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7.2 Covalent bonding

. . . contain a metal (i.e. Na) bonded to a nonmetal (i.e. Cl). The salt (NaCl) contains Na⁺ ions that are electrostatically attracted to Cl^- ions. result from the transfer of electrons while . . . contain a nonmetal (i.e. C) bonded to a nonmetal (i.e. O). Molecular compounds (i.e. CO₂) contain only covalent bonds. result from the sharing of electrons. Ionic compounds like table salt ( Na   Cl  ) do not exist as a single sodium ion attached to a single chloride ion. In the image to the right, each chloride ion is surrounded by six sodium ions and each sodium ion is surrounded by six chloride ions.

A good analogy would involve trying to identify unique pairs of black and white squares on a chess board.

Which black square is the "mate" of a specific white square?

Because ionic compounds don't form unique molecules, chemists refer to their "molecular weight" as There aren't any Na–Cl molecules . . . . but the formula, NaCl, has a formula mass of 58.443 g/mol (22.990 g/mol Na + 35.453 g/mol Cl)..

Formation of Covalent Bonds

Covalent bonds occur between . . . CS₂, F–F, CCl₄, NI₃, etc.) - atoms on the right side of the periodic table. Covalent bonds also occur between . . . H–Cl, H–O–H, NH₃, etc.). Because ionic bonds are stronger than covalent bonds, ionic compounds have higher . . . thermal energy breaks some intermolecular bonds allowing molecules to slide past each other. and . . . thermal energy breaks all intermolecular bonds allowing molecules to move freely inside their container. points than molecular compounds.

The two hydrogen atoms in a hydrogen molecule, H₂, are held together with a covalent bond. The animation to the right depicts the energy involved in the . . . energy must be added to break chemical bonds (an endothermic process, +ΔH), whereas forming chemical bonds releases energy (an exothermic process, –ΔH).. For the two separated hydrogen atoms to bond, their internuclear distance must decrease.

  As the hydrogen atoms approach each other, their valence electrons (1s¹) begin to interact with each other and establish an attractive force with the nuclei of both atoms  .   At an internuclear distance of 74 pm, the H–H bond has fully formed and both electrons freely move throughout the space around the H₂ molecule. The H–H bondlength of 74 pm corresponds to an energy minima of –7.24 × 10^-19 J  .   Decreasing the internuclear distance further causes a rise in the potential energy as the positive charges in the two nuclei begin to repel each other  . The length of any chemical bond is defined as the internuclear distance at which the lowest potential energy is achieved.

Pure and Polar Covalent Bonds

While a chemical bond consists of two electrons, it is the position of these two electrons that determine the type of bond. In the animation to the right,  the two atoms are too far apart for bonding to occur .

1.  Pure (nonpolar) covalent  - occurs when the two bonding atoms are identical nonmetals . . . . H₂, O₂, Cl₂, etc. Pure covalent bonds are formed from the  equal sharing of electrons .
2.  Polar covalent  - occurs when the two bonding atoms are different nonmetals . . . . SO₂, CCl₄, HCl, etc. Polar covalent bonds are formed from the  unequal sharing  of electrons .
3.  Ionic  - occurs when one bonding atom is a metal and the other is a nonmetal . . . . NaCl, BaI₂, CuSO₄, etc. Ionic bonds are formed from the  transfer  of electrons .

Electronegativity

Electronegativity is a measure of an atom's tendency to attract bonding electrons. If the electronegativity difference between the two bonding atoms is . . . .

1. < 0.4, the bond is pure covalent
2. < 1.8, the bond is polar covalent
3. > 1.8, the bond is ionic

In practice, chemists rarely "calculate" electronegativites. Instead, we use the periodic trends for electronegativity to qualitatively determine the bond type. In the periodic table, electronegativity increases from left to right (→) and from bottom to top (↑). This means that metals are less electronegative than nonmetals and that Fluorine is the most electronegative element. The assignment of bond types based on the quantitative ranges listed above can be qualitatively assigned based on the bonding atoms' position in the periodic table:

1. If both atoms are the same (i.e. Cl–Cl) or it's a C–H bond, the bond is pure covalent
2. If both atoms are nonmetals (i.e. the C–O bond in CO₂ or the C–Cl bond in CCl₄), the bond is polar covalent
3. If one atom is a metal and the other is a nonmetal (i.e. Na–Cl), the bond is ionic

Electronegativity versus Electron Affinity

It can be easy to confuse electronegativity with electron affinity. The electron affinity of an element is the measured energy that is released or absorbed when an atom acquires an electron. Electronegativity describes how tightly a bonding atom attracts the electrons in the bond.