Electron bonds are not the same, since each electron in a molecule has its own peculiar energy characteristics, dependent on the electrical fields that surround it from other electrons and nearby nuclei.
We can characterise bonds within a molecule by three properties that reflect the number of bonding electrons involved in each bond, including resonance bonds, the distance between nuclei, and the amount of work required to pull the bond apart.
Bond order is simply the number of electron pairs shared by any two atoms in a molecule. Suppose that you have a simple water molecule H-O-H. There is one pair of electrons shared between each hydrogen and the central oxygen, so the bond order of each bond is 1 pair per bond, or 1. In carbon monoxide, there is a double bond between the carbon and oxygen: C=O, so the bond order is 2.
When we calculate the bond order for resonances, we need to count up the number of bonds in which the resonance can occur and divide by the total electron pairs involved. For example, in sulfur dioxide, the bonds can move from O-S=O to O=S-O to O=S=O. In the 3-pair sharing, the order is 3 pairs of shared electrons to two pairs of bonded atoms, or 3:2 = 1.50. In the four-pair version, the bond order is 4:2 or 2.
We've already seen how this works: bond length is simply the distance between nuclei in a bond. Length will be greater if the atoms are larger, and smaller if the bonds are multiple. Determining exact bond length is a matter of experiment, but estimating relative bond length is possible from trends in atomic size and bond order.
We can always think of energy in terms of work. Bond energy is the amount of work we need to do to pull the atoms in the bond completely apart. But this is simply the enthalpy of the reaction to form the molecule in the first place! So bond energy is our old friend, enthalpy of formation, or delta_Ho, which is the difference between the enthalpy of formation of the reactants, which we will have to break apart, and the enthalpy of formation of the products, which we are putting together. Breaking bonds is always endothermic, forming bonds from atoms and radicals is always exothermic. Whether a complex rearrangement of atoms in molecules is endothermic or exothermic depends on the final state of the bonds. If we have to spend more energy breaking bonds than forming the new ones, the reaction will be endothermic.
As we continue with our investigation of electron bonding between atoms, keep in mind the fundamental principles of all electrical force, as demonstrated by the Coulomb formula:
Fe = kq1q2 / r2
The force will be positive and repulsive if the charges are both negative or both positive; it will be negative and attractive if one of the charges is positive and the other is negative. So the electrons will repel each other in their own atom and in other atoms, but be attracted to their own positive nuclei and other positive nuclei or areas of net positive charge, such as unprotected side of a hydrogen atom in a covalent bond.
All chemical bonds result from a balance where the net attractive charge outweighs the net repulsive charge, so that we have to put energy into the system (or, to put it another way, do work) in order to pull the electrons out of their position. Within any molecule, the electrons will try to assume a symmetrical distribution as they repel one another. Where a completely symmetrical distribution occurs in a molecule with the same number of protons and electrons, the molecule is not only neutral overall, but has permanent net areas of charge concentration. This condition is rarely met, however, because most of the time, the molecule contains atoms of different electronegativities, so that the attractive forces are not equally balanced. This results in local areas of negative and positive charge, or dipoles, within the molecule, and these dipoles affect both the shape and the chemical properties of the molecule.
For the moment, we consider only covalent molecules, which can be isolated from one another while remaining a molecular substance. The special problems of ionic bonds in lattice solids we will look at later.
We use the concept of formal charge to help us keep track of where electrons are in the molecule. An atom has a positive formal charge if it contributes more electrons to the bonds than it gets back by sharing, and it has a negative formal charge if it gets more electrons by sharing than it contributes to the bonds. Lone pairs belong to their own atom, shared electrons are split between the bonded atoms. We count up all the valence electrons for the atom in its isolated, unbonded state, and compare this to the electrons it has when bonded in the molecule.
Consider, for example, our old friend SO2, a resonance molecule with one possible structure being
The rightmost oxygen has two lone pairs, and half shares in a double bond, for a total of 6 electrons, which is equal to its normal valence shell electron count. So the formal charge on this oxygen is 0.
The sulfur has one lone pair and half-interest in three shared pairs, for a total of 5 electrons. This is one less than its normal valence count of 6 electrons, so the formal charge on the sulfur is +1.
The leftmost oxygen has three lone pairs and half-interest in one shared pair, for a total of 7 electrons, one over its normal valence count of 6, so its formal charge is -1.
