Chemical bonding is fundamental to the study of chemistry since chemical bonds describe how molecules and compounds are formed.
Chemical bonds describe an effect that occurs when one or more outer shell electrons are simultaneously attracted to two atomic nuclei. The forces holding a chemical bond together are electrostatic forces. A molecule is created when chemical bonds form. You can think about a molecule as an aggregate of atoms with distinct properties.
Most elements found on the periodic table readily form chemical bonds to create molecules.
Atoms often bond together chemically to form molecules. Since heat is released during this process, chemical bonding is an example of an exothermic reaction. The more exothermic the reaction, the more stable the product. Since the formation of chemical bonds creates a more stable product, there is a natural tendency toward bond formation. We describe the bonds as energetically favorable since energy is released.
In general, the more stable structure that results (molecules) also has lower potential energy than the original individual atoms. Remember their energy was released during the chemical bonding process.
Chemical bonds can take three forms depending on the atoms involved: ionic, covalent, and metallic.
In an ionic bond, ions of opposite charge are attracted to each other via electrostatic forces: an electron is donated from the negative ions to the positive ion to form the bond.
Table salt, or sodium chloride (NaCl), is an example of a compound that has an ionic bond. In the case of NaCl, the chloride (Cl-) ion donates its extra electron to the sodium ion (Na+) to form the bond. Ionic solids, such as NaCl are crystalline in form.
In a covalent bond, the electrons in the bond are shared. Covalent bonds generally form among nonmetal atoms: each atom in the bond usually contributes one electron to form the bonding electron pair.
We classify covalent bonds as
nonpolar covalent bonds or
polar covalent bonds. In a nonpolar covalent bond, the two atoms in the bond share the bonding electrons equally. In a polar covalent bond, the electrons are unevenly shared.
This means, there is a greater electron density near one of the atoms in the bond than the other. This is because one of the atoms has higher electronegativity than the other atom. Electronegativity is a measure of how much an atom pulls electrons toward itself in a covalent bond.
Metallic bonds form, not surprisingly, among metal atoms. Metal atoms have low electronegativities, and have empty, or near empty, outer electron shells. Consequently, metal atoms do not attract electrons, nor do they easily donate them.
A simple way to think about metal bonding is that positive metal ions are immersed in an electron fluid of free-flowing electrons. This leads to many properties of metals. For example, metals conduct electricity because the valence electrons are mobile throughout the material.
Intermolecular forces are the forces that hold molecules of the same type together in the condensed phase.
Van der Waals forces, also known as intermolecular forces, describe the forces that exist between molecules that hold molecules together in the liquid or solid state.
In dipole-dipole force interactions, molecules with permanent dipole moments interact. A molecule has a permanent dipole moment if it has an uneven distribution of negative charge within the overall neutrally charged molecule. We call molecules with a permanent dipole moment polar molecules. A polar molecule has a partially positive area and a partially negative area within the molecule.
When two polar molecules come together, the partially positive area of one molecule lines up with the partially negative area of the other molecule. The weak electrostatic attraction among the molecules leads to dipole-dipole interactions.
ion-induced dipole forces, an ion interacts with a
nonpolar molecule. In this case, the electrostatic charge on the ion induces or forces the nonpolar molecule to momentarily develop a dipole. This is a weaker interaction because the nonpolar molecule does not have a permanent dipole, and is therefore not strongly attracted to the ion.
ion-dipole forces, an ion interacts with a polar molecule. The polar molecule lines up, so the opposite partial charge is attracted to the ion. This is a stronger interaction than ion-induced dipole because the polar molecule has a permanent partially-positive and partially-negative charge.
In dipole-induced dipole interactions, a polar molecule interacts with a nonpolar molecule. The dipole in the polar molecule
induces or causes a momentary dipole in the nonpolar molecule. The resulting interaction is weak because the nonpolar molecule only has momentary and not permanent dipole.
London forces describe the forces between two nonpolar molecules. A momentary dipole that randomly forms in one molecule induces an induced dipole in the other nonpolar molecule. This is a very weak intermolecular force.
First, review how to write electron configurations for atoms in Section 3: The Aufbau Rules, in Electrons in Atoms. This will help you understand how to draw a Lewis Dot structure to apply VSEPR theory.
Next, read Section 2: Lewis Dot Structures to understand how Lewis dot structures work and review how to use VSEPR theory.
Section 3: How to Draw Lewis Dot Structures, provides a step-by-step description of how to draw Lewis dot structures in The Shared-Electron Covalent Bond.
To write a Lewis dot structure, first write the symbols for the elements in a simple diagram to show how the elements will be connected.
