CHEMICAL BONDS

CHEMICAL BONDS
Chemical bonds are formed when the electrons in an atom interact with the electrons in another atom. This allows for the formation of more complex molecules.

3 Types of Chemical bonds

*Covalent
These strong bonds form when two atoms share electrons.
Sometimes the electrons in an atom get shared. It's much like when you were a kid and got to sleep over at a friends house. Your friends parents were in charge of you both for one night and the next night you would sleep over at your house and your own parents would be in charge. This sharing of responsibility is functionally similar to the way covalent bonding works.

Normally this sharing is an equal proposition. Sometimes it's not equal (but that gets us into hydrogen bonding discussed below.)

*Ionic
Atoms gain or lose electron (opposites attract)
Ions have positive or negative charges. In dating situations, you may know that sometimes opposites attract. In Chemistry, opposites ALWAYS attract. This forms an ionic bond between two atoms.

*Hydrogen
Weakest bond between atoms
Occurs in molecules that have covalent bonds. Sometimes the electrons are not equally shared; one atom tends to have an electron more often than the other atom. In this situation one atom of the molecule becomes partly negative and the other then becomes partly positive.
Now we have positive and negative things becoming attracted to each other. (remember ionic bonds?) This is especially common between water molecules.

Shapes of Simple Molecules

Understanding the shapes of molecules is an important first step in being able to discuss and predict chemical properties. We shall discuss here some "simple" molecules. This topic, however, has important applications in understanding the behavior of much larger molecules. Much of biochemistry is now being discussed based on how macromolecules are shaped, and how different molecules "fit" together.
It is observed that the SF2 molecule is bent, the F-S-F bond angle being 98°. The BeCl2 molecule, however, is linear. Why are these two AX2 type molecules so different?

SF2

BeCl2

To understand any molecule, one must first complete a Lewis dot structure. It is then possible to predict the molecular shape using TWO BASIC PRINCIPLES:
1. The shapes of molecules are determined by the repulsion between electron pairs in the outer shell of the central atom. Both bond pairs (electron pairs shared by two atoms) and lone pairs (those located on a central atom but not shared) must be considered.
2. Lone pairs repel more than bond pairs.
The application of these principles is best seen by referring to specific examples. We shall start looking only at molecules with single bonds (for simplicity).

A. Central atoms with less than octet configurations: The BeCl2 molecule has a Lewis dot structure as shown. The central Be atom has two bond pairs in its outer shell. Repulsion between these two pairs (first principle above) causes the molecule to be linear (see above). If the molecule were bent in any direction, the two bond pairs would be brought closer together, increasing the repulsion.
The molecule BF3 has a dot symbol as:

BF3 :

Here the B atom has three bond pairs in its outer shell. Minimizing the repulsion causes this molecule to have a trigonal planar shape, with the F atoms forming an equilateral triangle about the B atom. The F-B-F bond angles are all 120°, and all the atoms are in the same plane.

B. Central atoms with octet (noble gas) configurations:
The CH4 molecule has a dot structure as shown. The shape of this molecule, however, is not planar, as is suggested by the way we draw this dot structure.

CH4 :

Carbon has 4 bond pairs. The four H atoms are arranged about the C atom in a tetrahedral shape . This shape minimizes the repulsion between the bond pairs. The 109.5° angle is the same for all H-C-H bond angles and is called the tetrahedral angle.

There are many molecules that have four bond pairs and this regular tetrahedral shape; CCl4, SiF4, and SnCl4 are just a few examples.

The molecule NH3 has a dot symbol much like that for BF3 (see above). Now, however, there is a lone pair in the outer shell of the central N atom.

NH3 :

In NH3 the N has 3 bond pairs and 1 lone pair, (4 total pairs). The shape is called trigonal pyramidal (approximately tetrahedral minus one atom).

WHACHACALLIT: A shape name is based on what is experimentally observed - the location of atomic nuclei. In NH3 the N, with 4 e pairs, will have a tetrahedral electron pair orientation. (The total number of e pairs determines the electron pair orientation.) The lone pair occupies one corner of the tetrahedron. It is difficult to "see" lone pairs experimentally. Looking only at the atoms, we see a short, rather distorted tetrahedron. This is called a pyramid. The pyramids of Egypt have square bases. The NH3 pyramid has a triangular base. Hence the shape is called trigonal pyramidal.

LONE PAIR DISTORTIONS: Due to the greater repelling character of lone pairs, (second principle above) the H atoms in NH3 are bent closer together than the normal tetrahedral angle of 109.5°. In NH3 the observed angle is 107.3°. Other molecules with this one-lone, three-bond-pair configuration (:NCl3 and :PCl3) have this same trigonal pyramidal shape, slightly different bond angles, but all less than 109.5°.

