Any exceptions? Re: Shorter bond length always leads to stronger bond? Atoms that are closer together are bonded more strongly to each other, and those that are far apart have a weak bond.
Triple and double bonds are much stronger than single bonds because they are shorter, allowing the atoms to be closer together and the bonds are harder to break apart. The strength of a bond between two atoms increases as the number of electron pairs in the bond increases.
Generally, as the bond strength increases, the bond length decreases. Thus, we find that triple bonds are stronger and shorter than double bonds between the same two atoms; likewise, double bonds are stronger and shorter than single bonds between the same two atoms.
Average bond energies for some common bonds appear in Table 9. When one atom bonds to various atoms in a group, the bond strength typically decreases as we move down the group. The bond energy is the difference between the energy minimum which occurs at the bond distance and the energy of the two separated atoms. This is the quantity of energy released when the bond is formed.
Conversely, the same amount of energy is required to break the bond. For the H 2 molecule shown in Figure 5. This may seem like a small number. However, as we will learn in more detail later, bond energies are often discussed on a per-mole basis. For example, it requires 7. A comparison of some bond lengths and energies is shown in Figure 5. We can find many of these bonds in a variety of molecules, and this table provides average values.
For example, breaking the first C—H bond in CH 4 requires As seen in Table 9. We can use bond energies to calculate approximate enthalpy changes for reactions where enthalpies of formation are not available. Calculations of this type will also tell us whether a reaction is exothermic or endothermic. This can be expressed mathematically in the following way:. The bond energy is obtained from a table like Table 9. Thus, in calculating enthalpies in this manner, it is important that we consider the bonding in all reactants and products.
Because D values are typically averages for one type of bond in many different molecules, this calculation provides a rough estimate, not an exact value, for the enthalpy of reaction. Cited by. Download options Please wait Supplementary information PDF K. Article type Communication. Submitted 16 Jun Accepted 23 Jun First published 27 Jun Download Citation. Request permissions. 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. Structures 1 and 2 are resonance structures of each other, but structure 2 is the lower energy structure, even though it violates the octet rule.
Sulfur can accommodate more than eight electrons, and the formal charges in structure 2 are all zero. Lone pairs go in the equatorial positions, since they take up more room than covalent bonds. The Lewis structures of the previous examples can be used to predict the shapes around their central atoms:.
Formula Lewis Structure Bonding Shape 1. CH 4 4 bonds. NH 3 3 bonds. HCN 2 bonds. CO 2 2 bonds. CCl 4 4 bonds. COCl 2 3 bonds. O 3 2 bonds. CO 3 2- 3 bonds. O: bent BF 3 3 bonds.
NO linear PCl 5 5 bonds. SF 6 6 bonds. SF 4 4 bonds. XeF 4 4 bonds. O: bent With Lewis structures involving resonance, it is irrelevant which structure is used to determine the shape, since they are all energetically equivalent. Polar and Nonpolar Covalent Bonds Electronegativity is a measure of the ability of an atom in a molecule to attract shared electrons in a covalent bond.
Electronegativity is a periodic property, and increases from bottom to top within a group and from left to right across a period:. When two atoms of the same electronegativity share electrons, the electrons are shared equally , and the bond is a nonpolar covalent bond — there is a symmetrical distribution of electrons between the bonded atoms.
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