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Glossary: Chemical bonds
- alkane. paraffin. Compare with hydrocarbon and alkene.
- A series of organic compounds with general formula CnH2n+2. Alkane names end with -ane. Examples are propane (with n=3) and octane (with n=8).
- antibonding orbital. antibonding; antibonding molecular orbital.
- A molecular orbital that can be described as the result of destructive interference of atomic orbitals on bonded atoms. Antibonding orbitals have energies higher than the energies its constituent atomic orbitals would have if the atoms were separate.
- average bond enthalpy. Compare with bond enthalpy.
- Average enthalpy change per mole when the same type of bond is broken in the gas phase for many similar substances
- 1. An atom, bond, or lone pair that is perpendicular to equatorial atoms, bonds, and lone pairs in a trigonal bipyramidal molecular geometry.
- bond energy. Compare with bond enthalpy.
- Energy change per mole when a bond is broken in the gas phase for a particular substance.
- bond enthalpy. Compare with average bond enthalpy.
- Enthalpy change per mole when a bond is broken in the gas phase for a particular substance.
- bond length.
- The average distance between the nuclei of two bonded atoms in a stable molecule.
- bond order.
- 1. In Lewis structures, the number of electron pairs shared by two atoms. 2. In molecular orbital theory, the net number of electron pairs in bonding orbitals (calculated as half the difference between the number of electrons in bonding orbitals and the number of electrons in antibonding orbitals.
- chemical bond. bond; bonding; chemical bonding.
- A chemical bond is a strong attraction between two or more atoms. Bonds hold atoms in molecules and crystals together. There are many types of chemical bonds, but all involve electrons which are either shared or transferred between the bonded atoms.
- covalent bond. covalent; covalently bound. Compare with covalent compound and ionic bond.
- A covalent bond is a very strong attraction between two or more atoms that are sharing their electrons. In structural formulas, covalent bonds are represented by a line drawn between the symbols of the bonded atoms.
- electric dipole. dipole.
- An object whose centers of positive and negative charge do not coincide. For example, a hydrogen chloride (HCl) molecule is an electric dipole because bonding electrons are on average closer to the chlorine atom than the hydrogen, producing a partial positive charge on the H end and a partial negative charge on the Cl end.
- electric dipole moment. (µ) dipole moment.
- A measure of the degree of polarity of a polar molecule. Dipole moment is a vector with magnitude equal to charge separation times the distance between the centers of positive and negative charges. Chemists point the vector from the positive to the negative pole; physicists point it the opposite way. Dipole moments are often expressed in units called Debyes.
- electronegativity Compare with ionization energy and electron affinity.
- Electronegativity is a measure of the attraction an atom has for bonding electrons. Bonds between atoms with different electronegativities are polar, with the bonding electrons spending more time on average around the atom with higher electronegativity.
- enthalpy of atomization. (Hat) atomization enthalpy; heat of atomization.
- The change in enthalpy that occurs when one mole of a compound is converted into gaseous atoms. All bonds in the compound are broken in atomization and none are formed, so enthalpies of atomization are always positive.
- free radical.
- A free radical is a molecule with an odd number of electrons. Free radicals do not have a completed octet and often undergo vigorous redox reactions. Free radicals produced within cells can react with membranes, enzymes, and genetic material, damaging or even killing the cell. Free radicals have been implicated in a number of degenerative conditions, from natural aging to Alzheimer's disease.
- geometric isomer.
- Geometric isomers are molecules that have the same molecular formula and bond connections, but distinctly different shapes.
- hydrogen bond. hydrogen bonding.
- An especially strong dipole-dipole force between molecules X-H...Y, where X and Y are small electronegative atoms (usually F, N, or O) and ... denotes the hydrogen bond. Hydrogen bonds are responsible for the unique properties of water and they loosely pin biological polymers like proteins and DNA into their characteristic shapes.
- incomplete octet.
- 1. An atom with less than eight electrons in its valence shell. 2. An atom with less than eight total bonding and nonbonding electrons in a Lewis structure, for example, B in BH3 has an incomplete octet.
- inductive effect. inductance effect.
- An inductive effect is the polarization of a chemical bond caused by the polarization of an adjacent bond. (Field effects are polarization caused by nonadjacent bonds).
- inert pair. inert pair effect.
- Valence electrons in an s orbital penetrate to the nucleus better than electrons in p orbitals, and as a result they're more tightly bound to the nucleus and less able to participate in bond formation. A pair of such electrons is called an "inert pair". The inert pair effect explains why common ions of Pb are Pb4+ and Pb2+, and not just Pb4+ as we might expect from the octet rule.