The total of the formal charges 0 + 1 + (-1) = 0, so the net charge on the molecule is also 0.
Electrons do not spend their bond equidistant from the nuclei of the atoms they are bonding unless the atoms are the same element. The electrons in the O2 bond are equally attracted by each nuclei, so this bond is truly covalent. But in any case where the atoms in the bond are not the same element, the one with the greater electronegativity will pull the electrons in the bond closer. [Where the electronegativity difference is great enough, the stronger element will detach the electron entirely from the weaker element, forming an ionic bond.] Any compound covalent bond between two dissimilar elements which is not actually ionic will therefore be polar. This can affect the shape of the molecule and its overall polarity, depending on the nature of the other bonds in the molecule.
There are some simple rules to memorize for molecular shapes. These are determined by a theory called VSEPR, or the valence shell electron pair repulsion theory, which is just the formal way of saying that electrons will distribute themselves as evenly and symmetrically as possible in any configuration, in response to their mutual repulsion.
The actual shape of the molecule depends on whether all available pairs are part of bonds, or whether there are lone pairs to repel the bonding pairs. If all the electrons of a central atom are involved in bonds, the terminal atoms will be symmetrically distributed. But if the central atom has lone pairs, these will take the place of a virtual terminal atom, and the actual terminal atoms will not be symmetrically distributed.
|Number of electron pairs||General arrangement||Lone pairs||Net Bonding Pairs||Molecular shape|
|3||trigonal (flat triangle)||0||3||linear|
|4||tetrahedral (three-sided pyramid)||0||4||tetrahedral|
|5||trigonal bipyramidal||0||5||trigonal bipyramidal|
Another way to think of this is that all positions are filled by an electron pair, but not all available positions are filled by a bonded atom. So in a situation where there are six pairs available, but two of the pairs are lone pairs, the lone pairs will take up positions on opposite sides of their central atom, and the remaining four pairs will get the four bonding atoms in a plane between the two lone pairs.
To determine whether a molecule is polar, you have to look at both the presence of polar bonds (where the electronegativity of the atoms involved is different) and the distribution of these bonds. It is possible for individual bonds to be polar, but the overall molecule to have a symmetrical distribution, as well as for an asymmetrical distribution to result in a polar molecule.
In the case of CCl4, the Cl pulls the electrons more tightly than the C, but since the molecule is symmetrical, with each Cl at a point of the tetrahedron, equidistant from the central C and from each other as much as possible, the electrons are still symmetrically distributed around the central C. Similarly, BF3 is trigonal planar, so the fluorine atoms are symmetrically distributed and their polar bonds with the B cancel out overall.
In the case of PCl3, however, the lone pair on the phosphorus makes the molecule trigonal pyramidal, and the lone pair takes the fourth position of a tetrahedron with the other points occupied by the chlorine atoms. The polar bonds between the phosphorus and each chlorine do not cancel out, and the molecule is polar —and that will be true of all trigonal pyramidal molecules.
Ions can be non-polar if the arrangement of atoms within the molecule cancels out any internal polar bonds. Some examples are NO3- and PO4-3. In these ions, the extra electrons fill out the bonding requirements to complete the octets, but do so in such a way that the charges are evenly distributed within the molecule. There is no internal dipole, but seen from outside, there is still a net charge, since there are more electrons than protons in the molecule.
Why do we care? Well, net polarity allows molecules to form temporary and weak but still effective bonds with other molecules. The net local charges hold sugar molecules together in sugar cubes. The fact that these local net charges are relatively weak accounts for the low melting points of sugar, making it possible to liquify corn syrup and dextrose on the kitchen stove and recombine them as they cool into Christmas candies.
In some cases, the polar nature of the molecular has significant consequences. The polarity of the water molecule creates bonds that provide for surface tension on water: boats float more easily, insects walk on it. The hydrogen bonds formed between the "exposed" side of the hydrogen atoms and the net negative side other water molecules where the oxygen atom resides means that more energy is necessary to pull water molecules apart, to heat them up, and to cause water to boil. This is why water has such a relatively high specific heat for a compound that is liquid at room temperature. This high heat capacity makes possible life on earth as we know it, since large bodies of water retain heat, making coastal regions more temperate.
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