Atoms bond by either transferring electrons (ionic bonding) or sharing them (covalent bonding). Whether a bond is ionic or covalent depends on the electronegativity difference between the two atoms: a difference greater than 1.6 generally means ionic, less than 1.6 means covalent. Lewis structures let you map out where the valence electrons sit in a molecule, which sets up everything you need for predicting molecular shape later.
Difficulty: Introductory to Intermediate
Prerequisites: a basic understanding of atomic structure (protons, neutrons, electrons), the periodic table layout, and the concept of valence electrons. If you are not comfortable identifying how many valence electrons an element has from its group number, review that first.
This topic sits at the heart of general chemistry. Everything from predicting physical properties (melting point, conductivity) to understanding chemical reactions depends on knowing how and why atoms bond. Topic 4 bridges atomic theory (how individual atoms behave) and molecular behaviour (how groups of atoms behave together). If you are behind, this is worth catching up on, because VSEPR theory, intermolecular forces, and reaction mechanisms all build directly on the bonding concepts here. You should already know how to read the periodic table and count valence electrons before diving in.
Cation
A positively charged ion formed when a metal atom loses one or more valence electrons. Think of it as a metal that has given away electrons and now has more protons than electrons.
Anion
A negatively charged ion formed when a nonmetal atom gains one or more electrons. In simple terms, a nonmetal picks up extra electrons to fill its valence shell.
Octet rule
The principle that atoms tend to bond in ways that give them eight electrons in their valence shell, achieving the electron configuration of a noble gas. Think of it as atoms wanting a "full outer shell" of eight.
Electronegativity
A measure of how strongly an atom attracts shared electrons in a bond. In simple terms, it tells you how greedy an atom is for electrons. Fluorine has the highest electronegativity (4.0) on the periodic table.
Ionic bond
A bond formed by the electrostatic attraction between a cation and an anion, typically when the electronegativity difference between two atoms is greater than 1.6. Think of it as one atom taking an electron from another, then the opposite charges holding them together.
Covalent bond
A bond formed when two atoms share one or more pairs of electrons, typically when the electronegativity difference is less than 1.6. Think of it as two atoms holding onto the same electrons at the same time.
Lewis structure (Lewis dot diagram)
A diagram showing the arrangement of valence electrons around atoms in a molecule, using dots for lone pairs and lines for bonding pairs. In simple terms, it is a map of where every valence electron lives in the molecule.
Lone pair
A pair of valence electrons on an atom that is not shared in a bond. Think of it as electrons that belong to one atom only, sitting there taking up space.
Bonding pair
A pair of electrons shared between two atoms in a covalent bond. Each single bond is one bonding pair; a double bond is two bonding pairs; a triple bond is three.
Valence electrons
The electrons in the outermost shell of an atom, responsible for bonding behaviour. The number of valence electrons for main-group elements matches the group number on the periodic table.
Metals lose valence electrons to form cations (positive ions).
Nonmetals gain electrons to form anions (negative ions).
The electrostatic attraction between the cation and anion holds them together as an ionic bond.
Ionic bonds typically form when the electronegativity difference between two atoms is greater than 1.6.
Example: Be (1.6) and Cl (3.2) have a difference of 1.6, which sits right at the boundary. In practice, 1.6 is often still treated as covalent; the cutoff is a guideline, not a hard rule.
Ionic compounds form crystal lattice structures, not discrete molecules.
Covalent bonds form when two nonmetals share electron pairs.
The electronegativity difference between the atoms is less than 1.6.
Example: B (2.0) and F (4.0) have a difference of 2.0, which is above 1.6, yet BF₃ is considered covalent. This is a well-known exception worth remembering.
A single bond shares one pair of electrons (one line in a Lewis structure).
A double bond shares two pairs (two lines). A triple bond shares three pairs (three lines).
Electronegativity is the key predictor for bond type.
Difference > 1.6: ionic bond (electron transfer).
Difference < 1.6: covalent bond (electron sharing).
There are exceptions (BF₃ being the classic one in this course), so treat the 1.6 threshold as a guideline.
Step 1: Count the total number of valence electrons for all atoms in the molecule.
Each bond counts as one electron towards the total for each atom it connects.
Hydrogen always has a maximum of 2 electrons (one bond, no lone pairs).
Step 2: Draw single bonds between the central atom and each surrounding atom.
Step 3: Distribute remaining electrons as lone pairs, starting with the outer atoms, to satisfy the octet rule.
Step 4: If the central atom lacks an octet, convert lone pairs on outer atoms into double or triple bonds.
