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In chemistry, polarity refers to a separation of electric charge leading to a molecule or its chemical groups having an electric dipole or multipole moment. Polar molecules interact through dipole–dipole intermolecular forces and hydrogen bonds. Molecular polarity is dependent on the difference in electronegativity between atoms in a compound and the asymmetry of the compound's structure. For example, a molecule of water is polar because of the unequal sharing of its electrons between oxygen and hydrogen in which the former has larger electronegativity than the latter, resulting in a "bent" structure, whereas methane is considered nonpolar because the carbon shares the electrons with the hydrogen atoms almost uniformly. Polarity underlies a number of physical properties including surface tension, solubility, and melting- and boiling-points.
Electrons are not always shared equally between two bonding atoms; one atom might exert more of a force on the electron cloud than the other. This "pull" is termed electronegativity and measures the attraction for electrons a particular atom has. The unequal sharing of electrons within a bond leads to the formation of an electric dipole: a separation of positive and negative electric charge. Partial charges are denoted as δ+ (delta plus) and δ− (delta minus). These symbols were introduced by Christopher Ingold and his wife Hilda Usherwood in 1926.
Atoms with high electronegativities — such as fluorine, oxygen, and nitrogen — exert a greater pull on electrons than atoms with lower electronegativities. In a bond, this can lead to unequal sharing of electrons between atoms, as electrons will be drawn closer to the atom with the higher electronegativity.
Bonds can fall between one of two extremes — being completely nonpolar or completely polar. A completely nonpolar bond occurs when the electronegativities are identical and therefore possess a difference of zero. A completely polar bond is more correctly termed ionic bonding and occurs when the difference between electronegativities is large enough that one atom takes an electron from the other. The terms "polar" and "nonpolar" bonds usually refer to covalent bonds. To determine the polarity of a covalent bond using numerical means, the difference between the electronegativity of the atoms is taken. If the result is between 0.4 and 1.7 then, generally, the bond is polar covalent.
While the molecules can be described as "polar covalent", "nonpolar covalent", or "ionic", it must be noted that this is often a relative term, with one molecule simply being more polar or more nonpolar than another. However, the following properties are typical of such molecules.
A molecule is composed of one or more chemical bonds between molecular orbitals of different atoms. A molecule may be polar either as a result of polar bonds due to differences in electronegativity as described above, or as a result of an asymmetric arrangement of nonpolar covalent bonds and non-bonding pairs of electrons known as a full molecular orbital.
Examples of common household polar molecules include sugar, for instance the sucrose sugar variety. Sugars have many polar oxygen–hydrogen (-OH) groups and are overall highly polar.
Due to the polar nature of the water molecule (H2O) itself, polar molecules are generally able to dissolve in water.
Ammonia: the two lone electrons are shown in yellow, the hydrogen atoms in white
A molecule may be nonpolar either because there is (almost) no polarity in the bonds (when there is an equal sharing of electrons between two different atoms) or because of the symmetrical arrangement of polar bonds.
Examples of household nonpolar compounds include fats, oil, and petrol/gasoline. Therefore (per the "oil and water" rule of thumb), most nonpolar molecules are water-insoluble (hydrophobic) at room temperature. However, many nonpolar organic solvents, such as turpentine, are able to dissolve polar substances. When comparing a polar and nonpolar molecule with similar molar masses, the polar molecule in general has a higher boiling point, because of the dipole–dipole interaction between their molecules. The most common form of such an interaction is the hydrogen bond, which is also known as the H-bond.
Methane: the bonds are arranged symmetrically so there is no overall dipole
Boron trifluoride: trigonal planar arrangement of three polar bonds results in no overall dipole
Large molecules that have one end with polar groups attached and another end with nonpolar groups are good surfactants. They can aid in the formation of stable emulsions, or blends, of water and fats. Surfactants reduce the interfacial tension between oil and water by adsorbing at the liquid–liquid interface.
Phospholipids are effective natural surfactants that have important biological functions
|HAx||Molecules with a single H||HF|
|AxOH||Molecules with an OH at one end||C2H5OH|
|OxAy||Molecules with an O at one end||H2O|
|NxAy||Molecules with an N at one end||NH3|
|Nonpolar||A2||Diatomic molecules of the same element||O2|
|CxAy||Most carbon compounds||CO2|
Any molecule with an centre of inversion ( "i" ) or a horizontal mirror plane ( "σh ") will not possess dipole moments. Likewise, a molecule with more than one Cn axis will not possess dipole moment because dipole moments cannot lie in more than one dimension. As a consequence of that constraint, all molecules with D symmetry (Schönflies notation) will, therefore, not have dipole moment because, by definition, D point groups have two or multiple Cn axis.
Since C1, Cs,C∞h Cn and Cnv point groups do not have a centre of inversion, horizontal mirror planes or multiple Cn axis, molecules in one of those point groups will have dipole moment.
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