Because the electron bones in our analogy have a negative charge, the puppy thief becomes negatively charged due to the additional bone. The puppy that lost its electron bone becomes positively charged.
Because the puppy who lost his bone has the opposite charge of the thief puppy, the puppies are held together by electrostatic forces, just like sodium and chloride ions!
In our analogy, each puppy again starts out with an electron bone. Some covalently bonded molecules, like chlorine gas Cl2 , equally share their electrons like two equally strong puppies each holding both bones. Other covalently bonded molecules, like hydrogen fluoride gas HF , do not share electrons equally. The fluorine atom acts as a slightly stronger puppy that pulls a bit harder on the shared electrons see Fig. Even though the electrons in hydrogen fluoride are shared, the fluorine side of a water molecule pulls harder on the negatively charged shared electrons and becomes negatively charged.
The hydrogen atom has a slightly positively charge because it cannot hold as tightly to the negative electron bones. Covalent molecules with this type of uneven charge distribution are polar. Molecules with polar covalent bonds have a positive and negative side. In this analogy, each puppy represents an atom and each bone represents an electron. Water H2O , like hydrogen fluoride HF , is a polar covalent molecule.
When you look at a diagram of water see Fig. 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 In general, if the electronegativity difference between two atoms is less than 0.
Nonpolar molecules also form when atoms sharing a polar bond arrange such that the electric charges cancel each other out. Examples of nonpolar molecules include:. If you know the polarity of molecules, you can predict whether or not they will mix together to form chemical solutions. The general rule is that "like dissolves like", which means polar molecules will dissolve into other polar liquids and nonpolar molecules will dissolve into nonpolar liquids.
This is why oil and water don't mix: oil is nonpolar while water is polar. It's helpful to know which compounds are intermediate between polar and nonpolar because you can use them as an intermediate to dissolve a chemical into one it wouldn't mix with otherwise.
For example, if you want to mix an ionic compound or polar compound in an organic solvent, you may be able to dissolve it in ethanol polar, but not by a lot. Then, you can dissolve the ethanol solution into an organic solvent, such as xylene. Actively scan device characteristics for identification. Use precise geolocation data.
Select personalised content. Create a personalised content profile. Measure ad performance. Select basic ads. Create a personalised ads profile. Dipoles that are directly opposite one another cancel each other out. What electronegativity difference would indicate a polar bond? Is a molecule with symmetric polar bonds a polar molecule?
Review What is a dipole? How does shape affect the polarity of a molecule? What is the difference between a polar bond and a polar molecule? Show References References Courtesy of G. CK Foundation — Christopher Auyeung. CK Foundation — Zachary Wilson. CK Foundation — Joy Sheng. Licenses and Attributions. CC licensed content, Shared previously.
0コメント