Name: metallic bonding and crystal packing
Uploaded: 2020-09-24T15:06:49.000Z
Duration: 2 min 37 s
Channel: JoVE

K12 Standards

ohios-learning-standards-and-model-curriculum

Ohio’s Learning Standards and Model Curriculum

9-12

Chemistry

Structure And Properties Of Matter

Intermolecular forces of attraction

Metallic Solids

Metallic solids, such as copper, aluminum, and gold, are solids that have metal atoms as their constituent particles, which typically are held together by strong metallic bonds.
In a metallic bond, the electron clouds of neighboring metal atoms overlap, such that these valence electrons are delocalized throughout the solid. Metallic solids are sometimes described as a lattice of stationary metal cations in a sea of delocalized electrons. 
This results in a more flexible internal structure than that of molecular or ionic solids. Metallic solids generally can bend and deform without breaking to a greater extent than other solids. 
As exhibited by gold and mercury, their melting points vary widely depending on the strength of the metallic bonds. The mobile electrons also make these solids excellent conductors of heat and electricity. 
The atoms of pure metallic solids are typically arranged, or &#8216;packed&#8217;, in one of the two most efficient possible configurations for spheres of a uniform size: hexagonal close-packing or cubic close-packing.
In hexagonal close-packing, the metal atoms are arranged hexagonally in two repeating layers, where the atoms in the second layer occupy the depressions

metallic solids

Metallic solids such as crystals of copper, aluminum, and iron are formed by metal atoms. The structure of metallic crystals is often described as a uniform distribution of atomic nuclei within a &#8220;sea&#8221; of delocalized electrons. The atoms within such a metallic solid are held together by a unique force known as metallic bonding that gives rise to many useful and varied bulk properties.
All metallic solids exhibit high thermal and electrical conductivity, metallic luster, and malleability. Many are very hard and quite strong. Because of their malleability (the ability to deform under pressure or hammering), they do not shatter and, therefore, make useful construction materials. The melting points of the metals vary widely. Mercury is a liquid at room temperature, and the alkali metals melt below 200 &#176;C. Several post-transition metals also have low melting points, whereas the transition metals melt at temperatures above 1000 &#176;C. These differences reflect differences in the strengths of metallic bonding among metals.
Properties of Metallic Solids
Owing to their crystalline structure, metallic solids exhibit few unique properties associated with the structure and have been tabulated in the following table.
<table style="margin-left: auto; margin-right: auto;">
	<tbody>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Type of Solid</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Type of Particles</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Type of Attractions</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Properties</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Examples</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Metallic</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Atoms of electropositive elements</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Metallic bonds</td>
			<td style="color: black; border: 1px solid black; text-align: center;">shiny, malleable, ductile, conducts heat and electricity well, variable hardness and melting temperature</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Cu, Fe, Ti, Pb, U</td>
		</tr>
	</tbody>
</table>
Crystal Structure of Metallic Solids: Close-packing
Solids that are made of identical atoms can have two types of arrangements: square or close-packed (Figure 1). Since close-packing maximizes the overall attractions between atoms and minimizes the total intermolecular energy, the atoms in most metals pack in this manner.
<img alt="Atomic packing in crystal structures; face-centered cubic, hexagonal close-packed; diagram." src="/files/ftp_upload/11356/11356_Image_1.png" style="text-align: center; height: 110px;" data-altupdatedat="1754904545000" data-alt="Atomic packing factor in crystalline structures; APF=0.74 for FCC, 0.68 for BCC; diagram." />
Figure 1. Square vs close-packed arrangement.
We find two types of closest packing in simple metallic crystalline structures: hexagonal closest packing (HCP), and cubic closest packing (CCP). Both consist of repeating layers of hexagonally arranged atoms. In both types, a second layer (B) is placed on the first layer (A) so that each atom in the second layer is in contact with three atoms in the first layer. The third layer is positioned in one of two ways.
In HCP, atoms in the third layer are directly above atoms in the first layer (i.e., the third layer is also a type A), and the stacking consists of alternating type A and type B close-packed layers (i.e., ABABAB&#8943;) (Figure 2a).
In CCP, atoms in the third layer are not above atoms in either of the first two layers (i.e., the third layer is type C), and the stacking consists of alternating type A, type B, and type C close-packed layers (i.e., ABCABCABC&#8943;) (Figure 2b). Cubic face-centered (FCC) and CCP arrangements are actually the same structures with compact packing of atoms, occupying 74% of the volume.
<img alt="Crystalline structure diagram; substitutional and interstitial alloys with atomic layers; structural analysis." src="/files/ftp_upload/11356/11356_Image_2.png" style="text-align: center; height: 170px;" data-altupdatedat="1754904545000" data-alt="Atomic arrangement diagram of substitutional alloys, illustrating lattice disturbances." />
Figure 2. (a) Hexagonal close-packing consists of two alternating layers (ABABAB&#8230;). (b) Cubic close-packing consists of three alternating layers (ABCABCABC&#8230;).
In both types of packing, each atom contacts six atoms in its own layer, three in the layer above, and three in the layer below. Thus each atom touches 12 near neighbors and therefore has a coordination number of 12.
About two&#8211;thirds of all metals crystallize in closest-packed arrays with coordination numbers of 12. Metals that crystallize in an HCP structure include Cd, Co, Li, Mg, Na, and Zn, and metals that crystallize in a CCP structure include Ag, Al, Ca, Cu, Ni, Pb, and Pt.
This text has been adapted from <a href="https://openstax.org/books/chemistry-2e/pages/10-5-the-solid-state-of-matter">Openstax, Chemistry 2e, Sections 10.5 The Solid State of Matter</a>, and <a href="https://openstax.org/books/chemistry-2e/pages/10-6-lattice-structures-in-crystalline-solids">10.6 Lattice Structures in Crystalline Solids</a>.

