Covalent Bonds

Chapter 4 described atoms that hand electrons over and then cling together as oppositely charged ions. Plenty of familiar substances are not like that. The oxygen you breathe is an element, as Chapter 1 listed it, yet it is not a crowd of separate atoms: it is made of atoms joined in pairs. Wikipedia's Covalent bond article defines the kind of link that does the joining: "a chemical bond that involves the sharing of electrons to form electron pairs between atoms". This chapter covers sharing instead of transferring, how to count and draw shared pairs, why some bonds pull harder to one side than the other, and why the shape of a molecule matters. It uses the valence electrons introduced in Chapter 2 and the group rule from Chapter 4.

Sharing Instead of Handing Over

The Covalent bond article calls the shared electron pairs "shared pairs" or "bonding pairs". It explains why atoms do this with Lewis's idea: "Lewis proposed that an atom forms enough covalent bonds to form a full (or closed) outer electron shell." That is the octet rule of Chapter 4 again, reached by a different route. Its example is hydrogen: "the hydrogen atoms share the two electrons via covalent bonding", so each hydrogen atom in an H₂ unit has two electrons around it, a full first shell (the full first shell of two is my addition, not a quote from the page (check)).

Single, Double and Triple Bonds

Atoms can share more than one pair. The Covalent bond article describes the three cases: "A single bond is usually a σ bond", "a double bond between two given atoms consists of one σ and one π bond", and "a triple bond is one σ and two π bonds". The Greek-letter names describe how the electron clouds overlap, which this course does not go into. What matters here is the count: a single bond is one shared pair, a double bond two pairs, a triple bond three. Counting pairs gives the bond order. The page has no general definition of the term, but for oxygen it writes "The bond order is 1+0.5+0.5=2", and the Diatomic molecule article says that diatomic oxygen has a double bond.

The shared-electron column is my own arithmetic, two electrons per pair, run in Python: 1 x 2 = 2, 2 x 2 = 4 and 3 x 2 = 6.

Molecules and Molecular Formulae

Wikipedia's Molecule article defines a molecule as "a group of two or more atoms that are held together by attractive forces known as chemical bonds". A molecular formula "reflects the exact number of atoms that compose the molecule". The article's examples are water, "two hydrogen atoms and one oxygen atom; H2O", and the oxygen molecule, "two atoms in the oxygen molecule (O2)". The subscript counts atoms of the element in front of it, and no subscript means one. This is the difference from Chapter 4: the article says "No typical molecule can be defined for salts", and that atoms joined by ionic bonds "are typically not considered single molecules". Sodium chloride was a lattice with a formula ratio; water is a countable unit of three atoms.

Diatomic Elements: The Oxygen Promise

Chapter 1 gave oxygen gas as an example of an element, which it defined as a substance made of one kind of atom. Here is how that fits with atoms and molecules. The Diatomic molecule article says such molecules "are molecules composed of only two atoms, of the same or different chemical elements". It lists the elements that are diatomic at 1 bar and 25 °C: "the gases hydrogen (H2), nitrogen (N2), oxygen (O2), fluorine (F2), and chlorine (Cl2), and the liquid bromine (Br2)". It adds that iodine forms a diatomic gas at slightly higher temperatures. So the oxygen in air is O₂: two oxygen atoms joined by a double bond, still one element because both atoms are the same kind. Chlorine is on the list too: the chlorine atom that took an electron from sodium in Chapter 4 is not how chlorine exists on its own, which is Cl₂ (the link to Chapter 4 is my own (check)).

Drawing a Lewis Structure

The Lewis structure article describes these as "diagrams that show the bonding between atoms of a molecule", where "excess electrons that form lone pairs are represented as pairs of dots" and "lines are drawn between atoms that are bonded to one another". The page's own full worked example is the nitrite ion, and it shows a water molecule only as a picture caption, so the steps below are my own application of the page's rules, using the valence-electron group rule from Chapter 4 (check).

      ··
  H — O — H
      ··

  Each line   = one shared pair (2 electrons)
  Each ··     = one lone pair (2 electrons)
  Total       = 2 bonds x 2 + 2 lone pairs x 2 = 8 electrons

Everything in the drawing is accounted for: the 8 valence electrons counted in step 1 equal the 8 electrons drawn. Counting back like this is the best check that a Lewis structure is right. Note that the drawing shows which atoms are joined, not the angle between the bonds; the shape section below comes to that.

Electronegativity and Polar Bonds

In a shared pair, the two atoms do not necessarily pull equally. Wikipedia's Electronegativity article defines electronegativity as "the tendency for an atom of a given chemical element to attract shared electrons". The Covalent bond article says covalent character "is greatest between atoms that have similar electronegativities", and "an unequal relationship creates a polar covalent bond such as with H−Cl". Equal sharing, as in H–H, gives a nonpolar bond.

The Chemical polarity article gives a rough guide: polar bonds "generally occur when the difference in electronegativity between the two atoms is roughly between 0.5 and 2.0", with smaller differences nonpolar and larger ones ionic. The Electronegativity article itself gives no numeric cut-offs and says bonding falls along a continuous scale from covalent to ionic, so treat the numbers as a guide, not a law. Below are Pauling-scale values from the Electronegativity article, and the differences I calculated in Python.

The last row ties back to Chapter 4: the sodium chloride example sits on the ionic side of the guide, and the article's idea of a continuous scale means it is not a sharp line. In a polar bond the more electronegative atom is at the negative end of the dipole, so in H–Cl the chlorine end is the negative one (the Electronegativity article).

Shape: Why Water Is Bent

Water is the standard example. The Properties of water article gives its shape as "bent", not linear, and the Chemical polarity article says the bent geometry means the O–H bond dipoles do not cancel, so the molecule has a net dipole and is polar. The Properties of water article states the angle twice: "The hydrogen–oxygen–hydrogen angle is 104.45°" in one place, and a gas-phase bend angle of 104.48° in another, which differ by 0.03° (Python). So say "about 104.5°" and note the page is not consistent to the second decimal place. The methane article calls the molecule "a tetrahedral molecule with four equivalent C–H bonds"; it did not give the bond angle in the part I read, so no angle is given here.

Ionic and Covalent Compounds Compared

Chapter 4 gave the properties of ionic compounds. The Covalent bond article describes covalent substances in two groups and does not compare them directly with ionic ones, so the table below sets the two side by side using only wording from the pages I read; a blank cell means the pages did not say.

The table shows that "covalent" covers two quite different kinds of material, and that one set of properties does not fit them all. The pages give no melting points for network solids, so the comparison stays qualitative. For two molecular examples, the Properties of water article gives 0.00 °C for melting and 99.98 °C for boiling; the Methane article gives −182.456 °C and −161.49 °C. Compare these with sodium chloride's 801 °C melting point from Chapter 4. Three substances are a small sample, not proof of a pattern.

Where to Read More

• Encyclopaedias. The Wikipedia articles Covalent bond, Molecule, Diatomic molecule, Lewis structure, Electronegativity, Chemical polarity, Properties of water and Methane are the sources used here. See the note above for how much of each was read.
• A school textbook. Look for worked Lewis structures of ammonia, carbon dioxide and nitrogen, and for the VSEPR method of predicting shape, which this chapter leaves out.
• Practical. A molecular model kit (balls and sticks) makes the bent shape of water and the tetrahedral shape of methane easy to see.

Exercises