Chemical bonds are the attractions that hold atoms together in substances, from table salt to water to copper wire. Atoms bond because a bonded arrangement often has lower energy than separate atoms. The type of bond depends mainly on how valence electrons are transferred, shared, or allowed to move.
Understanding bonding helps explain melting points, electrical conductivity, hardness, solubility, and many other material properties.
Ionic bonds form when electrons transfer from one atom to another, creating oppositely charged ions that attract in a crystal lattice. Covalent bonds form when atoms share electron pairs, often making molecules with specific shapes. Metallic bonds form when metal atoms share a sea of mobile electrons throughout a solid.
These different electron arrangements create different properties, such as brittle ionic crystals, low-melting molecular covalent substances, and conductive, malleable metals.
Key Facts
- Ionic bonding: metal + nonmetal, electron transfer, attraction between cations and anions.
- Covalent bonding: nonmetal + nonmetal, electron sharing, often forms molecules.
- Metallic bonding: metal atoms in a lattice with delocalized electrons that move freely.
- Ion charge comes from electron loss or gain: losing electrons makes a positive ion, gaining electrons makes a negative ion.
- Electronegativity difference helps predict bond type: large difference is often ionic, small difference is often covalent.
- Coulomb attraction increases when charges are larger and distance is smaller: F = kq1q2/r^2.
Vocabulary
- Valence electron
- A valence electron is an electron in the outer energy level of an atom that can participate in bonding.
- Ionic bond
- An ionic bond is the electrostatic attraction between oppositely charged ions formed after electron transfer.
- Covalent bond
- A covalent bond is a bond in which atoms share one or more pairs of electrons.
- Metallic bond
- A metallic bond is the attraction between positive metal ions and delocalized electrons that move through the metal.
- Electronegativity
- Electronegativity is a measure of how strongly an atom attracts shared electrons in a chemical bond.
Common Mistakes to Avoid
- Calling all compounds molecules is wrong because ionic compounds are usually crystal lattices, not separate molecules.
- Thinking ionic bonds share electrons is wrong because ionic bonding is based on electron transfer followed by attraction between ions.
- Assuming covalent bonds always share electrons equally is wrong because polar covalent bonds have unequal sharing due to electronegativity differences.
- Saying metals conduct because atoms move through the solid is wrong because electrical conductivity mainly comes from mobile delocalized electrons.
Practice Questions
- 1 Sodium has 1 valence electron and chlorine has 7 valence electrons. What ions form when they bond, and what is the formula of the ionic compound?
- 2 A bond forms between atoms with electronegativities 3.5 and 2.1. Calculate the electronegativity difference and predict whether the bond is more likely nonpolar covalent, polar covalent, or ionic.
- 3 Explain why solid copper conducts electricity well but solid sodium chloride does not, even though both contain charged particles.