chemical bonding

A chemical bond is an interaction that holds molecules and compounds together by the sharing or exchanging of electrons. When an atom comes into proximity to another atom and its valence (outer) electrons are attracted to the positive (nuclear) charge of the other atom, a bond between the two atoms can be formed. The strength of the bond, how tightly the chemical species are held together, is determined by the interactions between the repulsive and attractive forces of the nuclei and electrons. When atoms come into proximity and interact with each other they arrange themselves into the lowest energy formation possible. In this way, the chemical bond forms molecules, liquids, gases, crystals, and the objects we use everyday. For instance, sodium is a highly reactive metal and chlorine is a poisonous gas, but when sodium and chlorine react with each other they form sodium chloride, table salt, a unique compound with properties that are very different from the individual elements.

Chemical bonding tends to be of two types; covalent, in which electrons are shared between atoms, and ionic in which two oppositely charged ions attract one another. An ion is a chemical species that possesses a charge due to the loss or gain of one or more electrons.

covalent

In a covalent bond, sometimes called a molecular bond, valence electron pairs are shared between atoms in a stable balance of attractive and repulsive forces. Atoms are most stable when their valence electron (electrons located in the outermost orbital shell) shell is full. If atoms can’t fill their valence shell by transferring electrons, they share to achieve stability. In addition, atoms share electrons to achieve a charge balance. The positive charge on a proton is attracted to the nearest negative charged species – usually, electrons. If a pair of atoms has a slight positive charge, sharing electrons lets them balance their charge and a bond is formed.

Covalent bonding occurs between nonmetallic atoms with similar electronegativities. If the atoms are identical, like two hydrogens forming \(\ce{H_{2(g)}}\), then the bond is purely covalent. If the atoms are different, like hydrogen and chlorine forming HCl, then the difference in electronegativity will affect the polarity of the bond meaning the electrons have a higher probability to be closer to one atom than the other creating an imbalance of charge. This is called a polar covalent bond. Molecules with polar covalent bonds often dissolve in polar solvents, such as water. Differences in electronegativities can give rise to dipole-dipole interactions which are interactions between polar regions of two molecules. The difference in electronegativity results in unequal sharing of electrons since the more electronegative atom holds the shared electrons more tightly creating a dipole moment.

covalent bond, HCl

ionic

An ionic bond is an electrostatic attraction between atoms with opposite charge. This type of bond is very strong and has a high level of energetic stability that comes from the interaction of the positively-charged nuclei with the negatively-charged electrons. In an ionic bond there is a transfer of valence electrons between atoms resulting in two oppositely charged ions. This type of bond requires an electron donor and an electron acceptor; the atoms that ‘lose’ electrons become positively-charged and are called cations; atoms that ‘gain’ electrons are negatively-charged and are called anions.

Ionic bonds form because the valence shells of metal atoms are not full. By losing the few valence electrons they do have, metals can achieve a stable, zero net charge, noble gas configuration and satisfy the octet rule. The octet rule is the observation that main-group elements are likely to form bonds where each atom has eight electrons in its valence shell. Main-group elements are those in periodic table groups 1, 2 (s-block), and 13-18 (p-block); not including H, He, Li, and Be. Low atomic weight elements (atomic number 20 and below) are the most likely to follow this rule. When they have a full octet their s- and p-orbitals are completely filled.

periodic table showing how shells are filled for main group elements

EXAMPLE

Sodium chloride, table salt, is a most classic example of ionic bonding. You may think that one Na attaches to one Cl but in reality networks of ions are formed by electrostatic interactions. The oppositely-charged ions are held together by electrostatic attraction and the like-charged ions are repelled. As the strength of the attraction is stronger than the repulsion in the lattice, solid NaCl forms a very ordered, rigid, ionic structure. In NaCl, each \(\ce{Na^{+}}\) ion is surrounded by 6 \(\ce{Cl^{-}}\) ions.

These solid ionic lattices have high melting and boiling points and act as insulators; they do not conduct electricity. To conduct electricity, you need charged particles that are free to move. So, upon dissolving in water, NaCl separates into ions, \(\ce{Na^{+}}\) and \(\ce{Cl^{-}}\), allowing electricity to be conducted. The lattice properties no longer hold. Once the lattice is dissolved the strong electrostatic attractions are no longer in place and the high melting and boiling points are reduced.

ionic lattice

NaCl dissolved in water

Ionic and covalent bonds are not the only types of chemical bonds, there are many other types; intermolecular – interactions between molecules, metallic – attractions between metal atoms; and vibrational – a lightweight element oscillating between much heavier atoms and holding them together. A full description of these is outside the scope of this particular wiki.

