AP Chemistry - Chemical Bonding
Molecular Architecture
Chemical Bonds
Chemical bonds are the electrostatic force of attraction which holds atoms, ions and molecules together. Clues to the type of bond can be obtained by studying the properties of substances and the structural characteristics of the particle. Understanding the nature and origin of chemical bonds is an important part of understanding chemistry, because changes in these bonding forces constitute the underlying basis for all chemical reactions. Old bonds break and new bonds form when chemicals react.
Ionic Bonds
Ionic bonds are formed when metals react with non-metals. For example, when sodium reacts with chlorine, the sodium loses one electron while the chlorine gains one electron. The Na atom which was electrically neutral takes on a positive charge and the chlorine atom which was also electrically neutral takes on a negative charge.
Nao -----> Na+ + e-
Clo + e- -----> Cl-
Nao + Cl- -> Na+Cl-
The reason for the attraction is the fact that opposite charges attract. But why are electrons transferred between these two atoms? Why does Nao form Na+ ions and not Na2+ or Na-? Why does Clo form Cl- and not Cl2- or Cl+ ions? The fact is it depends upon an energy change. In order for these ions to form there is a net energy decrease to a more stable energy level. Making a Na2+ ion is not possible because it would require to much energy. The same is true for the Cl-. Making a Cl- ion is easy. You would have to force it to become Cl2-.
Three factors affect the energy involved in the formation of an ionic compound. One is the removal of electrons from the atoms that become cations (eg. sodium). Formation of a cation requires an input of energy - the ionization energy. (The amount of energy it takes to move an electron out of orbit in a neutral atom and remove it to some infinite point away from the nucleus is the ionization energy). You have a table of ionization energies in your databook. A second factor is the energy change that accompanies the addition of one or more electrons to the atoms that become anions. (eg. chlorine). This energy is the electron affinity. The ionization energy and the electron affinity are energies associated with the changes of isolated gaseous atoms. A crystal of salt, however, does not consist of isolated atoms. A crystal of salt is a group of ions packed tightly into a regular pattern. This pattern is referred to as a lattice, and it has a lower energy than the isolated ions.
To understand this, imagine that we want the vaporize a salt crystal. In order to do this we must add heat energy in order to get the crystal vibrating fast. In the crystal the forces of attraction exceed the forces of repulsion, so to accomplish our vaporization we have to add enough energy to overcome these forces of attraction. This would of course require work, so vaporizing the crystal increases the ions' potential energy and is endothermic. The reverse process - the imaginary process that forms the lattice form from isolated ions - must therefore lead to a lowering of the potential energy of the system and be exothermic. The amount that the energy of the system is lowered because of these mutual attractions of its ions is the lattice energy.
The lattice energy is the major stabilizing factor for ionic compounds. In almost every case, the energy input required by the ionization energy is larger than the energy recovered by the electron affinity, so the IE and EA combined have a net energy-raising effect. It is were not for the large energy-lowering effect of the lattice energy, formation of ionic compounds would be endothermic and they simply wouldn't be formed.
Now why do atoms react? Right from the beginning, you where told that metals tend to form positive ions and non-metals tend to form negative ions. At the left of the period table are the metals - elements with small IE and EA. Relatively little energy is needed to remove electrons from them to produce positive ions. At the upper right of the periodic table are the non-metals - elements with large IE and EA. It is very difficult to remove electrons from these elements, but sizeable amounts of energy are released when they gain electrons. On an energy basis, it is least "expensive" to form a cation from a metal and an anion from a non-metal.
Formation of Ions by the Representative Elements
What happens when sodium loses an electron. The electronic structure of Na is
Na 1s22s22p63s1
The electron that is lost is the one least tightly held. For sodium that is the single outer 3s electron. The electronic structure of the Na+ ion, then is
Na+ 1s22s22p6
The removal of the first electron from Na does not require much energy because the first IE of Na is so small.
For Na 1st  IE = 496 kJ/mol
            2nd IE = 4563 kJ/mol
Therefore, an input of energy equal to the first IE can be easily recovered by the exothermic lattice energy of ionic compounds containing the Na+ ion. However, removal of a second electron from sodium is very difficult - the second IE of Na is enormous. The amount of energy that must be invested to create a Na2+ ion is therefore much greater then the amount of energy that can be recovered by the lattice energy, so overall the formation of a compound containing Na2+ is very energetically unfavourable. This is why we never observe compounds that contain this ion, and why sodium stops losing electrons once it has achieved a noble gas configuration.