Metallic solids such as crystals of copper, aluminum, and iron are formed by metal atoms. The structure of metallic crystals is often described as a ...

Education

JoVE Core

Liquids, Solids, and Intermolecular Forces

KEY TERMS AND CONCEPTS

metallic bonding and crystal packing

Metallic solids are made of metal atoms arranged in crystals. Copper, aluminum, and iron are common examples. These solids are often described as metal nuclei spread through a sea of delocalized electrons. The atoms are held together by metallic bonding, and that bonding gives metals many of their useful bulk properties.

Metallic solids are shiny, conduct heat and electricity well, and can be shaped without breaking. This shaping ability is called malleability, which means a metal can deform under pressure or hammering. Because of that, metals do not shatter easily and are useful as construction materials. Many metallic solids are also hard and strong, although their hardness can vary.

Their melting points also vary widely. Mercury is a liquid at room temperature, and the alkali metals melt below 200 °C. Several post-transition metals also have low melting points. In contrast, transition metals melt above 1000 °C. These differences show that metallic bonding is not equally strong in every metal.

Most metals pack their atoms in close-packed crystal structures. Close-packing helps maximize attractions between atoms and lowers the total energy. Two common forms are hexagonal close packing, or HCP, and cubic close packing, or CCP. Both use repeating layers of hexagonally arranged atoms.

In HCP, the stacking pattern follows ABABAB... In CCP, the pattern follows ABCABCABC... The cubic face-centered, or FCC, arrangement is the same as CCP and packs atoms very tightly, filling about 74% of the volume. In both structures, each atom touches 12 near neighbors, so the coordination number is 12. About two-thirds of all metals crystallize in these closest-packed arrays.

Metallic Bonding and Crystal Packing

molecular comparison of gases, liquids, and solids

molecular-comparison-of-gases-liquids-and-solids

How Intermolecular Forces Shape States

Intermolecular forces and particle motion shape the states of matter: solids, liquids, and gases. In a solid, particles are tightly packed and often arranged in a regular pattern. A solid has a fixed shape, and its particles vibrate in place rather than flowing past one another. In a liquid, particles stay close together but do not have a regular arrangement. Liquids have no fixed shape, yet their particles move past each other while remaining in contact. That is why liquids can flow and take...

