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Types of Bonds and Orbitals

Types of Bonds and Orbitals

Types of Bonds and Orbitals Terms

Angular Momentum Quantum Number

The second quantum number used to determine the type of sublevel or subshell that a particular electron is occupying.

Anion

A negatively charged ion.

Atom

Aufbau Prinicple

States that electrons in an atom must fill the lowest energy level first. Poor little electrons…so many rules and regulations…fill this way and fill that way. Why can't an electron ever fill the way it wants to fill?

Cation

A positively charged ion.

Coulomb's Law

Explains the electrostatic interaction between charged particles. If the two charges have the same sign, the electrostatic force between them is repulsive. If they have different signs, the force between them is attractive. Charges close together mean stronger bonds, while charges far apart mean weaker bonds. It ain't fiction just a natural fact.

Covalent Bonding

A bond in which one or more electron pairs are shared between two atoms, typically between two nonmetals. Caring is sharing.

Diamagnetism

Repulsion of a molecule from a magnetic field due to the presence of all paired electrons.

Dipole Moment

A measure of the polarity in a chemical bond or molecule, equal to the product of one charge and the distance between the charges.

Electron

A negatively charged subatomic particle found in the electron "cloud" or volume of space around the nucleus. Electrons are thought to have essentially no mass, but actually their mass is just really super duper small: approximately 1/1836 that of a proton.

Electrostatic Attraction

The attraction between atoms that have opposite charges, and hold the atoms together in ionic bonds.

Electronegativity

Electron Configuration

A notation that shows details about the likely locations of the electrons in a given atom.

Electron Pool Theory

This theory states that each atom present in a metallic crystal loses all of its valence electrons. As a result a "pool" of electrons is formed. It is believed that positively charged metal ions are held together by this "pool" of electrons. Do your electrons want to go for a swim?

Hund's Rule

States that the most stable arrangement for electrons in subshells is the one with the greatest number of parallel spins. Therefore electrons are added to the orbitals in the following way: filling them all half first before any pairing occurs in each subshell.

Hybrid Orbitals

Atomic orbitals formed as a result of mixing the atomic orbitals of the atoms involved in the covalent bond. The type of hybrid orbital formed is dependent on the types of orbitals mixed (for example s, p, and d orbitals). The number of hybrid orbitals formed is the same as the number of atomic orbitals mixed.

Hydrogen Bonding

Considered not a true chemical bond, but it occurs when a hydrogen atom has an electromagnetic attractive interaction with an electronegative atom. The hydrogen bond is typically stronger than Van der Waals attractions but weaker than covalent or ionic bonds.

Ion

An atom or molecule that has obtained a charge by either gaining or losing one or more electrons. This means that the proton number does not match the number of electrons giving the atom either a net negative or positive charge.

Ionic Bond

A bond formed from the electrostatic attraction between ions with different charges. This typically consists of a metal (cation) and a nonmetal (anion). A metal loses electrons to form a cation and some non-metal gains those electrons to form an anion. Opposites attract.

Isoelectronic

Describes molecular entities that share the same number of valence electrons or have the same electron configuration.

Isotope

Atoms of the same element with different masses. Since neutrons have no charge associated with them they don't affect the atom's identity. The atomic number or the number of protons defines that characteristic. However, since neutrons have a mass associated with them they do affect the element's mass number. That's why the same element can have different masses. Carbon-12 says to Carbon-11, "Looking good...what's your secret?" Carbon-11 replies, "Just got a rid of a pesky neutron that was weighing me down."

Lewis Electron Dot Structure

A notation that shows information about an atom's outermost shell, depicting the element and its valence electrons as dots around the element. Lewis sure loved his dots.

Magnetic Quantum Number

The third quantum number used to determine the spatial orientation of the orbital that the electron is occupying.

Mass Number

Adding the total number of protons and neutrons in the nucleus of an atom gives you the atom's mass number. Since the mass of an electron is negligible it doesn't contribute to the overall mass number of an atom.

Metallic Bonding

A bond between metals where the valence electrons are donated to a vast electron "pool", so that the valence electrons are free to move throughout the entire metallic solid.

Molecular Geometry

A type of geometry that defines the shape of a molecule, in which the nonbonding electrons become "invisible" and only the geometry of the atomic nuclei are considered.

Molecular Orbital (MO) Theory

A model that represents the bonding that takes place in covalent compounds. This theory states that atomic orbitals on individual atoms will combine to form molecular orbitals that encompass the entire molecule.

Molecule

A group of atoms that are held together by covalent bonds.

Neutron

A subatomic particle with no charge that resides in the nucleus of an atom. The mass of a neutron is essentially the same as the mass of a proton. Nothing but a chargeless copycat.

Non-polar Covalent Bond

A bond formed when atoms share one or more pairs of electrons relatively equally. Each of the atoms attract the bonding pair of electrons equally.

Octet Rule

States that atoms will lose, gain, or share electrons in order to achieve a filled valence shell, to complete their octet. Anybody up for a game of Crazy Eights?