For example, for ammonia, NH3, or hydroxylamine, you would write the following:
Then, using the Aufbau principle, draw electron dot structures for each of the elements. These drawings tell us how many
valence shell electrons are in each atom, based on their group on the periodic table. The group number on the periodic table equals the number of valence electrons.
For example, hydrogen is in group 1A, so it has one valence electron. Nitrogen is in group 5A so it has five valence electrons, and oxygen is in group 6A so it has six valence electrons.
Finally, bring the atoms together in a way that places eight electrons around each atom wherever possible. Note that hydrogen is an exception to this rule and will only ever have two electrons.
After studying the basics of Lewis electron dot structures, we can begin to understand VSEPR Theory.
VSEPR Theory is an acronym for Valence Shell Electron Pair Repulsion Theory. It describes the three-dimensional shapes of molecules.
VSEPR Theory focuses on the valence electron pairs in the outermost electron shells of the atoms involved in bonding. These pairs are the electrons that can form chemical bonds.
We assume that electrons involved in bonding exist in between the two atoms being bonded. We also know that similar or “like” charges repel each other via electrostatic forces. Therefore, the lone pairs of electrons, or nonbonding valence electrons in the molecule are repelled by the bonding pair and by all of the other lone pairs of electrons in the molecule.
Because they are repelled, all of the valence electron pairs will adopt shapes that make them as far apart from each other as possible in a three-dimensional space. This description provides the basis of VSEPR Theory and predicts the shapes different molecules will take based on the number of bonding and nonbonding valence electron pairs.
Remember that we use VSEPR Theory to predict the shape molecules and polyatomic ions will take on based on their Lewis electron dot structure.
Again, it is important to have a strong foundation in Lewis electron dot structures, so be sure to review Sections 2 and Section 3, in The Shared-Electron Covalent Bond.
Molecules with the formula AX2 have two bonding electron pairs that are 180o apart to maximize the distance between them. We call them linear molecules.
Molecules with the formula AX3 have three bonding electron pairs that are 120o apart to maximize distance between them. We call them trigonal planar molecules.
The most common configuration we see in chemistry is the AX
4 molecule. These molecules have bonding electron pairs 109.5
o apart in three dimensional space. These molecules are called
Methane, CH4, is an example of a tetrahedral molecule.
Sometimes, one or more of the valence electron pairs in a molecule are lone pairs of electrons. This alters the geometry of the molecule since the lone pairs take up a bit more space than bonding electrons do.
A molecule of the form AX3E is based in tetrahedral geometry, but has one lone pair of electrons (E). Ammonia, NH3, is an example of this type of molecule. Here, the molecule takes a trigonal pyramidal shape and the bonding angles are approximately 107o.
20, is an example of a molecule with AX
2 geometry. In this molecule, there are two bonding pairs of valence electrons and two lone pairs. This molecule has a
bent geometry with a bond angle around 104.5
Occasionally, we see molecules with five or six valence electron pairs that adopt other, more complex geometric shapes.
To review how VSEPR theory describes different molecular geometries, see the descriptions in Molecular Geometry.
Hybridization theory accounts for molecular shapes that have been seen in experiments when combining atomic orbitals into molecular orbitals.
Hybridization theory presents the limitations of VSEPR theory. In VSEPR theory, we assumed the electrons were in their atomic orbitals. If this were the case, many molecules we know to exist should never form.
For example, beryllium hydride, BeH2, is a known compound, but beryllium (Be) has a set of paired electrons in its valence shell (the 2s atomic orbital). The paired electrons in the atomic orbitals are stable and should not combine with other atomic orbitals to form bonds.
However, we know beryllium hydride, BeH2, exists. The only way for this to happen is for an electron from the beryllium valence shell to move to the next energy level, 2p, so there are unpaired electrons. We know this does not happen because it would require too much energy. Therefore, a new model of bonding is needed to explain molecules such as BeH2 which cannot be explained by VSEPR.
In hybridization theory, the electrons from the atoms involved in the bond are hybridized or combined into new molecular orbitals which are combinations of the input atomic orbitals. This is a mathematical construct that gives us molecular orbitals that possess properties that are consistent with what we observe about molecules.
Hybrid orbitals are constructed by combining the wavefunctions, ψ, of the atomic orbitals involved in the bond. Because the wavefunction describes a wave, the wavefunctions interfere constructively and destructively, to create a new shape for the hybrid molecular orbital.
To use a hybrid orbital diagram, we first fill the electron configurations for the atomic orbitals.
Review how to write electron configurations for atoms in Section 3: The Aufbau Rules, in Electrons in Atoms.
