Name: atomic orbital hybridization and bond shapes
Uploaded: 2020-09-03T13:38:29.000Z
Duration: 3 min 24 s
Channel: JoVE

K12 Standards

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Ohio’s Learning Standards and Model Curriculum

9-12

Chemistry

Structure And Properties Of Matter

Chemical bonding

Hybridization of Atomic Orbitals I

s and p orbitals overlapping to form covalent bonds cannot yield the various molecular shapes in the VSEPR model. Valence bond theory helps to explain this molecular geometry through the hybridization, or mixing, of atomic orbitals.&nbsp;

Some atomic orbitals involved in bonding recombine to form new orbitals whose shapes are a hybrid of the originals. The initial number of atomic orbitals and the number of hybrid orbitals generated is always the same.&nbsp;

Beryllium fluoride is a linear molecule. The beryllium atom has two valence electrons found in its 2s orbital in the ground state. &nbsp;

The s orbital mixes with one of the empty p orbitals to create orbitals that contain unpaired electrons available for bonding. This leaves two unhybridized p orbitals and produces two sp hybrid orbitals, named for the original atomic orbitals.

The hybrid orbitals have a different shape from their constituent atomic orbitals with one lobe that is significantly larger than the other. Thus, the electron probability density is highly concentrated in a directional lobe, which leads to a more effective overlap with the orbitals of other atoms. For clarity, these orbitals are often shown without the minor lobes.&nbsp;

hybridization of atomic orbitals i

The mathematical expression known as the wave function, &psi;, contains information about each orbital and the wavelike properties of electrons in an isolated atom. When atoms are bound together in a molecule, the wave functions combine to produce new mathematical descriptions that have different shapes. This process of combining the wave functions for atomic orbitals is called hybridization and is mathematically accomplished by the linear combination of atomic orbitals. The new orbitals that result are called hybrid orbitals.

<h4>Understanding Atomic Orbital Hybridization</h4>

The following ideas are important in understanding hybridization:

<ol>
	<li>Hybrid orbitals do not exist in isolated atoms. They are formed only in covalently bonded atoms.</li>
	<li>Hybrid orbitals have shapes and orientations that are very different from those of the atomic orbitals in isolated atoms.</li>
	<li>A set of hybrid orbitals is generated by combining atomic orbitals. The number of hybrid orbitals in a set is equal to the number of atomic orbitals that were combined to produce the set.</li>
	<li>All orbitals in a set of hybrid orbitals are equivalent in shape and energy.</li>
	<li>The type of hybrid orbitals formed in a bonded atom depends on its electron-pair geometry, as predicted by the VSEPR theory.</li>
	<li>Hybrid orbitals overlap to form &sigma; bonds. Unhybridized orbitals overlap to form &pi; bonds.</li>
</ol>

In the following sections, we shall discuss the common types of hybrid orbitals.

<h4>sp Hybridization</h4>

The beryllium atom in a gaseous BeCl2 molecule is an example of a central atom with no lone pairs of electrons in a linear arrangement of three atoms. There are two regions of valence electron density in the BeCl2 molecule that correspond to the two covalent Be&ndash;Cl bonds. To accommodate these two electron domains, two of the Be atom&rsquo;s four valence orbitals will mix to yield two hybrid orbitals. This hybridization process involves the mixing of the valence s orbital with one of the valence p orbitals to yield two equivalent sp hybrid orbitals that are oriented in a linear geometry. The set of sp orbitals appears similar in shape to the original p orbital, but there is an important difference. The number of atomic orbitals combined always equals the number of hybrid orbitals formed. The p orbital is one orbital that can hold up to two electrons. The sp set is two equivalent orbitals that point 180&deg; from each other. The two electrons that were originally in the s orbital are now distributed to the two sp orbitals, which are half-filled. In gaseous BeCl2, these half-filled hybrid orbitals will overlap with orbitals from the chlorine atoms to form two identical &sigma; bonds.

When atomic orbitals hybridize, the valence electrons occupy the newly created orbitals. The Be atom had two valence electrons, so each of the sp orbitals gets one of these electrons. Each of these electrons pairs up with the unpaired electron on a chlorine atom when a hybrid orbital and a chlorine orbital overlap during the formation of the Be&ndash;Cl bonds.

