Polar and Nonpolar Covalent Bonds H. Molecular Shape and Polarity I. Covalent Bonds and Lewis Structures When elements combine, there are two types of bonds that may form between them:. Lewis theory Gilbert Newton Lewis, focuses on the valence electrons, since the outermost electrons are the ones that are highest in energy and farthest from the nucleus, and are therefore the ones that are most exposed to other atoms when bonds form.
Lewis dot diagrams for elements are a handy way of picturing valence electrons, and especially, what electrons are available to be shared in covalent bonds. The valence electrons are written as dots surrounding the symbol for the element: one dot is place on each side first, and when all four positions are filled, the remaining dots are paired with one of the first set of dots, with a maximum of two dots placed on each side.
Lewis-dot diagrams of the atoms in row 2 of the periodic table are shown below:. Unpaired electrons represent places where electrons can be gained in ionic compounds, or electrons that can be shared to form molecular compounds. The valence electrons of helium are better represented by two paired dots, since in all of the noble gases, the valence electrons are in filled shells, and are unavailable for bonding. Covalent bonds generally form when a nonmetal combines with another nonmetal.
Both elements in the bond are attracted to the unpaired valence electrons so strongly that neither can take the electron away from the other unlike the case with ionic bonds , so the unpaired valence electrons are shared by the two atoms, forming a covalent bond :. The shared electrons act like they belong to both atoms in the bond, and they bind the two atoms together into a molecule. The shared electrons are usually represented as a line — between the bonded atoms.
In Lewis structures, a line represents two electrons. Atoms tend to form covalent bonds in such a way as to satisfy the octet rule , with every atom surrounded by eight electrons. Hydrogen is an exception, since it is in row 1 of the periodic table, and only has the 1 s orbital available in the ground state, which can only hold two electrons.
The shared pairs of electrons are bonding pairs represented by lines in the drawings above. The unshared pairs of electrons are lone pairs or nonbonding pairs. All of the bonds shown so far have been single bonds , in which one pair of electrons is being shared. It is also possible to have double bonds , in which two pairs of electrons are shared, and triple bonds , in which three pairs of electrons are shared:.
Examples 1. This uses up all of the valence electrons. The octet rule is satisfied everywhere, and all of the atoms have formal charges of zero. This uses up six of the eight valence electrons. All of the valence electrons have now been used up, the octet rule is satisfied everywhere, and all of the atoms have formal charges of zero. This uses up four of the valence electrons. This uses up six of the valence electrons. The remaining two valence electrons must go on the oxygen:.
All of the valence electrons have been used up, and the octet rule is satisfied everywhere. The remaining six valence electrons start out on the N:. The octet rule can be satisfied if we move two pairs of electrons from the N in between the C and the N, making a triple bond:. This uses up the sixteen valence electrons The octet rule is not satisfied on the C, and there are lots of formal charges in the structure:.
The octet rule can be satisfied, and the formal charges diminished if we move a pair of electrons from each oxygen atom in between the carbon and oxygen atoms:. Now, all of the valence electrons have been used up, the octet rule is satisfied everywhere, and all of the atoms have formal charges of zero. Place the remaining valence electrons on the O and Cl atoms:.
Making a carbon-chlorine double bond would satisfy the octet rule, but there would still be formal charges, and there would be a positive formal charge on the strongly electronegative Cl atom structure 2. Making a carbon-oxygen double bond would also satisfy the octet rule, but all of the formal charges would be zero, and that would be the better Lewis structure structure 3 :.
Examples continued from section B 9. We can satisfy the octet rule on the central O by making a double bond either between the left O and the central one 2 , or the right O and the center one 3 :. In this example, we can draw two Lewis structures that are energetically equivalent to each other — that is, they have the same types of bonds, and the same types of formal charges on all of the structures. The actual molecule is an average of structures 2 and 3 , which are called resonance structures.
Structure 1 is also a resonance structure of 2 and 3 , but since it has more formal charges, and does not satisfy the octet rule, it is a higher-energy resonance structure, and does not contribute as much to our overall picture of the molecule.
The real molecule does not alternate back and forth between these two structures; it is a hybrid of these two forms. The ozone molecule, then, is more correctly shown with both Lewis structures, with the two-headed resonance arrow between them:.
In contrast, the lone pairs on the oxygen in water are localized — i. Resonance delocalization stabilizes a molecule by spreading out charges, and often occurs when lone pairs or positive charges are located next to double bonds. Resonance plays a large role in our understanding of structure and reactivity in organic chemistry.