intermolecular vs intramolecular forces

How Forces Shape Liquids and Solids

Intermolecular forces are electrostatic attractions between molecules. Their strength depends on how far apart the molecules are. These forces matter most in solids and liquids, where particles are close together. In gases, they become important mainly under high pressure because the molecules are then closer together. Intermolecular forces help control key physical properties. They affect melting point, boiling point, density, and the enthalpies of fusion and vaporization. When a liquid is...

intermolecular forces

How Molecules Stick Together

Molecules and ions stick together through intermolecular forces. These are electrostatic attractions and repulsions caused by permanent, partial, or temporary charges. They act between ions, polar molecules, nonpolar molecules, and neutral atoms, and they vary in strength. Some intermolecular forces act between neutral particles and are called van der Waals forces. These include dipole-dipole forces, hydrogen bonding, and dispersion forces. Another important interaction is ion-dipole force,...

comparing intermolecular forces: melting point, boiling point, and miscibility

comparing-intermolecular-forces-melting-point-boiling-point

How Molecule Shape Affects Intermolecular Forces

Intermolecular forces are the attractions between molecules, and they help explain melting point, boiling point, and miscibility. Their strength depends on molar mass, molecular shape, and polarity. These forces shape the bulk behavior of substances that students see in chemistry labs. Dispersion forces are temporary attractions that occur in every substance, whether it is polar or nonpolar. They form when electrons are spread unevenly for a moment, creating a temporary dipole. Stronger...

surface tension, capillary action, and viscosity

surface-tension-capillary-action-and-viscosity

How Liquids Behave on Surfaces

Liquids behave in different ways on surfaces because of surface tension, adhesive forces, capillary action, and viscosity. These properties help explain why some liquids bead up, some spread out, and some flow more easily than others. Surface tension comes from cohesive forces, which are the attractions between identical molecules. Molecules inside a liquid are pulled equally in all directions, but molecules at the surface are pulled by fewer neighbors. That imbalance makes the liquid contract...

phase transitions

Melting, Boiling, and State Changes

Melting, boiling, and other state changes depend on how particles are arranged in a substance. In solids, particles pack closely in an ordered pattern. They vibrate in fixed positions but do not move past nearby particles. In liquids, particles are still close together, but they are randomly arranged and can move past one another. Solids and liquids are called condensed phases because their particles remain close and their intermolecular forces are relatively strong. Gases behave differently...

phase transitions: vaporization and condensation

phase-transitions-vaporization-and-condensation

Heat and Mass Transfer in Phase Changes

Heat and mass transfer drive phase changes in liquids. When a liquid is heated, its molecules move faster. Some surface molecules can then overcome intermolecular forces and escape into the gas phase. This process is called vaporization, and it happens at the surface of the liquid. Vaporization can also happen in a closed container. In that case, gas molecules cannot leave the container. The gas molecules move randomly and may collide with the surface of the condensed phase. Some of those...

vapor pressure

Vapor Pressure and Boiling Point

Vapor pressure and boiling point describe how a liquid behaves in a closed container and when it starts to boil. As a liquid vaporizes, some gas molecules stay above the liquid and can collide with the surface again. When those collisions turn gas back into liquid, the process is called condensation. If vaporization and condensation happen at the same rate, the liquid and vapor amounts stay constant. The vapor is then in equilibrium with the liquid. This is a dynamic equilibrium, which means...

clausius-clapeyron equation

Predicting Vapor Pressure with Temperature

Predicting vapor pressure with temperature is a key use of the Clausius-Clapeyron equation. It describes how the equilibrium between a liquid and its vapor changes as temperature changes. When temperature rises, the vapor pressure of the liquid also rises. The equation gives a quantitative link between vapor pressure, P, and temperature, T. It shows the rate at which vapor pressure increases for each unit increase in temperature. In this relationship, ΔHvap is the enthalpy of vaporization, R...

phase transitions: melting and freezing

Melting Point and Freezing Equilibrium

Melting point and freezing equilibrium describe how a crystalline solid changes between solid and liquid states. As a solid is heated, the average energy of its atoms, molecules, or ions increases, and the solid becomes hotter. At a certain point, the added energy becomes large enough to partly overcome the forces that hold the particles in fixed positions. The solid then starts melting, or changing into a liquid. During melting, the temperature stops rising even though heat is still being...

