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Chemical Bonding

Chemical bonding involves the joining of two atoms. The oxidation numbers that we learned about in the first semester have to do with whether the atoms involved in a bond “prefer” to lose, gain or share electrons. The ability to lose electrons can be determined by an element’s ionization energy whereas the ability of an element to accept electrons is its electron affinity. Both ionization energy and electron affinity apply to atoms by themselves. We're more interested in how atoms of one element behave in the presence of atoms of other elements.

An element’s electronegativity relates the ability of an element to attract electrons when they combine with another element. The same 4 factors affect electronegativity as electron affinity and ionization energy. We observe the same trends on the periodic table for each. Elements with the highest ionization energy, electron affinity and electronegativity are at the top, right on the periodic table (Noble Gases excluded). Elements with the lowest ionization energy, electron affinity and electronegativity are at the bottom, left of the periodic table.

Metals and Nonmetals and Electronegativity

Since metals lose electrons when they react, they have low electronegativities. Nonmetals can gain or lose electrons, and have higher electronegativities.

Bond Character in a Molecule

One factor that determines the nature of a bond between two atoms is the electronegativity difference (DEN). A rule of thumb most textbooks go by is that if the DEN is less than 1.67, the bond is covalent and electrons are shared. A DEN value greater than or equal to 1.67 results in a bond which is mainly ionic in nature.

An example:

BaO          EN of O = 3.37                                           CH4        EN of C = 2.56
            -  EN of Ba = 0.93                                               - EN of H = 2.20
                                  2.44        ionic bond                                            0.36 covalent bond                                                                                                           

This is a bit of a simplification as a bond can have some sharing and some electron transfer going on.

What Do Ionic and Covalent Bonds Look Like?

What Affects The properties of a Molecule?

When atoms combine to form molecules, the compounds have different properties than any of the atoms. The properties of a molecule are dependent on two factors. The are:

1) the type of chemical bonding taking place.

2) how the atoms arrange themselves in space.

There are several ways of explaining the shape of molecules. We will start with a 2 dimensional model called Lewis dot diagrams.

Lewis Dot Diagrams

We were first introduced to Lewis dot figures when we discussed electron arrangement. For the A groups, the number of dots (electrons) surrounding the symbol of the element is simply equal to the group number. We also know that atoms react in order to acheive a stable octet, 8 outer electrons. Atoms can either share or transfer electrons. In this section we will focus on molecules, which consist of atoms that share electrons.

Knowing that atoms combine to achieve a stable electron configuration, we can use the Lewis dot diagrams to construct the molecules. For example, two hydrogens combine with an oxygen atom to form water.

Another example would be the molecule methane, CH₄

While it is easy to see how the atoms go together in Lewis Dot figures, the main weakness is that the dot diagrams are a 2 dimensional representation.

VSEPR - Valence shell electron pair repulsion

Because molecules occupy three dimensions not two, we can easily modify our dot diagram model to account for the electrostatic repulsion that occurs between the electron clouds and visualize molecules in 3-D. Each electron represented as a dot is actually an electron cloud that repels other electron clouds. What shapes do the clouds make?

In an atom like Beryllium, there are two outer electrons. Since these electron clouds repel each other, they orient themselves as far away from each other as possible into a linear configuration. In a molecule of Beryllium hydride, BeH₂, the H-Be-H bond angle is 180 degrees. The shape of the BeH₂ molecule is linear.

H : Be : H

Boron on the other hand has 3 outer electrons, which orient themselves in the 12, 4, 8 O'clock positions. In this arrangement, the three electron clouds orient themselves along a plane at angles of 120 degrees from each other.

Carbon has four outer electrons, which arrange themselves in a 3 dimensional configuration. If you can visualize the carbon atom in the center of a tetrahedron, the electrons would be at the corners of the tetrahedron. We cannot represent a tetrahedron using Lewis dot diagrams. The best we can do is to draw the dots at 12, 3, 6 and 9 O'clock.

If you use your imagination, the electrons at 12 and 6 O'clock come out of the screen slightly while the 3 and 9 O'clock electrons go into the screen. A molecule of methane has hydrogens where each unpaired electron is located. The shape of the CH₄ molecule is a tetrahedral. The H-C-H bond angle in a tetrahedral shaped molecule is 109.5 degrees.

Molecules like PCl₅ and SF₆ have 3 dimensional arrangements that really cannot be represented in a Lewis Dot figure, but can be drawn out.

 Knowing the shape of the electron clouds can help us predict the shape of a molecule. If all the surrounding electron clouds are bonded to atoms, the resulting shape of the molecule matches the shape of the electron clouds. However, if there are some unshared pairs of electrons, the shape of the molecule is different.

Characterizing A Chemical Bond: A Closer Look

Summarizing, we have identified 3 kinds of bonds. They are:

1) metallic bonds - see above discussion

2) ionic bonds - a large difference in electronegativities between the two atoms results in a transfer of electrons from the atom with low EN to the atom with high EN. This results not in a molecule, but a positively charged ion attracted to a negatively charged ion. For example salt, NaCl is not really a molecule, but a collection of Na⁺¹ ions and Cl^-1 ions stuck together in a repeating pattern.