Any central atom surrounded by just two regions of valence electron density in a molecule will exhibit sp hybridization. Other examples include the mercury atom in the linear HgCl2 molecule, the zinc atom in Zn(CH3)2, which contains a linear C&ndash;Zn&ndash;C arrangement, and the carbon atoms in HCCH and CO2.

<h4>sp2 Hybridization</h4>

The valence orbitals of a central atom surrounded by three regions of electron density consist of a set of three sp2 hybrid orbitals and one unhybridized p orbital. This arrangement results from sp2 hybridization, the mixing of one s orbital and two p orbitals to produce three identical hybrid orbitals oriented in a trigonal planar geometry.

The observed structure of the borane molecule, BH3, suggests sp2 hybridization for boron in this compound. The molecule is trigonal planar, and the boron atom is involved in three bonds to hydrogen atoms. The three valence electrons of the boron atom in the three sp2 hybrid orbitals are redistributed, and each boron electron pairs with a hydrogen electron when B&ndash;H bonds form.

Any central atom surrounded by three regions of electron density will exhibit sp2 hybridization. This includes molecules with a lone pair on the central atom, such as ClNO, or molecules with two single bonds and a double bond connected to the central atom, as in formaldehyde, CH2O, and ethene, H2CCH2.

<h4>sp3 Hybridization</h4>

The valence orbitals of an atom surrounded by a tetrahedral arrangement of bonding pairs and lone pairs consist of a set of four sp3 hybrid orbitals. The hybrids result from the mixing of one s orbital and all three p orbitals, which&nbsp;produces four identical sp3 hybrid orbitals. Each of these hybrid orbitals points toward a different corner of a tetrahedron.

A molecule of methane, CH4, consists of a carbon atom surrounded by four hydrogen atoms at the corners of a tetrahedron. The carbon atom in methane exhibits sp3 hybridization. The four valence electrons of the carbon atom are distributed equally in the hybrid orbitals, and each carbon electron pairs with a hydrogen electron when the C&ndash;H bonds form.

In a methane molecule, the 1s orbital of each of the four hydrogen atoms overlaps with one of the four sp3 orbitals of the carbon atom to form a sigma (&sigma;) bond. This results in the formation of four strong, equivalent covalent bonds between the carbon atom and each of the hydrogen atoms to produce the methane molecule, CH4.

An sp3 hybrid orbital can also hold a lone pair of electrons. For example, the nitrogen atom in ammonia is surrounded by three bonding pairs and a lone pair of electrons directed to the four corners of a tetrahedron. The nitrogen atom is sp3 hybridized with one hybrid orbital occupied by the lone pair.

The molecular structure of water is consistent with a tetrahedral arrangement of two lone pairs and two bonding pairs of electrons. Thus we say that the oxygen atom is sp3 hybridized, with two of the hybrid orbitals occupied by lone pairs and two by bonding pairs. Since lone pairs occupy more space than bonding pairs, structures that contain lone pairs have bond angles slightly distorted from the ideal. Perfect tetrahedral structures have angles of 109.5&deg;, but the observed angles in ammonia (107.3&deg;) and water (104.5&deg;) are slightly smaller. Other examples of sp3 hybridization include CCl4, PCl3, and NCl3.

This text has been adapted from <a href="https://openstax.org/books/chemistry-2e/pages/8-2-hybrid-atomic-orbitals">Openstax, Chemistry 2e, Section 8.2: Hybrid Atomic Orbitals</a>.

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

Education

JoVE Core

Chemical Bonding: Molecular Geometry and Bonding Theories

chemical-bonding-molecular-geometry-and-bonding-theories

KEY TERMS AND CONCEPTS

atomic orbital hybridization and bond shapes

Atomic orbital hybridization explains how electron orbitals change shape when atoms form covalent bonds. The wave function, ψ, describes the behavior of electrons in an isolated atom. When atoms bond, their wave functions combine through the linear combination of atomic orbitals, and the new orbitals are called hybrid orbitals.

