Basic Concepts of Chemical Bonding

Basic Concepts of Chemical Bonding

Lewis Symbols and the Octet Rule

Why are some substances chemically bonded molecules and others are an association of ions?

depends upon the electronic structures of the atoms
nature of the chemical forces within the compounds

A broad classification of chemical forces:

Ionic bonds
Covalent bonds
Metallic bonds

Ionic bonds - electrostatic forces that exist between ions of opposite charge

typically involves a metal with a nonmetal

Covalent bonds - results from the sharing of electrons between two atoms

typically involves one nonmetallic element with another

Metallic bonds

found in solid metals (copper, iron, aluminum)
each metal bonded to several neighboring groups
bonding electrons free to move throughout the 3-dimensional structure

Lets look at the preferred arrangements of electrons in atoms when they form chemical compounds

Lewis Symbols and the Octet Rule

Valence electrons reside in the outer shell and are the electrons which are going to be involved in chemical interactions and bonding (valence comes from the Latin valere, "to be strong").

Electron-dot symbols (Lewis symbols):

convenient representation of valence electrons
allows you to keep track of valence electrons during bond formation
consists of the chemical symbol for the element plus a dot for each valence electron

Sulfur

Electron configuration is [Ne]3s²3p⁴, thus there are six valence electrons. Its Lewis symbol would therefore be:

Note:

The dots (representing electrons) are placed on the four sides of the atomic symbol (top, bottom, left, right)
Each side can accommodate up to 2 electrons
The number of valence electrons in the table below is the same as the column number of the element in the periodic table (for representative elements only)

Atoms often gain, lose, or share electrons to achieve the same number of electrons as the noble gas closest to them in the periodic table

Because all noble gasses (except He) have filled s and p valence orbitals (8 electrons), many atoms undergoing reactions also end up with 8 valence electrons. This observation has led to the Octet Rule:

Atoms tend to lose, gain, or share electrons until they are surrounded by 8 valence electrons

Note: there are many exceptions to the octet rule (He and H, for example), but it provides a useful model for understanding the basis of chemical bonding.

Ionic Bonding

Sodium metal reacts with chlorine gas in a violently exothermic reaction to produce NaCl (composed of Na⁺ and Cl^- ions):

2Na(s) + Cl₂(g) -> 2NaCl(s)

These ions are arranged in solid NaCl in a regular three-dimensional arrangement (or lattice):

The chlorine has a high affinity for electrons, and the sodium has a low ionization potential. Thus the chlorine gains an electron from the sodium atom. This can be represented using electron-dot symbols (here we will consider one chlorine atom, rather than Cl₂):

The arrow indicates the transfer of the electron from sodium to chlorine to form the Na⁺ metal ion and the Cl^- chloride ion. Each ion now has an octet of electrons in its valence shell:

Na⁺ 2s²2p⁶

Cl^- 3s²3p⁶

Covalent Bonding

Ionic substances:

usually brittle
high melting point
organized into an ordered lattice of atoms, which can be cleaved along a smooth line

the electrostatic forces organize the ions of ionic substances into a rigid, organized three-dimensional arrangement

The vast majority of chemical substances are not ionic in nature

gases and liquids, in addition to solids
low melting temperatures

G.N. Lewis

reasoned that an atom might attain a noble gas electron configuration by sharing electrons

A chemical bond formed by sharing a pair of electrons is called a covalent bond

The diatomic hydrogen molecule (H₂) is the simplest model of a covalent bond, and is represented in Lewis structures as:

The shared pair of electrons provides each hydrogen atom with two electrons in its valence shell (the 1s) orbital.

In a sense, it has the electron configuration of the noble gas helium

When two chlorine atoms covalently bond to form Cl₂, the following sharing of electrons occurs:

Each chlorine atom shared the bonding pair of electrons and achieves the electron configuration of the noble gas argon.

In Lewis structures the bonding pair of electrons is usually displayed as a line, and the unshared electrons as dots:

The shared electrons are not located in a fixed position between the nuclei. In the case of the H₂ compound, the electron density is concentrated between the two nuclei:

The two atoms are bound into the H₂ molecule mainly due to the attraction of the positively charged nuclei for the negatively charged electron cloud located between them

For the nonmetals (and the 's' block metals) the number of valence electrons is equal to the group number:

Element	Group	Valence electrons	Bonds needed to form valence octet
F	7A	7	1
O	6A	6	2
N	5A	5	3
C	4A	4	4

Examples of hydride compounds of the above elements (covalent bonds with hydrogen:

Thus, the Lewis bonds successfully describe the covalent interactions between various nonmetal elements

Multiple bonds

The sharing of a pair of electrons represents a single covalent bond, usually just referred to as a single bond

In many molecules atoms attain complete octets by sharing more than one pair of electrons between them.