In 1916 Gilbert N. Lewis described the covalent sharing of electron pairs between atoms and he introduced a notation in which valence electrons are represented as dots around atomic symbols. These drawings are known as Lewis dot structures or electron dot structures. The most common covalent bond is a single bond in which two atoms share two electrons (represented as two dots or one line). A single bond is called a σ-bond. It follows that there are double bonds (two atoms share four electrons), and triple bonds (two atoms share six electrons). A double bond consists of one σ–bond and one so-called π-bond, and a triple bond is one σ–bond and two π-bonds. As the orbitals overlap, molecular stability is increased.

It was Lewis who proposed the octet rule for main-group elements. Hydrogen is a notable exception to the octet rule with only one valence electron so it’s outer (and, only) shell can hold just two electrons. There are other exceptions, nitric oxide NO, as discussed earlier that has an odd number of electrons. The elements B and Al typically form compounds in which they have six electrons instead of eight. B, atomic number 5, and Al, atomic number 13, both have only 3 valence electrons which are not enough to fill an octet.

Lewis dot single, double, and triple bonds

Lewis dot poly-atomic molecule, \(\ce{CH_{4}}\). For methane we see that the carbon has fulfilled it’s required octet. This satisfies the number of electrons corresponding to full shells in the quantum mechanical model of the atom: the outer shell of a carbon atom has a principal quantum number of n = 2 and this shell can hold eight electrons. The hydrogen is an exception as mentioned earlier and need only follow a duet rule.

Lewis dot poly-atomic molecule, \(\ce{CO_{2}}\). The octet rule for oxygen and the most energetically stable configuration for CO2 is achieved by forming double bonds with each carbon.

EXAMPLE. You Try It. Draw the Lewis structure for water \(\ce{H_{2}O}\)

Write the letters of the elements you will draw electrons around. Usually the element that will likely have the most bonds should be put in the center:

\[\ce{H\space O\space H}\]

Draw Lewis symbols; electrons can be represented by dots or you can use a dash for a pair of electrons, around the atoms. Arrange the atoms in a way so that there are eight electrons (if possible) around each atom, or two electrons for H:

This arrangement fills the octet of O and gives H it’s two maximum valence electrons. Note that the electron pairs repulse the H’s and push them away.

EXAMPLE. You Try It. Draw the Lewis structure for acetic acid \(\ce{CH_{3}COOH}\)

Write the letters of the elements you will draw electrons around. Both C and O will need octets of electrons, H needs two electrons. Clues: There are three ligands in this molecule. A ligand is a functional group of atoms that will position itself as a unit in such a way as to give the molecule the lowest possible energy. The \(\ce{CH_{3}}\), the carbonyl (a carbon double bonded to an oxygen), and the OH (an alcohol) are ligands.

That’s a start but both O and the carbonyl C don’t have octets, so

OK, now we have octets around both C and the alcohol O but not the O bonded to the carbonyl C. It will need to have some additional free electrons two of which will be able to form a double bond with the carbonyl C, as such

Show the double bond in a clear manner and you have the Lewis structure for acetic acid:

There are other types of bonds as well, including one- and three-electron bonds, but they are not as commonly found. These bonds occur only in radicals, compounds with odd numbers of electrons. A one-electron bond can form when the nuclei have the same charge, such as \(\ce{H{_{2}^{+}}}\). One-electron bonds are sometimes called half bonds. An example of a three-electron bond is nitric oxide, NO.

When drawing a Lewis structure, something called formal charge (FC) must be considered to know if the structure is stable. When a bond is formed atoms gain or lose electrons in an attempt to fulfill the octet rule. Formal charge is the difference between the number of valence electrons of each atom and the number of electrons the atom is associated with. Formal charge assumes that shared electrons are equally shared between the bonded atoms. You can calculate formal charge for each atoms using the relationship:

\[formal\, charge = ev - en - \frac{eb}{2}\]

where ev = number of valence electrons of the isolated atom, en = number of unbound valence electrons on the atom in the molecule, and eb = number of electrons shared in bonds with other atoms in the molecule. Lewis structures are drawn in such a way that the formal charge is as small as possible. As a general guideline, if FC = 0, that’s good and the structure is stable and possible in nature; FC = -1 or 1, not ideal; FC < -2 or > 2, unstable and not possible in nature.

EXAMPLE Resonance Structures

Electrons have no memory of where they have been or where they belong, they are distributed over the molecule. To draw a good Lewis representation, sometimes resonance structures must be drawn to show possible electron (not atom) configurations. This may cause formal charge on an atom to change.

resonance structures of nitrogen dioxide \(\ce{NO{_{2}}}\)

EXAMPLE Draw the Lewis structure for \(\ce{BF{_{3}}}\)

Write the letters of the elements you will draw electrons around. Usually the element that will likely have the most bonds should be put in the center. That implies B will be the central atom with the three F around it.