Hybrid orbitals do not exist in isolated atoms. They form only in covalently bonded atoms. Their shapes and directions differ from those of the original atomic orbitals. A set of hybrid orbitals contains the same number of orbitals as the atomic orbitals used to make it. The orbitals in one set are equivalent in shape and energy. The kind of hybrid orbitals formed depends on the electron-pair geometry predicted by VSEPR theory. Hybrid orbitals overlap to make σ bonds, while unhybridized orbitals can overlap to make π bonds.

The simplest case is sp hybridization. In gaseous BeCl2, the beryllium atom has two regions of valence electron density and a linear arrangement of three atoms. To fit this geometry, one s orbital and one p orbital mix to form two equivalent sp hybrid orbitals. These orbitals point 180° apart. The two valence electrons of beryllium are placed into the two sp orbitals, and each half-filled orbital overlaps with a chlorine orbital to form a σ bond. The same pattern appears in other linear species, including HgCl2, Zn(CH3)2 with a linear C–Zn–C arrangement, and the carbon atoms in HCCH and CO2.

A central atom with three regions of electron density often shows sp2 hybridization. In this case, one s orbital and two p orbitals combine to make three identical sp2 hybrid orbitals in a trigonal planar arrangement. Borane, BH3, fits this model because boron forms three bonds to hydrogen atoms in a trigonal planar shape. The three valence electrons of boron occupy the three sp2 orbitals, and each one pairs with a hydrogen electron when B–H bonds form. At the same time, one unhybridized p orbital remains. This pattern also appears in atoms with a lone pair, such as in ClNO, and in molecules with a double bond on the central atom, such as formaldehyde, CH2O, and ethene, H2CCH2.

Atoms with a tetrahedral arrangement of bonding pairs and lone pairs usually exhibit sp3 hybridization. Here, one s orbital and all three p orbitals mix to form four equivalent sp3 hybrid orbitals. Each orbital points toward a different corner of a tetrahedron. Methane, CH4, is a clear example. Its carbon atom is surrounded by four hydrogen atoms, and each C–H bond forms when a hydrogen 1s orbital overlaps with one sp3 orbital on carbon. The result is four strong, equivalent covalent bonds.

An sp3 hybrid orbital can also hold a lone pair. In ammonia, the nitrogen atom has three bonding pairs and one lone pair arranged toward the corners of a tetrahedron, so nitrogen is sp3 hybridized. Water shows a similar pattern. Its oxygen atom has two lone pairs and two bonding pairs, which also matches sp3 hybridization. Lone pairs take up more space than bonding pairs, so the bond angles in these molecules are slightly smaller than the ideal 109.5°. The observed angles are 107.3° in ammonia and 104.5° in water. Other examples of sp3 hybridization include CCl4, PCl3, and NCl3.

Atomic Orbital Hybridization and Bond Shapes

atomic-orbital-hybridization-and-bond-shapes

vsepr theory and the basic shapes

Predicting Molecular Shape with VSEPR

VSEPR theory helps predict molecular shape and approximate bond angles around a central atom. VSEPR stands for valence shell electron-pair repulsion. The model uses the Lewis structure to count bonding pairs and lone pairs in the valence shell of the central atom. The idea behind the theory is simple. Electron pairs in the valence shell arrange themselves to reduce repulsion by staying as far apart as possible. Bonding pairs are found mainly between bonded atoms, while lone pairs stay on the...

vsepr theory and the effect of lone pairs

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

How Lone Pairs Change Molecular Shape

Lone pairs of electrons can change molecular shape and bond angles. Electron-pair geometry includes every region of electron density around a central atom, while molecular structure shows only where the atoms are located. These two ideas match when the central atom has no lone pairs, but they differ when lone pairs are present. Methane, CH4, is a simple example. It has four bonding pairs around carbon, so both the electron-pair geometry and the molecular structure are tetrahedral. Ammonia,...

predicting molecular geometry

VSEPR Rules for Molecular Shapes

VSEPR theory helps predict electron-pair geometry and molecular shape from a Lewis structure. First, write the Lewis structure of the molecule or polyatomic ion. Then count the electron groups around the central atom. A single, double, or triple bond each counts as one region of electron density. Next, match the number of electron groups to the electron-pair geometry. The main geometries are linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral. Then use the number of lone...

