Types of Chemical Bonds: Study Guide (Chemistry 2.1)

What decides the type of bond

In Topic 1.8 you saw that bonding is a tug of war for valence electrons. Each nucleus attracts its own valence electrons and those of its neighbour. How the electrons end up depends on how strongly each atom pulls:

• Both atoms pull about equally hard. They share the electrons equally: a nonpolar covalent bond.
• One atom pulls harder. The shared electrons spend more time near it: a polar covalent bond.
• One atom pulls much harder, and the other holds its electrons weakly. The electrons are, in effect, transferred and ions form: an ionic bond.
• Many metal atoms, all holding their valence electrons weakly. The electrons are shared by the whole solid: metallic bonding.

The property that measures “how hard an atom pulls” on bonding electrons is electronegativity.

Electronegativity and its trends

Electronegativity is the ability of an atom in a bond to attract the shared electrons towards itself. The most common scale is the Pauling scale, which has no unit. Fluorine is the highest, at 3.98.

Period 2	Li	C	N	O	F
Electronegativity	0.98	2.55	3.04	3.44	3.98

Group 17	F	Cl	Br	I
Electronegativity	3.98	3.16	2.96	2.66

The two trends for the main-group (representative) elements are:

• Across a period, electronegativity increases. Each step adds a proton to the nucleus, but the new electron joins the same shell, so shielding by core electrons hardly changes. The effective nuclear charge rises while the distance stays similar. By Coulomb’s law the nucleus attracts the bonding electrons more strongly. From Li to F the value rises by 3.00.
• Down a group, electronegativity decreases. Each step adds a shell. The bonding electrons are further from the nucleus and better shielded, so the attraction is weaker. From F to I the value falls by 1.32.

These are the same reasons that explain the ionization energy trend in Topic 1.7. That is why metals (low ionization energy) also have low electronegativity, and nonmetals on the upper right have high electronegativity. Noble gases are usually left out, because they form few bonds.

Nonpolar covalent bonds

When two atoms of the same element bond, as in H₂ or Cl₂, the shared pair is attracted equally. The electron density is symmetric and neither end of the bond is charged. This is a nonpolar covalent bond.

Atoms with similar electronegativities behave in almost the same way. The standard example is the carbon–hydrogen bond. Carbon (2.55) is slightly more electronegative than hydrogen (2.20), but the difference of 0.35 is small, and C–H bonds are treated as effectively nonpolar. This is why hydrocarbons such as methane and octane behave as nonpolar substances. The P–H bond (difference 0.01) is another near-perfect example.

Polar covalent bonds

When the electronegativities are unequal, the shared electrons are pulled towards the more electronegative atom. They are still shared, but not evenly. This is a polar covalent bond.

• The more electronegative atom carries a partial negative charge, written δ−. The other atom carries a partial positive charge, δ+. “Partial” means less than the charge of one electron: no ion has formed.
• The separation of charge is called a bond dipole. You may see it drawn as an arrow along the bond, pointing towards the δ− atom, with a small cross on the tail at the δ+ end.
• For single bonds, a bigger electronegativity difference gives a bigger bond dipole. H–Cl (difference 0.96) is more polar than N–H (0.84).

In H–Cl, chlorine (3.16) is more electronegative than hydrogen (2.20), so hydrogen is δ+ and chlorine is δ−: H(δ+)–Cl(δ−).

Every polar bond has some ionic character: the electron density is shifted, so the bond behaves partly as if the atoms were ions. The bigger the shift, the more ionic character.

A continuum, not three boxes

Because every polar bond is partly ionic, there is no sharp boundary between “covalent” and “ionic”. Bonds sit along a sliding scale.

Ionic bonds: more than a number

The electronegativity difference is a useful guide, but it is not the only factor in calling a bond ionic or covalent. Two rules of thumb work well:

• A metal with a nonmetal usually forms an ionic compound (NaCl, MgO, CaF₂). Metals have low ionization energies and low electronegativities, so their valence electrons are easily pulled away.
• Two nonmetals usually form covalent bonds (H₂O, CO₂, NH₃), even when the difference is large.

Hydrogen fluoride shows why a single cut-off number fails. Its difference is 1.78, bigger than the “about 1.7” that some books quote for ionic bonding. Yet HF is made of discrete covalent molecules, as you would expect for two nonmetals.

So the best way to decide the type of bonding in a compound is to examine its properties.

Metallic bonding

In a piece of metal there is no partner atom with a high electronegativity to take the electrons. Instead, each metal atom releases its valence electrons into the whole solid. The valence electrons are delocalized: they are not tied to any one atom or any one bond. What remains is a regular array of metal cations held together by their attraction to this mobile “sea” of electrons.

Using properties to identify bonding

Each type of bonding leaves fingerprints on the properties of a substance. You will study the reasons in Topics 2.3, 2.4 and 3.2; for now, use them as evidence.

