I. Chemical Bonding. First we will review the two most important chemical bond types: Ionic and Covalent Bonding.
a. Ionic Bonding
Key: Transfer of electrons. Occurs in metal-non metal compounds. For example, in sodium fluoride NaF, that is used by the dentist to strengthen the enamel covering of the apatite lattice of your teeth, the sodium atom, 1s2 2s22p6 3s1 has transferred one electron, its 3s electron, to the fluorine atom, 1s2 2s22p5. In the process the sodium atom became an Na+ ion: electron configuration 1s2 2s22p6. The F atom became an F-ion with the electron configuration: 1s2 2s22p6.You notice that both of these ions have the stable noble gas configuration of neon. They are by far not as reactive (=more stable) as the Na atom and the F atom. Therefore NaF is a stable ionic compound. We do not draw Lewis structures for positive single ions, because the outer shell electron(s) is/are transferred, but we do draw a Lewis structure for the F- ion: an F with 8 electrons drawn as dots around it.
Try to write the electron configurations of calcium atoms and chlorine atoms, followed by the electron transfer when calcium chloride is formed. What noble gas configurations do the Ca2+ ion and the Cl- ion have respectively?Also, draw the Lewis structures of the two ions in calcium chloride.
b. Covalent Bonding. As opposed to transfer of electrons, we encounter here sharing of electrons. That is the key! This occurs typically in nonmetal-nonmetal compounds, such as N2, CO2, NO, S8, CH3OH. For those molecules we can draw Lewis structures. It also occurs in complex ions that are combinations of non metals, such as nitrate, NO3-; ammonium, NH4+; carbonate, CO32-; hydronium, H3O+. Try to draw the Lewis structures of these ions and label their covalent bonds, which represent the shared electrons.
To review Lewis structures click here.
In the molecules we encounter non polar or polar covalent bonds depending on the difference in electro-negativity. Read the following text between the purple horizontal lines for an explanation of the shapes of molecules and the bond angles in them. Also, electro-negativity is explained here and more important, how to apply it.The correct shape of molecules and ions is determined using the so called VSEPR theory, the vesper theory if you will. The valence shell electron pair repulsion, in other words the repulsion between all the electrons around the central atom, bonding pairs and lone pairs, will determine the shape. Recall, in the space around the central atom electrons tend to move as far apart as possible from each other. Therefore CO2 is a linear molecule with a 1800 bond angle and not a bent molecule with a 900 bond angle. Water, H2O is a bent molecule with a bond angle of 1050, because the two lone pairs and the two bonding pairs around oxygen repel each other into a bent shape. Here is an overview with many examples.
| Draw the Lewis structures of all the examples mentioned in the table below and check the repulsion of the lone pairs and bonding pairs around the central atom that leads to the correct shapes. |
|
Shape | Bond Angle |
Examples |
| linear | 1800 | N2; HBr; CO2; HCN |
| angular bent | < 1200 | H2O; SO2 |
| triangular planar | 1200 | NO3-; SO3; BH3;H2CO |
| tetrahedral | 109.50 | CH4; SiF4; SO42- |
| pyramidal | < 109.50 | NH3; H3O+ |
| trigonal bipyramidal | 900 and 1200 | PCl5 |
| octahedral | 900 | SF6; Fe(CN)6- |
Using the difference in electronegativity, ΔEN between two atoms on either side of a bond we are able to determine what type of bond it is: non polar covalent, polar covalent or ionic. We define this electronegativity difference as an absolute value. For example, for the bond between hydrogen and fluor in H-F it would be:
n.p.c. = non-polar covalent bond
This scale means that the C-H bond present in many organic compounds is essentially non-polar covalent, because the electronegativity difference, the ΔEN C-H = |2.55-2.10| = 0.45Non-polar molecules and polar molecules. Considering each polar covalent bond as a bond moment, a vector in the direction of the most electronegative atom with a magnitude of the ΔEN value, we are able to determine if a molecule is non-polar or polar. Let's first take a look at non-polar molecules. In CO2 ΔEN= 0.89 we have two bond moments starting at C and going in the direction of the most electronegative atom O. Since CO2 is a linear molecule the two vectors cancel each other, because they work in opposite directions and are equal in magnitude, both 0.89. When the vectors cancel each other the electrons are evenly distributed over the whole CO2 molecule and the molecule has therefore no dipole, does not contain two poles. It is a non-polar molecule. Other examples of non-polar molecules are N2, SO3, CCl4, C3H8.
In
review sessions, lab, and lecture we will show and often build models of these molecules, so that it is easier for you to check the cancellation of the bond moments.
In general you can say that symmetrical molecules are non-polar. Take the
trigonal planar sulfur trioxide, SO3 molecule. The three S-O bond
moments are perfectly canceling each other so that a non-polar molecule results.
The S-O double bond and S-O single bond exert the same bond moments, there is no
difference.
It is evident from the examples above that in polar molecules no cancellation of bond moments occurs and that therefore polar molecules contain a positive and a negative side, they contain two poles, and are therefore called di-polar molecules or simpler: polar molecules. Examples are the bent molecules H2O and SO2, the pyramidal molecule NH3, and the tetrahedral molecule CHCl3.
The consequence of all this is that non-polar carbon dioxide does not dissolve very well in polar water, making our soda's go flat real quick! SO2 on the contrary dissolves much better in water, because it is a polar molecule. The ground rule for good solubility is: "Like dissolves like".
Check this out on sticky gecko feet! Great stuff!
Between the molecules (in liquids and solids) we find the van der Waals forces (pure LDF or LDF + DDF) and hydrogen bonds (HB). Strength LDF + DDF: approximately 5-10 kJ and HB: 10-40 kJ! These forces determine if a compound is a gas, a liquid or a solid at room temperature. The greater the attraction force the greater the chance that the compound is a solid. Look at the examples hereunder and see if you understand their physical states, (g), (l) and (s) based on the strength of the attraction force between their molecules.
LDF, the London Dispersion Force, is determined by the number of electrons in the molecule and the shape of the molecule. It is an induced dipole force. The attraction occurs, because tiny dipoles are formed between moving electrons on neighboring atoms. The more electrons the stronger the attraction between the molecules.
DDF, the Dipole-Dipole Force, is determined by the presence of a permanent dipole moment.
Here are some examples to clarify this. Exclusively LDF occurs in non dipole molecules, also called non polar molecules: Ne(g); O2(g); SiF4(l); SO3(l); C3H8(g).
LDF + DDF attraction forces occur between all molecules that contain a dipole, also called polar molecules: SO2(g); CH3Cl(l); CH2O(l).
| Draw the Lewis structures of the examples above to check this out! |
Hydrogen bonds (=HB), by far the most important intermolecular force, occur only between compounds that contain an OH, an NH or an HF group, because the electronegativity of O, N and F is large enough to create a significant delta positive charge on the small hydrogen atom. Examples of molecules with intermolecular forces that are a sum of their LDF + DDF + HB's are: H2O(l); NH3(l); HF(l); CH3OH(l) (=methanol); C2H5OH (l)(=ethanol, the alcohol in beer, wine, whiskey).
Here is a good picture of the hydrogen bonds in water. Note the distances: An O-H bond in a water molecule is 100 pm, while a hydrogen bond between two water molecules is 180 pm. 1 pm = 1.0 x 10-12 m. (pm= pico meter)
