Chemical Bonding and Intermolecular Forces
Students compare ionic, covalent, and metallic bonding and use particle-level attractions to explain observable properties of substances.
Illustrations are auto-generated and may be placeholders. They can be refreshed to match the narration.
Why Atoms Bond
Atoms bond when interactions among their charged particles produce a system with lower potential energy than the separated atoms. Positively charged nuclei attract negatively charged electrons, while like charges repel one another. At a particular distance, these attractions and repulsions balance, creating a stable bond length. Energy is released when a bond forms, and energy must be added to break that bond. For example, as two hydrogen atoms approach, each nucleus attracts both electrons. Their potential energy decreases until the H–H bond length is reached. If the atoms move too close, nucleus-nucleus and electron-electron repulsions sharply increase the energy. Bond formation does not mean atoms become inactive; instead, their valence electrons reach a more stable arrangement through electron transfer, electron sharing, or collective movement among many atoms.
Ionic and Covalent Bonds
Ionic bonding results from attraction between oppositely charged ions after electrons are transferred. In sodium chloride, sodium loses one valence electron to become Na+, while chlorine gains that electron to become Cl−. The ions form an extended crystal lattice rather than separate NaCl molecules. Covalent bonding occurs when atoms share electron pairs. In a water molecule, each hydrogen shares electrons with oxygen, producing two O–H covalent bonds. Metallic bonding occurs in metals such as copper, where positive metal ions are attracted to mobile, delocalized valence electrons. These models explain contrasting properties. Ionic solids are often brittle and conduct electricity only when melted or dissolved because their ions can then move. Molecular covalent substances often have lower melting points and poor conductivity. Metals conduct electricity as solids and can bend because mobile electrons maintain attraction as metal ions shift.
Lewis Structures and Polarity
A Lewis structure represents valence electrons as dots and shared pairs as lines. To draw one, count the total valence electrons, connect atoms with single bonds, complete outer-atom octets, and place remaining electrons on the central atom. Multiple bonds may be needed if the central atom lacks an octet. Molecular polarity depends on both bond polarity and molecular geometry. Oxygen attracts shared electrons more strongly than hydrogen, so each O–H bond in water is polar. Because water has a bent shape, its bond dipoles do not cancel, making the molecule polar. Carbon dioxide also has polar C=O bonds, but its linear, symmetrical shape makes the two bond dipoles cancel. Therefore, CO2 is nonpolar overall. Correct Lewis structures help predict geometry, charge distribution, intermolecular attractions, and physical behavior.
Intermolecular Forces
Intermolecular forces are attractions between separate particles, not the bonds holding atoms together within a particle. London dispersion forces arise from temporary uneven electron distributions and occur in all atoms and molecules. They generally become stronger as electron clouds grow larger and more polarizable. Dipole-dipole forces occur between permanent partial charges in polar molecules. Hydrogen bonding is an especially strong dipole-dipole attraction when hydrogen is covalently bonded to nitrogen, oxygen, or fluorine and is attracted to a lone pair on a nearby N, O, or F atom. For example, hydrogen bonds connect neighboring water molecules. More energy is required to separate these molecules, helping water have a much higher boiling point than similarly sized methane. Ion-dipole attractions also form when ions dissolve in polar solvents, such as Na+ and Cl− interacting with water.
Structure and Properties
Bulk properties provide evidence about particle structure and the strength of electrical forces. A useful investigation can compare equal samples of sodium chloride, sugar, paraffin wax, and copper. Students can observe hardness and brittleness, test solubility in equal volumes of water, and measure conductivity of each solid and resulting mixture with the same circuit. Melting behavior may be compared using published melting points or carefully controlled heating with appropriate safety equipment. Sodium chloride’s high melting point and solution conductivity support a strong ionic lattice. Sugar dissolves but its solution does not conduct well because neutral molecules remain intact. Wax melts at a relatively low temperature, indicating weaker attractions between nonpolar molecules. Copper conducts as a solid because it contains mobile electrons. Repeated trials, controlled sample amounts, and recorded quantitative measurements make the evidence more reliable and allow particle-level forces to be inferred from observable properties.