CO₂: C = 4 valence electrons, O = 6 each. Total = 4 + (2 × 6) = 16 electrons. Carbon forms two double bonds with oxygen. Each oxygen also carries two lone pairs.
CF₄: C = 4, F = 7 each. Total = 4 + (4 × 7) = 32 electrons. Carbon forms four single bonds to fluorine. Each fluorine carries three lone pairs.
NH₄⁺: N = 5, H = 1 each, minus 1 for the positive charge. Total = 5 + (4 × 1) − 1 = 8 electrons. Nitrogen forms four single bonds to hydrogen. No lone pairs on nitrogen in this ion.
NCl₃: N = 5, Cl = 7 each. Total = 5 + (3 × 7) = 26 electrons. Nitrogen forms three single bonds to chlorine and retains one lone pair.
Total valence electrons formula:
Sum of valence electrons for each atom, adjusted for charge. Add electrons for negative charges, subtract for positive charges.
Example: NH₄⁺ = 5 (N) + 4(1) (H) − 1 (charge) = 8 electrons
Electronegativity difference rule of thumb:
ΔEN > 1.6 → ionic bond (electron transfer)
ΔEN < 1.6 → covalent bond (electron sharing)
Exceptions exist (e.g. BF₃ with ΔEN = 2.0 is covalent).
Ionic compounds like NaCl (table salt) dissolve in water and conduct electricity in solution, which is why electrolyte drinks work: dissolved ions carry electrical signals your muscles need. Covalent compounds like water (H₂O) and carbon dioxide (CO₂) are everywhere in biological systems. Understanding bond types helps explain why some substances dissolve in water and others do not, and why metals conduct electricity but most molecular compounds do not.
Students often think the 1.6 electronegativity cutoff is absolute. It is a guideline. BF₃ has a difference of 2.0 but is covalent. Always check for known exceptions.
Students frequently confuse cations and anions. Cations are positive (metals lose electrons). Anions are negative (nonmetals gain electrons). A useful mnemonic: cation has a "t" that looks like a plus sign.
Some students believe every atom must follow the octet rule. Hydrogen is satisfied with just 2 electrons, and some elements (like Be in BeCl₂ and B in BF₃) are stable with fewer than 8.
Students sometimes count a double bond as two separate electron domains. A double bond (and a triple bond) counts as one electron domain for geometry purposes.
⚠️ You will almost certainly be asked to draw Lewis structures. Practise counting valence electrons quickly.
⚠️ Know the electronegativity difference rule (>1.6 = ionic, <1.6 = covalent) and the BF₃ exception.
⚠️ Be ready to identify whether a bond is ionic or covalent from a pair of elements and their electronegativity values.
⚠️ Expect questions that ask you to count total valence electrons for a molecule or polyatomic ion (remember to add or subtract for charges).
True or False: An electronegativity difference of 2.0 always means the bond is ionic.
Fill in the blank: Metals form ______ ions by losing electrons.
True or False: A double bond counts as two electron domains in VSEPR theory.
Fill in the blank: The total valence electrons in CO₂ are ______.
True or False: Hydrogen follows the octet rule.
Q: What type of bond forms between sodium (Na, EN = 0.9) and chlorine (Cl, EN = 3.0)?
A: Ionic bond. The electronegativity difference is 2.1, which is greater than 1.6, so sodium transfers its electron to chlorine.
Q: Draw the Lewis structure of water (H₂O) and state the total number of valence electrons.
A: Oxygen has 6 valence electrons, each hydrogen has 1. Total = 6 + 2(1) = 8. Oxygen forms two single bonds to hydrogen and retains two lone pairs.
Q: Why is BF₃ considered covalent despite having an electronegativity difference of 2.0?
A: BF₃ is a well-known exception to the 1.6 rule. Boron and fluorine share electrons rather than transferring them, partly because boron is a metalloid and forms stable covalent bonds with fluorine.
Q: How many valence electrons does the ammonium ion (NH₄⁺) have?
A: 8. Nitrogen contributes 5, four hydrogens contribute 1 each, and you subtract 1 for the positive charge: 5 + 4 − 1 = 8.
Q: Explain why hydrogen does not follow the octet rule.
A: Hydrogen's first (and only) electron shell holds a maximum of 2 electrons. It achieves a stable configuration by filling that shell, not by reaching 8.
This material connects directly to VSEPR theory and molecular geometry (Topic 4, Part 2), because you need Lewis structures to determine electron domains before predicting shape. It also feeds into intermolecular forces: the polarity of bonds (from electronegativity differences) determines whether a molecule is polar or nonpolar, which controls properties like boiling point and solubility. If you go on to organic chemistry, Lewis structures are the foundation for understanding reaction mechanisms.
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