phase transitions: sublimation and deposition

phase-transitions-sublimation-and-deposition

Sublimation, Frost, and Enthalpy Changes

Sublimation is a phase change in which a solid turns directly into a gas. It skips the liquid state. Dry ice is a clear example at room temperature and standard pressure, where solid carbon dioxide seems to disappear as it sublimes without forming a liquid first. Some other solids also sublime under the right conditions. Snow and ice can sublimate below water’s melting point, and wind or the lower air pressure at high altitudes can speed up that slow process. Solid iodine does the same when...

heating and cooling curves

Phase Changes on Heating and Cooling Curves

Phase changes on heating and cooling curves show how a substance responds when it absorbs or releases heat while isolated from its surroundings. The graph tracks changes in temperature and state as the substance is heated or cooled. When a solid is heated, its temperature rises first. The amount of heat absorbed depends on the solid’s heat capacity, which is written as q = mcΔT. In this equation, m is the mass and c is the specific heat. This relationship applies when matter changes...

phase diagrams

Water and CO2 Phase Behavior

Water and carbon dioxide phase behavior is shown with phase diagrams that map pressure against temperature. These plots show where a substance exists as a solid, liquid, or gas. They also show how melting points, sublimation points, and boiling points change with pressure. In a phase diagram, the labeled regions for solid, liquid, and gas represent single phases. The lines or curves between them show two phases in equilibrium. The triple point is the one condition of pressure and temperature...

structures of solids

Crystalline and Amorphous Solid Patterns

Crystalline and amorphous solids differ in how their particles are arranged. In crystalline solids, atoms, ions, or molecules follow a definite repeating pattern. Metals and ionic compounds typically form these ordered solids. Because the same particles are held in place by the same forces or energy, crystalline solids have a precise melting temperature. Amorphous solids, also called non-crystalline solids or glasses, do not have an ordered internal structure. Their particles are arranged...

molecular and ionic 
solids

Molecular and Ionic Crystal Properties

Crystalline solids are grouped into four types: molecular, ionic, metallic, and covalent network solids. This page focuses on molecular and ionic crystals and how their particle type and attractions shape their properties. Molecular crystalline solids are made of neutral molecules. Common examples include ice, sucrose, and iodine. The molecules are held together by weak intermolecular forces, including London dispersion forces, dipole-dipole interactions, and hydrogen bonds. These forces...

ionic crystal structures

Ionic Crystal Packing and Ion Holes

Ionic crystal packing depends on ion size, charge, and the ratio of cations to anions. In these solids, the ions usually behave like charged spheres. Stable structures form when each ion is surrounded by as many oppositely charged ions as possible and the ions remain in contact. In simple ionic crystals, the larger anions usually form a close-packed array. The smaller cations fit into the spaces between them. Two common spaces are tetrahedral holes, which are surrounded by four anions, and...

Metallic solids are made of metal atoms arranged in crystals. Copper, aluminum, and iron are common examples. These solids are often described as metal nuclei spread through a sea of delocalized electrons. The atoms are held together by metallic bonding, and that bonding gives metals many of their useful bulk properties. Metallic solids are shiny, conduct heat and electricity well, and can be shaped without breaking. This shaping ability is called malleability, which means a metal can deform...

band theory

Band Gaps in Conductors and Insulators

Band gaps in solids help explain why some materials conduct electricity and others do not. When many atoms come together, their atomic orbitals combine into molecular orbitals with distinct energies. In a solid, so many orbitals are involved that these energies merge into continuous regions called bands. The energy difference between bands is called the band gap. Electrons in a solid must move across orbitals of different energies to travel through the material. In conductors, the valence...

network covalent solids

Why Network Covalent Solids Are So Hard

Network covalent solids are made of atoms joined in a three-dimensional network of covalent bonds. They appear in the crystal structures of nonmetals such as diamond, graphite, and silicon. Some covalent compounds also form this kind of structure, including silicon dioxide, which is found in sand, and silicon carbide, the abrasive used on sandpaper. Many minerals also contain networks of covalent bonds. These solids are hard to melt or break because their covalent bonds must be broken first.

