What a Chemical Bond Really Is
A chemical bond is an attraction that holds atoms together in a stable arrangement. When atoms bond, they form substances with new properties: gases can become solids, reactive elements can become safe compounds, and materials can gain strength, flexibility, or electrical conductivity.
Bonding happens because atoms interact through electric forces: positive regions attract negative regions, and like charges repel. In practice, bonding is about how atoms arrange their electrons (especially the outer, “valence” electrons) and how the nuclei of neighboring atoms attract those electrons.
Think of bonding as a balance between attractions and repulsions. Two atoms approach: each nucleus attracts electrons, but the nuclei repel each other and the electrons repel each other. A bond forms when the overall attractions outweigh the repulsions at a particular distance, creating a stable “sweet spot” called the bond length.
Bond Formation as an Energy Story
Bonding is strongly connected to energy. When a bond forms, energy is usually released (the system becomes more stable). When a bond breaks, energy must be absorbed. This is why many reactions that form strong bonds give off heat, and why breaking bonds often requires heating, light, or electricity.
- Bond formation: typically releases energy (exothermic step).
- Bond breaking: requires energy input (endothermic step).
Even if you do not calculate energies yet, you can use this idea to reason about stability: structures with stronger, well-matched attractions tend to be lower in energy and more stable.
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Three Big Bonding Models You Will Use
For beginners, chemical bonding is usually organized into three main models: ionic bonding, covalent bonding, and metallic bonding. These are models—useful ways to describe what electrons are doing. Real substances can show mixed behavior, but these categories help you predict properties.
Ionic Bonding: Attraction Between Ions
Ionic bonding describes the attraction between positively charged ions and negatively charged ions. Instead of forming a single “pair bond” between two atoms, ionic substances typically form large, repeating crystal structures (lattices) where each ion is surrounded by many oppositely charged ions.
In the ionic model, electrons are transferred (or, more realistically, shifted so strongly) that one species becomes a cation and the other becomes an anion. The bond is the electrostatic attraction between these ions.
Common property patterns of ionic substances (useful for identification):
- Often solid at room temperature.
- Often have high melting points.
- Often dissolve in water (many do, not all).
- Conduct electricity when melted or dissolved (because ions can move), but not as solids (ions locked in place).
Covalent Bonding: Sharing Electrons
Covalent bonding describes atoms sharing electrons. The shared electrons are attracted to both nuclei, which can hold the atoms together as a molecule (a discrete unit) or as a network (a giant covalent structure).
Covalent bonds are directional: the arrangement in space matters, which is why molecules have shapes. Covalent substances vary widely: some are gases (like many small molecules), some are liquids, and some are solids (including very hard network solids).
Common property patterns of molecular covalent substances:
- Often lower melting and boiling points than ionic solids (especially for small molecules).
- Often poor electrical conductors (no free ions or electrons).
- Solubility depends on polarity (some dissolve in water, some do not).
Metallic Bonding: A Lattice with Mobile Electrons
Metallic bonding describes metal atoms in a lattice where valence electrons are not tied to one specific atom. A helpful picture is “positive metal ions in a sea of delocalized electrons.” The attraction between the positive metal centers and the mobile electrons holds the metal together.
Common property patterns of metals:
- Conduct electricity and heat well (mobile electrons carry charge and energy).
- Malleable and ductile (layers can slide without shattering because bonding is not strongly directional).
- Often shiny (electrons interact with light).
How to Decide Which Bonding Model Fits
When you meet a new substance, you can often choose a bonding model by looking at what kinds of elements are present and what the structure seems to be.
Step-by-Step: Classifying Bond Type from a Formula
Step 1: Identify the elements. Are they metals or nonmetals? (You do not need a full periodic table lesson here—just the basic idea that metals are typically on the left/center and nonmetals on the right.)
Step 2: Apply the quick rules.
- Metal + nonmetal → usually ionic (ions in a lattice).
- Nonmetal + nonmetal → usually covalent (molecules or networks).
- Metal + metal → metallic (metal lattice with mobile electrons).
Step 3: Watch for exceptions and special cases. Some compounds have mixed bonding (for example, a compound can be ionic overall but contain covalent bonds inside a polyatomic ion). Also, some nonmetal combinations form networks rather than discrete molecules.
Practice Examples
- NaCl: sodium (metal) + chlorine (nonmetal) → ionic model fits well.
- H2O: hydrogen and oxygen are nonmetals → covalent bonds inside water molecules.
- Cu: copper is a metal element → metallic bonding.
- CO2: carbon and oxygen are nonmetals → covalent bonds; discrete molecules.
Covalent Bonds in More Detail: Single, Double, Triple
Covalent bonds can involve one shared pair of electrons (single bond), two shared pairs (double bond), or three shared pairs (triple bond). More shared pairs generally mean a shorter and stronger bond, though the exact values depend on the atoms involved.