Two electron pairs shared a double bond

Three electron pairs shared a triple bond

Because each nitrogen contains 5 valence electrons, they need to share 3 pairs to each achieve a valence octet.

N₂ is fairly inert, due to the strong triple bond between the two nitrogens
The N - N bond distance in N₂ is 1.10 Å (fairly short)

From a study of various Nitrogen containing compounds bond distance as a function of bond type can be summarized as follows:

N-N 1.47Å
N=N 1.24Å
N=N 1.10Å

As a general rule, the distance between bonded atoms decreases as the number of shared electron pairs increases

Drawing Lewis Structures

The general procedure...

1. Sum the valence electrons from all atoms

Use the periodic table for reference
Add an electron for each indicated negative charge, subtract an electron for each indicated positive charge

2. Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond

You may need some additional evidence to decide bonding interactions
If a central atom has various groups bonded to it, it is usually listed first: CO₃^2-, SF₄
Often atoms are written in the order of their connections: HCN

3. Complete the octets of the atoms bonded to the central atom (H only has two)

4. Place any leftover electrons on the central atom (even if it results in more than an octet)

5. If there are not enough electrons to give the central atom an octet, try multiple bonds (use one or more of the unshared pairs of electrons on the atoms bonded to the central atom to form double or triple bonds)

Draw the Lewis structure of phosphorous trichloride (PCl₃)

This is an example of a central atom, P, surrounded by chlorine atoms

1. We will have 5(P) plus 21 (3*7, for Cl), or 26 total valence electrons

2. The general symbol, starting with only single bonds, would be:

3. Completing the octets of the Cl atoms bonded to the central P atom:

4. This gives us a total of (18 electrons) plus the 6 in the three single bonds, or 24 electrons total. Thus we have 2 extra valence electrons which are not accounted for. We will place them on the central element:

5. The central atom now has an octect, and there is no need to invoke any double or triple bonds to achieve an octet for the central atom. We are finished.

Draw the Lewis structure for the NO+ ion

1. We will have 5 (N) plus 6 (O) minus 1 (1+ ion), or 10 valence electrons

2. The general structure starting only with single bonds would be:

3. Completing the octet of the O bonded to N:

4. This gives us a total of 6 plus 2 for the single bond, or 8 electrons. There are 2 unaccounted for electrons and we will place them on the N:

5. There are only 4 atoms on the N atom, not enough for an octet, so lets try a double bond between the N and O:

The oxygen still has an octet, but the N only has 6 valence electrons, so lets try a triple bond:

Each atom now has a valence octet. We are finished.

The brackets with the + symbol are used to indicate that this is an ion with a net charge of 1+

Formal Charge

In some cases we can draw several different Lewis structures which fulfill the octet rule for a compound. Which one is the most reasonable?

One method is to tabulate the valence electrons around each atom in a Lewis structure to determine the formal charge. The formal charge is the charge that an atom in a molecule would have if we considered each atom to have the same electronegativity in a compound.

To calculate formal charge, assign electrons to their respective atoms as follows:

All of the unshared electrons are assigned to the atom on which they are found
The bonding electrons are divided up equally between each atom involved in the bond
The number of valence electrons expected in the isolated atom is compared to the actual number of electrons assigned from the Lewis structure:

The formal charge equals the number of valence electrons in the isolated atom, minus the number of electrons assigned in the Lewis structure

Example: Carbon Dioxide (CO₂)

Carbon has 4 valence electrons

Each oxygen has 6 valence electrons, therefore our Lewis structure of CO₂ will have 16 electrons:

One way we could draw the Lewis structure is:

Another way we could draw the Lewis structure is:

Both structures fulfill the octet rule. But what are the formal charges?

Which structure is correct? In general, when several Lewis structures can be drawn the most stable structure is the one in which:

The formal charges are the smallest
Any negative charge is found on the most electronegative atom

In the above case, the second structure is the one with the smallest formal charges (i.e. 0 on all the atoms).

Furthermore, in the first possible Lewis structure the carbon has a formal charge of 0 and one of the oxygens it is bonded to has a formal charge of +1.
Oxygen is more electronegative than Carbon, so this situation would seem unlikely.

It is important to remember that formal charges do not represent the actual charges on the atoms. Actual charges are determined by the electronegativity of the atoms involved.

1996 Michael Blaber