Name: how lone pairs change molecular shape
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 4 min 1 s
Channel: JoVE

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Ohio’s Learning Standards and Model Curriculum

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Chemistry

Structure And Properties Of Matter

Chemical bonding

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VSEPR Theory and the Effect of Lone Pairs

Predicting molecular geometry has more steps if the central atom has both bonding pairs and lone pairs of electrons.&nbsp;

The nitrogen atom in ammonia has four electron groups arranged in a tetrahedral fashion: three bonding pairs and one lone pair.&nbsp;

The lone pair of electrons occupies a larger space than the bonding pairs. This is because a lone pair is bound to only one nucleus, whereas a bonding electron group is shared by two nuclei.&nbsp;

The H-N-H bond angles are smaller than the expected tetrahedral angle of 109.5&deg; as observed in methane. This compression of the bond angle is attributed to the repulsive force exerted by a lone pair on the adjacent bonding electron groups.

The arrangement of electron pairs is called electron-pair geometry. The molecular geometry describes the arrangement of the atoms and differs from the electron-pair geometry. The electron pair geometry for ammonia is tetrahedral, whereas the molecular shape is trigonal pyramidal.

A water molecule also has four electron groups around the central atom. The electron pair geometry is also tetrahedral with two bonding electron groups and two lone pairs.

Lone pair-lone pair repulsions are greater than lone pair-bonding pa

vsepr theory and the effect of lone pairs

<h4>Effect of Lone Pairs of Electrons&nbsp;on Molecule Geometry</h4>

It is important to note that electron-pair geometry around a central atom is not the same thing as its molecular structure. Molecular structure describes the location of the atoms, not the electrons. The geometry that includes all electron pairs is the electron-pair geometry. The electron-pair geometries describe all regions where electrons are located, bonds as well as lone pairs. The structure that includes only the placement of the atoms in the molecule is called the molecular structure. The electron-pair geometries will be the same as the molecular structures when there are no lone electron pairs around the central atom, but they will be different when there are lone pairs present on the central atom.

For example, the methane molecule, CH4, which is the major component of natural gas, has four bonding pairs of electrons around the central carbon atom; the electron-pair geometry is tetrahedral, as is the molecular structure. On the other hand, the ammonia molecule, NH3, also has four electron pairs associated with the nitrogen atom and thus has a tetrahedral electron-pair geometry. One of these regions, however, is a lone pair, which is not included in the molecular structure, and this lone pair influences the shape of the molecule.

<h4>Angle Distortions Based on VSEPR Theory</h4>

The small distortions from the ideal angles can result from differences in repulsion between various regions of electron density. VSEPR theory predicts these distortions by establishing an order of repulsions and an order of the amount of space occupied by different kinds of electron pairs. The order of electron-pair repulsions from greatest to least repulsion is:

lone pair-lone pair &gt; lone pair-bonding pair &gt; bonding pair-bonding pair

This order of repulsions determines the amount of space occupied by different regions of electrons. A lone pair of electrons occupies a larger region of space than the electrons in a triple bond; in turn, electrons in a triple bond occupy more space than those in a double bond, and so on. The order of sizes from largest to smallest is:

lone pair &gt; triple bond &gt; double bond &gt; single bond

In the ammonia molecule, the three hydrogen atoms attached to the central nitrogen are not arranged in a flat, trigonal planar molecular structure, but rather in a three-dimensional trigonal pyramid with the nitrogen atom at the apex and the three hydrogen atoms forming the base. The ideal bond angles in a trigonal pyramid are based on the tetrahedral electron pair geometry. Again, there are slight deviations from the ideal because lone pairs occupy larger regions of space than do bonding electrons. The H&ndash;N&ndash;H bond angles in NH3 are slightly smaller than the 109.5&deg; angle in a regular tetrahedron because the lone pair-bonding pair repulsion is greater than the bonding pair-bonding pair repulsion.