x-ray crystallography

X-ray Diffraction in Crystal Structure

X-ray diffraction can reveal the unit cell size and atom arrangement in a crystal. This method is called X-ray crystallography. It uses diffraction, which is the bending of an electromagnetic wave when it meets a barrier about the same size as its wavelength. X-rays work well for this method because their wavelengths are about the same as the spacing between neighboring atoms in crystals. They are measured on the scale of a few angstroms. When monochromatic X-rays, or X-rays with one...

lattice centering and coordination number

lattice-centering-and-coordination-number

Crystal Packing in Cubic Unit Cells

Crystal packing in cubic unit cells shows how atoms are arranged in a repeating solid structure. A crystal is best described by its unit cell, which is the smallest repeating part of the lattice. Lattice points mark where atoms or ions are located. The full crystal is built when this unit cell repeats in three dimensions. The cubic lattice has three common unit cells. A simple cubic structure, also called a primitive cubic structure, can be pictured as identical spheres stacked in layers with...

key terms and concepts

Explore how metallic bonding, malleability, conductivity, and close-packed crystal structures give metals their key properties in high school chemistry.

Metallic Bonding and Crystal Packing | JoVE K12

chemistry

JoVE Core Chemistry brings general chemistry concepts to life using concise and easy-to-understand animated video lessons. Plus, in the scientist-in-action videos, we demonstrate classical and original research experiments performed in today's laboratories worldwide.

Physical Geology

Minerals

properties-minerals-hardness-luster-cleavage-streak-crystal-shape

Properties of minerals (hardness, luster, cleavage, streak, crystal shape, fluorescence, flammability, density/specific gravity, malleability)

Virginia  Standards of Learning

HIgh School

Physics

PH.9

PH.9.e

south-carolina-college-and-career-ready-science-standards

South Carolina College and Career Ready Science Standards

High School

Matter and Its Interactions

C-PS1-1

Georgia Standards of Excellence

SC2.

Molecular Structures in Material Design

The Atomic Theory of Matter

The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers. Leucippus and Democritus argued that all matter ...

The Bohr Model

Following the work of Ernest Rutherford and his colleagues in the early twentieth century, the picture of atoms consisting of tiny dense nuclei surrounded ...

The Quantum-Mechanical Model of an Atom

Shortly after de Broglie published his ideas that the electron in a hydrogen atom could be better thought of as being a circular standing wave instead of ...

Subatomic Particles

Dalton was only partially correct about the particles that make up matter. All matter is composed of atoms, and atoms are composed of three smaller ...

Quantum Numbers

It is said that the energy of an electron in an atom is quantized; that is, it can be equal only to certain specific values and can jump from one energy ...

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

The arrangement of electrons in the orbitals of an atom is called its electron configuration. We describe an electron configuration with a symbol that ...

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 ...

Thomson's e/m Experiment

In a beam of charged particles created by a heated cathode, the particles move at different speeds. However, many applications need a beam with uniform ...

Atomic structure
• Evolution of atomic models/theory
• Electrons
• Electron configurations

