalentes generalmente tienen puntos de fusión y ebullición mucho más bajos que los compuestos iónicos. Además, mientras que los compuestos iónicos son buenos conductores de la electricidad cuando se disuelven en el agua, la mayoría de los compuestos covalentes son insolubles en agua; ya que son eléctricamente neutros, son malos conductores de la electricidad en cualquier estado.
 
 <h4>Enlaces metálicos</h4>
 
 Los enlaces metálicos se forman entre dos átomos de metal. Paul Dr&uumml;de ha desarrollado un modelo simplificado para describir el enlace metálico llamado 'Modelo del mar de electrones'. Basado en las bajas energías de ionización de los metales, el modelo establece que los átomos de metal pierden fácilmente sus electrones de valencia y se convierten en cationes. Estos electrones de valencia crean un grupo de electrones deslocalizados alrededor de las cationes sobre todo el metal. 
 
 Los sólidos metálicos, como los cristales de cobre, aluminio y hierro, están formados por átomos metálicos, y todos presentan alta conductividad térmica y eléctrica, brillo metálico y maleabilidad. Muchos son muy duros y bastante fuertes. Debido a su maleabilidad (la capacidad de deformarse bajo presión o martilleo) no se rompen y, por lo tanto, son materiales de construcción útiles. Los puntos de fusión de los metales varían ampliamente. El mercurio es un líquido a temperatura ambiente y los metales alcalinos se funden por debajo de 200 °C. Varios metales post transicionales también tienen puntos de fusión bajos, mientras que los metales de transición se funden a temperaturas superiores a 1000 °C. Estas diferencias reflejan diferencias en la fuerza del enlace metálico entre los metales. 
 
 Este texto es adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/7-1-ionic-bonding">Openstax, Química 2e, Sección 7.1: Enlace Iónico</a>, <a href="https://openstax.org/books/chemistry-2e/pages/7-2-covalent-bonding">Openstax, Química 2e, Sección 7.2: Enlace Covalente</a>, y <big><a href="https://openstax.org/books/chemistry-2e/pages/10-5-the-solid-state-of-matter">Openstax, Química 2e, Sección 10.5: El Estado Sólido de la Materia</a>. </big>

Tipos de Enlaces Químicos

Las teorías de los enlaces químicos fueron iniciadas por el químico estadounidense Gilbert N. Lewis. Desarrolló un modelo llamado el modelo de Lewis para explicar el tipo y ...

Education

Educación

JoVE Core

Química

Chemical Bonding: Basic Concepts

Enlace químico: conceptos básicos

KEY TERMS AND CONCEPTS

TÉRMINOS Y CONCEPTOS CLAVE

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

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

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

chemistry

JoVE Core Química da vida a los conceptos de química general mediante lecciones de vídeo basadas en animaciones, concisas y fáciles de entender. Además, en los vídeos de científicos en acción, se muestran experimentos de investigación clásicos y originales realizados en laboratorios de todo el mundo.

Physical Science

PS.PSc.2.1

ohios-learning-standards-and-model-curriculum

Ohio’s Learning Standards and Model Curriculum

Study Of Matter

Bonding and compounds

Structure And Properties Of Matter

Chemical bonding

Physical Geology

Minerals

Chemical bonding (ionic, covalent, metallic)

Tennessee Academic Standards

High School

Matter and Its Interactions

CHEM1.PS1.14

motion-and-stability-forces-and-interactions

Motion and Stability: Forces and Interactions

CHEM1.PS2.1

CHEM1.PS2.2

Virginia  Standards of Learning

HIgh School

CH.3

CH.3.d

PS.3

PS.3.c

Minnesota Academic Standards in Science

1.2.1 9C.1.2.1.2

Arizona Science Standards

HS.P1

HS.P1U1.2

HS+C.P1

HS+C.P1U1.4

Texas Essential Knowledge and Skills

CHEM.6

CHEM.6.E

Alabama Course of Study Science

9-12.C.2.2.a

9-12.C.2.2.c

Florida Department of Education

SC.912.P.8

SC.912.P.8.6

Chemistry and Physics

IPC.7

IPC.7.A

new-york-state-science-learning-standards

New York State Science Learning Standards

HS. Structure and Properties of Matter

HS-PS1-3.

HS. Chemical Reactions

HS-PS1-2.

HS-PS1-4.

HS-PS1-12.

Georgia Standards of Excellence

SPS2.

Chemical Bonds

SC2.

bonding-types-from-nonpolar-covalent-to-ionic

Bonding Types: From Nonpolar Covalent to Ionic

Scientific Laws and Theories

Scientific Laws

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The Scientific Method

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Classifying Matter by State

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Classifying Matter by Composition

Matter: Pure Substances and Mixtures
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The Atomic Theory of Matter

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Subatomic Particles

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Elements: Chemical Symbols and Isotopes

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PS.Chm.1.1

Radioactivity and Nuclear Equations

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Types of Radioactivity

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Nuclear Binding Energy

The difference between the calculated and experimentally measured masses is known as the mass defect of the atom. In the case of helium-4, the mass defect ...

Radioactive Decay and Radiometric Dating

Radioactivity is a spontaneous disintegration of an unstable nuclide and is a random process, as all the nuclei in the sample do not decay simultaneously. ...

Nuclear Fission

Many heavier elements with smaller binding energies per nucleon can decompose into more stable elements that have intermediate mass numbers and larger ...

Nuclear Power

Controlled nuclear fission reactions are used to generate electricity. Any nuclear reactor that produces power via the fission of uranium or plutonium by ...

Nuclear Fusion

The process of converting very light nuclei into heavier nuclei is also accompanied by the conversion of mass into large amounts of energy, a process ...

Nuclear Transmutation

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Use models to compare nuclear reactions including alpha decay, beta decay and gamma decay; nuclear fusion and nuclear fission.

PS.Chm.1.2

Atomic Orbitals

An atomic orbital represents the three-dimensional regions in an atom where an electron has the highest probability to reside. The radial distribution ...

The Pauli Exclusion Principle

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The Energies of Atomic Orbitals

In an atom, the negatively charged electrons are attracted to the positively charged nucleus. In a multielectron atom, electron-electron repulsions are ...

The Aufbau Principle and Hund's Rule

To determine the electron configuration for any particular atom, we can build the structures in the order of atomic numbers. Beginning with hydrogen, and ...

electron-configuration-of-multielectron-atoms

Electron Configuration of Multielectron Atoms

The alkali metal sodium (atomic number 11) has one more electron than the neon atom. This electron must go into the lowest-energy subshell available, the ...

Use models to explain how electrons are distributed in atoms.

PS.Chm.1.3

Periodic Classification of the Elements

The periodic table arranges atoms based on increasing atomic number so that elements with the same chemical properties recur periodically. When their ...

The Periodic Table

As early chemists discovered more elements, they realized that various elements could be grouped by their similar chemical behaviors. One such grouping ...

Atomic Mass

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Molar Mass

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Use the Periodic Table as a model to predict the relative properties of elements based on the pattern of valence electrons in the outermost energy levels of atoms.

PS.Chm.2.1

atomic-radii-and-effective-nuclear-charge

Atomic Radii and Effective Nuclear Charge

The elements in groups of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number ...

Ionic Radii

Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom; it ...

Ionization Energy

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Electron Affinity

The electron affinity (EA) is the energy change for adding an electron to a gaseous atom to form an anion (negative ion).



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Construct an explanation to infer the atomic size, reactivity, electronegativity, and ionization energy of an element based on its position in the Periodic Table.

PS.Chm.2.2

Analyze the structure of atoms and isotopes.

Chemical Bonding and Bond Properties

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