According to VSEPR theory, the terminal atom locations are equivalent within the linear, trigonal planar, and tetrahedral electron-pair geometries. It does not matter which X is replaced with a lone pair because the molecules can be rotated to convert positions. For trigonal bipyramidal electron-pair geometries, however, there are two distinct X positions, an axial position (if we hold a model of a trigonal bipyramid by the two axial positions, we have an axis around which we can rotate the model) and an equatorial position (three positions form an equator around the middle of the molecule). The axial position is surrounded by bond angles of 90&deg;, whereas the equatorial position has more space available because of the 120&deg; bond angles. In a trigonal bipyramidal electron-pair geometry, lone pairs always occupy equatorial positions because these more spacious positions can more easily accommodate the larger lone pairs.

When a central atom has two lone electron pairs and four bonding regions, we have an octahedral electron-pair geometry. The two lone pairs are on opposite sides of the octahedron (180&deg; apart), giving a square planar molecular structure that minimizes lone pair-lone pair repulsions.

This text has been adapted from <a href="https://openstax.org/books/chemistry-2e/pages/7-6-molecular-structure-and-polarity">Openstax, Chemistry 2e, Section 7.6: Molecular Structure and Polarity.</a>

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 ...

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KEY TERMS AND CONCEPTS

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, NH3, also has four electron pairs around nitrogen, so its electron-pair geometry is tetrahedral too, but one region is a lone pair. That lone pair is not part of the molecular structure, and it changes the shape.

VSEPR theory explains why bond angles can shift away from ideal values. It predicts the order of repulsion as lone pair-lone pair greater than lone pair-bonding pair greater than bonding pair-bonding pair. It also shows how much space each region takes up. A lone pair occupies more space than a triple bond, a triple bond occupies more space than a double bond, and a double bond occupies more space than a single bond.

Ammonia shows this effect clearly. The three hydrogen atoms around nitrogen form a three-dimensional trigonal pyramid, not a flat trigonal planar shape. Its bond angles come from a tetrahedral electron-pair geometry, but the H-N-H angles are slightly smaller than 109.5° because lone pair-bonding pair repulsion is stronger than bonding pair-bonding pair repulsion.

In linear, trigonal planar, and tetrahedral electron-pair geometries, the terminal atom positions are equivalent. It does not matter which X position is replaced by a lone pair because the molecule can be rotated. Trigonal bipyramidal electron-pair geometry is different. It has axial positions and equatorial positions. Axial positions have 90° bond angles, while equatorial positions have more space because the bond angles are 120°.

For that reason, lone pairs in a trigonal bipyramidal arrangement always go to equatorial positions. These positions can better hold the larger lone pairs. When a central atom has two lone pairs and four bonding regions, the electron-pair geometry is octahedral. The two lone pairs sit opposite each other, 180° apart, which gives a square planar molecular structure and helps reduce lone pair-lone pair repulsion.

How Lone Pairs Change Molecular Shape

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...

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,...

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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...

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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...

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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...

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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.

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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...

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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...

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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

Explore how lone pairs affect electron-pair geometry, molecular shape, and bond angles using examples like methane, ammonia, trigonal bipyramidal, and square planar structures.

How Lone Pairs Change Molecular Shape | JoVE K12

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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

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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 ...

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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

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Periodic Classification of the Elements

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Atomic Radii and Effective Nuclear Charge

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Ionic Radii

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Ionization Energy

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Electron Affinity

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Alkali Metals

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Halogens

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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
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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

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The Born-Haber Cycle

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Trends in Lattice Energy: Ion Size and Charge

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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 ...

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Electronegativity values of ...

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Bond Polarity, Dipole Moment, and Percent Ionic Character

Bond Polarity

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Lewis Structures of Molecular Compounds and Polyatomic Ions

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Resonance

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Formal Charges

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Exceptions to the Octet Rule

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	...

Bond Energies and Bond Lengths

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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 ...

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

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Hybridization of Atomic Orbitals I

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Hybridization of Atomic Orbitals II

sp3d and sp3d&nbsp;2 Hybridization

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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

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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
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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

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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
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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.

Measurement: Standard Units

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Measurement: Derived Units

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Uncertainty in Measurement: Accuracy and Precision

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Uncertainty in Measurement: Reading Instruments

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Uncertainty in Measurement: Significant Figures

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Dimensional Analysis

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Elements: Chemical Symbols and Isotopes

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Atomic Mass

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Formula Mass and Mole Concepts of Compounds

Formula Mass of Covalent Compounds

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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

How Lone Pairs Change Molecular Shape

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