Covalent Bonding: Sharing Electrons, Molecules, and Polarity

What Covalent Bonding Means

In a covalent bond, two atoms are held together because they share one or more pairs of electrons. Instead of one atom fully taking electrons from another (as in ionic bonding), both atoms “use” the shared electrons. You can picture the shared electrons as spending time in the space between the nuclei, creating an attraction that pulls both atoms toward that shared region.

Covalent bonding is most common between nonmetals. Many everyday substances are made of covalent molecules: water (H₂O), carbon dioxide (CO₂), methane (CH₄), oxygen gas (O₂), and countless organic compounds.

Bonding language you will use

• Bonding pair: a pair of electrons shared between two atoms.
• Single bond: one shared pair of electrons.
• Double bond: two shared pairs of electrons.
• Triple bond: three shared pairs of electrons.
• Lone pair (nonbonding pair): a pair of electrons located on one atom that is not shared in a bond.
• Molecule: a discrete group of atoms connected by covalent bonds (for many substances).

How Sharing Electrons Creates Molecules

When atoms share electrons, they form a stable arrangement at a lower potential energy than the separated atoms. The shared electrons are attracted to both nuclei, and the nuclei are attracted to the shared electrons. At the same time, the nuclei repel each other, and the electrons repel each other. A covalent bond forms at a particular distance where attractions and repulsions balance. That distance is the bond length, and the strength of the bond is related to the bond energy (how much energy is required to break it).

Single, double, and triple bonds (and what changes)

As you go from single to double to triple bonds, two trends are especially useful for beginners:

• Bond length decreases: triple bonds are generally shorter than double bonds, which are shorter than single bonds.
• Bond energy increases: triple bonds are generally stronger than double bonds, which are stronger than single bonds.

Example comparisons (qualitative):

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• N–N (single) is longer and weaker than N≡N (triple) in nitrogen gas.
• C–C (single) is longer and weaker than C=C (double) in ethene.

Lewis Structures: A Practical Step-by-Step Method

Lewis structures are simple drawings that show which atoms are connected and where the bonding and lone-pair electrons are. They are not perfect pictures of the real electron distribution, but they are extremely useful for predicting formulas, connectivity, and many properties (including polarity).

Step-by-step: drawing a Lewis structure

Use this procedure for many common covalent molecules:

• Step 1: Count total valence electrons. Add the valence electrons from each atom. If the species is an ion, add electrons for a negative charge or subtract for a positive charge.
• Step 2: Choose a central atom. Often the least electronegative atom (except hydrogen) is central. Hydrogen is never central because it forms only one bond.
• Step 3: Connect atoms with single bonds. Each single bond uses 2 electrons.
• Step 4: Distribute remaining electrons as lone pairs. First complete the outer atoms’ typical valence (hydrogen gets 2; many main-group atoms commonly reach 8). Put leftover electrons on the central atom.
• Step 5: If the central atom lacks a typical valence, form multiple bonds. Convert lone pairs on outer atoms into bonding pairs to make double or triple bonds until the central atom is satisfied.
• 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 1: Water, H₂O

Step 1: Valence electrons: O has 6, each H has 1. Total = 6 + 1 + 1 = 8.

Step 2: Oxygen is central.

Step 3: Connect O to two H atoms with single bonds (2 bonds × 2 e⁻ = 4 electrons used).

Step 4: Remaining electrons: 8 − 4 = 4. Place them as two lone pairs on oxygen.

  ..      (two lone pairs on O)  ..
H-O-H

This structure explains why water has two O–H bonds and why oxygen carries lone pairs that strongly affect shape and polarity.

Worked example 2: Carbon dioxide, CO₂

Step 1: Valence electrons: C has 4, each O has 6. Total = 4 + 6 + 6 = 16.

Step 2: Carbon is central.

Step 3: Connect O–C–O with single bonds (2 bonds = 4 electrons).

Step 4: Remaining: 16 − 4 = 12. Complete octets on the oxygens first (each O gets 6 more electrons as lone pairs, total 12). Now carbon has only 4 electrons around it (two single bonds), which is short of a typical valence.

Step 5: Convert one lone pair from each oxygen into a bonding pair to form two double bonds.

O=C=O

CO₂ ends up with two C=O double bonds.

Worked example 3: Ammonia, NH₃

Step 1: Valence electrons: N has 5, each H has 1. Total = 5 + 3 = 8.

Step 2: N is central.

Step 3: Make three N–H single bonds (6 electrons).

Step 4: Remaining: 8 − 6 = 2 electrons, placed as one lone pair on N.

  ..
H-N-H
  |
  H

The lone pair is crucial for the molecule’s shape and polarity.