The H2O molecule has this dot structure:

H2O:



The O in H2O has 2 bond pairs and 2 lone pairs (again, 4 total pairs). The electron pair orientation around O is tetrahedral. Two corners of the tetrahedron are "missing" because they are occupied by lone pairs, not atoms. The shape is called bent. The H-O-H bond angle is 104.4°. This angle is less than that in NH3, due in part to the greater repulsions felt with two lone pairs

Other molecules with 2 bond plus 2 lone pairs include OF2, H2S, and SF2. Bond angles vary, but all are significantly less than 109.5°.

C. Central atoms with expanded octet
configurations:
A number of molecules have more electrons in the outer shell than in a noble gas configuration. This involves use of one or more d orbital (so as to not violate the Pauli Exclusion Principle). We will not worry, for now, about the types or shapes of the orbitals involved in the bonding, but will only consider electron pair repulsions.

The PCl5 molecule has 5 bond pairs in the outer shell of P. This molecule has a symmetrical trigonal bipyramidal shape.

PCl5 :

Note that the Cl atoms occupy two types of positions. The two Cl atoms which are on a straight line which passes through the P nucleus are said to occupy axial positions. The other three Cl are in equatorial positions.
NOTE:The axial Cl to P bond distance is slightly longer than the equatorial Cl to P bond distance.

Valence-Shell Electron-Pair Repulsion Theory (VSEPR)

There is no direct relationship between the formula of a compound and the shape of its molecules. The shapes of these molecules can be predicted from their Lewis structures, however, with a model developed about 30 years ago, known as the valence-shell electron-pair repulsion (VSEPR) theory.
The VSEPR theory assumes that each atom in a molecule will achieve a geometry that minimizes the repulsion between electrons in the valence shell of that atom. The five compounds shown in the figure below can be used to demonstrate how the VSEPR theory can be applied to simple molecules.


There are only two places in the valence shell of the central atom in BeF2 where electrons can be found. Repulsion between these pairs of electrons can be minimized by arranging them so that they point in opposite directions. Thus, the VSEPR theory predicts that BeF2 should be a linear molecule, with a 180o angle between the two Be-F bonds.
There are three places on the central atom in boron trifluoride (BF3) where valence electrons can be found. Repulsion between these electrons can be minimized by arranging them toward the corners of an equilateral triangle. The VSEPR theory therefore predicts a trigonal planar geometry for the BF3 molecule, with a F-B-F bond angle of 120o.
BeF2 and BF3 are both two-dimensional molecules, in which the atoms lie in the same plane. If we place the same restriction on methane (CH4), we would get a square-planar geometry in which the H-C-H bond angle is 90o. If we let this system expand into three dimensions, however, we end up with a tetrahedral molecule in which the H-C-H bond angle is 109o28'.
Repulsion between the five pairs of valence electrons on the phosphorus atom in PF5 can be minimized by distributing these electrons toward the corners of a trigonal bipyramid. Three of the positions in a trigonal bipyramid are labeled equatorial because they lie along the equator of the molecule. The other two are axial because they lie along an axis perpendicular to the equatorial plane. The angle between the three equatorial positions is 120o, while the angle between an axial and an equatorial position is 90o.
There are six places on the central atom in SF6 where valence electrons can be found. The repulsion between these electrons can be minimized by distributing them toward the corners of an octahedron. The term octahedron literally means "eight sides," but it is the six corners, or vertices, that interest us. To imagine the geometry of an SF6 molecule, locate fluorine atoms on opposite sides of the sulfur atom along the X, Y, and Z axes of an XYZ coordinate system.



Polarity of molecules

A compound is comprised of one or more chemical bonds between atoms. The polarity of each bond within the compound determines the overall polarity of the compound: how polar or non-polar it is. A polar molecule usually contains polar bonds - bonds which have unequal sharing of electrons between the two atoms involved in bonding. A non-polar compound usually contains non-polar bonds - bonds which have identical or similar sharing of electrons.Besides bond polarity, the other factor that decides if a molecule is polar is the molecule's symmetry. Even if a compound contains only polar bonds, it may be non-polar overall as the direction of the polarities cancel each other out, giving the molecule a net polarity of zero. This occurs in boron trifluoride, which contains three identical polar bonds all canceling each other out due to their symmetrical arrangement. Trigonal planar, tetrahedral and linear bonding arrangements often lead to symmetrical, non-polar molecules which contain polar bonds. On the other hand, even if a compound contains only non-polar bonds, it may be polar overall if it is a non-symmetric shape; for example, all the bonds in ozone are non-polar (between atoms of the same element), but the ozone molecule is nevertheless polar, because of its bent shape and the resulting asymmetry in electron distribution.



Intermolecular Forces

Water is the only substance we routinely encounter as a solid, a liquid, and a gas. At low temperatures, it is a solid in which the individual molecules are locked into a rigid structure. As we raise the temperature, the average kinetic energy of the molecules increases, which increases the rate at which these molecules move.
There are three ways in which a water molecule move: (1) vibration, (2) rotation, and (3) translation. Water molecules vibrate when H--O bonds are stretched or bent. Rotation involves the motion of a molecule around its center of gravity. Translation literally means to change from one place to another. It therefore describes the motion of molecules through space.

