1902-1916 - G.N. Lewis

Described atoms with a cube model, explaining that the electron in an atom were arranged in positions at the corners of a cube.  This model was based upon the following assumptions:  The number of electrons in an atom's outermost cube is equal to the number of electrons lost when an atom forms a positive ion.  Each successive neutral atom on the periodic table (as the atomic number increases) has one more electron it its outermost cube.  To fill the corners of the cube, it takes eight electrons, or an octet.  When a cube has filled its octet, it becomes the center about which another cube can be built.

Lewis explained simple ionic compounds by assuming that atoms with less than four electrons in their outermost cube transferred those electrons to another atom which needed to gain that many electrons to fill its octet.  Both atoms, then, would end up with an outer cube that was either completely filled or completely empty.

e.g.  Na loses one electron to Cl to form Na⁺ and Cl^-.

It was later understood why Lewis's "octet" theory was valid.  It requires eight electrons to fill the s (1) and p (3) orbitals in one shell which are the outermost orbitals in a shell.  The outermost shell of an atom came to be known as the valence shell, and the electrons which reside in the valence shell are now known as valence electrons.

Valence electrons - The electrons on an atom that are not present in the previous atom with a filled-shell electron configuration, ignoring filled d and f orbitals.

e.g.  Bromine has the electron configuration:  [Ar] 4s²3d¹⁰4p⁵.  If we add up the electrons in the s and p orbitals only, we find that bromine has 7 valence electrons.

Lewis later found that atoms could fill their outermost cubes by sharing one or more pairs of electrons.

e.g.  One Cl atom can share one of its outermost electrons with another Cl atom to form the diatomic molecule Cl₂.

e.g.  One O atom can share two of its outermost electrons with another O atom to form the diatomic molecule O₂.

Since atoms always seemed to share pairs of electrons when they formed bonds, Lewis changed his cube model with the eight electrons in the corners to a model with pairs of electrons.  These models are known as Lewis Structures (or Lewis Dot Structures).

e.g.  The bonding of two Cl atoms for form Cl₂ would now be written as:

These Lewis Structures models are still in use today.  The only major difference between Lewis's system, and the system in use today is that the covalent bonds between the atoms today are represented by a line instead of two dots.  Double bonds are represented by a double line and triple bonds are represented by triple lines.

Intramolecular Forces of Attraction  - The forces of attraction that exist between bonds within a molecule.

• Ionic bond - A bond between two ions which are formed from the TRANSFER or electrons from one (which forms a positive ion) to another (which forms a negative ion).  The bond is held together by the force of attraction from the opposite charges of the ions formed.
• Ionic compounds dissociate into their ions when they dissolve in water.  This makes them GOOD CONDUCTORS of electricity in aqueous solution.
NaCl_(s) Na⁺_(aq) + Cl^-_(aq)
• Covalent bond - A bond between two atoms formed by the SHARING of electrons.
• Covalent compounds cannot dissociate when they dissolve in water.  This makes them POOR CONDUCTORS of electricity in aqueous solution.
C₆H₁₂O_6(s) C₆H₁₂O_6(aq) Electronegativity and Polarity
• Electronegativity - The tendency of an atom to draw electrons in a bond toward itself.
• There are two periodic trends concerning electronegativity.
• As you move down a group, electronegativity decreases.
• As you move across a period, electronegativity increases.

Bonds can be classified according to the difference in electronegativities of the atoms (EN).  Bonds are:

• ionic if EN > 1.8.
• polar covalent if 1.8 EN  0.4
• nonpolar covalent if EN < 0.4
The larger the difference in the electronegativities of the atoms in a bond, the stronger the strength of the bond.  As the bond becomes stronger, melting and boiling points generally increase as well.

When comparing bonds with the same atoms but with different oxidation states, we have to consider another factor besides the difference in the electronegativities of the atoms since this will be the same if the atoms are the same.  As the oxidation number on an atom increase, its ability to draw electrons in a bond toward itself increases as well.  For example, the electronegativity of an Mn atom in Mn₂O₇ (oxidation number of +7) is much greater than the electronegativity of an Mn atom in MnO (oxidation number of +2).

Polar Versus Nonpolar Bonds

Nonpolar covalent - Covalent bond in which the electrons are shared equally.

e.g.  The diatomic elements (Br, I, N, Cl, H, O, and F) are all examples of nonpolar covalent bonds.

Two general relationships between single, double and triple bonds

1. BOND STRENGTH increases from single to double to triple

 

 
 




SINGLE < DOUBLE < TRIPLE

2. BOND LENGTH decreases from single to double to triple

 

 
 




SINGLE > DOUBLE > TRIPLE

Polar covalent - Covalent bond in which the electrons are NOT shared equally.  The charged ends are called dipoles and are represented by the symbol.

The end of the bond with the larger electronegativity is the slightly charged negative end.

The end of the bond with the smaller electronegativity is the slightly charged positive end.

In this H-Cl bond, for example, the bond is covalent since the EN < 1.8, but the electrons are shared unequally, so the bond is polar.  The Cl atom has the higher EN so the electrons are drawn more toward the Cl atom and away from the H atom.  Since electrons carry a negative charge, this gives the Cl atom a slightly negative charge.  The H atom, consequently, has a slightly positive charge.

Polar Versus Nonpolar Molecules

Molecules that contain polar covalent bonds may or may not be polar molecules.  The polarity of a molecule is determined by measuring the dipole moment which is represented by the symbol .  This depends on two factors:

The degree of the overall separation of charge between the atoms in the bond

The distance between the positive and negative poles

If there are equal polar bonds that balance each other around the central atom, then the overall molecule will be NONPOLAR with no dipole moment, even though the bonds within the molecule may be polar.

In this molecule, CH4, the C atom has a higher EN than the surrounding H atoms.  This makes the individual C-H bonds polar with a slightly negative charge on the C atom and a slightly positive charge on the H atoms.  However, the polarities on these bonds balance each other out around the central C atom.  The overall dipole moment of the molecule is therefore 0 and the molecule is nonpolar.

If there are unequal polar bonds around the central atom, then the overall molecule will be POLAR with an a dipole moment.

In the this molecule, CH3Cl, all of the bonds around the central C atom are not equal.  The the Cl atom has a higher EN than the  central C atom, so the electrons are more drawn toward the Cl atom.  The H atoms have a lower EN than the central C atom and cannot balance out the large EN of the Cl atom.  Therefore, the overall molecule has a dipole moment with a slightly negative charge on the Cl end and a slightly positive charge on the H ends, making the molecule polar.

In this molecule, H2O, the two O-H bonds are polar but are equal to each other.  However, the central O atom has a higher EN than the H atoms so the electrons are drawn toward the O atom.  The O atom also has two pairs of nonbonding electrons, and combined the nonbonding electrons and the electrons drawn toward the O from the O-H bonds result in an overall dipole moment.  Therefore, this molecule is polar.

Check out more about bonding

• Chemical Bonds in the MIT Biology Hypertextbook