Content elaboration: Physical Science included properties and locations of protons, neutrons and electrons, atomic number, mass number, cations and anions, isotopes and the strong nuclear force which holds the nucleus together. In this course, the historical development of the atomic model and the positions of electrons are explored in greater detail. Atomic models are constructed to explain experimental evidence and make predictions. The changes in the atomic model over time exemplify how scientific knowledge changes as new evidence emerges and how technological advancements like electricity extend the boundaries of scientific knowledge. Thompson's study of electrical discharges in cathode-ray tubes led to the discovery of the electron and the development of the plum pudding model of the atom. Rutherford's experiment, in which he bombarded gold foil with a-particles, led to the discovery that most of the atom consists of empty space with a relatively small, positively charged nucleus. Bohr used data from atomic spectra to propose a planetary model of the atom in which electrons orbit the nucleus, like planets around the sun. Later, Schrödinger used the idea that electrons travel in waves to develop a model in which electrons travel randomly in regions of space called orbitals (quantum mechanical model). Based on the quantum mechanical model, it is not possible to predict exactly where electrons are located but there is a region of space surrounding the nucleus in which there is a high probability of finding an electron (electron cloud or orbital). Data from atomic spectra (emission and absorption) gives evidence that electrons can only exist at certain discrete energy levels and not at energies between these levels. Atoms are usually in the ground state where the electrons occupy orbitals with the lowest available energy. However, the atom can become excited when the electrons absorb a photon with the precise amount of energy (indicated by the frequency of the photon) to move to an orbital with higher energy. Any photon without this precise amount of energy will be ignored by the electron. The atom exists in the excited state for a very short amount of time. When an electron drops back down to the lower energy level, it emits a photon that has energy equal to the energy difference between the levels. The amount of energy is indicated by the frequency of the light that is given off and can be measured. Each element has a unique emission and absorption spectrum due to its unique electron configuration and specific electron energy jumps that are possible for that element. Being aware of the quantum mechanical model as the currently accepted model for the atom is important for science literacy as it explains and predicts subatomic interactions, but details should be reserved for more advanced study. Electron energy levels consist of sublevels (s, p, d and f), each with a characteristic number and shape of orbitals. Orbital diagrams and electron configuration can be constructed to show the location of the electrons in an atom using established rules. Valence electrons are responsible for most of the chemical properties of elements. In this course, electron configuration (extended and noble gas notation) and orbital diagrams can be shown for any element in the first three periods. Although the quantum mechanical model of the atom explains the most experimental evidence, other models can still be helpful. Thinking of atoms as indivisible spheres is useful in explaining many physical properties of substances, such as the state (solid, liquid or gas) of a substance at room temperature. Bohr's planetary model is useful to explain and predict periodic trends in the properties of elements.

Atomic structure

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 ...

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 ...

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

The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy ...

Electron Affinity

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



This process can be either ...

Alkali Metals

Group 1 elements are soft and shiny metallic solids. They are malleable, ductile, and good conductors of heat and electricity. The melting points of the ...

Halogens

Group 17 elements, known as halogens, are nonmetals. At room temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine a solid. ...

Noble Gases


The elements in group 18 are noble gases (helium, neon, argon, krypton, xenon, and radon). They earned the name &ldquo;noble&rdquo; because they were ...

Periodic Table
• Properties
• Trends

Content elaboration: In the Physical Science course, the concept that elements are placed in order of increasing atomic number in the periodic table such that elements with similar properties are placed in the same column is introduced. How the periodic table is divided into groups, families, periods, metals, nonmetals and metalloids is also included and will be revisited here. In this course, with more information about the electron configuration of elements, similarities in the configuration of the valence electrons for a particular group can be observed. The electron configuration of an atom can be determined from the position on the periodic table. The repeating pattern in the electron configuration for elements on the periodic table explains many of the trends in the properties observed. Atomic theory is used to describe and explain trends in properties across periods or down columns including atomic radii, ionic radii, first ionization energies, electronegativities and whether the element is a solid or gas at room temperature. Additional ionization energies, electron affinities and periodic properties of the transition elements, and the lanthanide and actinide series are reserved for more advanced study.

Periodic Table

Types of Chemical Bonds

Chemical bonding theories were pioneered by American chemist Gilbert&#xA0;N. Lewis. He developed a model called the Lewis model to explain the type and ...

Lewis Symbols and the Octet Rule

Chemical bonds are complex interactions between two or more atoms or ions, which reduce the potential energy of the molecule. Gilbert N. Lewis developed a ...

Ionic Bonding and Electron Transfer

Ions are atoms or molecules bearing an electrical charge. A cation (a positive ion) forms when a neutral atom loses one or more electrons from its valence ...

The Born-Haber Cycle

Lattice Energy&#xA0;

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a ...

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

Trends in Lattice Energy: Ion Size and Charge

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a compound is a measure ...

Covalent Bonding and Lewis Structures

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

Electronegativity

Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity.&#xA0;

Electronegativity values of ...