To understand the effect of this motion, we need to differentiate between intramolecular and intermolecular bonds. The covalent bonds between the hydrogen and oxygen atoms in a water molecule are called intramolecular bonds. (The prefix intra- comes from the Latin stem meaning "within or inside." Thus, intramural sports match teams from the same institution.) The bonds between the neighboring water molecules in ice are called intermolecular bonds, from the Latin stem meaning "between." (This far more common prefix is used in words such as interface, intercollegiate, and international.)
The intramolecular bonds that hold the atoms in H2O molecules together are almost 25 times as strong as the intermolecular bonds between water molecules. (It takes 464 kJ/mol to break the H--O bonds within a water molecule and only 19 kJ/mol to break the bonds between water molecules.)
All three modes of motion disrupt the bonds between water molecules. As the system becomes warmer, the thermal energy of the water molecules eventually becomes too large to allow these molecules to be locked into the rigid structure of ice. At this point, the solid melts to form a liquid in which intermolecular bonds are constantly broken and reformed as the molecules move through the liquid. Eventually, the thermal energy of the water molecules becomes so large that they move too rapidly to form intermolecular bonds and the liquid boils to form a gas in which each particle moves more or less randomly through space.
The difference between solids and liquids, or liquids and gases, is therefore based on a competition between the strength of intermolecular bonds and the thermal energy of the system. At a given temperature, substances that contain strong intermolecular bonds are more likely to be solids. For a given intermolecular bond strength, the higher the temperature, the more likely the substance will be a gas.
The kinetic theory assumes that there is no force of attraction between the particles in a gas. If this assumption were correct, gases would never condense to form liquids and solids at low temperatures. In 1873 the Dutch physicist Johannes van der Waals derived an equation that not only included the force of attraction between gas particles but also corrected for the fact that the volume of these particles becomes a significant fraction of the total volume of the gas at high pressures.
The van der Waals equation is used today to give a better fit to the experimental data of real gases than can be obtained with the ideal gas equation. But that wasn't van der Waals's goal. He was trying to develop a model that would explain the behavior of liquids by including terms that reflected the size of the atoms or molecules in the liquid and the strength of the bonds between these atoms or molecules. The weak intermolecular bonds in liquids and solids are therefore often called van der Waals forces. These forces can be divided into three categories: (1) dipole-dipole, (2) dipole-induced dipole, and (3) induced dipole-induced dipole.

Dipole-Dipole Forces
Many molecules contain bonds that fall between the extremes of ionic and covalent bonds. The difference between the electronegativities of the atoms in these molecules is large enough that the electrons aren't shared equally, and yet small enough that the electrons aren't drawn exclusively to one of the atoms to form positive and negative ions. The bonds in these molecules are said to be polar, because they have positive and negative ends, or poles, and the molecules are often said to have a dipole moment.
HCl molecules, for example, have a dipole moment because the hydrogen atom has a slight positive charge and the chlorine atom has a slight negative charge. Because of the force of attraction between oppositely charged particles, there is a small dipole-dipole force of attraction between adjacent HCl molecules.



The dipole-dipole interaction in HCl is relatively weak; only 3.3 kJ/mol. (The covalent bonds between the hydrogen and chlorine atoms in HCl are 130 times as strong.) The force of attraction between HCl molecules is so small that hydrogen chloride boils at -85.0oC.

Dipole-Induced Dipole Forces
What would happen if we mixed HCl with argon, which has no dipole moment? The electrons on an argon atom are distributed homogeneously around the nucleus of the atom. But these electrons are in constant motion. When an argon atom comes close to a polar HCl molecule, the electrons can shift to one side of the nucleus to produce a very small dipole moment that lasts for only an instant.

By distorting the distribution of electrons around the argon atom, the polar HCl molecule induces a small dipole moment on this atom, which creates a weak dipole-induced dipole force of attraction between the HCl molecule and the Ar atom. This force is very weak, with a bond energy of about 1 kJ/mol.

Induced Dipole-Induced Dipole Forces
Neither dipole-dipole nor dipole-induced forces can explain the fact that helium becomes a liquid at temperatures below 4.2 K. By itself, a helium atom is perfectly symmetrical. But movement of the electrons around the nuclei of a pair of neighboring helium atoms can become synchronized so that each atom simultaneously obtains an induced dipole moment.

These fluctuations in electron density occur constantly, creating an induced dipole-induced dipole force of attraction between pairs of atoms. As might be expected, this force is relatively weak in helium -- only 0.076 kJ/mol. But atoms or molecules become more polarizable as they become larger because there are more electrons to be polarized. It has been argued that the primary force of attraction between molecules in solid I2 and in frozen CCl4 is induced dipole-induced dipole attraction.