Difficulty: Introductory | Prerequisites: Basic atomic structure (protons, neutrons, electrons), electron shells, valence electrons, periodic table groups.
Electron dot diagrams sit at the boundary between understanding atomic structure and predicting how atoms bond. If you are comfortable with electron configuration and valence electrons, you are ready. If "2, 8, 8" means nothing to you yet, review atomic structure first.
This worksheet covers how atoms share electrons to form covalent bonds, how to represent those bonds with Lewis (electron dot) diagrams, and how to distinguish covalent bonding from ionic and metallic bonding. It also introduces the octet rule, bonding versus non-bonding pairs, and the relationship between bond order, bond length and bond strength. The material is foundational for everything that follows in molecular geometry, polarity and intermolecular forces.
Atoms form covalent bonds by sharing valence electrons so that each atom (usually) ends up with eight electrons in its outer shell. Electron dot (Lewis) diagrams show every valence electron around each atom, making it easy to count bonding pairs (shared) and lone pairs (not shared). The number of bonds an atom forms, and whether those bonds are single, double or triple, follows directly from how many electrons it needs to complete its octet.
Covalent bond
A chemical bond formed when two non-metal atoms share one or more pairs of electrons. In simple terms, both atoms contribute electrons to a shared pool so each can fill its outer shell.
Ionic bond
A bond formed by the electrostatic attraction between a positively charged ion (cation) and a negatively charged ion (anion), typically between a metal and a non-metal. Think of it as one atom handing its electron to another, creating opposite charges that stick together.
Metallic bond
Bonding in metals where valence electrons are delocalised across a lattice of positive metal ions. In simple terms, the outer electrons float freely in a "sea" shared by all the metal atoms, which is why metals conduct electricity.
Electron dot diagram (Lewis structure)
A diagram showing all valence electrons of each atom in a molecule, with dots for lone electrons and lines (or paired dots) for shared bonding pairs. Think of it as a map of who is sharing what.
Octet rule
The principle that atoms tend to gain, lose or share electrons until they have eight electrons in their outermost shell. Hydrogen is the main exception: it is satisfied with just two (a "duet").
Bonding pair
A pair of electrons shared between two atoms in a covalent bond. Each line in a Lewis structure represents one bonding pair.
Non-bonding pair (lone pair)
A pair of valence electrons on an atom that is not involved in bonding. These are the dots that sit on one atom only, not between two atoms.
Single bond
One shared pair of electrons between two atoms (represented by one line or two dots).
Double bond
Two shared pairs of electrons between two atoms (represented by two lines or four dots). Shorter and stronger than a single bond.
Triple bond
Three shared pairs of electrons between two atoms (represented by three lines or six dots). Shorter and stronger than a double bond.
Valence electrons
The electrons in the outermost shell of an atom. These are the electrons involved in bonding. The group number on the periodic table tells you how many valence electrons an element has (for main-group elements).
Molecule
A discrete group of atoms held together by covalent bonds. Not all substances are molecular: metals and ionic compounds form extended lattices, not individual molecules.
Substances fall into three broad bonding categories. The type depends on which elements are involved:
Metallic: Pure metals and alloys (e.g. copper, brass). Electrons are delocalised across a lattice of metal cations.
Ionic: Compounds of a metal and a non-metal (e.g. calcium sulfide, potassium carbonate). Electrons are transferred from the metal to the non-metal.
Covalent: Compounds of two or more non-metals (e.g. O₂, NH₃, H₂O). Electrons are shared.
Some compounds contain both ionic and covalent bonding. Potassium carbonate (K₂CO₃) is ionic between K⁺ and CO₃²⁻, but the carbonate ion itself has covalent bonds between carbon and oxygen. Calcium carbonate (CaCO₃) works the same way: Ca²⁺ is ionically bonded to the covalent CO₃²⁻ ion.
A molecule is a discrete particle of covalently bonded atoms. When deciding whether something is "composed of molecules":
Glucose (C₆H₁₂O₆): yes, a covalent molecular compound.
Nitrogen monoxide (NO): yes, a covalent molecule.
Helium (He): no, it is monatomic (single atoms, no bonds).
Brass: no, it is an alloy (metallic bonding, no discrete molecules).
Aluminium chloride (AlCl₃): no in the solid state, it forms an ionic lattice. (In the gas phase it can exist as molecular Al₂Cl₆, but the default answer treats it as ionic.)
The method, step by step:
Count the total valence electrons for every atom in the molecule. For polyatomic ions, add one electron per negative charge or subtract one per positive charge.