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

Bond Polarity, Dipole Moment, and Percent Ionic Character

Bond Polarity

The absolute value of the difference in electronegativity (&Delta;EN) of two bonded atoms provides a rough measure of the polarity to be ...

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

Lewis Structures of Molecular Compounds and Polyatomic Ions

To draw Lewis structures for complicated molecules and molecular ions, it is helpful to follow a step-by-step procedure as outlined:


	Determine the ...

Resonance

The Lewis structure of a nitrite anion (NO2&minus;) may actually be drawn in two different ways, distinguished by the locations of the N-O and N=O ...

Formal Charges

In some cases, there are seemingly more than one valid Lewis structures for molecules and polyatomic ions. The concept of formal charges can be used to ...

Exceptions to the Octet Rule

Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:


	...

Bond Energies and Bond Lengths

Stable molecules exist because covalent bonds hold the atoms together. The strength of a covalent bond is measured by the energy required to break it, ...

VSEPR Theory and the Basic Shapes

Overview of VSEPR Theory

Valence shell electron-pair repulsion theory (VSEPR theory) enables us to predict the molecular structure, including approximate ...

vsepr-theory-and-the-effect-of-lone-pairs

VSEPR Theory and the Effect of Lone Pairs

Effect of Lone Pairs of Electrons&nbsp;on Molecule Geometry

It is important to note that electron-pair geometry around a central atom is not the same ...

Predicting Molecular Geometry

VSEPR Theory for Determination of Electron Pair Geometries

The following procedure uses VSEPR theory to determine the electron pair geometries and the ...

Molecular Shape and Polarity

Dipole Moment of a Molecule

Polar covalent bonds connect two atoms with differing electronegativities, leaving one atom with a partial positive charge ...

Valence Bond Theory

Overview of Valence Bond Theory

Valence bond theory describes a covalent bond as the overlap of half-filled atomic orbitals (each containing a single ...

Hybridization of Atomic Orbitals I

The mathematical expression known as the wave function, &#x3C8;, contains information about each orbital and the wavelike properties of electrons in an ...

Hybridization of Atomic Orbitals II

sp3d and sp3d&nbsp;2 Hybridization

To describe the five bonding orbitals in a trigonal bipyramidal arrangement, we must use five of the valence shell ...

Molecular Orbital Theory I

Overview of Molecular Orbital Theory

Molecular orbital theory describes the distribution of electrons in molecules in the same way as the distribution of ...

Molecular Orbital Theory II

Molecular Orbital Energy Diagrams

The relative energy levels of atomic and molecular orbitals are typically shown in a molecular orbital diagram. For a ...

Chemical bonding
• lonic
• Polar/covalent

Content elaboration: Content in the Physical Science course included recognizing that atoms with unpaired electrons tend to form ionic and covalent bonds with other atoms, forming molecules, ionic lattices or network covalent structures. In this course, electron configuration, electronegativity values and energy considerations will be applied to bonding and the properties of materials with different types of bonding. Atoms of many elements are more stable when they are bonded to other atoms. In such cases, as atoms bond, energy is released to the surroundings, resulting in a system with lower energy. An atom's electron configuration, particularly the valence electrons, determines how an atom interacts with other atoms. Molecules, ionic lattices and network covalent structures have different, yet predictable, properties that depend on the identity of the elements and the types of bonds formed. Differences in electronegativity values can be used to predict where a bond fits on the continuum between ionic and covalent bonds. The polarity of a bond depends on the electronegativity difference and the distance between the atoms (bond length). Polar covalent bonds are introduced as an intermediary between ionic and pure covalent bonds. The concept of metallic bonding is also introduced to explain many of the properties of metals (e.g., conductivity). Since most compounds contain multiple bonds, a substance may contain more than one type of bond. Carbon atoms can bond together and with other atoms, especially hydrogen, oxygen, nitrogen and sulfur, to form chains, rings and branching networks that are present in a variety of important compounds, including synthetic polymers, fossil fuels and the large molecules essential to life.

Chemical bonding

Metallic Bonding and Crystal Packing

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