POLARITY OF THE MOLECULES                                                    

Bond Polarity

Polarity in organic chemistry refers to a separation of charge and can describe a bond or an entire molecule. Experimentally, bond polarity is measured by its dipole moment. Bonds connecting atoms of different electronegativity are polar with a higher density of bonding electrons around the more electronegative atom giving it a partial negative charge (designated as d-). The less electronegative atom has some of its electron density taken away giving it a partial positive charge (d+).
This polarization of charge in the H-Cl bond is due to different electronegativities of chlorine and hydrogen.

The polarity in the bond can also be represented by a arrow indicating a dipole (two charges separated by a distance). The tip of the arrow points toward the more electronegative atom.

Molecular Polarity

The polarity of the molecule is the sum of all of the bond polarities in the molecule. Since the dipole moment (m, measured in Debyes (D)) is a vector (a quantitiy with both magnitude and direction), the molecular dipole moment is the vector sum of the individual dipole moments. If we compare the molecular dipole moments of formaldehyde and carbon dioxide, both containing a polar carbonyl (C=O) group, we find that formaldehyde is highly polar while carbon dioxide is nonpolar . Since CO₂ is a linear molecule, the dipoles cancel each other.

Water is a bent molecule with polar O-H bonds. The bond dipole moments add to give a resultant dipole (m = 1.85 D) directed toward the more electronegative oxygen.

The polarity of chloromethanes reveals the importance of symmetry. All of these compounds contain polar C-Cl bonds but the tetrahedral symmetry of CCl₄ causes the bond dipoles to cancel giving a nonpolar molecule.

Chloromethane The top image show the bond electron density and the bottom image the molecular dipole. m = 1.87 D	Dichloromethane The top image show the bond electron density and the bottom image the molecular dipole m = 1.54 D	Trichloromethane The top image show the bond electron density and the bottom image the molecular dipole m = 1.02 D	Tetrachloromethane The top image show the bond electron density and the bottom image the molecular dipole m = 0 D

Dipoles and Intermolecular Attraction

Melting points and boiling points are important physical properties. These properties reveal something about the forces that hold molecules together in condensed phases (liquids and solids). Chemists recognize three major kinds of attractive forces in covalent molecules, all of which are related to dipoles.
Polar molecules have a permanent dipole moment. Since opposite charges attract, when polar molecules approach each other they orient themselves in a head-to-tail manner. The following example shows the dipole-dipole attraction in chloroform (trichloromethane, bp 61^oC).

Carbon tetrachloride (tetrachloromethane, bp 77^oC) is a nonpolar molecule but it is a liquid at room temperature, indicating that some attraction between molecules must exist. The molecule has no permanent dipole but an instantaneous dipole is formed when two CCl4 molecules approach each other. The electron cloud in one molecule repulses the electrons in the second molecule breaking the symmetry. These temporary dipoles exist for only a short time and fluctuate from one molecule to another. The result is a weak dipole-dipole attraction called the London dispersive force (van der Waals force). The more contact area between molecules the stronger the van der Waals forces. We will see examples of this trend when we examine the boiling points of hydrocarbons.

Tetrachloromethane molecules far apart.  No dipole moment.	Induced dipole moment of two tetrachloromethane molecules close together.

Hydrogen bonding is the result of strong dipole-dipole attraction in molecules containing O-H and N-H bonds. HF also undergoes hydrogen bonding but since F is monovalent, this bond is not found in organic molecules. These hydrogen bonds are strongly polarized with the hydrogen atom carrying a partial positive charge. Although hydrogen bonding is only about 10% as strong as covalent bonds it is responsible for the unusual high boiling points of water and alcohols which contain O-H bonds. Ammonia and amines contain N-H bonds which are less polar than O-H bonds and the resulting hydrogen bonding is weaker.

chemicalbonding

Though the periodic table has only 118 or so elements, there are obviously more substances in nature than 118 pure elements. This is because atoms can react with one another to form new substances called compounds (see our Chemical Reactions module). Formed when two or more atoms chemically bond together, the resulting compound is unique both chemically and physically from its parent atoms.

Let's look at an example.  The element sodium is a silver-colored metal that reacts so violently with water that flames are produced when sodium gets wet.  The element chlorine is a greenish-colored gas that is so poisonous that it was used as a weapon in World War I.  When chemically bonded together, these two dangerous substances form the compound sodium chloride, a compound so safe that we eat it every day - common table salt!

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In 1916, the American chemist Gilbert Newton Lewis proposed that chemical bonds are formed between atoms because electrons from the atoms interact with each other. Lewis had observed that many elements are most stable when they contain eight electrons in their valence shell. He suggested that atoms with fewer than eight valence electrons bond together to share electrons and complete their valence shells.

While some of Lewis' predictions have since been proven incorrect (he suggested that electronsoccupy cube-shaped orbitals), his work established the basis of what is known today about chemical bonding. We now know that there are two main types of chemical bonding; ionic bonding and covalent bonding.

Ionic bonding

In ionic bonding, electrons are completely transferred from one atom to another. In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely charged ions are attracted to each other by electrostatic forces, which are the basis of the ionic bond.

For example, during the reaction of sodium with chlorine:

sodium (on the left) loses its one valence electron to chlorine (on the right),

resulting in

a positively charged sodiumion (left) and a negatively charged chlorine ion (right).

The reaction of sodium with chlorine

Concept simulation - Reenacts the reaction of sodium with chlorine.