Place the least electronegative atom in the centre (hydrogen is always on the outside).
Draw a single bond (one pair of electrons) between each pair of bonded atoms.
Distribute the remaining electrons as lone pairs, starting with the outer atoms, to satisfy the octet rule.
If an atom (typically the central atom) still lacks an octet, convert one or more lone pairs on a neighbouring atom into additional bonding pairs to form double or triple bonds.
Carbon has an electron configuration of 2, 4, meaning four valence electrons. It needs four more to reach eight. Each hydrogen atom has one valence electron and needs one more to fill its shell (duet rule).
Four hydrogen atoms each contribute one electron, giving carbon a share of eight. That is why the formula is CH₄, not CH₃ (which would leave carbon with only seven) or CH₅ (which would give carbon nine, exceeding the stable octet).
Once you draw the Lewis structure, count:
Bonding pairs: electron pairs sitting between two atoms (shared). Each single bond = 1 bonding pair. Each double bond = 2 bonding pairs. Each triple bond = 3 bonding pairs.
Non-bonding (lone) pairs: electron pairs sitting on one atom only.
Worked examples from the worksheet:
F₂: One single bond (1 bonding pair), six lone pairs (three on each fluorine) = 6 non-bonding pairs.
CO₂: Two double bonds (4 bonding pairs total, counted as "2 double"), four lone pairs (two on each oxygen) = 4 non-bonding pairs.
H₂S: Two single bonds (2 bonding pairs), two lone pairs on sulfur = 2 non-bonding pairs.
More shared pairs means a shorter, stronger bond:
Single bond: longest, weakest.
Double bond: intermediate.
Triple bond: shortest, strongest.
Comparing O₂ and N₂:
O₂ has a double bond. N₂ has a triple bond.
O₂ therefore has the longer bond length.
N₂ has the stronger bond.
From the worksheet's list (ethyne C₂H₂, trichloromethane CHCl₃, sulfur dioxide SO₂, phosphorus trichloride PCl₃):
Contains three single bonds: phosphorus trichloride (PCl₃). Phosphorus forms one single bond to each of three chlorine atoms.
Nine lone pairs of electrons: trichloromethane (CHCl₃). Each Cl has three lone pairs (3 × 3 = 9).
Contains a double bond: sulfur dioxide (SO₂). Sulfur forms a double bond with each oxygen (or one double and one single with resonance, depending on the model used at this level).
Contains a triple bond: ethyne (C₂H₂). The two carbon atoms share three pairs of electrons.
Electron dot diagrams are difficult to render in text, but here is a summary of the key structures from the worksheet. Use these as a reference when drawing your own.
F₂ (fluorine gas)
Each fluorine has 7 valence electrons. They share one pair to form a single bond. Each F retains three lone pairs. Bonding pairs: 1. Non-bonding pairs: 6.
CO₂ (carbon dioxide)
Carbon has 4 valence electrons, each oxygen has 6. Carbon forms a double bond with each oxygen. Each oxygen retains two lone pairs. Bonding pairs: 4 (two double bonds). Non-bonding pairs: 4.
H₂S (hydrogen sulfide)
Sulfur has 6 valence electrons. It forms a single bond with each of two hydrogen atoms, leaving two lone pairs on sulfur. Bonding pairs: 2. Non-bonding pairs: 2.
CH₄ (methane)
Carbon forms four single bonds, one to each hydrogen. No lone pairs on carbon. Bonding pairs: 4. Non-bonding pairs: 0.
H₂O (water)
Oxygen forms two single bonds (one to each hydrogen) and retains two lone pairs. Bonding pairs: 2. Non-bonding pairs: 2.
H₂O₂ (hydrogen peroxide)
Each oxygen bonds to one hydrogen and to the other oxygen. Each oxygen has two lone pairs. Structure: H-O-O-H. Bonding pairs: 3. Non-bonding pairs: 4.
OH⁻ (hydroxide ion)
Oxygen bonds to one hydrogen. The extra electron (from the negative charge) gives oxygen three lone pairs. Enclose the structure in brackets with a minus sign. Bonding pairs: 1. Non-bonding pairs: 3.
NO₃⁻ (nitrate ion)
Nitrogen is the central atom, bonded to three oxygens. One N-O bond is a double bond; the other two are single bonds (with resonance, the double bond can be drawn to any of the three oxygens). The ion carries a negative charge. Enclose in brackets with a minus sign. Total valence electrons: 5 (N) + 3 × 6 (O) + 1 (charge) = 24.