A similar situation exists for other metals too. Consider calcium, for example. We known that this metal forms ions with a 2+ charge. This means that when it reacts, a calcium atom loses its two outermost electrons.
Cao 1s22s22p63s23p64s2

Ca2+ 1s22s22p63s23p6

The two 4s electrons of Ca are not held too tightly, so the amount of energy that must be invested to remove them (the sum of the first and second IE) can be recovered easily by the lattice energy of a Ca2+ compound.
For calcium    1st IE = 590 kJ/mol
                    2nd IE = 1140 kJ/mol
                     3rd IE = 4912 kJ/mol
However, the removal of yet another electron from calcium to form Ca3+ requires breaking into the noble gas core. A tremendous amount of energy is needed to accomplish this - much more than would be regained by the lattice energy of Ca3+ compound. Therefore, a calcium atom loses just two electrons when it reacts.
For sodium and calcium, we find that the stability of the noble gas core that lies below the outer shell of electrons effectively limits the number of electrons that they lose, and that the ions that are formed have a noble gas electron configuration. A similar configuration also tends to be the fate of nonmetals when they form anions.
Chlorine and oxygen are typical non-metals that form anions when they react with metals such as sodium or calcium. When a chlorine atom reacts, it gains one electron. For chlorine we have
Clo 1s22s22p63s23p5
and when an electron is gained, its configuration becomes
Cl- 1s22s22p63s23p6
At this point, electron gain ceases, because if another electron were to be added, it would have to enter an orbital in the next higher shell.
With oxygen, a similar situation exists. The formation of the oxide ion, O2-, gives oxygen a noble gas configuration without much difficulty,
Oo 1s22s22p4
O2- 1s22s22p6
and the large lattice energies of metal oxides leads to stable compounds. However, we never see the formation of O3- because, once again, the last electron would have to enter an orbital in the next higher shell, and this is very energetically unfavourable.
The energy factors cause many atoms to form ions that have a noble gas electron configuration. This leads us to the useful generalization that when they form ions, atoms of most of the representative elements tend to gain or lose electrons until they have obtained a configuration that is that of the nearest noble gas.
Go to the Ions Worksheet #1
Go to the Atoms and Ion Worksheet #2
Ions of Transition Metals
The transition elements are metals that form cations when they react. However, the situation is a bit more complex. Many transition elements are able to form more than one cation because they have a partially filled d subshell that is just slightly lower in energy than the outer s subshell. When a transition metal forms a positive ion, it always loses electrons from its outer s subshell first. Once these are gone, any further electron loss takes place from the partially filled d subshell. Iron is a typical example. Its electron configuration is
Fe [Ar]3d64s2
When iron reacts, it loses its 4s electrons fairly easily to give Fe2+. But because the 3d subshell is close in energy to the 4s, it is not very difficult to remove still another electron to give Fe3+.
Fe3+ [Ar]3d5
Because so many transition elements are able to form ions in a similar way, the ability to form more than one positive ion is usually cited as one of the characteristic properties of the transition elements. Frequently, one of the ions formed has a 2+ charge, which arises from the loss of the outer two s electrons. Ions with larger positive charges result when additional d electrons are lost.
Electron Bookkeeping with Lewis Symbols
Ionic bonds use the principle of electron transfer and electrostatic attraction to hold together. Some compounds are held together with covalent bonds which exist because of mutual sharing of bonding valence electrons. It is useful to be able to keep track of valence electrons. It is recommended that you use a simple bookkeeping device called a Lewis symbol, named after their inventor, G.N.Lewis (1875-1946).
To draw the Lewis symbol for an element, start with the chemical symbol surrounded by a number of dots (or some other symbol), which represent the atom's valence electrons. For example, the element lithium, which has one valence electron in its 2s subshell, has the Lewis symbol
In fact, each element in Group IA has a similar Lewis symbol, because each has only one valence electron. The Lewis symbols for all of the Group IA elements are
Li·     Na·      K·     Rb·   Cs·
The Lewis symbols for the eight A-group elements in period 2 are1
GROUP   IA     IIA     IIIA     IVA     VA      VIA     VIIA      O
                                      ·           ·          ·                       ..         ..
