Bond polarity

(noun)

A covalent bond is polar if one atom is more electronegative than its bonding partner, resulting in a net dipole moment between the two atoms.

Related Terms

Examples of Bond polarity in the following topics:

  • Bond Polarity

    • Bond polarity exists when two bonded atoms unequally share electrons, resulting in a negative and a positive end.
    • Bonds can fall between one of two extremes, from completely nonpolar to completely polar.
    • The terms "polar" and "nonpolar" usually refer to covalent bonds.
    • To determine the polarity of a covalent bond using numerical means, find the difference between the electronegativity of the atoms; if the result is between 0.4 and 1.7, then, generally, the bond is polar covalent.
    • The hydrogen fluoride (HF) molecule is polar by virtue of polar covalent bonds; in the covalent bond, electrons are displaced toward the more electronegative fluorine atom.

  • Bond Polarity

    • Molecular polarity is dependent on the presence of polar covalent bonds and the molecule's three-dimensional structure.
    • Bond polarity: when atoms from different elements are covalently bonded, the shared pair of electrons will be attracted more strongly to the atom with the higher electronegativity.
    • Such bonds are said to be 'polar' and possess partial ionic character.
    • In molecules containing more than one polar bond, the molecular dipole moment is just the vector addition of the individual bond dipole moments.
    • Apply knowledge of bond polarity and molecular geometry to identify the dipole moment of molecules

  • Types of Bonds

    • Pure ionic bonding cannot exist: all ionic compounds have some degree of covalent bonding.
    • The larger the difference in electronegativity between the two atoms involved in the bond, the more ionic (polar) the bond is.
    • Bonds with partially ionic and partially covalent character are called polar covalent bonds.
    • This difference in charge is called a dipole, and when the covalent bond results in this difference in charge, the bond is called a polar covalent bond.
    • Similarly, the higher the difference in electronegativity, the more unequal the sharing of electrons is between the nuclei, and the higher the polarity of the bond.

  • Charge Distribution in Molecules

    • Such a covalent bond is polar, and will have a dipole (one end is positive and the other end negative).
    • The degree of polarity and the magnitude of the bond dipole will be proportional to the difference in electronegativity of the bonded atoms.
    • Thus a O–H bond is more polar than a C–H bond, with the hydrogen atom of the former being more positive than the hydrogen bonded to carbon.
    • Methane is essentially non-acidic, since the C–H bond is nearly non-polar.
    • As noted above, the O–H bond of water is polar, and it is at least 25 powers of ten more acidic than methane.

  • Ionic vs Covalent Bond Character

    • The bond formed between any two atoms is not a purely ionic bond.
    • In the conventional presentation, bonds are designated as ionic when the ionic aspect is greater than the covalent aspect of the bond.
    • Bonds that fall in between the two extremes, having both ionic and covalent character, are classified as polar covalent bonds.
    • This bond is considered to have characteristics of both covalent and ionic bonds.
    • Discuss the idea that, in nature, bonds exhibit characteristics of both ionic and covalent bonds

  • Hydrogen Bonding

    • If we compare the boiling points of methane (CH4) -161ºC, ammonia (NH3) -33ºC, water (H2O) 100ºC and hydrogen fluoride (HF) 19ºC, we see a greater variation for these similar sized molecules than expected from the data presented above for polar compounds.
    • Hydrogen forms polar covalent bonds to more electronegative atoms such as oxygen, and because a hydrogen atom is quite small, the positive end of the bond dipole (the hydrogen) can approach neighboring nucleophilic or basic sites more closely than can other polar bonds.
    • The molecule providing a polar hydrogen for a hydrogen bond is called a donor.
    • Also, O–H---O hydrogen bonds are clearly stronger than N–H---N hydrogen bonds, as we see by comparing propanol with the amines.
    • Although the hydrogen bond is relatively weak (ca. 4 to 5 kcal per mole), when several such bonds exist the resulting structure can be quite robust.

  • Covalent Bonds

    • Covalent bonding interactions include sigma-bonding (σ) and pi-bonding (π).
    • Double bonds occur when four electrons are shared between the two atoms and consist of one sigma bond and one pi bond.
    • For atoms with equal electronegativity, the bond between them will be a non-polar covalent interaction.
    • In non-polar covalent bonds, the electrons are equally shared between the two atoms.
    • For atoms with differing electronegativity, the bond will be a polar covalent interaction, where the electrons will not be shared equally.

  • Introduction to Bonding

    • These bonds include both strong intramolecular interactions, such as covalent and ionic bonds.
    • When there is a greater electronegativity difference than between covalently bonded atoms, the pair of atoms usually forms a polar covalent bond.
    • Again, polar covalent bonds tend to occur between non-metals.
    • Bonds, especially covalent bonds, are often represented as lines between bonded atoms.
    • Acetylene has a triple bond, a special type of covalent bond that will be discussed later.

  • Ion-Dipole Force

    • The ion-dipole force is an intermolecular attraction between an ion and a polar molecule.
    • However, ion-dipole forces involve ions instead of solely polar molecules.
    • Ion-dipole bonding is also stronger than hydrogen bonding.
    • Ion-dipole forces are generated between polar water molecules and a sodium ion.
    • These intermolecular ion-dipole forces are much weaker than covalent or ionic bonds.

  • Properties of Aldehydes and Ketones

    • Because of the greater electronegativity of oxygen, the carbonyl group is polar, and aldehydes and ketones have larger molecular dipole moments (D) than do alkenes.
    • The resonance structures in the first diagram below illustrate this polarity, and the relative dipole moments of formaldehyde, other aldehydes and ketones confirm the stabilizing influence that alkyl substituents have on carbocations (the larger the dipole moment the greater the polar character of the carbonyl group).
    • The polarity of the carbonyl group also has a profound effect on its chemical reactivity, compared with the non-polar double bonds of alkenes.
    • Since a C–C σ-bond has a bond energy of 83 kcal/mole, the π-bond energy may be estimated at 63 kcal/mole (i.e. less than the energy of the sigma bond).
    • The C–O σ-bond is found to have an average bond energy of 86 kcal/mole.