Bonding
Our job in describing the bonding of a substance is to describe: 1) what atoms are connected to what other atoms 2) what types of bonds hold these atoms together 3) how much energy would it take to break the bonds 4) what is the bond length (what distance separates the two nuclei) 5) what is the shape of the molecule (what are the angles between each bond) 6) how is the electron density distributed around the molecule (polarity, reactivity) Most of these things can be measured empirically Bond energy can be measured by shining light of increasing frequency into a container of a substance in the gas state until the substance starts to break down. If more than one bond is present, one must determine from the identity of the products which bond has broken. Bond length and angle can be measured by shining electrons or X-rays through a crystalline sample of a substance and calculating the bond attributes from the diffraction pattern that results (makes more sense in you have taken physics). Bond angle and length also can be calculated by observing the response of the substance to microwave radiation or by seeing how difficult it is to orient the molecules in an electric field. (more about some of these methods later)Ionic Bonding
An ionic bond is the electrostatic attraction between two ions of opposite charge. The actual trading of electrons to form ions is not actually part of the ionic bonding process. Ionic bonds can form between ions as they precipitate out of a solution. Example: Ionization (e- trading): Na + Energy ---> Na+ + e- Cl + e- ---> Cl- + Energy Ionic Bond Formation: Na+ + Cl- ---> NaCl
The ionic formula gives only the simplest ratio of the ions involved.
In solid form, ionic substances form crystal lattices with alternating positive and negative charged ions
with a ratio and arrangement specific to the substance.
Molecular Bonding
Molecules are held together by covalent bonds. A covalent bond is the mutual attraction of two nuclei (atoms) for a shared pair of electrons. Molecular Lewis structures are constructed (in simple cases) by circling unpaired electrons between two atoms. The actual Lewis structure shows bonded electron pairs as a short line.
Co-ordinate covalent bonds form when one atom contributes both of the electrons in the bond. For this to
work, one of the atoms in the bond must have a lone pair and the other must have a high electronagativity
and be short two electrons.
Bond Order
Multiple bonds are more than just the sharing of more electrons. Because more electrons are being shared, there is a greater attraction between the atoms. For this reason, the atoms are drawn closer together, so the bond length is less. The bond energy, on the other hand, is greater. Table 1: Average Bond Lengths (pm) (averaged over a number of different molecules) ------------------------------------------------------------------------------------------------ Bond Length Bond Length Bond Length H-H 74.14 C-C 154 N-N 145 H-C 110 C=C 134 N=N 123 H-N 100 C=C 120 N=N 109.8 H-O 97 C-N 147 N-O 136 H-S 132 C=N 128 N=O 120 H-F 91.7 C=N 116 O-O 145 H-Cl 127.4 C-O 143 O=O 121 H-Br 141.4 C=O 120 F-F 143 H-I 160.9 C-Cl 178 Cl-Cl 199 Br-Br 228 I-I 266 ------------------------------------------------------------------------------------------------ If a bond between two different elements is not in the table, you can calculate it as in this example: the O-Cl bond is not listed, but O-O and Cl-Cl are at 145 pm and 199 pm respectively. The O-Cl bond length can be predicted as (145/2) + (199/2) = 172 pm Source: Petrucci, R., Harwood, W., & Herring, F.G., (2002), General Chemistry: Principles and Modern Applications, Prentice Hall ------------------------------------------------------------------------------------------------ You can see that for a given element, double bonds are shorter than single bonds. Triple bonds are even shorter. However, the difference between double and single bonds is greater than that between double and triple. We also can see from this chart, that the elements involved in the bond also have an effect on the bond length (and energy). This makes sense, since the nuclei of different elements have different charges and, thus, will have different attractions for the bonding electron pairs. Bond Order Bond order is the number of shared electron pairs in the bond: a single bond has a bond order of 1 a double bond has a bond order of 2 a triple bond has a bond order of 3 Why introduce yet another term when we already had the terms "single," "double," ect. The answer will be clear in a few days when we encounter substances in which the bonds act as if they are sharing fractions of an electron pair ("behave" you say? Well, for example, they have bond lengths that are inbetween that of the single and double bond).
