Chemical polarity

In chemistry, polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole moment, with a negatively charged end and a positively charged end.This amphiphilic molecule has several polar groups (hydrophilic, water-loving) on the right side and a long nonpolar chain (lipophilic, fat-loving) at the left side. This gives it surfactant propertiesA micelle – the lipophilic ends of the surfactant molecules dissolve in the oil, while the hydrophilic charged ends remain outside in the water phase, shielding the rest of the hydrophobic micelle. In this way, the small oil droplet becomes water-soluble.Phospholipids are effective natural surfactants that have important biological functionsCross section view of the structures that can be formed by phospholipids. They can form a micelle and are vital in forming cell membranes In chemistry, polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole moment, with a negatively charged end and a positively charged end. Polar molecules must contain polar bonds due to a difference in electronegativity between the bonded atoms. A polar molecule with two or more polar bonds must have a geometry which is asymmetric in at least one direction, so that the bond dipoles do not cancel each other. Polar molecules interact through dipole–dipole intermolecular forces and hydrogen bonds. Polarity underlies a number of physical properties including surface tension, solubility, and melting and boiling points. Not all atoms attract electrons with the same force. The amount of 'pull' an atom exerts on its electrons is called its electronegativity. Atoms with high electronegativities – such as fluorine, oxygen, and nitrogen – exert a greater pull on electrons than atoms with lower electronegativities such as alkali metals and alkaline earth metals. In a bond, this leads to unequal sharing of electrons between the atoms, as electrons will be drawn closer to the atom with the higher electronegativity. Because electrons have a negative charge, the unequal sharing of electrons within a bond leads to the formation of an electric dipole: a separation of positive and negative electric charge. Because the amount of charge separated in such dipoles is usually smaller than a fundamental charge, they are called partial charges, denoted as δ+ (delta plus) and δ− (delta minus). These symbols were introduced by Sir Christopher Ingold and Dr. Edith Hilda (Usherwood) Ingold in 1926. The bond dipole moment is calculated by multiplying the amount of charge separated and the distance between the charges. These dipoles within molecules can interact with dipoles in other molecules, creating dipole-dipole intermolecular forces. 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 called an ionic bond, and occurs when the difference between electronegativities is large enough that one atom actually takes an electron from the other. The terms 'polar' and 'nonpolar' are usually applied to covalent bonds, that is, bonds where the polarity is not complete. To determine the polarity of a covalent bond using numerical means, the difference between the electronegativity of the atoms is used. Bond polarity is typically divided into three groups that are loosely based on the difference in electronegativity between the two bonded atoms. According to the Pauling scale: Pauling based this classification scheme on the partial ionic character of a bond, which is an approximate function of the difference in electronegativity between the two bonded atoms. He estimated that a difference of 1.7 corresponds to 50% ionic character, so that a greater difference corresponds to a bond which is predominantly ionic.