A substance is a pure form of matter, existing as a single element or a compound, possessing a constant chemical composition. The properties of a substance are its characteristic, measurable traits, such as color, density, melting point, and electrical conductivity. These observable properties are dictated by the substance’s internal structure and the specific forces holding its constituent particles together.
The Foundation: Atomic Structure
Every element is defined by the number of protons in its nucleus. The element’s chemical behavior is governed by the arrangement of its electrons, particularly valence electrons in the outermost shell.
The number of valence electrons dictates an atom’s tendency to gain, lose, or share electrons to achieve a stable, full outer shell configuration. Atoms with one or two valence electrons tend to lose them, forming positive ions and acting as highly reactive metals. Conversely, atoms that are close to having a complete shell are more likely to gain electrons, forming negative ions.
Primary Determinants: Chemical Bonding
The formation of strong chemical bonds becomes the primary determinant of a substance’s fundamental nature. These bonds are the intramolecular forces that link atoms together within a compound or structure.
Ionic Bonding
Ionic bonding involves the complete transfer of one or more valence electrons from a metal atom to a non-metal atom. This transfer creates positively charged ions (cations) and negatively charged ions (anions), held together by electrostatic attraction. This strong attraction results in the formation of a rigid crystal lattice structure.
Covalent Bonding
Covalent bonding occurs when non-metal atoms share pairs of valence electrons to achieve a stable electronic configuration. This sharing can result in the formation of small, discrete molecules, such as water or carbon dioxide. Alternatively, the sharing can extend throughout the entire sample, creating a giant covalent network structure, as seen in materials like diamond or silicon dioxide.
Metallic Bonding
Metallic bonding is unique to metals, where valence electrons are delocalized. These electrons form a mobile “sea” of charge that is shared among a lattice of positively charged metal ions. This arrangement is responsible for the unique bulk properties of metals, which are distinct from those formed by ionic or covalent bonds.
Secondary Determinants: Molecular Arrangement and Intermolecular Forces
For substances composed of discrete covalent molecules, the properties are further refined by their specific three-dimensional arrangement and the weak forces between them. Molecular geometry dictates whether the electron sharing is equal or unequal, thereby determining the molecule’s polarity. A molecule with an uneven distribution of charge is considered polar, having distinct positive and negative ends.
These temporary or permanent charges lead to the formation of intermolecular forces (IMFs), which are the relatively weak attractions between separate molecules. The weakest are London Dispersion Forces, which arise from temporary imbalances in electron distribution and are present in all substances. Stronger IMFs include Dipole-Dipole interactions between polar molecules and Hydrogen bonds, which form when hydrogen is bonded to a highly electronegative atom like oxygen or nitrogen.
The total strength of these IMFs determines many physical properties, as energy is required to overcome these attractions for a substance to change state. Substances with strong hydrogen bonding, like water, have relatively higher boiling points than similar-sized molecules held together only by weaker dispersion forces. These weak forces also influence a liquid’s internal resistance to flow, with stronger IMFs leading to higher viscosity.
Linking Structure to Macroscopic Properties
The type of bonding and the nature of the molecular forces directly translate to the macroscopic properties we observe. Electrical conductivity requires the presence of mobile charged particles. Metals are excellent conductors because their “sea” of delocalized electrons is free to move when a voltage is applied.
Ionic solids are poor conductors in their solid state because their ions are locked into the crystal lattice. However, when an ionic compound is melted or dissolved in water, the ions become mobile, allowing the substance to conduct electricity. Conversely, most molecular covalent compounds are non-conductors because their electrons are tightly held in specific bonds and there are no free ions.
Melting and boiling points are a direct measure of the energy required to break the forces holding the particles together. Ionic compounds and giant covalent networks, like diamond, have extremely high melting points because the entire structure is held by strong chemical bonds that must be broken. In contrast, molecular substances have low melting points because only the relatively weak intermolecular forces between the molecules need to be overcome to transition into the liquid or gas phase.
Solubility is also determined by the compatibility of the forces between particles, often summarized by the rule “like dissolves like”. Polar solvents, such as water, can break the strong attractions in other polar substances through favorable dipole-dipole or hydrogen bonding interactions. Non-polar substances, like oil, dissolve in non-polar solvents because the weak London Dispersion Forces between the solvent and solute molecules are similar in strength.