Symbol     Li·     ·Be·     ·B·       ·C·      ·N:       · :        ·F:       :Ne:
                                                  ·          ·           ·           ..          ..
The elements in each group below those given have Lewis symbols identical to those above except of course for the chemical symbol of the element. Notice that when an atom has more than four valence electrons, the additional electrons are shown to be paired with others. Also note that the group number is also equal to the number of valence electrons.
1.   For beryllium, boron, and carbon, the number of unpaired electrons in the Lewis symbol doesn't agree with the number predicted from the atom's electron configuration. Boron, for example, has two electrons paired in its 2s orbital and a third electron in one of its 2p orbitals; therefore, there is actually only one unpaired electron in a boron atom. The Lewis symbols are drawn as shown, however, because when beryllium, boron, and carbon form bonds, they behave as if they have two, three, and four unpaired electrons, respectively.
Electron Sharing
Most of the substances we encounter are not ionic. Rather than existing as collections of electrically charged particles (ions), they occur as electrically neutral combinations of atoms called molecules. Water, H2O, consists of molecules, table sugar C12H22O11 and gasoline C8H18 are also examples.
Ionic bonding results because of the energy-lowering lattice energy and the energy raising IE and EA. Many times this is not possible, particulary when the IE of all the atoms involved is large, as happens when non-metals combine with other non-metals. In such instances, nature uses a different way to lower the energy - electron sharing.
What happens when two hydrogen atoms join together to form an H2 molecule? As the two atoms approach each other, the electron of each atom begins to feel the attraction of both nuclei. This causes the electron density around each nucleus to shift toward the region between the two atoms. Therefore, as the distance between the nuclei decreases, there is an increase in the probability of finding either electron near either nucleus. In effect, then, each of the hydrogen atoms in this H2 molecule now has a share of two electrons.
When the electron density shifts to the region between the two hydrogen atoms, it attracts each of the nuclei and pulls them together. Being of the same charge, however, the two nuclei also repel each other, as do the two electrons. In the molecule that forms, therefore, the atoms are held at a distance at which these attractions and repulsions are balanced. Overall, the nuclei are kept from separating, and the net force of attraction produced by the sharing of the pair of electrons is called a covalent bond.
Every covalent bond is characterized by two quantities, the average distance between the nuclei held together by the bond, and the energy needed to separate the two atoms to produce neutral atoms again. In the hydrogen molecule, the attractive forces pull the nuclei to a distance of 75 pm, and this distance is called the bond length or bond distance. Because a covalent bond holds atoms together, work must be done to separate them. When a bond is formed, an equivalent amount of energy is released. The amount of energy released when the bond is formed is called the bond energy.
Formation of a covalent bond releases the bond energy, which means that as the bond forms, the energy of the atom decreases. In the diagram below you can see how the energy changes when two hydrogen atoms form H2. The minimum energy occurs at a bond distance of 75 pm, and that 1 mol of hydrogen molecules is more stable than 2 mol of hydrogen atoms by 435 kJ. In other words the bond energy of H2 is 435 kJ/mole.

Before joining, each of the separate hydrogen atoms has one electron in its 1s orbital. When these electrons are shared, the 1s orbital of each atom has, in a sense, become filled. The electrons have also become paired, as required by Pauli's Exclusion Principle, (each pair of atoms has a different spin and hence opposite magnetic poles). For this reason a covalent bond is sometimes referred to as an electron pair bond.
Covalent bonds which are sharing an electron pair bond are indicated with a dash, just like what was used in the organic unit.

H· + ·H ---> H:H which is shown as H-H. (the line indicates a covalent bond pair of electrons).

Covalent Bonding and the Octet Rule
Hydrogen, with just one electron in its 1s orbital, can achieve a noble gas configuration (that of helium) by obtaining a share of one electron from another atom. The Lewis structure of H2 implies that both atoms have access to both electrons in the bond.