Right away there is a problem. If you want to give all atoms an octet you’ll have to draw resonance structures like

However, these resonance structures do not represent a workable structure for \(\ce{BF{_{3}}}\). Remember what we learned about boron – it has only three valence electrons and doesn’t obey the octet rule. That give a much simpler Lewis structure as

The Lewis approach does not give any indication about the arrangement of the atoms, free electrons, or molecules in space as there is no direct relationship between molecular formula and molecular shape. To predict three-dimensional molecular geometry of simple, symmetric molecules we look to the valence-shell electron-pair repulsion (VSEPR) model. Shape is predicted using the number of electrons around the central atom in the Lewis structure. VSEPR is instead a method of counting to predict 3D structures.

Single atoms or functional groups of atoms, called ligands, are positioned around the central atom to give a molecular structure with the lowest possible energy. Electrostatic repulsion, in addition to electron-electron repulsion due to the Pauli exclusion principle, make the most stable geometry the one that minimizes these repulsions. For example, a simple molecule like carbon dioxide \(\ce{CO{_{2}}}\) is linear because the valence electron pairs on the C repel each other forcing the two O to opposite sides of the C as

the bond angle is 180°. Consider another example, methane \(\ce{CH{_{4}}}\). To get the hydrogen atoms are as far apart as possible you need a 109.5° bond angle which forms a tetrahedron:

AXE method

To use VSEPR theory, the "AXE method" of electron counting is commonly employed. “A” represents the central atom, “X” represents each of atoms bonded to “A”, and “E” represents the number of electron pairs surrounding the central atom. X + E gives the steric number that is used to predict the molecular geometry that will be formed.

Electron pairs and atoms are counted and the molecule or ion is represented as \(\ce{AX{_{m}E{_{n}}}}\), where m and n are integers telling how many atoms or electron pairs X and E, respectively, represent.

To predict VSEPR molecular geometry:

1. Draw the Lewis structure,
2. Determine the electron group arrangement around the central atom that
   minimizes electronic repulsions by assigning an \(\ce{AX{_{m}E{_{n}}}}\) designation to describe geometry; for quick reference (schematics shown in table below):

1. Find the electron bonding-pair and lone (nonbonding) electrons to identify and deviations from ideal bond angles.

EXAMPLES. Predicting VSEPR Geometry

\(\ce{CO{_{2}}}\)

● has a central atom C with two O atoms bonded to it

● The Lewis structure would be

● To find an \(\ce{AX{_{m}E{_{n}}}}\) designation we look at the Lewis structure and see that m = 2 because 2 O atoms are bonded to the C and that n = 0
because there are no non-bonding electrons. This gives the designation as \(\ce{AX{_{2}E{_{0}}}}\)

● m + n = 2 meaning the molecular is linear

\(\ce{CH{_{4}}}\)

● has a central atom C with four H atoms bonded to it

● The Lewis structure would be

● To find an \(\ce{AX{_{m}E{_{n}}}}\) designation we look at the Lewis structure and see that m = 4 because four H atoms are bonded to the C and that n = 0 because there are no non-bonding electron pairs. This gives the designation as \(\ce{AX{_{4}E{_{0}}}}\)

● m + n = 4 meaning the molecular is tetrahedral or trigonal pyramidal; in this case it is tetrahedral because all four H atoms are equally distributed around the central C as in a tetrahedron.

\(\ce{NH{_{3}}}\)

● has a central atom N with three H atoms bonded to it

● The Lewis structure would be

● To find an \(\ce{AX{_{m}E{_{n}}}}\) designation we look at the Lewis structure and see that m = 3 because three H atoms are bonded to the N and that n = 2 because there is one non-bonding electron pair. This gives the designation as \(\ce{AX{_{3}E{_{1}}}}\)

● m + n = 4 meaning the molecular is tetrahedral or trigonal pyramidal; in this case it is trigonal pyramidal because the non-bonded pair repels the three H atoms slightly away from it.

VSEPR geometry designations. This is a quick reference guide for commonly encountered structural types; there are many more designations. A (red), X (blue), and E (greenish-white)

VSEPR theory gives no information about the presence of multiple bonds, the effects of orbital symmetries, or bond length (the distance between atoms at the most stable position where electrostatic forces are a minimum). Some contend that Bent's rule is capable of replacing VSEPR as a simple model for explaining molecular structure. Bent’s rule states that “atomic s character (spherical) tends to concentrate in orbitals that are directed toward electropositive groups and atomic p character (dumbbell) tends to concentrate in orbitals that are directed toward electronegative groups ”. Otherwise stated, electron distribution around ligands are generally electronegative and thus tend to have more p character since the s character concentrated on the central atom.

Despite these shortcomings, VSEPR theory is a useful visualisation tool of electron distribution for symmetric molecules and continues to be utilized before more sophisticated models need to be invoked. Molecular orbital theory, discussed in a separate wiki entitled Chemical Bonding – Molecular Orbital Theory, is now being used as a more accurate way to visualize distribution of bonding electrons and molecular shape when simpler models like Lewis and VSEPR no longer suffice.