3) covalent bonds - a difference in electronegativities of less than 1.67 results in sharing of electrons between two atoms. We can identify two kinds of covalent bonds. They are:

The Dipole Moment (m) is the measure of the polarity of a bond. It is dependent on the size of the partial charge and the distance between them. A larger difference in EN can result in a larger dipole moment as can a larger bond length.

What Determines the Characteristics of a Substance?

We have learned that both shape of the molecule and bonding type affect the properties of a molecule. We can categorize substances as belonging to 4 different classes. They are metals, ionic substances, nonpolar covalent and polar covalent molecules.

1) Metals are malleable, ductile, shiny and conduct heat and electricity. They tend to lose electrons when they react.

2) Ionic substances are hard, brittle materials with fairly high melting points. They are nonconductors in the solid state but will conduct electricity in the molten phase and in the aqueous state.

3) Non-polar Covalent Molecules are mostly low boiling liquids or gases, though there are some solid non-polar substances. They tend not to be soluble in water and do not conduct electricity in any state.

4) Polar Covalent Molecules are mostly higher boiling liquids or solids. They are soluble in water and are nonconductors in any state.

What Determines Polarity of a Molecule

A substance with non-polar covalent bonds will always be non-polar. A molecule with polar covalent bonds can be either polar or non-polar depending on how the bonds are arranged in space.

Molecules with similar dipole moments will be more likely to dissolve in each other. Substances with very different dipole moments tend to be insoluble or immiscible in each other. The term like dissolves like can be useful when comparing the relative polarities of two materials.

For example:

In general, most polar and ionic solutes will dissolve in water, which is a polar solvent. Most nonpolar solutes are either insoluble or barely soluble in water. Remember, this is a generalization.

Intermolecular Forces - van der waals Forces

Molecules must interact with their neighbors, otherwise all matter would exist only in the gas state. What holds molecules together in the solid and liquid states?

Intermolecular forces provide the "glue" that holds molecules of a solid or liquid together. There are three types of intermolecular forces. They are:

These intermolecular forces are much weaker than covalent bonds and can be broken by heating to the boiling point.

The intermolecular forces vary with the size of the molecules. In general, the larger the molecule, the larger the intermolecular forces. This explains why fluorine (F₂) exists as a gas whereas iodine (I₂) exists as a solid. The I₂ molecule is much larger than the F₂ molecule which results in stronger intermolecular forces holding the iodine together as a solid.

A similar pattern is observed in compounds containing hydrogen and VIIA and hydrogen and VA. HF and NH₃, being molecules with the lowest molar mass in their respective series all boil at unexpectedly high temperatures. What causes this unexpected deviation?

All 3 series of molecules contain hydrogen. When we look at the electronegativities of F, O and N we see that they are the three most electronegative elements on the periodic table. There must be something about hydrogen bonded to a highly electronegative element that results in a large intermolecular force and therefore a higher boiling point.

Water: Some Interesting Properties Due to Hydrogen Bonding

Most substances expand when heated and contract when cooled. This should make sense based on our understanding of kinetic theory. Any heat added will result in more movement of the particles which would result in a spreading out. The same should hold true for melting or freezing. Most substances expand when melted and contract when cooled. Water behaves like most materials except between 0 and approximately 4 C. When ice is heated to its melting temperature it begins to melt. After melting, we notice that it has contracted instead of expanded. Furthermore, when liquid water at 0 C is heated to 4 C, it contracts. After 4 C water behaves like most other substances expanding when heated and contracting when cooled.

The most interesting consequence of this behavior is that unlike most substances solid water is less dense than liquid water. We can see why this is when we look at the hydrogen bonding between neighboring molecules in ice. The result of the hydrogen bonding is much empty space between the molecules, which explains the lower density. This explains why ice floats on water. The layer of ice on lakes and ponds in the winter provides insulation and prevents the entire lake or pond from freezing over. If water was like most substances, ice would be more dense than water and it would sink to the bottom of the lake. With no insulation, the lake would freeze solid.

Surface Tension

Because of the hydrogen bonding, water molecules exert forces on their neighbors and are also pulled at by their neighbors. Molecules below the surface are pulled equally in all directions but surface molecules are not. Since surface molecules do not have any molecules above them, the result is a net force downwards. This net force acting on the surface molecules contributes to surface tension. Some interesting results of surface tension are the beading up of water on cars, water striders being able to walk on water and the spherical shape of water droplets created from a splash.

Surface tension is due to cohesive forces holding the water molecules together. Attraction of water to other objects is referred to as adhesion. The concave meniscus created by water is due to the adhesive forces between the water and glass being greater than the cohesive forces between water molecules. This also explains capillary rise. Mercury has a much stronger surface tension, which cannot be overcome by adhesive forces of glass in a graduated cylinder. This explains the convex meniscus.

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