Unit 2 · Topic 2.1

2.1 Types of Chemical Bonds

Chemical bonds fall on a range from equal sharing of electrons (nonpolar covalent) through unequal sharing (polar covalent) to transfer (ionic). Where a bond falls depends mostly on the difference in electronegativity between the two atoms, while metals hold their atoms together with electrons spread over the whole solid.

Key terms

electronegativity
nonpolar covalent bond
polar covalent bond
ionic bond
metallic bond
partial charge

Electronegativity

Electronegativity is how strongly an atom pulls shared electrons in a bond toward itself. It increases from left to right across a period and decreases down a group, for the same Coulomb's law reasons as ionization energy (topic 1.7): more effective nuclear charge and a smaller radius mean a stronger pull. Fluorine is the most electronegative element (3.98 on the Pauling scale); oxygen (3.44), chlorine (3.16) and nitrogen (3.04) come next.

When two atoms share electrons, the difference in their electronegativities decides how evenly the electrons are shared.

Noble gases are usually left out of electronegativity tables, since most of them rarely form bonds.

Three kinds of covalent and ionic bonding

Nonpolar covalent: the atoms have equal or very similar electronegativity, so electrons are shared about equally. H–H and Cl–Cl are perfectly nonpolar. C–H bonds (2.55 vs 2.20) are treated as essentially nonpolar too.
Polar covalent: electrons are shared unequally. The more electronegative atom gets a partial negative charge (δ−) and the other gets a partial positive charge (δ+). In H–Cl, chlorine is δ− and hydrogen is δ+. This separation of charge is a bond dipole.
Ionic: the difference is so large that electrons are essentially transferred, making cations and anions that attract each other. NaCl (0.93 vs 3.16) is the classic example.

Bond dipoles

For single bonds, a bigger electronegativity difference means a bigger bond dipole. H–F (difference 1.78) is more polar than H–Cl (0.96).

You can show a bond dipole with an arrow pointing toward the δ− end, with a small cross on the δ+ end. In H–F the arrow points at F. Bond dipoles become important in topic 2.7, where you decide whether a whole molecule's dipoles add up or cancel.

A continuum, not three boxes

There's no sharp line between polar covalent and ionic. Every polar bond has some ionic character, and the change from covalent to ionic is gradual. Some textbooks quote cutoff values, but the AP exam treats bonding as a continuum.

In practice, use this rule of thumb: a metal bonded to a nonmetal is ionic, and two nonmetals bonded together are covalent. When the rule and the electronegativity difference disagree, the best evidence is how the substance behaves. A compound that has a high melting point and conducts electricity when melted is behaving like an ionic compound.

Metallic bonding

In a metal, the atoms' valence electrons are delocalized: they don't belong to any one atom but move freely through the whole solid. You can picture positive metal ions sitting in a 'sea' of electrons. Topic 2.4 builds on this model, and topic 3.2 uses it to explain why metals conduct electricity.

Worked examples

Try each one yourself first, then open the solution.

Example 1
Ranking bond polarity
Rank these bonds from least polar to most polar: O–H, C–H, H–F, N–H. Electronegativities: H 2.20, C 2.55, N 3.04, O 3.44, F 3.98.
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1. Step 1: Find each difference: C–H = 0.35, N–H = 0.84, O–H = 1.24, H–F = 1.78.
2. Step 2: A larger difference means a larger bond dipole.
3. Step 3: In each polar bond, the atom other than H carries the δ− charge.
Answer: C–H < N–H < O–H < H–F
Example 2
Classifying with evidence
Solid X melts at 801 °C. It does not conduct electricity as a solid, but it does when melted or dissolved in water. Is X more likely ionic or covalent (molecular)? Justify.
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1. Step 1: A high melting point means the particles are held by strong attractions.
2. Step 2: Conducting when melted or dissolved but not as a solid means charged particles are present that can move only once the solid breaks down. Those are ions.
3. Step 3: A molecular substance would have a low melting point and wouldn't conduct in any state.
Answer: Ionic. The high melting point and conductivity only when molten or dissolved show it is made of ions held in a rigid lattice (this is sodium chloride's melting point).
Example 3
Where are the partial charges? (classic trap)
A student says that in a C–Cl bond, carbon is partially negative because carbon has more electrons than chlorine. Correct the statement.
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1. Step 1: Partial charges depend on electronegativity, not on how many electrons each atom has.
2. Step 2: Chlorine (3.16) is more electronegative than carbon (2.55), so the shared electrons spend more time near chlorine.
3. Step 3: Also, carbon has 6 electrons and chlorine has 17, so the student's premise is wrong as well.
Answer: Chlorine is δ− and carbon is δ+, because chlorine is more electronegative.

Common mistakes

Treating a fixed electronegativity cutoff as the rule for ionic bonding. The exam treats bonding as a continuum and uses the metal/nonmetal rule plus properties.
Putting the δ− on the less electronegative atom. Electrons are pulled toward the more electronegative atom.
Saying a C–H bond is strongly polar. It's treated as nonpolar.
Describing metallic bonding as atoms sharing electron pairs with neighbors. In metals, valence electrons are delocalized across the whole solid.

On the exam

Expect to compare bond polarity using given electronegativity values, or to explain the electronegativity trend using effective nuclear charge and atomic radius.
When asked to classify bonding from data, cite a specific property, such as melting point, conductivity in each state or solubility, and connect it to the particles present.

Connected topics

Videos

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