Since hydrogen obtains a stable valence shell configuration when it shares just one pair of electrons with another atom, hydrogen atoms only form one covalent bond. (This should be easily confirmed from the organic unit as well).
The Octet Rule
The valence shell of the noble gases other than helium all contain eight electrons, and the tendency of many atoms to acquire such an outer shell electron configuration forms the basis of the octet rule.
Octet Rule: When atoms react, they tend to achieve an outer shell having eight electrons.
Many of the representative elements (sodium and chlorine for instance) follow this rule when they form ions. In the ionic case they achieve the noble gas shell configuration by either a gain or loss of electrons. When atoms other than hydrogen form covalent bonds, an octet is accomplished by sharing. The octet rule can be used to explain the number of covalent bonds an atom forms. This number normally equals the number of electrons that atom needs to have a total of eight electrons (an octet) in its outer shell. For example, the halogens (Group VIIA), all have seven valence electrons.
The Lewis symbol for a typical member of this group, chlorine, is
There is only one electron needed to complete an octet. Of course, chlorine can actually gain this electron when it becomes and ion. When chlorine combines with any other nonmetal, the transfer of electrons is not energetically favourable. Therefore, in forming such compounds as HCl or Cl2, chlorine gets the one electron it needs by forming a covalent bond.
               ..                   ..              ..
     H· + :Cl: ------> H:Cl: --> H-Cl:
               .                    ..             ..
Multiple Bonds
The bond produced by the sharing of one pair of electrons between two atoms is called a single bond. This first bond is designated a sigma bond. There are however many molecules in which more than one pair of electrons are shared between two atoms. For example, we that nitrogen, is diatomic, N2. Each N atom has the Lewis symbol of
and each nitrogen needs three electrons to complete its octet. When the N2 molecule is formed, each nitrogen atom shares three electrons with the other.
The result is called a triple bond. Notice that in the Lewis formula for the molecule, the three shared pairs of electrons are placed between the two atoms. We count all of these electrons as though they belong to both of the atoms. Each nitrogen therefore has an octet
           :N:::N: OR :NN:
In the triple bond above the middle bond is a sigma type bond and the other two are Pi bonds. The Pi bonds are slightly different because they have to be bent out of shape in order to connect properly. But more on this later.
Double bonds are also possible. eg. CO2
      ..           ..          ..            ..
   : O :: C :: O : or : O = C = O :
      ..           ..          ..            ..
When the Octet Rule Fails
Sometimes it is just impossible to write a Lewis structure in which all of the atoms in a molecule obey the octet rule. This happens most often when an atom forms more than four bonds. Examples are PCl5 and SF6, in which there are five P-Cl bonds and six S-F bonds, respectively. Since each covalent bond requires the sharing of a pair of electrons, P and S must exceed eight electrons in their outer shells. The Lewis formula of these two molecules are shown below.
Elements in period 2 such as carbon or nitrogen, never exceed an octet simply because their valence shell, having n=2, can hold a maximum of only 8 electrons. Elements in periods below period 2, however, sometimes do exceed an octet, because their valence shells can hold more than 8 electrons. For example, the valence shell for elements in period 3, for which n=3, can hold a maximum of 18 electrons, and the valence shell for period 4 elements can hold as many as 32 electrons.
In some molecules (but not many), an atom has less than an octet. The most common examples are compounds of beryllium and boron.
                ..               ..         ..          ..         ..
·Be· + 2 ·Cl: -----> :Cl··Be··Cl: --> :Cl-Be-Cl:
                ..              ..          ..          ..         ..
Note: there is only 4e- around Be
                                    :Cl:                  ..
                                      ·                   :Cl:
 ·            ..                ..    ·     ..          ..  ..   ..
·B· + 3 ·Cl: ------> :Cl··B··Cl: --> :Cl-B-Cl:
             ..                ..        ..           ..       .. 

How to Write Lewis Structures
Step 1: Decide which atoms are bonded.
Step 2: Count all valence electrons
Step 3: Place two electrons in each bond.
Step 4: Complete the octets of the atoms attached to the central atom by adding e-'s in pairs.
Step 5: Place any remaining electrons on the central atom in pairs.
Step 6: If the central atom does not have an octet, form double bonds. If necessary form triple bonds.
