2.1 Types of Chemical Bonds

Chemical bonds form when atoms share or transfer valence electrons to reach a lower-energy, more stable state. Bond type (ionic, covalent, or metallic) depends mostly on electronegativity and whether the atoms are metals or nonmetals, but compound properties give the stronger evidence. For AP Chemistry, connect bond type to melting point, conductivity, polarity, and Coulombic attractions.

Why This Matters for the AP Chemistry Exam

Topic 2.1 is your foundation for explaining how the elements in a bond connect to the properties you observe. The main skill here is making a scientific claim: you predict bond type from electronegativity and periodic position, then back it up with evidence like melting point, conductivity, and atomic size.

This shows up in both multiple-choice and free-response thinking. You may need to compare melting points using Coulomb's law, justify why one bond is more polar than another, or explain why a metal conducts electricity. Getting comfortable with electronegativity trends and the ionic-covalent continuum sets you up for the rest of Unit 2, where you build Lewis structures, analyze geometry, and predict polarity.

Key Takeaways

• Electronegativity increases left to right across a period and decreases down a group; use the shell model and Coulomb's law to explain why.
• Bonds between atoms of similar electronegativity are nonpolar covalent; unequal electronegativity makes a bond polar covalent, with the more electronegative atom getting the partial negative charge.
• All polar bonds have some ionic character. Ionic versus covalent is a continuum, not two separate boxes.
• Metal plus nonmetal usually means ionic; two nonmetals usually means covalent, but examining the compound's properties is the most reliable way to identify bond type.
• Coulomb's law explains ionic bond strength: larger charges and smaller ions produce stronger attractions and higher melting points.
• In metallic bonding, valence electrons are delocalized into a shared "sea" rather than belonging to any single atom.

Principles of Bonding

Atoms bond to reach a more stable, lower-energy state. The electrons that do this work are valence electrons, not core electrons.

Valence Electrons

Valence electrons are the outermost electrons and they determine how an atom bonds. Quick review from Unit 1:

• Valence electrons are found in the s and p orbitals of the outermost occupied shell.
• A large jump in successive ionization energies signals how many valence electrons an element has.
• Elements in the same group have the same number of valence electrons, so they tend to form similar compounds.

Electronegativity

Electronegativity measures an atom's ability to attract shared electrons in a bond. The two trends you need:

• Across a period (left to right): electronegativity increases. More protons in the nucleus means a stronger pull on electrons.
• Down a group: electronegativity decreases. A larger atomic radius puts valence electrons farther from the nucleus, weakening the attraction.

You can explain both trends qualitatively using the shell model and effective nuclear charge.

Coulomb's Law

Coulomb's law describes the strength of attraction between charged particles and depends on two things:

1. Magnitude of charge: greater charges produce stronger attraction.
2. Distance between nuclei: closer particles produce stronger attraction.

This directly explains the electronegativity trend down a group. As atoms get larger, valence electrons sit farther from the nucleus, the attractive force drops, and electronegativity decreases.

Ionic Bonding

Ionic bonds form by the transfer of valence electrons from one atom to another, usually from a metal to a nonmetal. A common example is sodium and chlorine:

$\text{Na}(s) + \frac{1}{2}\text{Cl}_2(g) \rightarrow \text{NaCl}(s)$

NaCl is a brittle salt with a high melting point. Ionic compounds are held together by strong electrostatic attractions between positive and negative ions, not by shared electrons. Those attractions take a lot of energy to break, which is why ionic solids have high melting and boiling points.

Ionic compounds form a crystal lattice: ions arranged in a repeating, three-dimensional pattern that maximizes attractions between oppositely charged ions while minimizing repulsions. When an ionic compound is melted or dissolved in water, the ions can move freely and conduct electricity. In the solid state, the ions are locked in place and do not conduct.

When sodium and chlorine react, sodium transfers a valence electron to chlorine. The atom that loses an electron (sodium) becomes a positively charged cation, and the atom that gains an electron (chlorine) becomes a negatively charged anion.

Applying Coulomb's Law to Ionic Compounds

Coulomb's law tells you that larger charges and smaller ions create stronger attractions. When comparing melting points of ionic compounds, look at charge and ionic size.

Worked Comparisons

1. Higher melting point: MgF2 or NaF? Mg has a +2 charge while Na has only +1. The larger charge means stronger attraction, so MgF2 has the higher melting point.
2. Higher melting point: LiF or NaBr? Both have +1 / -1 charges, so size decides it. Li and F are in period 2 (smaller ions), while Na is in period 3 and Br is in period 4 (larger ions). Smaller ions sit closer together, so LiF has the higher melting point. You may be asked to justify this trend on a free-response question.

Covalent Bonding

In covalent bonds, electrons are shared between atoms, usually two nonmetals. There are two types based on the electronegativities involved:

1. Polar covalent bond: unequal sharing, so charge is distributed unevenly.
2. Nonpolar covalent bond: roughly equal sharing, so charge is distributed evenly.

You will analyze polarity in more depth when you reach molecular geometry later in this unit.

Telling Nonpolar and Polar Apart

The deciding factor is electronegativity difference.