A more accurate picture of bonding in molecules like this is found in Molecular Orbital theory, but this theory is more advanced, and mathematically more complex topic, and will not be dealt with here.
Examples We can satisfy the octet rule and make the formal charges smaller by making a carbon-oxygen double bond. Once again, structure 1 is a resonance structure of 2 , 3 , and 4 , but it is a higher energy structure, and does not contribute as much to our picture of the molecule. Multi-Center Molecules Molecules with more than one central atoms are drawn similarly to the ones above. The octet rule and formal charges can be used as a guideline in many cases to decide in which order to connect atoms.
C 2 H 6 ethane C 2 H 4 ethylene The octet rule is not satisfied on the B, but the formal charges are all zero. In fact, trying to make a boron-fluorine double bond would put a positive formal charge on fluorine; since fluorine is highly electronegative, this is extremely unfavorable. In this structure, the formal charges are all zero, but the octet rule is not satisfied on the N.
Since there are an odd number of electrons, there is no way to satisfy the octet rule. Nitric oxide is a free radical, and is an extremely reactive compound. In the body, nitric oxide is a vasodilator, and is involved in the mechanism of action of various neurotransmitters, as well as some heart and blood pressure medications such as nitroglycerin and amyl nitrite.
Notice that the formal charge on the phosphorus atom is zero. Notice that the formal charge on the sulfur atom is zero. Notice that the formal charge on the xenon atom is zero. Movies in 3D pop out at us. Before, we see movies that are just on the screen and that's good. What's better? For bond angles, 3D is better. Therefore, tetrahedrals have a bond angle of How scientists got that number was through experiments, but we don't need to know too much detail because that is not described in the textbook or lecture.
Using the example above, we would add that H 2 O has a bond angle of A molecule is polar when the electrons are not distributed equally and the molecule has two poles. The more electronegative end of the molecule is the negative end and the less electronegative end is the positive end. A common example is HCl.
Using the cross bow arrow shown below we can show that it has a net dipole. The net dipole is the measurable, which is called the dipole moment. Dipole moment is equal to the product of the partial charge and the distance. The equation for dipole moment is as follows. The units for dipole is expressed in debye which is also known as Coulombs x meter C x m. The cross base arrow demonstrates the net dipole.
On the cross-base arrow, the cross represents the positive charge and the arrow represents the negative charge. Here's another way to determine dipole moments. We need to comprehend electronegativity which is abbreviated EN. What is EN? Well, EN is how much an element really wants an electron. Think about basketball and how two players pass the ball to each other. Each player represent an element and the ball represents the electron. Let's say one player is a ball hog.
The player that is the ball hog is more electronegative because he or she wants the ball more. What if we are not given EN? Luckily, there is a trend in the periodic table for EN. From bottom to the top, EN will increase. From left to right, EN will increase. The most electronegative element is Flourine with 4. Now, we are ready to apply EN to determine whether or not molecules are polar.
We look back at the picture of H 2 O above. The EN is given. What do we do with all the EN? We compare the EN between each bond. Oxygen has a greater EN than Hydrogen. Therefore, we can draw a cross bow arrow towards Oxygen. We have two arrows because Oxygen is bonded to two Hydrogens. Since both arrows point toward Oxygen, we can say that there is a net EN. We added the arrows that point to Oxygen and we end up with a new, bigger arrow.
This is examplified in the picture above. Refer back to the Lewis dot diagram of CO 2. The shape is linear and the EN arrows point towards Oxygen. The arrows are opposite of each other and have the same EN difference.
Therefore, we have no net charge and the molecule is non-polar. To recap, when a molecule is polar it means that the electron is not distributed evenly and there is a difference in the electronegativity of the atoms. If a molecule is polar, it means that it had a net dipole which results in having a dipole moment. If the molecule has a net dipole, then it is polar. If the structure is symmetric, then it is non-polar C.
There are three rules to this part: 1. When there are no lone pairs on the center atom, then the molecule is non-polar 2. If it is linear or square planar, then it is non-polar.
This rule is more important than rule 1, so it overrules it because it has lone pairs. If it has different terminal atoms, then it is polar. This rule overrules rule 1 and 2 because it is more important. Draw the Lewis Structure and name the shape of each compound. Also determine the polarity and whether or not it has a dipole moment.
Introduction To determine the shapes of molecules, we must become acquainted with the Lewis electron dot structure. Valence-Shell Electron-Pair Repulsion Theory Now that we have a background in the Lewis electron dot structure we can use it to locate the the valence electrons of the center atom. Electron-group geometry is determined by the number of electron groups. However, the structures of some compounds and ions cannot be represented by a single formula.