Go to the Molecular Architecture Lewis Structures Worksheet
Resonance: When Lewis Structures Fail
There are some molecules and ions for which we cannot write Lewis structures that agree with experimental measurements of bond length and bond energy. One example is the formate ion, CHO2-, which is produced by neutralizing formic acid, HCHO2.
The skeletal structure for this ion is:
and, following the usual steps, we would write its Lewis structure as
This structure suggests that one carbon-oxygen bond should be longer than the other, but experimental evidence shows that they are in fact identical. In fact, the C-O bond lengths are about halfway between that expected for a single bond and that expected for a double bond. The Lewis structure doesn't match the experimental evidence, and there's no way to write one that does unless we use 1.5 electrons per bond.
We get around this problem with resonance. The Lewis structure can be shown as two structures
The term resonance is often misleading. The word itself suggests that the actual structure flip-flops back and forth between the two structures shown. This is not the case.
There is a simple way to determine when resonance should be applied to Lewis structures. If you find that you must move electrons to create a double bond while following the procedure, the number of resonance structures is equal to the number of equivalent choices for the location of the double bond.
eg 2. An SO3 molecule showing resonance structures. 
Go to the Molecular Architecture Worksheet #2
Coordinate Covalent Bonding
Sometimes one atom supplies both of the electrons that are shared in a covalent bond. When ammonia, NH3, is placed in an acidic solution, it picks up a hydrogen ion, H+, and becomes NH4+.
The H+ ion has a vacant valence shell that can accommodate two electrons. When the H+ is bonded to the nitrogen of NH3, the nitrogen donates both of the electrons to the bond. This type of bond, in which both electrons of the shared pari come from one of the atoms, is called a coordinate covalent bond. Even though we are making a distinction about where the electrons come from, once the NH4+ forms all four of the N-H bonds are identical.
Another example of a coordinate bond occurs when a molecule having an incomplete valence shell reacts with a molecule having valence shell electrons that aren't being used in bonding.
VSEPR Theory:The Shapes of Molecules
VSEPR Theory Tutorial  -  an excellent on-line tutorial from Purdue University
One of the important ways that molecular compounds differ from ionic compounds is in their structures. Ionic bonding is nondirectional in the sense that, at a given distance, an ion will attract others of the opposite sign with the same force regardless of where these other ions are located around the ion in question - there is no preferred direction. Because of this, the way ions arrange themselves in an ionic solid is determined just by the tendency to maximize attractions between ions of opposite charge and to minimize repulsions between like-charged ions.
In molecular compounds, quite a different situation exists. Covalent bonds are highly directional. That is, for a given central atom in a molecule there are preferred orientations for the atoms attached to it, because covalent bonds are not formed with equal ease in all directions. As a result, in polyatomic molecules the atoms remain in the same relative orientation, regardless of whether the substance is a solid, liquid or a gas and we can say that the molecules have a definite structure or shape.
The shapes of molecules are very important because many of their physical and chemical properties depend upon the three-dimensional arrangements of their atoms. For example, the functioning of enzymes, which are substances that control how fast biochemical reactions occur, requires that there be a very precise fit between one molecule and another. Even slight alterations in molecular geometry can destroy this fit and deactivate the enzyme, which in turns prevents the biochemical reaction involved from occurring. Nerve poisons work this way.
There is a very simple theory that is remarkably effective in predicting the shapes of molecules formed by the representative elements. It is called the valence shell electron pair repulsion theory (VSEPR theory, for short). The theory is based on the idea that valence shell electron pairs, being negatively charged, stay as far part from each other as possible so that repulsions between them at a minimum.
Consider the BeCl2 molecule. We've seen that its Lewis structure is
      ..       ..
     ..        ..
But how are these atoms arranged? Is BeCl2 linear or is it nonlinear, that is, do the atoms lie in a straight line, or do they form some angels less than 180o?
According to VSEPR theory, we can predict the shape of a molecule by looking at the electron pairs in the valence shell of the central atom. For BeCl2, there are two pairs of electrons around the central beryllium atom. The question is, how can they locate themselves to be as far apart as possible? The answer is that the minimum repulsions will occur when the electron pairs are on opposite sides of the nucleus. (180o apart).
