Source: Purdue University, CHEM 101 lecture notes
Tags: chemical bonding, molecular structures, ionic bonds, covalent bonds, metallic bonds, electronegativity, polarity, Lewis structures, Lewis dot diagrams, valence electrons, octet rule, bond character, molecular geometry, VSEPR
Difficulty: Introductory
Prerequisites: Basic atomic structure (protons, neutrons, electrons), periodic table familiarity, understanding of valence electrons.
Chemical bonding is the reason matter holds together. Every molecule you encounter, from water to DNA, exists because atoms form bonds by transferring or sharing electrons. This topic sits at the heart of general chemistry: once you understand why and how atoms bond, you can predict molecular shapes, explain physical properties, and make sense of chemical reactions. You should already be comfortable locating elements on the periodic table and identifying how many valence electrons an atom has before working through this material.
Atoms bond in three main ways: by sharing electrons (covalent), by transferring electrons (ionic), or through a "sea" of delocalised electrons (metallic). The type of bond depends on the electronegativity difference between atoms. Lewis structures are the standard tool for mapping out how electrons are distributed in a molecule, and getting them right is the gateway to predicting molecular shape and polarity.
Chemical bond
A lasting attraction between atoms that allows the formation of molecules and compounds. In simple terms, it is the glue holding atoms together.
Metallic bond
A bond between metal atoms in which a lattice of positive ions is surrounded by a "sea" of delocalised electrons. Think of it as metal atoms pooling their outer electrons into a shared cloud that all the ions sit in.
Ionic bond
A bond formed when one atom transfers one or more electrons to another, creating oppositely charged ions that attract each other. In simple terms, one atom gives and the other takes, and the resulting plus/minus attraction holds them together.
Covalent bond
A bond formed when two nonmetal atoms share one or more pairs of electrons. Think of it as two atoms each contributing electrons to a shared pool between them.
Nonpolar covalent bond
A covalent bond in which electrons are shared equally between atoms, typically because the atoms have identical or very similar electronegativities. The classic example is H₂.
Polar covalent bond
A covalent bond in which electrons are shared unequally because one atom is more electronegative than the other. In simple terms, one atom hogs the electrons a bit more, creating a slight charge imbalance. HCl is a standard example.
Electronegativity
A measure of an atom's ability to attract bonding electrons towards itself within a covalent bond. Think of it as how greedy an atom is for shared electrons. Fluorine is the most electronegative element.
Polar molecule
A molecule with an uneven distribution of electron density, resulting in a net dipole moment. This happens when polar bonds are arranged asymmetrically or when lone pairs create an imbalance.
Nonpolar molecule
A molecule with no net dipole moment, either because it contains only nonpolar bonds or because its polar bonds are arranged symmetrically so that dipoles cancel out.
Lewis structure (Lewis dot diagram)
A diagram showing the bonding between atoms in a molecule and any lone pairs of electrons that exist. Think of it as a map of where all the valence electrons live in a molecule.
Valence electrons
The electrons in the outermost shell of an atom, which participate in bonding. The number of valence electrons for main-group elements matches the group number on the periodic table.
Octet rule
The principle that atoms tend to bond in ways that give each atom eight electrons in its valence shell (hydrogen is the exception, needing only two). In simple terms, atoms want a full outer shell, and bonding is how they get there.
Formal charge
The charge assigned to an atom in a Lewis structure, calculated as: valence electrons minus (lone pair electrons + half of bonding electrons). It is a bookkeeping tool for choosing the most stable Lewis structure.
Metallic bonds form between metal atoms. The outer electrons detach from individual atoms and move freely through a lattice of positive ions.
This "electron sea" model explains why metals conduct electricity and heat, and why they are malleable.
Example: steel, an alloy of iron and carbon.
Ionic bonds form between metals and nonmetals through electron transfer.
The metal loses electrons to become a cation; the nonmetal gains electrons to become an anion.
The electrostatic attraction between the oppositely charged ions holds the compound together.
Example: sodium chloride (NaCl). Sodium gives one electron to chlorine.
Covalent bonds form between nonmetal atoms through electron sharing.
Nonpolar covalent: electrons shared equally. Occurs between identical atoms (H₂, O₂, N₂) or atoms with the same electronegativity.
Polar covalent: electrons shared unequally. The more electronegative atom pulls the shared pair closer. Example: HCl, where chlorine pulls harder than hydrogen.
Electronegativity increases across a period (left to right) and decreases down a group on the periodic table.
The electronegativity difference (ΔEN) between two bonded atoms determines bond type:
ΔEN = 0 (or very small): nonpolar covalent
ΔEN moderate (roughly 0.4 to 1.7): polar covalent
ΔEN large (roughly above 1.7): ionic
These boundaries are guidelines, not hard cutoffs. Some textbooks use slightly different thresholds.
A molecule's polarity depends on two things: whether it contains polar bonds, and how those bonds (and any lone pairs) are arranged in three-dimensional space.
Polar molecules have polar bonds arranged asymmetrically, or lone pairs that create an uneven charge distribution. Water (H₂O) is the standard example: its bent shape means the two O–H dipoles do not cancel.
Nonpolar molecules either have no polar bonds at all, or have polar bonds arranged symmetrically so that the individual dipoles cancel out. CO₂ is linear and symmetric, so despite having polar C=O bonds, the molecule overall is nonpolar.
Step 1: Identify the central atom. The central atom is usually the least electronegative element (hydrogen and fluorine are never central).
