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A bond dipole moment is a measure of the polarity of a chemical bond between two atoms in a molecule. It involves the concept of electric dipole moment, which is a measure of the separation of negative and positive charges in a system.
Dipole moments occur when there is a separation of charge. They can occur between two ions in an ionic bond or between atoms in a covalent bond; dipole moments arise from differences in electronegativity. The larger the difference in electronegativity, the larger the dipole moment.
Bond angles also contribute to the shape of a molecule. Bond angles are the angles between adjacent lines representing bonds. The bond angle can help differentiate between linear, trigonal planar, tetraheral, trigonal-bipyramidal, and octahedral.
Highlights. Learning Objectives. By the end of this section, you will be able to: Predict the structures of small molecules using valence shell electron pair repulsion (VSEPR) theory. Explain the concepts of polar covalent bonds and molecular polarity. Assess the polarity of a molecule based on its bonding and structure.
The magnitude of a bond dipole moment is represented by the Greek letter mu (µ) and is given by the formula shown below, where Q is the magnitude of the partial charges (determined by the electronegativity difference) and r is the distance between the charges: μ= Qr μ = Qr.
To determine if a molecule is polar we have to look at all the bond dipoles present within the molecule. If no bond dipoles exist, the molecule is nonpolar. If we have bond dipoles present in the molecule we have to look at the addition of these and their orientation in 3D space.
The dipole moment, μ (lowercase Greek letter mu), is defined as the magnitude of the charge Q at either end of the molecular dipole times the distance r between the charges, μ = Q × r. Dipole moments are expressed in debyes (D), where 1 D = 3.336 × 10 –30 coulomb meters (C · m) in SI units.