Making Realistic Lewis Structures
I usually is possible to draw several valid Lewis structures for a molecular formula. How do you know which one(s) represent(s) the real configuration of the molecule as it exists in nature. In some cases, none of the seemingly valid structures representents the real molecule and the empirical data shows that the real structure is quite different. We could use evidence from bond lengths and angles (X-ray diffraction), decomposition products, or products from other chemical reactions. All of these are time consuming and expensive. Luckily, there are a few sets of rules, which if followed, usually result in a realistic structure.Skeletal Structure
The first step is to correctly arrange the atoms before you start bonding them together. After this is done, you can start connecting the atoms with bonds. Rules 1) With the exception of organic polymers, most molecules are compact and symmetrical. e.g. H3PO4 H H O O P O O H O H H O P O H O not realistic realistic 2) a) The central atom in a molecule usually is the atom with the lowest electronegativity and, less important, the greatest bonding capacity (usually the least electronegativity). e.g. in H3PO4, the Phosphorous atom is the best candidate to place at the center of the structure. b) Carbon always is a central atom. 3) Hydrogen never is a central atom, rather it is a terminal atom as it makes only one bond. (in some Boron hydride molecules, it seems to have more than one bond).
Placing Bonds
We went over the rules for bonding atoms in grade 11 chemistry. 1) Add up all the valence electrons of each atom in the molecule. Add one for each negative charge (e.g., add 2 for SO42-) Subtract one for each positive charge. 2) Place one bond (2 electrons) between each atom in the skeletal structure to join atoms with covalent bonds. The bonds should join atoms towards the center of the molecule. Ring-shaped structures are rare outside of organic molecules as well. Subtract the number of electrons used from the total. 3) Use the remaining electrons to add lone pairs where needed to all but the central atom. Subtract the electrons used from the total. 4) If any electrons remain, use them to add lone pairs to the central atom. 5) If any atom does not have a stable octet, promote a lone pair on a more peripheral atom to be a bond. 6) Change all bonded electron pairs to dashes and the Lewis structure is complete. In many cases, however, these two sets of rules can still produce more than one valid stucture. Which is the more likely to be realistic. One way to determine this is to look at how the electrons are distributed around the molecule. To do this, we can use a technique called formal charges.
Formal Charges
A formal charge is an apparent charge on an atom in a molecule as a result of unequal contributions of electrons to covalent bond electron pairs. They are not actual charges, they are more of a bookeeping system to keep track of electrons. To calculate the formal charge of an atom in a molecule: the number of valence the number of lone pair one half the number FC = electrons that the atom - electrons that the atom - of bond pair electrons has when on its own has in the molecule that the atom has in a (not the number of lone molecule (not the number pairs) of bonds) Examples: Oxygen with one lone pair and a triple bond
FC = 6 - 2 - .5(6) = 1+
Oxygen with three lone pairs and a single bond 
FC = 6 - 6 - .5(2) = 1-
The H3PO4 molecule from today's lesson.
The left-most oxygen: FC = 6 - 6 - .5(2) = 1-
The phosphorous: FC = 5 - 0 - .5(8) = 1+
Any other oxygen: FC = 6 - 4 - .5(4) = 0
Any hydrogen: FC = 1 - 0 - .5(2) = 0
Indicate formal charges on the structural diagram with charges inside small circles.
Formal charges help us decide when a particular structure is likely to occur or whether some other structure
for the same formula is the real structure for a compound. Thus, if one cannot decide between several
valid structures for a formula, determine the formal charges.
Formal Charge Rules
Natural molecules usually follow these rules. Once you have produced a molecule using the steps
outlined in class today, check it out using these rules of formal charges.
1) In a neutral molecule, the formal charges of the component atoms must add up to zero.
In a polyatomic ion, the formal charges of the component atoms must add up to the charge.
2) Formal charges should be as small as possible.
3) Negative formal charges are usually on the most electronegative atom.
4) Atoms with the same sign of formal charge are unlikely to be bonded together.
Example
One of these two valid structures for NO2+ is a better candidate than the other as it has
smaller formal charges. Which one is it?