Resonance: When One Lewis Structure Is Not Enough

Sometimes you can draw more than one valid Lewis structure by placing double bonds in different positions while keeping the same arrangement of atoms. These are resonance structures. The real molecule is best thought of as a blend (a resonance hybrid) rather than switching back and forth between drawings.

Example: Ozone, O₃

Ozone can be drawn with a double bond on the left or on the right:

O=O-O   and   O-O=O

In reality, both O–O bonds are the same length, intermediate between a single and a double bond. Resonance is a way to represent electron delocalization (electrons spread over more than two atoms).

Exceptions and Useful Rules of Thumb

The octet guideline and common exceptions

Lewis structures often aim for 8 electrons around many main-group atoms, but there are important exceptions you should recognize:

• Hydrogen is stable with 2 electrons (one bond).
• Boron often forms compounds with only 6 electrons around B (for example, BF₃).
• Expanded valence: some atoms in period 3 and beyond can have more than 8 electrons around them in a Lewis structure (for example, PCl₅, SF₆). In beginner work, you mainly need to recognize that these exist and that the central atom can have more than four bonds.

Coordinate (dative) covalent bonds

In a typical covalent bond, each atom contributes one electron to the shared pair. In a coordinate covalent bond, both electrons in the shared pair come from the same atom (a lone pair is donated). After the bond forms, it behaves like a normal covalent bond.

A common pattern: a molecule with a lone pair (like NH₃) can donate that pair to an electron-poor species (like H⁺) to form a bond. You may see this represented with an arrow in Lewis structures to show the direction of donation.

From Lewis Structures to 3D Shapes (VSEPR Basics)

Lewis structures show connectivity, but molecules are three-dimensional. A practical way to predict shape is to count regions of electron density around the central atom. Each single bond, double bond, triple bond, or lone pair counts as one region. These regions arrange themselves to minimize repulsion.

Step-by-step: predicting a basic molecular shape

• Step 1: Draw the Lewis structure.
• Step 2: Identify the central atom (for simple molecules).
• Step 3: Count electron regions around the central atom (bonding regions + lone pairs).
• Step 4: Use the count to predict the electron-region geometry (arrangement of regions).
• Step 5: Determine the molecular shape by ignoring lone pairs (they affect angles but are not “seen” as atoms).

Common shapes you will meet often

• 2 regions: linear (example: CO₂ is linear around carbon).
• 3 regions: trigonal planar (example: many molecules like BF₃).
• 4 regions: tetrahedral electron geometry (example: CH₄ is tetrahedral; NH₃ becomes trigonal pyramidal due to one lone pair; H₂O becomes bent due to two lone pairs).

Lone pairs repel more strongly than bonding pairs, so they compress bond angles. That is why H₂O has a smaller H–O–H angle than the ideal tetrahedral angle.

Electronegativity and Bond Polarity

Even though covalent bonding involves sharing, the sharing is not always equal. If one atom attracts the shared electrons more strongly, the electron density shifts toward that atom. This creates a polar covalent bond, meaning one end of the bond is slightly negative and the other is slightly positive.

The tendency of an atom to attract bonding electrons is called electronegativity. When two atoms have similar electronegativity, the bond is nonpolar covalent (roughly equal sharing). When the difference is larger, the bond becomes more polar.

How to label partial charges and dipoles

In a polar bond, the more electronegative atom is labeled δ− (partial negative), and the less electronegative atom is labeled δ+ (partial positive). You may also see a dipole arrow pointing toward the δ− end, with a small “plus” at the δ+ end.

Examples:

• In H–Cl, chlorine is more electronegative, so Cl is δ− and H is δ+.
• In O–H, oxygen is δ− and hydrogen is δ+.
• In C–H, the bond is often treated as nearly nonpolar in many basic contexts (small electronegativity difference), though it can have slight polarity depending on the situation.

Molecular Polarity: When Bond Dipoles Add Up (or Cancel)

A molecule can have polar bonds but still be nonpolar overall if the bond dipoles cancel due to symmetry. Molecular polarity depends on two things together:

• Bond polarity (are there polar bonds?)
• Molecular shape (do the dipoles cancel or reinforce?)

Step-by-step: deciding whether a molecule is polar

• Step 1: Draw (or imagine) the Lewis structure and determine the 3D shape around the central atom.
• Step 2: Identify which bonds are polar by comparing electronegativities (qualitatively is usually enough).
• Step 3: Consider the directions of the bond dipoles in the 3D shape.
• Step 4: Decide whether the dipoles cancel (nonpolar molecule) or produce a net dipole (polar molecule).

Example: CO₂ is nonpolar overall

Each C=O bond is polar (oxygen pulls electron density). But CO₂ is linear: O=C=O. The two bond dipoles are equal and opposite, so they cancel. Result: CO₂ has no net dipole and is nonpolar overall.