Notice that when sodium loses its one valence electron it gets smaller in size, while chlorine grows larger when it gains an additional valence electron. This is typical of the relative sizes ofions to atoms. Positive ions tend to be smaller than their parent atoms while negative ions tend to be larger than their parent. After the reaction takes place, the charged Na+ and Cl- ions are held together by electrostatic forces, thus forming an ionic bond. Ionic compounds share many features in common:

• Ionic bonds form between metals and nonmetals.
• In naming simple ionic compounds, the metal is always first, the nonmetal second (e.g., sodium chloride).
• Ionic compounds dissolve easily in water and other polar solvents.
• In solution, ionic compounds easily conduct electricity.
• Ionic compounds tend to form crystalline solids with high melting temperatures.

This last feature, the fact that ionic compounds are solids, results from the intermolecular forces(forces between molecules) in ionic solids. If we consider a solid crystal of sodium chloride, the solid is made up of many positively charged sodium ions (pictured below as small gray spheres) and an equal number of negatively charged chlorine ions (green spheres). Due to the interaction of the charged ions, the sodium and chlorine ions are arranged in an alternating fashion as demonstrated in the schematic. Each sodium ion is attracted equally to all of its neighboring chlorine ions, and likewise for the chlorine to sodium attraction. The concept of a single moleculedoes not apply to ionic crystals because the solid exists as one continuous system. Ionic solids form crystals with high melting points because of the strong forces between neighboring ions.

	Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1
Sodium Chloride Crystal	NaCl Crystal Schematic

Covalent bonding

The second major type of atomic bonding occurs when atoms share electrons. As opposed to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when two (or more) elements share electrons. Covalent bonding occurs because the atoms in thecompound have a similar tendency for electrons (generally to gain electrons). This most commonly occurs when two nonmetals bond together. Because both of the nonmetals will want to gain electrons, the elements involved will share electrons in an effort to fill their valence shells. A good example of a covalent bond is that which occurs between two hydrogen atoms. Atoms of hydrogen (H) have one valence electron in their first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H2. Because the hydrogen compound is a combination of equally matched atoms, the atoms will share each other's single electron, forming one covalent bond. In this way, both atoms share the stability of a full valence shell.

Covalent bonding between hydrogen atoms

Concept simulation - Recreates covalent bonding between hydrogen atoms.

Unlike ionic compounds, covalent molecules exist as true molecules. Because electrons are shared in covalent molecules, no full ionic charges are formed.  Thus covalent molecules are not  strongly attracted to one another.  As a result, covalent molecules move about freely and tend to exist as liquids or gases at room temperature.  

Multiple Bonds: For every pair of electrons shared between two atoms, a single covalent bond is formed.  Some atoms can share multiple pairs of electrons, forming multiple covalent bonds.  For example, oxygen (which has six valence electrons) needs two electrons to complete its valence shell.  When two oxygen atoms form the compound O2, they share two pairs of electrons, forming two covalent bonds.  

Lewis Dot Structures: Lewis dot structures are a shorthand to represent the valence electrons of an atom. The structures are written as the element symbol surrounded by dots that represent the valence electrons. The Lewis structures for the elements in the first two periods of the periodic table are shown below.

	Lewis Dot Structures

Lewis structures can also be used to show bonding between atoms. The bonding electrons are placed between the atoms and can be represented by a pair of dots or a dash (each dash represents one pair of electrons, or one bond). Lewis structures for H2 and O2 are shown below.

H2	H:H	or	H-H
O2

Polar and nonpolar covalent bonding

There are, in fact, two subtypes of covalent bonds. The H2 molecule is a good example of the first type of covalent bond, the nonpolar bond. Because both atoms in the H2 molecule have an equal attraction (or affinity) for electrons, the bonding electrons are equally shared by the two atoms, and a nonpolar covalent bond is formed. Whenever two atoms of the same element bond together, a nonpolar bond is formed.

A polar bond is formed when electrons are unequally shared between two atoms. Polar covalent bonding occurs because one atom has a stronger affinity for electrons than the other (yet not enough to pull the electrons away completely and form an ion). In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule.

H2O: a water molecule

Water molecules contain two hydrogen atoms (pictured in red) bonded to one oxygen atom (blue). Oxygen, with six valence electrons, needs two additional electrons to complete its valence shell. Each hydrogen contains one electron. Thus oxygen shares the electrons from two hydrogen atoms to complete its own valence shell, and in return shares two of its own electrons with each hydrogen, completing the H valence shells.

Polar covalent bonding simulated in water

The primary difference between the H-O bond in water and the H-H bond is the degree ofelectron sharing. The large oxygen atom has a stronger affinity for electrons than the small hydrogen atoms. Because oxygen has a stronger pull on the bonding electrons, it preoccupies their time, and this leads to unequal sharing and the formation of a polar covalent bond.  

The dipole

Because the valence electrons in the water molecule spend more time around the oxygen atomthan the hydrogen atoms, the oxygen end of the molecule develops a partial negative charge (because of the negative charge on the electrons). For the same reason, the hydrogen end of the molecule develops a partial positive charge. Ions are not formed; however, the molecule develops a partial electrical charge across it called a dipole. The water dipole is represented by the arrow in the pop-up animation (above) in which the head of the arrow points toward the electron dense (negative) end of the dipole and the cross resides near the electron poor (positive) end of the molecule.