- Single bond: longest and weakest of the three (in general).
- Double bond: shorter and stronger than a single bond.
- Triple bond: shortest and strongest (in general).
These differences matter in reactions. For example, breaking a triple bond typically requires more energy than breaking a single bond, and molecules with multiple bonds often react in characteristic ways because the electron distribution is different.
Step-by-Step: Drawing a Simple Lewis Structure (Practical Skill)
Lewis structures are a beginner-friendly way to show which atoms are connected and where valence electrons are placed. They are not perfect, but they help you reason about bonding and electron pairs.
Step 1: Count total valence electrons. Add the valence electrons from each atom. If the species has a charge, add electrons for a negative charge or subtract for a positive charge.
Step 2: Choose a central atom. Often the least electronegative atom (not hydrogen) goes in the center. Hydrogen is never central.
Step 3: Connect atoms with single bonds. Each single bond uses 2 electrons.
Step 4: Distribute remaining electrons to complete outer shells. Place lone pairs on outer atoms first, then on the central atom.
Step 5: If the central atom lacks electrons, form multiple bonds. Convert lone pairs from outer atoms into bonding pairs to make double or triple bonds.
Step 6: Check formal charges (optional but powerful). Prefer structures with minimal formal charges and with negative formal charge on more electronegative atoms.
Worked Example: CO2
Step 1: Valence electrons: C has 4, each O has 6 → total = 4 + 6 + 6 = 16.
Step 2: Carbon is central: O–C–O.
Step 3: Two single bonds use 4 electrons, leaving 12.
Step 4: Complete octets on oxygens: each O gets 6 more electrons (3 lone pairs) → uses 12 electrons, leaving 0.
Step 5: Carbon now has only 4 electrons around it (two single bonds). To give carbon an octet, convert one lone pair from each oxygen into a bonding pair, making two double bonds.
Final picture: O=C=O with two lone pairs on each oxygen.
Polarity: When Sharing Is Unequal
Not all covalent bonds share electrons equally. If one atom attracts the shared electrons more strongly than the other, the bond becomes polar: one end is slightly negative (δ−) and the other slightly positive (δ+). This is not a full transfer like the ionic model; it is an imbalance in electron density.
Polarity matters because it affects solubility, boiling point, and how molecules interact. Polar molecules tend to mix with polar substances (like water), while nonpolar molecules tend to mix with nonpolar substances (like many oils).
Bond Polarity vs. Molecular Polarity
A molecule can contain polar bonds but still be nonpolar overall if the bond dipoles cancel due to symmetry. For example, CO2 has polar C=O bonds, but the molecule is linear and symmetric, so the overall molecule is nonpolar.
Water (H2O) has polar O–H bonds and a bent shape, so the dipoles do not cancel; the molecule is polar.
Step-by-Step: Predicting Whether a Molecule Is Polar
Step 1: Decide whether bonds are polar. Use electronegativity difference as a guide: larger difference → more polar bond.
Step 2: Consider the shape (symmetry). If the molecule is symmetric, dipoles may cancel.
Step 3: Conclude molecular polarity. If dipoles cancel → nonpolar; if they do not → polar.
Even without exact shapes, you can often use simple symmetry cues: identical outer atoms arranged evenly around a center often cancel.
Intermolecular Forces: Why Molecules Attract Each Other
Chemical bonds (ionic, covalent, metallic) hold atoms together within a substance’s basic units (ions in a lattice, atoms in a molecule, atoms in a metal). But many everyday properties—like boiling point, viscosity, and why some substances are liquids—depend on attractions between molecules. These are called intermolecular forces (IMFs). They are weaker than typical covalent or ionic bonds, but they are crucial.
Three Common Intermolecular Forces
- London dispersion forces: present in all molecules and atoms; caused by temporary fluctuations in electron density. Stronger in larger, more easily polarized electron clouds.
- Dipole–dipole forces: occur between polar molecules; positive end of one attracts negative end of another.
- Hydrogen bonding: a strong type of dipole–dipole attraction when hydrogen is bonded to N, O, or F and is attracted to a lone pair on N, O, or F in a neighboring molecule.
These forces help explain why some small molecules are gases while others are liquids, and why water has an unusually high boiling point for its size.
Step-by-Step: Using Intermolecular Forces to Compare Boiling Points
Step 1: Ask if hydrogen bonding is possible (H bonded to N/O/F). If yes, expect higher boiling point than similar-sized molecules without it.
Step 2: If no hydrogen bonding, ask if the molecule is polar. If yes, dipole–dipole forces raise boiling point compared with nonpolar molecules of similar size.
Step 3: Consider size and shape for dispersion forces. Larger molar mass and more surface area generally increase dispersion forces and raise boiling point.
Example comparison: CH4 (nonpolar, small) has very low boiling point; NH3 (hydrogen bonding) boils much higher; C3H8 (nonpolar but larger than CH4) boils higher than CH4 due to stronger dispersion forces.
