Dipole
Dipole
Dipole, literally, means “two poles.” It refers to two electrical charges, one negative and the other positive. Dipoles are common in atoms whenever electrons (negatively charged) are unevenly distributed around nuclei (positively charged), and in molecules whenever electrons are unevenly shared between two atoms in a covalent bond. A dipole can also be created artificially by charging two ends of an object with opposite charges. One class of radio antennas, used for both transmitting and receiving, is the dipole antenna.
When a dipole is present, the atom or covalent bond is said to be polarized, or divided into negative and positive regions. This is indicated by the use of partial negative (δ–) and partial positive (δ+) signs. The magnitude and direction of the electrical charge separation is indicated by using an arrow, drawn from the positive pole in a molecule to the negative pole.
In covalent bonds, permanent dipoles are caused when two different atoms share their electrons unevenly. The atom that is more electronegative— the one that holds electrons more tightly—pulls the electrons closer to itself, creating a partial negative charge there. The less electronegative atom becomes partially positive as a result because it has lost partial possession of the electrons. The electric strength of a dipole generally increases as the electronegativity difference between the atoms in the bond increases. This strength, called a dipole moment, can be measured experimentally. The size of a dipole moment is expressed in Debye units in honor of the Dutch chemist, Peter Debye (1884–1966).
The dipole moments of a series of molecules are listed below:
| Molecule | Dipole moment (in Debye units, D) |
|---|---|
| HF | 1.91 D |
| HC1 | 1.03 D |
| HBr | 0.78 D |
| HI | 0.38 D |
The measurement of dipole moments can help determine the shape of a molecule. The net dipole moment of a water molecule (H2 O) represents the overall electrical charge distribution in that molecule. (See Figure 1.)
The H2 0 molecule is bent; its dipole vectors do not cancel. The water molecule therefore has a net resultant dipole moment of 1.87 D. If the molecule were linear, the measured dipole moment would be zero; its individual dipoles in the two oxygen-hydrogen covalent bonds would cancel each other out.
Individual atoms (and ions) are naturally polarized if their electrons happen to move irregularly about their nuclei, creating, at least temporarily, lopsided charge distributions with δ+ and δ- portions. Natural collisions occurring between atoms can induce this temporary deformity from an atom’s normal spherical, symmetric shape. Larger atoms are considered to be “softer” than smaller, “harder” atoms—more easily disturbed. Larger atoms are then more likely to be polarized or to have stronger dipoles than smaller atoms.
The presence of dipoles helps to explain how atoms and molecules attract each other. Figure 2 shows how the electrically positive side of one xenon atom (Xe) lines up and pulls towards the negative side of another xenon atom. Likewise, the positive side of one H-C1 molecule is attracted to the negative side of another
H-C1 molecule. When many atoms and molecules are present in matter, these effects continue on indefinitely from atom to atom and molecule to molecule.
Dipole forces tend to organize matter and pull it together. Atoms and molecules most strongly attracted to each other will tend to exist as solids. Weaker interactions tend to produce liquids. The gaseous state of matter will tend to exist when the atoms and molecules are non-polar, or when virtually no dipoles are present.
Dipole
Dipole
Dipole, literally, means "two poles," two electrical charges, one negative and one positive. Dipoles are common in atoms whenever electrons (-) are unevenly distributed around nuclei (+), and in molecules whenever electrons are unevenly shared between two atoms in a covalent bond.
When a dipole is present, the atom or covalent bond is said to be polarized, or divided into negative and positive regions. This is indicated by the use of partial negative (δ-) and partial positive (δ+) signs. The magnitude and direction of the electrical charge separation is indicated by using an arrow, drawn from the positive pole in a molecule to the negative pole.
In covalent bonds, permanent dipoles are caused when two different atoms share their electrons unevenly. The atom that is more electronegative—the one that holds electrons more tightly—pulls the electrons closer to itself, creating a partial negative charge there. The less electronegative atom becomes partially positive as a result because it has lost partial possession of the electrons. The electric strength of a dipole generally increases as the electronegativity difference between the atoms in the bond increases. This strength, called a dipole moment, can be measured experimentally. The size of a dipole moment is expressed in Debye units in honor of the Dutch chemist, Peter Debye (1884-1966).
The dipole moments of a series of molecules are listed below:
| Molecule | Dipole Moment (in Debye units, D) |
| HF | 1.91 D |
| HC1 | 1.03 D |
| HBr | 0.78 D |
| HI | 0.38 D |
The measurement of dipole moments can help determine the shape of a molecule. The net dipole moment of a water molecule (H2O) represents the overall electrical charge distribution in that molecule. (See Figure 1.)
The H20 molecule is bent. Its dipole vectors do not cancel. The water molecule has a net resultant dipole moment of 1.87 D. If the molecule were linear, the measured dipole moment would be zero . Its individual dipoles in the two oxygen-hydrogen covalent bonds would have cancelled each other out.
Individual atoms (and ions) will be naturally polarized if their electrons happen to move irregularly about their nuclei creating, at least temporarily, lopsided looking atoms with δ+ and δ-portions. Natural collisions occurring between atoms can induce this temporary deformity from an atom's normal spherical, symmetric shape. Larger atoms are considered to be "softer" than smaller, "harder" atoms. Larger atoms are then more likely to be polarized or to have stronger dipoles than smaller atoms.
The presence of dipoles helps to explain how atoms and molecules attract each other. Figure 2 shows how the electrically positive side of one xenon atom (Xe) lines up and pulls towards the negative side of another xenon atom. Likewise, the positive side of one H-C1 molecule is attracted to the negative side of another H-C1 molecule. When many atoms and molecules are present in matter , these effects continue on indefinitely from atom to atom and molecule to molecule.
Dipole forces tend to organize matter and pull it together. Atoms and molecules most strongly attracted to each other will tend to exist as solids. Weaker interactions tend to produce liquids. The gaseous state of matter will tend to exist when the atoms and molecules are nonpolar, or when virtually no dipoles are present.
dipole
di·pole / ˈdīˌpōl/ • n. Physics a pair of equal and oppositely charged or magnetized poles separated by a distance. ∎ an antenna consisting of a horizontal metal rod with a connecting wire at its center. ∎ Chem. a molecule in which a concentration of positive electric charge is separated from a concentration of negative charge.DERIVATIVES: di·po·lar / dīˈpōlər/ adj.