Example: H₂O is polar

Each O–H bond is polar, and the molecule is bent due to lone pairs on oxygen. The two bond dipoles do not point in opposite directions, so they do not cancel. Result: water has a net dipole and is polar.

Example: CH₄ is nonpolar overall

CH₄ is tetrahedral and symmetric. Even if you consider each C–H bond to have slight polarity, the symmetry causes the dipoles to cancel. Result: methane is nonpolar overall.

Example: NH₃ is polar

NH₃ has polar N–H bonds and a trigonal pyramidal shape due to a lone pair on nitrogen. The dipoles do not cancel. Result: ammonia is polar.

Polarity and Real-World Properties

Polarity strongly influences how substances behave, especially in mixing, boiling points, and interactions between molecules.

“Like dissolves like” (a practical mixing rule)

A very useful guideline is that substances with similar polarity tend to mix well:

• Polar substances tend to dissolve in polar solvents (for example, many alcohols dissolve in water).
• Nonpolar substances tend to dissolve in nonpolar solvents (for example, oils mix with other nonpolar liquids).

This is not a magic rule with no exceptions, but it is a strong first prediction tool. When a polar and a nonpolar substance are mixed, they often separate into layers because the molecules prefer interacting with similar molecules.

Why polarity affects boiling and melting points

To boil a liquid, you must separate molecules from each other. Polar molecules often have stronger intermolecular attractions than nonpolar molecules of similar size, so they often have higher boiling points. For example, water has a much higher boiling point than you might expect from its small molar mass because its polarity leads to strong attractions between molecules.

Practice Toolkit: Build, Check, and Predict

A repeatable workflow for beginners

When you are given a covalent compound and asked about structure or polarity, use this workflow:

• 1) Write the formula and identify the likely central atom.
• 2) Count valence electrons and draw a Lewis structure.
• 3) Add multiple bonds if needed to satisfy typical valence patterns.
• 4) Count electron regions and predict the 3D shape.
• 5) Mark polar bonds and decide whether dipoles cancel.
• 6) Use polarity to predict behavior (mixing tendency, relative boiling point compared to similar-sized molecules).

Guided practice: sulfur dioxide, SO₂ (structure and polarity)

Lewis structure (high-level): S is central, with two O atoms. Total valence electrons: S (6) + 2×O (12) = 18. After connecting and distributing electrons, you typically end with a structure that includes resonance (one S=O double bond and one S–O single bond in each resonance form) and a lone pair on sulfur. Many introductory treatments show two S=O bonds with a lone pair on S as a simplified resonance hybrid idea.

Shape: Around sulfur there are 3 electron regions (two bonding regions to O, one lone pair) giving a bent molecular shape.

Polarity: Each S–O bond is polar, and because the molecule is bent, the dipoles do not cancel. SO₂ is polar.

Guided practice: carbon tetrachloride, CCl₄ (why it is nonpolar)

Lewis structure: Carbon central with four single bonds to chlorine; each Cl has three lone pairs.

Shape: 4 electron regions around carbon → tetrahedral and symmetric.

Polarity: Each C–Cl bond is polar, but the tetrahedral symmetry cancels the dipoles. CCl₄ is nonpolar overall, which helps explain why it behaves more like a nonpolar solvent.

Common Mistakes and How to Avoid Them

Mistake 1: Assuming “covalent = nonpolar”

Covalent bonds can be nonpolar or polar. Many important molecules are covalent and strongly polar (water is the classic example). Always check electronegativity differences and shape.

Mistake 2: Forgetting lone pairs when predicting shape

Lone pairs count as electron regions and can change the molecular shape dramatically. For instance, CH₄ and NH₃ both have four electron regions around the central atom, but NH₃ is not tetrahedral in molecular shape because one region is a lone pair.

Mistake 3: Thinking resonance means the molecule flips between drawings

Resonance structures are bookkeeping tools. The real electron distribution is spread out (delocalized) and the real structure is a hybrid.

Mistake 4: Ignoring symmetry when deciding molecular polarity

Symmetry is the main reason polar bonds can cancel. If a molecule is highly symmetric (linear with identical ends, trigonal planar with identical outer atoms, tetrahedral with identical outer atoms), it is often nonpolar overall even if its bonds are polar.

Mini Reference: Quick Polarity Predictions for Common Shapes

• Linear: nonpolar if the two ends are identical (CO₂), polar if ends differ (HCN).
• Trigonal planar: nonpolar if all three outer atoms are identical, polar if not.
• Tetrahedral: nonpolar if all four outer atoms are identical (CH₄, CCl₄), polar if not (CH₃Cl).
• Bent: usually polar because dipoles rarely cancel (H₂O, SO₂).
• Trigonal pyramidal: usually polar (NH₃).