In order for the electrons to be in the Be-Cl bonds, the Cl atoms must be placed where the electrons are; the result is that we predict that a BeCl2 molecule should be linear.
Another example will be the BCl3 molecule. Its Lewis structure is
    ..  ..  ..
    ..     ..
Here, the central atom, B, has three electron pairs. What arrangement will lead to minimum repulsions? The electron pairs will be as far apart as possible when they are located at the corners of a triangle with the boron in the centre.
When we attach the Cl atoms we obtain a triangular molecule.
This shape is planar triangular, because all four atoms lie in the equatorial plane.
In some molecules, all of the central pairs of the central atom are not bonding pairs. These unbonded pairs are actually unshared electron pairs - also called lone pairs - and these affect the geometry of the molecule. An example is SnCl2.
     ..  ..   ..
   ..        ..
There are three pairs of electrons around the central Sn atom - the two in the bonds plus the lone pair. As in BCl3 , the mutual repulsions of the three pairs will place them at the corners of a triangle.
Adding on the two chlorine atoms gives
This is not a triangular shape, even though that is how the electron pairs are arranged. Molecular shape describes the arrangements of atoms, not the arrangement of electron pairs. Therefore, this shape is described as being bent.
Determining the Shape of a Molecule
The best way to determine the architecture of a molecule is to:
1. Determine what the central atom is.
2. Draw the Lewis structure of the molecule.
3. Determine the number of bonding pairs and lone pairs around the central atom.
4. Refer to the following chart and determine the shape of the molecule.
VSEPR CHART     A = central atom   X = bonding pair   E = lone pair
# of 
# of 
Name Diagram
2 0  AX2 Linear 
2 2 AX2E2 Bent 
2 3 AX2E3 Linear 
3 0        AX3 Planar triangular
3 1 AX3 Trigonal pyramidal
3 2 AX3E2  T-shaped 
4 0  AX4 Tetrahedral 
4 1 AX4 See-saw
4 2  AX4E2 Square planar 
5 0 AX5 Trigonal bipyramidal 
5 1 AX5 Square pyramidal
6 0 AX6  Octahedral 
Shapes of Molecules and Ions with Double or Triple Bonds
The presence of double or triple bonds does not complicate matters at all. In a double bond, both electron pairs must stay together between the two atoms; they can't wander off to different locations in the valence shell. This is also true for the three pairs of electrons in a triple bond. For the purposes of predicting molecular geometry, then, we can treat double and triple bonds just as we do single bonds. For example, the Lewis formula for CO2 is
         ..         ..
         ..         ..
The carbon atom has no lone pairs. Therefore, the two groups of electron pairs that make up the double bonds are located on opposite sides of the nucleus and a linear molecule is formed. Similarly, you should be able to predict the following shapes for SO2 and SO3
Go to the Molecular Architecture VESPR Structures Worksheet

Polar Bonds and Electronegativity
When two identical atoms form a covalent bond, as in H2 or Cl2, each has an equal share of the electron pair in the bond. The electron density at both ends of the bond is the same, because the electrons are equally attracted to both nuclei.
When different kinds of atoms combine, as in HCl, the attractions usually are not equal. Generally one of the nuclei attracts the electrons more strongly than the other.
The effect of unequal attractions for the bonding electrons is an unbalanced distribution of electron density within the bond. It has been found that a chlorine atom attracts the electrons more strongly than does a hydrogen. In the HCl molecule, therefore, the electron cloud is pulled more tightly around the Cl, and that end of the molecule experiences a slight buildup of negative charge. The electron density that shifts toward the chlorine atom is removed from the hydrogen, which causes the hydrogen end to acquire a slight positive charge.
In HCl the electron transfer is incomplete. The electrons are still shared, but unequally. The charges on either end of the molecule are less than full +1 and -1 charges; they are partial charges, normally indicated by the lowercase Greek letter delta.
A bond that carries partial positive and negative charges on opposite ends is called a polar bond, or polar covalent bond. The term polar comes from the notion of poles of opposite charge at either end of the bond. Because there are two poles of charge involved, the bond is said to be a dipole.
The polar bond in HCl causes the molecules as a whole to have opposite charges on either end, so we say that HCl is a polar molecule. The HCl molecule as a whole is also a dipole.