Orbital Notation

The notation used to designate the electrons and orbitals of a given element. Basically just another way of expressing the electron configuration of an atom in picture perfect way.

Paramagnetism

Attraction of a molecule to a magnetic field due to the presence of unpaired electrons.

Pi (π) Bonds

Overlapping of orbitals, where two lobes of one atomic orbital overlap two lobes of another atomic orbital. These overlapping orbitals share a nodal plane that passes through the two nuclei of the atoms involved in the covalent bond. A double bond always consists of one sigma bond and one pi bond. Mmm Pie.

Polar Covalent Bond

A case where the two atoms involved in the covalent bond are not sharing the electrons equally. The bonded electrons are pulled toward the atom with the greater attraction towards electrons. We've all been there.

Principle Quantum Number

The first quantum number used to determine the energy level or shell that a particular electron is occupying.

Proton

A positively charged atomic particle. The number of protons in an atom determines the identity of the atom; for example, carbon atoms always have 6 protons.

Quantum Numbers

Numbers given to a particular electron, each electron is given four. The numbers are used as a coordinate system to help identify and find the particular electron you want to study.

Resonance Structures

The average of all the possible Lewis structures of a given compound. Some molecules can have more than one possible written structure. No individual structure truly represents the actual structure. That would require an average of all the possible structures. The Resonance Theory is used to describe this situation and often uses this two-headed arrow notation (← → ) between each of the possible structures drawn.

Sigma (σ) Bonds

Overlap of orbitals occurs on a line between two atoms involved in the covalent bond.

Spin Quantum Number

The fourth quantum number used to determine the "spin" direction of the electron you are studying in a particular orbital.

Valence Electron

An electron found in the outermost shell of an atom...we're talking the highest energy level of the atom. These guys are definitely not afraid of heights.

Valence Bond Theory

Is one way in which the molecular geometry of molecules may be determined. It describes covalent bonding as a mixing of atomic orbitals to form a new kind of orbital…hybrid orbitals.

Van Der Waals Attractions

A weak attractive force between atoms or nonpolar molecules caused by a temporary change in dipole moment arising from a brief shift of orbital electrons to one side of one atom or molecule. This creates a similar shift in adjacent atoms or molecules.

Valence-Shell Electron-Pair Repulsion (VSEPR) Theory

A theory that predicts molecular geometry. The theory states that electron pairs around the central atom will try to stay as far away from each other as possible in order to minimize repulsive forces. Honestly, who wouldn't want to stay away from things that repulse them?

Angular Momentum Quantum Number

The second quantum number used to determine the type of sublevel or subshell that a particular electron is occupying.

Anion

A negatively charged ion.

Atom

Aufbau Prinicple

States that electrons in an atom must fill the lowest energy level first. Poor little electrons…so many rules and regulations…fill this way and fill that way. Why can't an electron ever fill the way it wants to fill?

Cation

A positively charged ion.

Coulomb's Law

Explains the electrostatic interaction between charged particles. If the two charges have the same sign, the electrostatic force between them is repulsive. If they have different signs, the force between them is attractive. Charges close together mean stronger bonds, while charges far apart mean weaker bonds. It ain't fiction just a natural fact.

Covalent Bonding

A bond in which one or more electron pairs are shared between two atoms, typically between two nonmetals. Caring is sharing.

Diamagnetism

Repulsion of a molecule from a magnetic field due to the presence of all paired electrons.

Dipole Moment

A measure of the polarity in a chemical bond or molecule, equal to the product of one charge and the distance between the charges.

Electron

A negatively charged subatomic particle found in the electron "cloud" or volume of space around the nucleus. Electrons are thought to have essentially no mass, but actually their mass is just really super duper small: approximately 1/1836 that of a proton.

Electrostatic Attraction

The attraction between atoms that have opposite charges, and hold the atoms together in ionic bonds.

Electronegativity

Electron Configuration

A notation that shows details about the likely locations of the electrons in a given atom.

Electron Pool Theory

This theory states that each atom present in a metallic crystal loses all of its valence electrons. As a result a "pool" of electrons is formed. It is believed that positively charged metal ions are held together by this "pool" of electrons. Do your electrons want to go for a swim?

Hund's Rule

States that the most stable arrangement for electrons in subshells is the one with the greatest number of parallel spins. Therefore electrons are added to the orbitals in the following way: filling them all half first before any pairing occurs in each subshell.

Hybrid Orbitals

Atomic orbitals formed as a result of mixing the atomic orbitals of the atoms involved in the covalent bond. The type of hybrid orbital formed is dependent on the types of orbitals mixed (for example s, p, and d orbitals). The number of hybrid orbitals formed is the same as the number of atomic orbitals mixed.

Hydrogen Bonding

Considered not a true chemical bond, but it occurs when a hydrogen atom has an electromagnetic attractive interaction with an electronegative atom. The hydrogen bond is typically stronger than Van der Waals attractions but weaker than covalent or ionic bonds.