Then, we combine the electrons from both atoms together to fill in the hybrid molecular orbitals created by the atomic orbitals. We always fill the molecular orbitals starting at the lowest energy level. For each energy level, we fill in one electron per molecular orbital at a time and then fill the second electrons per molecular orbital in the energy level before moving on to the next energy level.
In linear molecules, s and p atomic orbitals combine to form hybrid sp orbitals.
In our example of beryllium hydride, BeH
2, we can use hybrid orbitals to explain why this molecule exists and has a linear shape. Two
sp hybrid orbitals are formed in this molecule and overlap with the 1s orbitals of each hydrogen atom to form covalent bonds.
We can use hybridization theory to explain larger molecules as well. Here we will look at a trigonal planar molecule, BF
3. In this case, the atomic orbitals of B hybridize to become three
hybrid orbitals. These sp
2 hybrid orbitals have unpaired electrons that can combine with the unpaired electrons of fluorine to create covalent bonds.
When the sp
2 hybrid orbitals are formed, they create trigonal planar geometry:
Finally, we will describe the most important geometry in chemistry,
tetrahedral, and how
hybridization theory predicts the molecular geometry. We will use methane, CH
4, as our example.
Here, the electrons in carbon hybridize to form four sp3 hybrid orbitals. Each of the four sp3 hybrid orbitals contains an unpaired electron that can combine with a hydrogen electron to form a bond.
3 orbitals form the expected
tetrahedral molecular geometry.
Molecular polarity occurs in molecules where the electrons are not evenly distributed.
Electronegativity measures how well an atom attracts electrons toward itself in a bond.
An atom that is highly-electronegative attracts electrons toward itself more strongly. On the other hand, an atom that has low electronegativity and does not attract electrons toward itself.
When determining polarity, chemists examine the relative electronegativities of an atom in a bond. The location of each element in the periodic table indicates this trend.
Fluorine is the most electronegative atom. The elements that appear closest to fluorine in the periodic table (top right) are the most electronegative. Elements that are located far away from fluorine in the periodic table (bottom left) are the least electronegative and said to be electropositive.
polar covalent bond occurs among atoms with different electronegativities. The electrons in the bond are more attracted to the more electronegative atom and therefore spend more time closer to that atom. So, in these bonds, the electrons are not evenly shared.
We say polar covalent bonds have a dipole moment, which is a vector that points from the less electronegative atom to the more electronegative atom.
Vectors describe a mathematical quantity that also have a direction associated with them. Because dipole moments are vectors, we can sum the vectors to determine if a molecule is polar. If the vectors do not cancel out, the molecule is polar because one area of the molecule has a higher electron density than the rest. If the vectors do cancel out, the molecule is nonpolar because the electron distribution in the molecule is even.
Review a mathematical treatment of dipole moments in Section 2: Molecular Dipole Moments, in Polar Covalence.
In the figure below, we can see how dipole moments determine molecular polarity. The oxygen molecule, O2, is homonuclear, or made up of only one element type. Therefore, there cannot be a difference in electronegativity and it is a nonpolar molecule.
Carbon monoxide, CO, consists of two different types of atoms (carbon and oxygen). Oxygen is more electronegative and therefore the dipole moment goes from the carbon to the oxygen (NOTE: The figure from your text contains an error here). Because there is a dipole moment, carbon monoxide is a polar molecule.
Carbon dioxide, CO2, on the other hand, also contains carbon-oxygen bonds. However, the two dipole moments are 180 degrees opposite each other and cancel each other out. Consequently, CO2 is a nonpolar molecule that contains polar bonds.
We create resonance structures for molecules in which some electrons are delocalized. In general, resonance structures describe a molecule or polyatomic ion in which we could write more than one equivalent Lewis electron dot structure.
First, review the rules for writing Lewis electron dot structures in Section 2: Lewis Dot Structures, and, Section 3: How to Draw a Lewis Dot Structure, in The Shared-Electron Covalent Bond.
Resonance structures are necessary when there are multiple equivalent Lewis electron dot structures that could be written for a given molecule or polyatomic ion. Generally, resonance structures involve double bonds.
Let’s look at the example of the nitrate ion, NO3. In the nitrate ion Lewis electron dot structure, we make two single N-O bonds and one double N=O bond. When we draw this, it does not matter which nitrogen-oxygen bonds we make single or double. We say that all three possibilities are resonance structures and we denote these with double arrows between them:
In reality, the true structure of the nitrate ion is a
superposition, or combination of all three of these resonance structures. Each bond is really about 1 ⅓ of a bond, or
We write resonance structures because we cannot accurately draw the true structure of compounds with resonance.
Review the section for more examples of molecules and polyatomic ions with resonance structures, Multiple Equivalent Structures: Resonance, in The Shared-Electron Covalent Bond.