Step 2: Count the total number of valence electrons. Add up valence electrons for every atom in the molecule. For ions, add electrons for negative charges or subtract for positive charges.
Step 3: Draw single bonds from the central atom to each surrounding atom. Each single bond uses two electrons.
Step 4: Complete the octets on the outer atoms first (remember, hydrogen only needs two electrons), then place remaining electrons on the central atom.
Step 5: If the central atom lacks an octet, convert lone pairs on outer atoms into double or triple bonds.
Step 6: Check formal charges. The best Lewis structure minimises formal charges and places any negative formal charge on the more electronegative atom.
Central atom: sulfur (less electronegative than bromine)
Valence electrons: S has 6, each Br has 7 → total = 6 + 7 + 7 = 20
Draw two S–Br single bonds (uses 4 electrons, leaving 16)
Complete octets on each Br (6 electrons each as lone pairs, uses 12, leaving 4)
Place remaining 4 electrons on S as two lone pairs
S now has 8 electrons (2 bonding + 6 lone pair), satisfying the octet. All formal charges are zero.
Formal charge formula:
Formal charge = (valence electrons) – (lone pair electrons) – (½ × bonding electrons)
Electronegativity difference thresholds (approximate):
0 to ~0.4 → nonpolar covalent
~0.4 to ~1.7 → polar covalent
Above ~1.7 → ionic
These thresholds vary slightly by textbook. Your professor's version is the one to use on the exam.
The type of bonding in a material dictates its properties. Ionic compounds like table salt dissolve in water and conduct electricity when melted, which is why salt solutions carry current in batteries and biological systems. Metallic bonding explains why copper wire conducts electricity in your home. Understanding molecular polarity is essential in pharmacology: drug molecules need the right polarity profile to cross cell membranes and bind to their targets.
Students often assume that any molecule containing polar bonds must itself be polar. That is not the case. If the polar bonds are arranged symmetrically (as in CO₂ or CCl₄), the dipoles cancel and the molecule is nonpolar overall.
Students sometimes think ionic bonds involve sharing electrons. They do not. Ionic bonds involve a transfer of electrons, creating ions. The bond is the electrostatic attraction between those ions.
A common mistake when drawing Lewis structures is forgetting to check formal charges after completing octets. A structure can satisfy the octet rule but still not be the best representation if formal charges are unnecessarily large.
Students frequently place the most electronegative atom at the centre of a Lewis structure. The least electronegative atom (excluding hydrogen) should be the central atom.
⚠️ Electronegativity differences and bond classification (nonpolar, polar, ionic) are heavily tested. Be ready to look at two atoms and categorise the bond.
⚠️ Lewis structure drawing is a near-guaranteed exam question. Practise the five-step method until it is automatic.
⚠️ Molecular polarity versus bond polarity is a classic exam trap. Know the difference and be able to explain why CO₂ is nonpolar despite having polar bonds.
⚠️ Formal charge calculation comes up in questions asking you to choose the "best" Lewis structure from several options.
True or false: A molecule with polar bonds is always a polar molecule.
Fill in the blank: Ionic bonds form when the electronegativity difference between two atoms is __________.
True or false: Hydrogen can be the central atom in a Lewis structure.
Fill in the blank: In metallic bonding, metal atoms release their outer electrons into a delocalised __________.
True or false: The formal charge on an atom in the best Lewis structure should be as close to zero as possible.
Answers: 1. False (symmetric arrangements cancel dipoles). 2. Large (roughly above 1.7). 3. False (hydrogen bonds to only one other atom). 4. Electron sea. 5. True.
Q: What type of bond forms between sodium (Na) and chlorine (Cl), and why?
A: An ionic bond. Sodium has a low electronegativity and chlorine has a high electronegativity, giving a large ΔEN. Sodium transfers its one valence electron to chlorine, forming Na⁺ and Cl⁻ ions that attract each other.
Q: Draw the Lewis structure for H₂O and determine whether the molecule is polar or nonpolar.
A: Oxygen is the central atom with 6 valence electrons; each hydrogen has 1, giving 8 total. Draw two O–H single bonds (4 electrons used), then place the remaining 4 electrons as two lone pairs on oxygen. The bent molecular geometry (due to lone pairs) means the dipoles do not cancel, so H₂O is a polar molecule.
Q: Explain why CO₂ is nonpolar even though it contains polar bonds.
A: CO₂ is linear and symmetric. The two C=O bond dipoles point in exactly opposite directions and cancel each other out, giving the molecule no net dipole moment.
Q: Calculate the formal charge on sulphur in SBr₂ if S has two bonding pairs and two lone pairs.
A: Formal charge = 6 (valence) – 4 (lone pair electrons) – ½(4 bonding electrons) = 6 – 4 – 2 = 0. Sulphur has a formal charge of zero, which is ideal.
Q: A student draws a Lewis structure with the most electronegative atom in the centre. What is wrong with this approach?
A: The least electronegative atom (excluding hydrogen) should be the central atom, because the central atom shares the most electrons and a less electronegative atom is better able to do that. Placing the most electronegative atom at the centre typically produces unfavourable formal charges.
This material connects directly to VSEPR theory (Valence Shell Electron Pair Repulsion), which uses Lewis structures as the starting point for predicting three-dimensional molecular shapes. It also underpins intermolecular forces: whether a molecule is polar or nonpolar determines whether it experiences dipole-dipole interactions or only London dispersion forces, which in turn explains boiling points, solubility, and many physical properties covered later in the course.
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