Ion

An atom or molecule that has obtained a charge by either gaining or losing one or more electrons. This means that the proton number does not match the number of electrons giving the atom either a net negative or positive charge.

Ionic Bond

A bond formed from the electrostatic attraction between ions with different charges. This typically consists of a metal (cation) and a nonmetal (anion). A metal loses electrons to form a cation and some non-metal gains those electrons to form an anion. Opposites attract.

Isoelectronic

Describes molecular entities that share the same number of valence electrons or have the same electron configuration.

Isotope

Atoms of the same element with different masses. Since neutrons have no charge associated with them they don't affect the atom's identity. The atomic number or the number of protons defines that characteristic. However, since neutrons have a mass associated with them they do affect the element's mass number. That's why the same element can have different masses. Carbon-12 says to Carbon-11, "Looking good...what's your secret?" Carbon-11 replies, "Just got a rid of a pesky neutron that was weighing me down."

Lewis Electron Dot Structure

A notation that shows information about an atom's outermost shell, depicting the element and its valence electrons as dots around the element. Lewis sure loved his dots.

Magnetic Quantum Number

The third quantum number used to determine the spatial orientation of the orbital that the electron is occupying.

Mass Number

Adding the total number of protons and neutrons in the nucleus of an atom gives you the atom's mass number. Since the mass of an electron is negligible it doesn't contribute to the overall mass number of an atom.

Metallic Bonding

A bond between metals where the valence electrons are donated to a vast electron "pool", so that the valence electrons are free to move throughout the entire metallic solid.

Molecular Geometry

A type of geometry that defines the shape of a molecule, in which the nonbonding electrons become "invisible" and only the geometry of the atomic nuclei are considered.

Molecular Orbital (MO) Theory

A model that represents the bonding that takes place in covalent compounds. This theory states that atomic orbitals on individual atoms will combine to form molecular orbitals that encompass the entire molecule.

Molecule

A group of atoms that are held together by covalent bonds.

Neutron

A subatomic particle with no charge that resides in the nucleus of an atom. The mass of a neutron is essentially the same as the mass of a proton. Nothing but a chargeless copycat.

Non-polar Covalent Bond

A bond formed when atoms share one or more pairs of electrons relatively equally. Each of the atoms attract the bonding pair of electrons equally.

Octet Rule

States that atoms will lose, gain, or share electrons in order to achieve a filled valence shell, to complete their octet. Anybody up for a game of Crazy Eights?

Orbital Notation

The notation used to designate the electrons and orbitals of a given element. Basically just another way of expressing the electron configuration of an atom in picture perfect way.

Paramagnetism

Attraction of a molecule to a magnetic field due to the presence of unpaired electrons.

Pi (π) Bonds

Overlapping of orbitals, where two lobes of one atomic orbital overlap two lobes of another atomic orbital. These overlapping orbitals share a nodal plane that passes through the two nuclei of the atoms involved in the covalent bond. A double bond always consists of one sigma bond and one pi bond. Mmm Pie.

Polar Covalent Bond

A case where the two atoms involved in the covalent bond are not sharing the electrons equally. The bonded electrons are pulled toward the atom with the greater attraction towards electrons. We've all been there.

Principle Quantum Number

The first quantum number used to determine the energy level or shell that a particular electron is occupying.

Proton

A positively charged atomic particle. The number of protons in an atom determines the identity of the atom; for example, carbon atoms always have 6 protons.

Quantum Numbers

Numbers given to a particular electron, each electron is given four. The numbers are used as a coordinate system to help identify and find the particular electron you want to study.

Resonance Structures

The average of all the possible Lewis structures of a given compound. Some molecules can have more than one possible written structure. No individual structure truly represents the actual structure. That would require an average of all the possible structures. The Resonance Theory is used to describe this situation and often uses this two-headed arrow notation (← → ) between each of the possible structures drawn.

Sigma (σ) Bonds

Overlap of orbitals occurs on a line between two atoms involved in the covalent bond.

Spin Quantum Number

The fourth quantum number used to determine the "spin" direction of the electron you are studying in a particular orbital.

Valence Electron

An electron found in the outermost shell of an atom...we're talking the highest energy level of the atom. These guys are definitely not afraid of heights.

Valence Bond Theory

Is one way in which the molecular geometry of molecules may be determined. It describes covalent bonding as a mixing of atomic orbitals to form a new kind of orbital…hybrid orbitals.

Van Der Waals Attractions

A weak attractive force between atoms or nonpolar molecules caused by a temporary change in dipole moment arising from a brief shift of orbital electrons to one side of one atom or molecule. This creates a similar shift in adjacent atoms or molecules.

Valence-Shell Electron-Pair Repulsion (VSEPR) Theory

A theory that predicts molecular geometry. The theory states that electron pairs around the central atom will try to stay as far away from each other as possible in order to minimize repulsive forces. Honestly, who wouldn't want to stay away from things that repulse them?

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