Evidence	Ionic	Metallic	Covalent molecules
Melting point	high	varies, often high	usually low
Conducts electricity as a solid?	no	yes	no
Conducts when molten (or dissolved in water)?	yes	yes (molten)	usually no
Response to hammering	shatters (brittle)	bends or flattens (malleable)	soft or crumbly solids

A few covalent substances, such as diamond and quartz, are giant networks with very high melting points. You will meet them in Topic 3.2.

Worked example 1: polarity of carbon bonds

Question. Electronegativities: H 2.20, C 2.55, N 3.04, O 3.44, F 3.98. For the single bonds C–H, C–N, C–O and C–F, (a) identify the atom with the partial negative charge in each bond, and (b) rank the bonds in order of increasing bond dipole.

(a) Find the more electronegative atom in each pair.

Bond	Difference	Partial negative atom
C–H	2.55 − 2.20 = 0.35	C (but the bond is effectively nonpolar)
C–N	3.04 − 2.55 = 0.49	N
C–O	3.44 − 2.55 = 0.89	O
C–F	3.98 − 2.55 = 1.43	F

(b) For single bonds, a larger difference gives a larger bond dipole: C–H < C–N < C–O < C–F.

Interpretation. You could also reach the ranking without numbers. N, O and F lie in the same period, to the right of carbon, and electronegativity rises across a period. The further right the partner, the more it pulls the shared pair. Notice that carbon is δ+ in three of the bonds but slightly δ− in C–H.

Worked example 2: a claim from properties

Question. A student tests three white solids, labelled P, Q and R. The data are invented for this example.

Solid	Melting point (°C)	Conducts as solid?	Conducts when molten?	When hit with a hammer
P	95	no	no	crumbles
Q	860	no	yes	shatters
R	650	yes	yes	flattens

Make a claim about the type of bonding in each solid and support it with evidence.

1. P: covalent molecules. It melts at a low temperature and does not conduct in either state. With no ions and no delocalized electrons, there are no charged particles that can move.
2. Q: ionic. It does not conduct as a solid but does when molten. Ions are present but fixed in place in the solid; melting lets them move and carry charge. The high melting point and brittleness fit strong attractions between oppositely charged ions.
3. R: metallic. It conducts as a solid, which needs charged particles that can move inside the solid: delocalized valence electrons. It flattens instead of shattering, as metals do.

Why this is the strongest method. The test tells you about the actual substance, whatever its electronegativity values suggest.

Worked example 3: predicting bond types from the elements

Question. Electronegativities: N 3.04, O 3.44, S 2.58, Cl 3.16, K 0.82, Br 2.96. Predict the type of bonding in (a) potassium bromide, KBr; (b) the N–Cl bonds in NCl₃; (c) the S–O bonds in SO₂; (d) brass, an alloy of copper and zinc.

(a) K is a group 1 metal and Br is a nonmetal. The difference is 2.96 − 0.82 = 2.14. Both clues agree: ionic (K⁺ and Br⁻).

(b) N and Cl are both nonmetals. The difference is 3.16 − 3.04 = 0.12, which is tiny. The bonds are covalent and almost nonpolar; chlorine carries a very small δ−. (Students often guess that nitrogen is the more electronegative atom here, but on the Pauling scale chlorine is slightly higher.)

(c) S and O are both nonmetals, so covalent. The difference is 3.44 − 2.58 = 0.86, so the bonds are polar covalent, with O as δ− and S as δ+.

(d) Copper and zinc are both metals. Neither atom holds its valence electrons strongly enough to take electrons from the other, so the electrons are delocalized across the whole alloy: metallic bonding. Do not use the electronegativity difference here; two metals do not form ionic or covalent bonds with each other in this course.

Common misconceptions

• “Ionic and covalent are completely different kinds of bond.” They are two ends of one continuum. Every polar covalent bond has some ionic character.
• “A difference above 1.7 always means ionic.” There is no sharp cut-off. HF (1.78) is covalent. Use the type of elements and the properties as well.
• “The more electronegative atom becomes a negative ion.” In a polar covalent bond it only gains a partial charge, δ−. A full charge means an ion, which happens in ionic compounds.
• “C–H bonds are polar because carbon is more electronegative.” The difference (0.35) is so small that C–H bonds are treated as nonpolar.
• “Metal atoms share electron pairs in fixed bonds.” Metallic bonding has no fixed pairs. The valence electrons are delocalized over the whole solid.
• “A molecule with polar bonds must be polar.” Bond polarity and molecular polarity are different. Molecular shape decides whether bond dipoles cancel, which you will meet in Topics 2.7 and 3.1.
• “Electronegativity is a measured energy.” It is a relative scale with no unit. Electron affinity is the measured energy (Topic 1.7).

Where this leads

Next, Topic 2.2, Intramolecular Force and Potential Energy, looks inside a single bond: why atoms settle at a particular distance and what decides bond length and bond strength. Topics 2.3 and 2.4 then explain the properties of ionic solids and metals that you used as evidence here. Try the practice questions now, then use the revision notes and the checklist to consolidate.