Nonpolar covalent: atoms of similar electronegativity share electrons evenly. Two oxygen atoms in O2 pull on the shared electrons with equal strength, so the bond is nonpolar. A C-H bond is also treated as effectively nonpolar even though carbon is slightly more electronegative than hydrogen. Think "nonpolar = balanced."

Polar covalent: atoms of unequal electronegativity share electrons unevenly. In water, hydrogen has an electronegativity around 2.2 and oxygen around 3.44, so oxygen pulls the shared electrons toward itself and develops a partial negative charge while each hydrogen develops a partial positive charge.

This is where the continuum matters. All polar bonds have some ionic character because electrons are partially pulled toward one atom. The line between ionic and covalent is not sharp. As the electronegativity difference grows, a single bond becomes more polar and the bond dipole gets larger. The symbol δ (delta) marks a partial positive or partial negative charge.

For now, remember: greater electronegativity difference means a larger bond dipole and more ionic character.

Metallic Bonding

Besides electron transfer (ionic) and electron sharing (covalent), metals use a third type: metallic bonding. Here the valence electrons are delocalized, meaning they are not tied to any single atom. Instead they form a shared "sea" of electrons that moves freely throughout the metal.

Picture positively charged metal cations surrounded by a mobile electron sea. This arrangement explains several metal properties:

• Electrical conductivity: mobile electrons carry current easily.
• Malleability and ductility: atoms can slide past each other without breaking the bonding because the electron sea adjusts.
• Metallic luster: free electrons absorb and re-emit light.
• Thermal conductivity: mobile electrons transfer energy efficiently.

Which Bond Will Form?

Electronegativity differences give useful guidelines, but they are not the whole story.

• Ionic Bonds
  • Usually form between a metal and a nonmetal
  • Join a cation (positive ion) and an anion (negative ion)
  • Tend to appear with large electronegativity differences
• Covalent Bonds
  • Usually form between two nonmetals
  • Polar covalent: moderate electronegativity difference, uneven sharing
  • Nonpolar covalent: similar electronegativity, even sharing (for example, C-H)
• Metallic Bonds
  • Form between metal atoms of the same or different elements
  • Involve delocalized valence electrons
  • Produce the characteristic properties of metals

The most reliable way to identify bond type is to examine the compound's actual properties, not just the electronegativity numbers.

Using Properties to Identify Bond Type

• High melting point and conducts electricity when melted or dissolved (but not as a solid): most likely ionic. Ions are locked in place as a solid but free to move once melted or dissolved.
• Low melting point and does not conduct electricity in any state: most likely a molecular compound held together by covalent bonds.
• Conducts electricity as a solid, has a shiny appearance, and can be hammered into sheets or drawn into wires: a metal with metallic bonding. The delocalized electrons conduct even while atoms stay in fixed positions.
• High melting point but does not conduct electricity in any state: a covalent network solid. You will see more of these in Unit 3.

These categories are idealized. Real compounds often land somewhere between them, which reflects the continuous nature of bonding.

How to Use This on the AP Chemistry Exam

Make a Scientific Claim

The core skill in this topic is making a claim and supporting it. When asked to predict or compare bonding, state your claim (for example, "this bond is more polar") and back it with electronegativity, periodic position, or measured properties.

Problem Solving

• For melting point comparisons of ionic compounds, apply Coulomb's law. Check charge first, then ionic size. Larger charges and smaller ions mean stronger attraction and higher melting points.
• For polarity comparisons, compare electronegativity differences. The bond with the larger difference is more polar and has the larger dipole.

Free Response

• Be ready to justify a periodic trend, such as why LiF has a higher melting point than NaBr, using charge and distance.
• Be ready to explain why a metal conducts electricity using delocalized electrons.
• When asked to identify bond type from data, cite the property evidence (conductivity in different states, melting point, malleability) rather than just stating a rule.

Common Trap

Do not rely only on an electronegativity-difference cutoff to declare a bond ionic or covalent. The exam expects you to know that bond character is a continuum and that the compound's properties are the strongest evidence.

Common Misconceptions

• "Ionic and covalent are completely separate." They sit on a continuum. Polar covalent bonds carry partial ionic character, and the larger the electronegativity difference, the more ionic a bond behaves.
• "A fixed electronegativity-difference number always decides bond type." It is a guideline, not a rule. Metal-versus-nonmetal identity and the compound's actual properties are more reliable.
• "Ionic solids conduct electricity in every state." They conduct only when melted or dissolved, because the ions must be free to move. As a solid, the ions are locked in the lattice.
• "Nonpolar means the atoms must be identical." Bonds between different atoms with very similar electronegativity, like C-H, are treated as effectively nonpolar.
• "In metallic bonding, electrons belong to specific atoms." The valence electrons are delocalized across the whole metal, which is what gives metals their conductivity and malleability.
• "Higher charge and larger size both raise melting point." Higher charge raises it, but larger ions actually lower the attraction because the nuclei are farther apart. Smaller ions give stronger attraction.

Related AP Chemistry Guides

• Unit 2 Overview: Compound Structure and Properties
• 2.2 Intramolecular Force and Potential Energy
• 2.3 Structure of Ionic Solids
• 2.4 Structure of Metals and Alloys
• 2.7 VSEPR and Bond Hybridization
• 2.6 Resonance and Formal Charge