The degree to which a covalent bond is polar depends on the relative abilities of bonded atoms to attract electrons. The term that we use to describe this relative attraction of an atom for the electrons in a bond is called the electronegativity of the atom. In HCl the chlorine is more electronegative than hydrogen. The electron pair of the covalent bond spends more of its time around the more electronegative atom, which is why that end of the bond acquires a partial negative charge.
The concept of electronegativity has been put on a quantitative basis - as numerical values have been assigned for each element. Refer to your periodic table for these assigned values. This information is useful because the difference in electronegativity values provides an estimate of the degree of polarity of a bond. In addition, the relative magnitudes of the electronegativities indicate which end of the bond carries the negative charge. For instance, fluorine is more electronegative than chlorine. Therefore, we would expect an HF molecule to be more polar than an HCl molecule. In addition, hydrogen is less electronegative than either fluorine or chlorine, so in both of these molecules the hydrogen bears the positive charge.
     ..       ..
H-F: H-Cl:
    ..       ..
+   -     +    -
The concept of electronegativity shows that there is no sharp dividing line between ionic and covalent bonding.
Ionic bonding and nonpolar covalent bonding simply represent the two extremes. Ionic bonding occurs when the difference in electronegativity between two atoms is very large; the more electronegative atom acquires essentially complete control of the bonding electrons. In a nonpolar covalent bond, there is no difference in electronegativity, so the pair of bonding electrons is shared equally.
The degree of polarity, which is the amount of ionic character of the bond varies in a continuous way with changes in the electronegativity difference.   The bond becomes more than 50% ionic when the electro-negativity difference exceeds 1.7
Within the periodic table, electronegativity varies in a more of less systematic way, and the trends follow those for the ionization energy (IE). Atoms with large ionization energies also have large electronegativities. An atom that has a small IE will give away an electron more easily than an atom with a large IE, just as an atom with a small electronegativity will lose its share of an electron pair more readily than an atom with a large electronegativity.
Molecular Shapes and the Polarity of Molecules
One of the main reasons you should be concerned about the polarity of molecules is because many physical properties, such as melting point and boiling point, are affected by it. This is because polar molecules attract each other - the positive end of one polar molecule attracts the negative end of another.
The strength of the attraction depends both on the amount of charge on either end of the molecule and on the distance between the charges. There are many molecules that are not dipoles even though they contain bonds that are polar. The reason for this is that can be seen if you look at the key role that molecular structure plays in determining molecular polarity.
Lets start with the H-Cl molecule, which has only two atoms and therefore only one bond. This bond is polar, and opposite ends of the bond carry partial charges of opposite sign. Because there are only two atoms in this molecule, which are located at the ends of the bond, the molecule as a whole has ends with equal but opposite charges. A molecule with equal but opposite charges on opposite poles is polar, so HCl is a polar molecule. In fact, any molecule composed of just two atoms that differ in electronegativity must be polar.
For molecules that contain more than two atoms, consider the combined effects of all the polar bonds. Sometimes, when all the atoms attached to the central atom are the same, the effects of the individual polar bonds cancel and the molecule as a whole is non-polar. In the diagram below the +---> sign is used to show the direction of the dipole. The arrow head points to the negative end of the bond and the crossed end is the positive end. 
In CO2 both bonds are identical, so each bond dipole is of the same magnitude. Because CO2 is a linear molecule, these bond dipoles point in opposite directions and work against each other. The net result is that their effects cancel, and CO2 is a non-polar molecule. Although it is not as easy to visualize, the same thing also happens in BCl3 and CCl4. In each of these molecules the effects of one of the bond dipoles is cancelled by the effects of the others.
If all the atoms attached to the central atom are not the same, or if there are lone pairs in the valence shell of the central atom, the molecule will usually be polar.
Not every structure that contains lone pairs on the central atom produces polar molecules. The following are two exceptions.

Go to the Molecular Architecture Polarity Worksheet

Hybrid Orbitals
The mixing of orbitals of different energy levels gives rise to new orbital types of intermediate energy that are able to form stronger bonds.