Network Solids vs. Molecular Substances
Covalent bonding can produce two very different kinds of solids:
- Molecular solids: made of discrete molecules held together by intermolecular forces. They tend to have lower melting points and can be soft (like many organic solids).
- Network covalent solids: atoms are connected in a continuous covalent network (no separate molecules). They tend to be very hard and have very high melting points.
This distinction helps explain why some covalent substances are gases or liquids while others are extremely hard solids.
Why Bond Type Predicts Properties (A Practical Map)
Bonding models are powerful because they connect microscopic structure to macroscopic properties you can observe and measure.
Ionic Substances: Property Reasons
- High melting points: strong electrostatic attractions in a lattice require lots of energy to separate ions.
- Brittleness: when layers shift, like charges can line up and repel, causing the crystal to crack.
- Conductivity when molten or dissolved: ions are free to move and carry charge.
Molecular Covalent Substances: Property Reasons
- Lower melting/boiling points: molecules are held together by weaker intermolecular forces, not by a full ionic lattice.
- Poor conductivity: no mobile ions or delocalized electrons.
- Solubility depends on polarity: polar molecules interact well with polar solvents; nonpolar with nonpolar.
Metals: Property Reasons
- Conductivity: delocalized electrons move through the lattice.
- Malleability: non-directional bonding allows layers to slide while staying bonded.
- Alloys: mixing metals can change strength and hardness by disrupting layer sliding.
Bonding and Reactions: What Changes and What Stays
In chemical reactions, atoms are rearranged: some bonds break and new bonds form. The identity of atoms stays the same, but their connections change. Thinking in terms of bonds helps you interpret reactions as “before-and-after” changes in connectivity.
Step-by-Step: Describing a Reaction in Bond Terms
Step 1: Identify which bonds exist in the reactants.
Step 2: Identify which bonds exist in the products.
Step 3: Determine which bonds must break (present in reactants, absent in products).
Step 4: Determine which bonds must form (absent in reactants, present in products).
Step 5: Use the energy idea: breaking bonds absorbs energy; forming bonds releases energy. Whether the reaction overall releases or absorbs energy depends on the balance.
Example (conceptual): If a reaction replaces several weak bonds with fewer but much stronger bonds, it often releases energy overall.
Common Beginner Misunderstandings (and Fixes)
“Ionic bonds are not real bonds because they are just attraction.”
All bonds are based on attraction. The ionic model emphasizes electrostatic attraction between ions, while the covalent model emphasizes shared electron density. Both are real and measurable in their effects on structure and energy.
“Covalent means equal sharing.”
Covalent means sharing, but sharing can be unequal. Polar covalent bonds are still covalent bonds; they simply have an uneven electron distribution.
“A formula unit like NaCl is a molecule.”
In an ionic crystal, NaCl does not exist as separate NaCl molecules. The formula tells you the simplest ratio of ions in the lattice. Many ions surround each ion in a repeating structure.
“Stronger intermolecular forces mean stronger covalent bonds.”
Intermolecular forces are attractions between molecules; covalent bonds are within molecules. A substance can have strong covalent bonds but weak intermolecular forces (for example, small molecules), or strong intermolecular forces (like hydrogen bonding) without changing the covalent bonds inside each molecule.
Mini Skill Practice: From Bonding to Properties
Problem 1: Predict Conductivity
You have three samples: solid NaCl, liquid water (pure), and solid copper. Which conduct electricity well?
- Solid NaCl: poor conductor as a solid (ions fixed).
- Pure water: poor conductor (very few ions present).
- Solid copper: good conductor (mobile electrons).
Problem 2: Predict Melting Point Trend
Which is likely to have the highest melting point: CO2, NaCl, or SiO2 (silicon dioxide)?
- CO2: molecular substance with weak intermolecular forces → low melting point.
- NaCl: ionic lattice → high melting point.
- SiO2: network covalent solid → very high melting point (often higher than many ionic salts).
Problem 3: Predict Solubility in Water
Which is more likely to dissolve in water: ethanol (C2H5OH) or hexane (C6H14)?
- Ethanol: has an O–H group, can hydrogen bond with water → dissolves well.
- Hexane: nonpolar, only dispersion forces → does not dissolve well in water.
Optional Deepening: Resonance as “Multiple Valid Drawings”
Sometimes one Lewis structure is not enough to represent electron distribution. In resonance, you draw two or more valid structures that differ only in electron placement, not in atom positions. The real molecule is a blend of these drawings, with electron density spread out (delocalized) more than a single drawing suggests.
This idea helps explain why some bonds have lengths and strengths that are “in between” single and double bonds, and why certain ions and molecules are unusually stable.
Example idea (not a full drawing set): a structure may be shown with a double bond in two possible locations. The actual bonding is shared across those locations.