For clarity the two ambiguous bonds to oxygen are given different colors in these formulas. If only one formula for sulfur dioxide was correct and accurate, then the double bond to oxygen would be shorter and stronger than the single bond. This averaging of electron distribution over two or more hypothetical contributing structures canonical forms to produce a hybrid electronic structure is called resonance.
Likewise, the structure of nitric acid is best described as a resonance hybrid of two structures, the double headed arrow being the unique symbol for resonance. The above examples represent one extreme in the application of resonance.
Here, two structurally and energetically equivalent electronic structures for a stable compound can be written, but no single structure provides an accurate or even an adequate representation of the true molecule. In cases such as these, the electron delocalization described by resonance enhances the stability of the molecules, and compounds or ions composed of such molecules often show exceptional stability.
The electronic structures of most covalent compounds do not suffer the inadequacy noted above. Nevertheless, the principles of resonance are very useful in rationalizing the chemical behavior of many such compounds. For example, the carbonyl group of formaldehyde the carbon-oxygen double bond reacts readily to give addition products. The course of these reactions can be explained by a small contribution of a dipolar resonance contributor, as shown in equation 3.
Here, the first contributor on the left is clearly the best representation of this molecular unit, since there is no charge separation and both the carbon and oxygen atoms have achieved valence shell neon-like configurations by covalent electron sharing. If the double bond is broken heterolytically, formal charge pairs result, as shown in the other two structures. The preferred charge distribution will have the positive charge on the less electronegative atom carbon and the negative charge on the more electronegative atom oxygen.
Therefore the middle formula represents a more reasonable and stable structure than the one on the right. The application of resonance to this case requires a weighted averaging of these canonical structures. The double bonded structure is regarded as the major contributor, the middle structure a minor contributor and the right hand structure a non-contributor. Since the middle, charge-separated contributor has an electron deficient carbon atom, this explains the tendency of electron donors nucleophiles to bond at this site.
These are the canonical forms to be considered, and all must have the same number of paired and unpaired electrons. The following factors are important in evaluating the contribution each of these canonical structures makes to the actual molecule. The stability of a resonance hybrid is always greater than the stability of any canonical contributor.
Consequently, if one canonical form has a much greater stability than all others, the hybrid will closely resemble it electronically and energetically. This is the case for the carbonyl group eq. On the other hand, if two or more canonical forms have identical low energy structures, the resonance hybrid will have exceptional stabilization and unique properties.
This is the case for sulfur dioxide eq. To illustrate these principles we shall consider carbon monoxide eq. In each case the most stable canonical form is on the left. For carbon monoxide, the additional bonding is more important than charge separation. Furthermore, the double bonded structure has an electron deficient carbon atom valence shell sextet. A similar destabilizing factor is present in the two azide canonical forms on the top row of the bracket three bonds vs. The bottom row pair of structures have four bonds, but are destabilized by the high charge density on a single nitrogen atom.
All the examples on this page demonstrate an important restriction that must be remembered when using resonance. No atoms change their positions within the common structural framework. Only electrons are moved. A more detailed model of covalent bonding requires a consideration of valence shell atomic orbitals. The spatial distribution of electrons occupying each of these orbitals is shown in the diagram below.
If this were the configuration used in covalent bonding, carbon would only be able to form two bonds. Very nice displays of orbitals may be found at the following sites: J.
Gutow, Univ. Wisconsin Oshkosh , R. Spinney, Ohio State and M. Winter, Sheffield University. These hybrid orbitals have a specific orientation, and the four are naturally oriented in a tetrahedral fashion.
Just as the valence electrons of atoms occupy atomic orbitals AO , the shared electron pairs of covalently bonded atoms may be thought of as occupying molecular orbitals MO. It is convenient to approximate molecular orbitals by combining or mixing two or more atomic orbitals.
In general, this mixing of n atomic orbitals always generates n molecular orbitals. The hydrogen molecule provides a simple example of MO formation.
The bonding MO is occupied by two electrons of opposite spin, the result being a covalent bond. The notation used for molecular orbitals parallels that used for atomic orbitals. In the case of bonds between second period elements, p-orbitals or hybrid atomic orbitals having p-orbital character are used to form molecular orbitals. For example, the sigma molecular orbital that serves to bond two fluorine atoms together is generated by the overlap of p-orbitals part A below , and two sp 3 hybrid orbitals of carbon may combine to give a similar sigma orbital.
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