The approach that we have taken so far has worked well with some simple molecules. their shapes are explained very nicely by the overlap of simple atomic orbitals. It does note take long to find molecules with shapes and bond angles that fail to fir the model that has been developed. For example, CH4, has a shape that the VSEPR theory predicts to be tetrahedral. The H-C-H bond angles in this molecule are 109.5o. No simple atomic orbitals are orientated at this bond angle. So what kinds of orbitals are CH4 molecules using?
When some atoms form bonds, their simple s, p, and d orbitals often mix to form new atomic orbitals, called hybrid atomic orbitals. These new orbitals have new shapes and new directional properties. The reason for this mixing can be seen if we look at their shapes.
One kind of hybrid atomic orbital is formed by mixing a s orbital with a p orbital. This creates two new orbitals called sp hybrid orbitals (the sp is used to designate the kinds of orbitals from which the hybrid was formed).
Notice that each of the hybrid orbitals has the same shape - each has one large lobe and another much smaller lobe. The large lobe extends further from the nucleus than either the s or p orbital from which the hybrid orbital was formed. This allows the hybrid orbital to overlap more effectively with an orbital on another atom when a bond is formed. In general, the greater the overlap of two orbitals, the stronger the bond.
Another point to notice is that the large lobes of the two sp hybrid orbitals point in opposite directions - that is, they are 180o apart.
Let's look at a specific example, the linear beryllium hydride molecule, BeH2, as it would be formed in the gas phase.
The orbital diagram for the valence shell of beryllium is
Note that the 2s orbital is filled and the three 2p orbitals are empty. For bonds to form at a 180o angle between beryllium and the two hydrogen atoms, two conditions must be met:(1) the two orbitals that beryllium uses to for the Be-H bonds must be aligned oppositely at 180o, and (2) each of the beryllium orbitals must contain only one electron. The reason for the first requirement is obvious. The reason for the second is that each bond is that each bond must contain two electrons, one from the beryllium and one from the hydrogen. The net effect of all this is that when the Be-H bonds form, the electrons from the beryllium unpaired, and the resulting half-filled s and p atomic orbitals become hybridized.
Other Hybrid Orbitals
Hybrid orbitals can be formed by mixing more than just two simple atomic orbitals. If an s and two p orbitals combine, three hybrid orbitals, each similar in shape to the sp hybrids, are formed. they are called sp2 hybrid orbitals, the superscript 2 specifying the number of p orbitals taking part in the formation of the hybrids. Also, notice that the number of hybrids is the set is equal to the number of simple atomic orbitals from which the hybrids are formed.
BCl3 is a molecule in which the central boron atom uses sp2 hybrids for bonding. A boron atom has the valence shell configuration
To form three bonds, boron must have three half-filled orbitals, so its 2s electrons must become unpaired. The resulting half-filled s and p orbitals then become hybridized.
A chlorine atom has the valence shell configuration 1s2 2s2 2p5 , the half-filed 3pz orbital of each chlorine overlaps with one of the sp2 hybrids of boron to give the molecule.
These bonds are illustrated below. The geometry of the BCl3 molecules is planar triangular because all three hybrid orbitals fit nicely into the plane of the equator, 120o apart.
The list below shows important types of hybrid orbitals. The directional properties of the various hybrids are also shown.
sp  hybrids 
sp2 hybrids 
sp3 hybrids 
sp3d hybrids 
sp3d2 hybrids 
The bonds in the ethane molecule. Notice the overlap in the orbitals. The degree of overlap of the sp3 orbitals in the carbon-carbon bond does not appreciably affect the rotation of the two CH3- groups.
Using VSEPR Theory to Predict Hybridization
If we know the geometry of a molecule, we can predict the kind of hybrid orbitals that are used. Since VSEPR works so well in predicting geometry, we can use it to help us obtain information on the hybrid orbitals. For example, the Lewis structures of CH4 and SF6 are
In CH4, there are four electron pairs around carbon. The VSEPR model tells us that they should be arranged tetrahedrally. The only hybrid orbitals that are tetrahedral are sp3 hybrids, and we have seen that they explain the structure of this molecule well. Similarly, the VSEPR theory tells us that the six electron pairs around sulphur should be arranged octahedrally. The only octahedrally orientated hybrids are sp3d2 - the sulphur in the SF6 molecules must have these hybrids.

Go to the Molecular Architecture VSEPR Shapes Worksheet