molecular shape and polarity

Bond Dipoles and Molecular Polarity

Bond dipoles and molecular polarity explain why some molecules have a charged end and others do not. In a polar covalent bond, atoms with different electronegativities share electrons unevenly. The more electronegative atom pulls electron density toward itself and becomes partially negative, while the other atom becomes partially positive. This separation of charge creates a bond dipole moment. The size of a bond dipole moment is written with the Greek letter mu, or μ. It depends on the size...

valence bond theory

Covalent Bond Strength and Orbital Overlap

Covalent bond strength depends on how atomic orbitals overlap. In valence bond theory, a covalent bond forms when two half-filled orbitals, each with one electron, overlap and the two electrons pair up. The shared electron pair is then attracted to both positively charged nuclei, which holds the atoms together. The amount of overlap affects both bond strength and energy. When atoms are far apart, there is no overlap, so the energy is set to zero by convention. As the atoms move closer, their...

Atomic orbital hybridization explains how electron orbitals change shape when atoms form covalent bonds. The wave function, ψ, describes the behavior of electrons in an isolated atom. When atoms bond, their wave functions combine through the linear combination of atomic orbitals, and the new orbitals are called hybrid orbitals. Hybrid orbitals do not exist in isolated atoms. They form only in covalently bonded atoms. Their shapes and directions differ from those of the original atomic orbitals.

hybridization of atomic orbitals ii

Trigonal Bipyramidal and Octahedral Hybridization

Trigonal bipyramidal and octahedral hybridization explain how atoms form five or six bonding orbitals. These models use valence shell atomic orbitals to match the shape of the molecule. They apply to atoms that have d orbitals in their valence subshells, so they do not apply to atoms in the first or second period. For five bonding orbitals, an atom uses one s orbital, three p orbitals, and one d orbital. This creates five sp3d hybrid orbitals. In phosphorus pentachloride, PCl5, the five P-Cl...

molecular orbital theory i

Bonding and Antibonding Molecular Orbitals

Molecular orbital theory explains how electrons are arranged in molecules. It uses quantum mechanics and wave functions, written as Ψ, to describe electron behavior in a molecule. Like electrons in isolated atoms, electrons in molecules occupy discrete, quantized energy levels. A molecular orbital is the region where a valence electron is likely to be found. It is full when it contains two electrons with opposite spin. Molecular orbitals can be built by combining atomic orbitals through the...

molecular orbital theory ii

MO Diagrams, Bond Order, and Mixing

Molecular orbital diagrams show how atomic orbitals combine to make bonding and antibonding orbitals in diatomic molecules. In these diagrams, the atomic orbitals of one atom are placed on the left and the atomic orbitals of the other atom are placed on the right. The molecular orbitals appear in the center. Each horizontal line stands for one orbital, and each orbital can hold two electrons. Dashed lines show which atomic orbitals combine to form each molecular orbital. When atomic orbitals...

key terms and concepts

Learn how atomic orbitals hybridize into sp, sp2, and sp3 orbitals, and how these shapes help explain linear, trigonal planar, and tetrahedral bonding.

Atomic Orbital Hybridization and Bond Shapes | JoVE K12

chemistry

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Virginia  Standards of Learning

HIgh School

CH.3

CH.3.e

Physics

PH.9

PH.9.e

Alabama Course of Study Science

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Motion and Stability: Forces and Interactions

9-12.C.2.2.d

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New York State Science Learning Standards

High School

Physical Science

HS. Structure and Properties of Matter

HS-PS1-3.

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 ...

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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 ...

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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 ...

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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

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 ...

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 ...

Bonding in Metals

Metallic bonds are formed between two metal atoms. A simplified model to describe metallic bonding has been developed by Paul Dr&#xFC;de called the ...

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 Formulas

A chemical formula presents information about the proportions of atoms constituting a particular chemical compound or molecule, mainly using symbols of ...

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Ionic Compounds: Formulas and Nomenclature

An element composed of atoms that readily lose electrons (a metal) can react with an element composed of atoms that readily gain electrons (a nonmetal) to ...

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Molecular Compounds: Formulas and Nomenclature

Molecular compounds or covalent compounds result when atoms share electrons to form covalent bonds. Since there is no electron transfer, molecular ...