- infrared spectroscopy. IR spectroscopy.
- A technique for determining the structure (and sometimes concentration) of molecules by observing how infrared radiation is absorbed by a sample.
- ionic bond. ionically bound; ionic bonding. Compare with covalent bond.
- An attraction between ions of opposite charge. Potassium bromide consists of potassium ions (K+) ionically bound to bromide ions (Br-). Unlike covalent bonds, ionic bond formation involves transfer of electrons, and ionic bonding is not directional.
- ionic compound. salt. Compare with covalent compound and ionic bond.
- A compound made of distinguishable cations and anions, held together by electrostatic forces.
- Lewis structure. electron dot structure; dot structure.
- A model pioneered by Gilbert N. Lewis and Irving Langmuir that represents the electronic structure of a molecule by writing the valence electrons of atoms as dots. Pairs of dots (or lines) wedged between atoms represent bonds; dots drawn elsewhere represent nonbonding electrons.
- lone pair. nonbonding pair; unshared pair.
- Electrons that are not involved in bonding.
- molecular geometry.
- 1. The three-dimensional shape of a molecule. For example, methane (CH4) has a tetrahedral molecular geometry. 2. The study of molecular shapes.
- molecular orbital. Compare with atomic orbital and orbital.
- A wavefunction that describes the behavior of an electron in a molecule. Molecular orbitals are usually spread across many atoms in the molecule, and they are often described as a combination of atomic orbitals on those atoms.
- multiple bond.
- Sharing of more than one electron pair between bonded atoms. A double bond consists of two shared pairs of electrons; a triple bond consists of three shared pairs.
- A set of eight valence electrons.
- octet rule.
- A guideline for building Lewis structures that states that atoms tend to gain, lose, or share valence electrons with other atoms in a molecule until they hold or share eight valence electrons. The octet rule almost always holds for carbon, nitrogen, oxygen, and fluorine; it is regularly violated for other elements.
- pi bond. ( bond) Compare with sigma bond.
- In the valence bond theory, a pi bond is a valence bond formed by side-by-side overlap of p orbitals on two bonded atoms. In most multiple bonds, the first bond is a sigma bond and all of the others are pi bonds.
- polar bond. Compare with covalent bond and ionic bond.
- A bond involving electrons that are unequally shared. Polar bonds can be thought of as intermediate between the extremes represented by covalent bonds and ionic bonds.
- polar molecule. polar. Compare with covalent compound, ionic compound and polar bond.
- An asymmetric molecule containing polar bonds. H2O, NH3, and HCl are examples of polar molecules. Non-examples are CO2, CCl4, and BCl3 which contain polar bonds but are nonpolar because they have symmetric shapes. Alkanes are usually asymmetric but are nonpolar because they contain no polar bonds. Polar molecules are electric dipoles and they attract each other via dipole-dipole forces.
- Description of the ground state of a molecule with delocalized electrons as an average of several Lewis structures. The actual ground state doesn't switch rapidly between the separate structures: it is an average.
- resonance effect. mesomeric effect.
- If electron density at a particular point in a molecule is higher or lower than what you'd expect from a single Lewis structure, and various canonical structures can be drawn to show how electron delocalization will explain the discrepancy, the difference in electron density is called a "resonance effect" or "mesomeric effect".
- sigma bond. ( bond) Compare with pi bond.
- In the valence bond theory, a sigma bond is a valence bond that is symmetrical around the imaginary line between the bonded atoms. Most single bonds are sigma bonds.
- triple bond. ()
- A covalent bond that involves 3 bonding pairs. In the valence bond theory, one of the bonds in a triple bond is a sigma bond and the other two are pi bonds. For example, the central bond in acetylene is a triple bond: H-CC-H.
- The number of hydrogen atoms that typically bond to an atom of an element. For example, in H2O, oxygen has a valence of 2; carbon in CH4 has a valence of four.
- valence bond.
- In the valence bond theory, a valence bond is a chemical bond formed by overlap of half-filled atomic orbitals on two different atoms.
- valence electron.
- Electrons that can be actively involved in chemical change; usually electrons in the shell with the highest value of n. For example, sodium's ground state electron configuration is 1s2 2s2 2p6 3s1; the 3s electron is the only valence electron in the atom. Germanium (Ge) has the ground state electron configuration 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p2; the 4s and 4p electrons are the valence electrons.
- valence shell.
- The shell corresponding to the highest value of principal quantum number in the atom. The valence electrons in this shell are on average farther from the nucleus than other electrons; they are often directly involved in chemical reaction.