|Chapter Three, Section Two|
I. Bond Angles and Shapes
Here is a great site on a VSEPR tutorial. It does require the Chime plugin to visualize and manipulate the molecules. Go to this site to download Chime. It is free and will be a big help.
All diatomics are linear (2 points define a straight line) and all triatomics are planar (3 points define a plane). Beyond this basic information, we need another model. One model used is the Valence-Shell Electron-Pair Repulsion Theory (VSEPR) which assumes that electrons in bonds and unshared electron pairs repel one another so that they tend to stay as far apart as possible. So VSEPR model says that the resulting molecular structures are the result of the electron pairs trying to stay as far apart as possible. Remember that the predicted structure is determined by considering the paired and unpaired electrons, but the name of the type of structure does not include the lone pair electrons.
The other important item to help you in predicting molecular structures is the order of repulsion of the electron pairs:
lp-lp > lp-bp > bp-bp
Here is a table to help in predicting the structure of molecules considering the number of bond pairs and the number of lone pairs (Note that these are sets of electron pairs,, not numbers of electrons. Also a single bond, a double bond, and a triple bond are all one set of electron pairs.). No molecules with d-electrons are included in this table.
|Total shared sets||shared bonding sets||unshared bonding sets||shape||bond angle|
Examples of the first table entry (2,2,0) are HCN and CO2. Both of these molecules have multiple bonds, each of which counts as one set and they are both linear:
Here is HCN:
Here is CO2:
An example of the second table entry (3,3,0) is BH3:
An example of the third table entry (3,2,1) is SO2, which has one pair of nonbonding electrons:
An example of the fourth table entry (4,4,0) is CH4:
An example of the fifth table entry (4,3,1) is NH3, which has one pair of nonbonding electrons:
Finally, an example of the sixth table entry (4.2.2) is H2O, which has two pairs of nonbonding electrons:
You need to look at a number of examples to hopefully help you get comfortable with some predictions of molecular structures. Be sure to understand the examples in your text. Go to the referenced sites in "Assignments" and study those sites.
II. Polar and Nonpolar Molecules
As discussed in Section One, the Pauling scale of electronegativity is probably the most used scale and it is the one which is in your book. On this scale H has the value 2.1 and F has the value 4.0. F is the most electronegative element in the periodic table. So what does this have to do with the polarity of a bond? Good question. Now we define another variable given the symbol delta -- I will just use the word delta rather than the Greek symbol for which I have no HTML code. Delta is the difference in electronegativity values of the two atoms involved in the bond. The larger the delta value, the greater the polarity of the bond. For example:
Bond H---H S---H Cl---H O---H F---H EN 2.1 2.1 2.5 2.1 3.0 2.1 3.5 2.1 4.0 2.1 delta 0 0.4 0.9 1.4 1.9
As you see the delta increasing, the polarity of the bonds is increasing and the H-F bond is the most polar covalent bond.
Can a molecule containing polar bonds be a nonpolar molecule? Good question. If a molecule has a separtion of positive and negative charges in an unsymmetrical way then the molecule has a dipole moment and is said to be polar, like HF and water (water is not a linear molecule). the bonds in these molecules are also polar. But a molecule can have polar bonds and be a nonpolar molecule like carbon dioxide which is linear. Carbon dioxide looks like:
O = C = O
so that the partial negative charges on each oxygen are symmetrical about the partial negative charge on the carbon atom and the result is a nonpolar molecule. However the delta for the C-O bond is 1.0 so the bonds are polar.
Now consider water:
Water is a non-linear molecule and each bond is polar. So water is a polar molecule with polar bonds.
III. Naming Inorganic Compounds
A. Binary Ionic Compounds
1. Monatomic Ions
Here we are looking at ions combining with ions. Thus we are looking at metals (form cations) combining with nonmetals (form anions).
In these compounds the cation is named first. The monatomic cation has the name of the element and the anion has the root of the element name + ide. Some examples:
NaCl sodium chloride KCl potassium chloride KI potassium iodide CaS calcium sulfide CsBr cesium bromide MgO magnesium oxide Li3N lithium nitride
Why do we have 3 lithium atoms combining with 1 nitrogen atom to form the compound lithium nitride? Well, compounds have to have a neutral overall charge. Lithium forms a +1 cation (remember Group IA) and nitrogen tends to form a -3 anion. Thus to have a neutral charge we must have three +1 charges and one -3 charge. You will need to memorize Table 3.1 showing some common monatomic cations and anions (know the elements' names and charges).
Some common names are used for some compounds like H2O is water and NH3 is ammonia.
Unfortunately life is not even this simple. Some atoms form more than one type of cations. You will need to memorize the table below which shows these common cations and their names. Fortunately the names are easy. It used to be that the names were more complicated : the ion with the higher charge has a name ending in -ic, and the on with the lower charge has a name ending in -ous. You will still see this terminology in some literature, but we will use the Roman numeral method as below.
Note that mercury is a bit wierd. It's ionic +1 state is a diatomic.
So how do we name these ionic compounds whose cations can take on more than one form? We just use the naming system above for the particular cation. For example:
CuCl would be copper(I)chloride HgO would be mercury(II)oxide Fe2O3 would be iron(III)oxide (remember that the overall charge has to be zero so we need two +3 charged cations and three -2 charged anions.) FeO would be iron(II)oxide MnO2 would be manganese(IV)oxideDo not put the Roman numerals after cations which only form one type of ion.
2. Polyatomic ions
Now you have another table to memorize: Table 3.6 of common polyatomic ions. In that table you will see a number of ions which contain oxygen atoms. These are called oxyanions and their names depend upon the charge, kinda like above. However here we don't use the simpler Roman numeral system for stating which charge the ion has -- I don't know why not. Instead we use a system rather like the older method I talked about above. The polyatomic ion with the small number of oxygen atoms ends in -ite and the name of the one with the larger number of oxygen atoms ends in -ate. Examples would be:
SO32- and SO42- where the first would be called sulfite and the second would be called sulfate. When there are more than two oxyanions in the series, the prefix hypo (meaning less than) and per (meaning more than) are used. For example:
ClO-1 would be hypochlorite ClO2-1 would be chlorite ClO3-1 would be chlorate ClO4-1 would be perchlorateSo let's look at a few naming exercises combining some of the above.
Na2SO4 would be sodium sulfate Fe(NO3)3 would be iron(III)nitrate Mn(OH)2 would be manganese(II)hydroxide Na2CO3 would be sodium carbonate.B. Molecular Compounds
Now we are looking at the combining of two nonmetals to form a compound. The first element is named first. The second element is named as if it were an anion. Prefixes (mono, di, tri, etc) are used to denote the number of atoms present except mono is never used for the first element. Here are some examples:
N2O would be dinitrogen monoxide or nitrous oxide NO would be nitrogen monoxide or nitric oxide NO2 would be nitrogen dioxide N2O3 would be dinitrogen trioxide N2O4 would be dinitrogen tetroxide (not tetra oxide) N2O5 would be dinitrogen pentoxide (not penta oxide)
After you have studied this material and practiced some problems, take quiz two. If you score at least 80 on the test then you are ready to continue to the next section.
Web Author: Dr. Leon L. Combs
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