Organic Compounds

All living things are formed mostly of carbon compounds called organic compounds. The category of organic compounds includes both natural and synthetic ...

Classification of Elements and Compounds

Pure substances consist of only one type of matter. A pure substance can be an element or a compound. An element consists of only one type of atom, while ...

Molecular Models

Physical models representing molecular architectures of chemical compounds play essential roles in understanding chemistry. The use of molecular models ...

Representing compounds
• Formula writing
• Nomenclature
• Models and shapes (Lewis structures, ball and stick, molecular geometries)

Content elaboration: Using the periodic table, formulas of ionic compounds containing specific elements can be predicted. This can include ionic compounds made up of elements from groups 1, 2, 17, hydrogen, oxygen and polyatomic ions (given the formula and charge of the polyatomic ion). Given the formula, a compound can be named using conventional systems that include Greek prefixes and Roman numerals where appropriate. Given the name of an ionic or covalent substance, formulas can be written. Many different models can be used to represent compounds including chemical formulas, Lewis structures, and ball and stick models. These models can be used to visualize atoms and molecules and to predict the properties of substances. Each type of representation provides unique information about the compound. Different representations are better suited for particular substances. Lewis structures can be drawn to represent covalent compounds using a simple set of rules and can be combined with valence shell electron pair repulsion (VSEPR) theory to predict the three-dimensional electron pair and molecular geometry of compounds. Lewis structures and molecular geometries will only be constructed for the following combination of elements: hydrogen, carbon, nitrogen, oxygen, phosphorus, sulfur and the halogens.

Representing compounds

Measurement: Standard Units

Every measurement provides three kinds of information: the size or magnitude of the measurement (a number), a standard of comparison for the measurement ...

Measurement: Derived Units

The International System of Units or SI system, by international agreement, has fixed measurement units for seven fundamental properties: length, mass, ...

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Uncertainty in Measurement: Accuracy and Precision

Scientists typically make repeated measurements of a quantity to ensure the quality of their findings and to evaluate both the precision and the accuracy ...

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Uncertainty in Measurement: Reading Instruments

Counting is the type of measurement that is free from uncertainty, provided the number of objects being counted does not change during the process. Such ...

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Uncertainty in Measurement: Significant Figures

All the digits in a measurement, including the uncertain last digit, are called significant figures or significant digits. Note that zero may be a ...

Dimensional Analysis

Dimensional analysis, also known as the factor label method, is a versatile approach for mathematical operations. The main principle behind this approach ...

Elements: Chemical Symbols and Isotopes

A chemical symbol is an abbreviation used to indicate an element or an atom of an element. For example, the symbol for mercury is Hg. The same symbol is ...

Atomic Mass

Atoms &mdash; and the protons, neutrons, and electrons that compose them &mdash; are extremely small. For example, a carbon atom weighs less than 2 ...

Molar Mass

The identity of a substance is defined not only by the types of atoms or ions it contains but by the quantity of each type of atom or ion. For example, ...

Quantifying matter

Content elaboration: In earlier grades, properties of materials were quantified with measurements that were always associated with some error. In this course, scientific protocols for quantifying the properties of matter accurately and precisely are studied. Using the International System of Units (SI), significant digits or figures, scientific notation, error analysis and dimensional analysis are vital to scientific communication. There are three domains of magnitude in size and time: the macroscopic (human) domain, the cosmic domain and the submicroscopic (atomic and subatomic) domain. Measurements in the cosmic domain and submicroscopic domains require complex instruments and/or procedures. Matter can be quantified in a way that macroscopic properties such as mass can reflect the number of particles present. Elemental samples are a mixture of several isotopes with different masses. The atomic mass of an element is calculated given the mass and relative abundance of each isotope of the element as it exists in nature. Because the mass of an atom is very small, the mole is used to translate between the atomic and macroscopic levels. A mole is equal to the number of atoms in exactly 12 grams of the isotope carbon-12. The mass of one mole of a substance is equal to its molar mass in grams. The molar mass for a substance can be used in conjunction with Avogadro's number and the density of a substance to convert between mass, moles, volume and number of particles of a sample.

Quantifying matter

Atomic Orbital Hybridization and Bond Shapes

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