CaCO₃ (calcium carbonate)
This compound contains both ionic and covalent bonding. The calcium ion (Ca²⁺) is ionically bonded to the carbonate ion (CO₃²⁻). Within the carbonate ion, carbon is covalently bonded to three oxygens with one double bond and two single bonds (plus resonance).
Lewis structures are the starting point for predicting molecular shape (VSEPR theory), which in turn determines polarity, and polarity determines physical properties like boiling point and solubility. Pharmaceutical chemists use Lewis structures daily to understand how drug molecules will interact with receptors in the body.
The distinction between ionic and covalent bonding explains why table salt (NaCl) dissolves in water and conducts electricity when dissolved, while sugar (C₁₂H₂₂O₁₁) dissolves but does not conduct: one dissociates into ions, the other stays molecular.
Students often think ionic compounds are "molecules." They are not. Ionic compounds form extended lattices with no discrete molecular unit. NaCl is not a molecule.
Students sometimes forget that hydrogen only needs two electrons (a duet), not eight. Drawing four bonds to hydrogen is a common early mistake.
A double bond is sometimes counted as one bonding pair. It is two bonding pairs. A triple bond is three bonding pairs. Count the pairs, not the bonds.
Students confuse the total electron count with the number of bonding pairs. The total includes all valence electrons. Bonding pairs are only the shared ones.
Expect to be asked to draw Lewis structures for simple molecules and polyatomic ions, including those with double and triple bonds.
Classifying substances as ionic, covalent or metallic is a reliable exam question. Know the rule: metal + non-metal = ionic, non-metal + non-metal = covalent, pure metal or alloy = metallic.
Counting bonding pairs and lone pairs from a Lewis structure is frequently tested. Make sure you can do it quickly and accurately.
Bond order vs bond length vs bond strength: more bonds = shorter and stronger. O₂ (double) vs N₂ (triple) is a textbook comparison.
Compounds with both ionic and covalent character (like CaCO₃) are a favourite exam topic because they test whether you can see both bonding types in one substance.
True or false: Brass is composed of molecules. (False: brass is an alloy with metallic bonding.)
Fill in the blank: A double bond consists of ______ bonding pairs. (Two.)
True or false: N₂ has a longer bond length than O₂. (False: N₂ has a triple bond, which is shorter than O₂'s double bond.)
Fill in the blank: Carbon forms ______ bonds with hydrogen in CH₄ because it needs four more electrons to complete its octet. (Four.)
True or false: The carbonate ion (CO₃²⁻) contains only ionic bonds. (False: the bonds between C and O within the ion are covalent.)
Q: Classify the bonding in each of the following: copper, oxygen (O₂), ammonia (NH₃), calcium sulfide, potassium carbonate.
A: Copper: metallic. O₂: covalent. NH₃: covalent. Calcium sulfide: ionic. Potassium carbonate: ionic (between K⁺ and CO₃²⁻) with covalent bonding within the carbonate ion.
Q: Draw the electron dot diagram for H₂O and state the number of bonding and non-bonding pairs.
A: Oxygen is the central atom with two single bonds to hydrogen and two lone pairs. Bonding pairs: 2. Non-bonding pairs: 2.
Q: Explain why carbon forms CH₄ rather than CH₃ or CH₅.
A: Carbon has four valence electrons and needs four more to complete its octet. Each hydrogen provides one electron. Four hydrogens give carbon exactly eight outer-shell electrons. Three would leave it one short (seven), and five would exceed the stable octet (nine).
Q: Which has the longer bond length, O₂ or N₂? Explain your reasoning.
A: O₂ has the longer bond length. O₂ contains a double bond, while N₂ contains a triple bond. A triple bond pulls the nuclei closer together than a double bond does.
Q: From the list (ethyne, trichloromethane, sulfur dioxide, phosphorus trichloride), which substance contains a triple bond?
A: Ethyne (C₂H₂). The two carbon atoms are connected by a triple bond, with one hydrogen bonded to each carbon.
Q: Describe the bonding in calcium carbonate (CaCO₃).
A: There is ionic bonding between the Ca²⁺ ion and the CO₃²⁻ ion. Within the carbonate ion, the bonding between carbon and the three oxygen atoms is covalent (one double bond and two single bonds, with resonance delocalising the double bond across all three C-O positions).
Electron dot diagrams feed directly into VSEPR theory (Valence Shell Electron Pair Repulsion), which predicts the three-dimensional shape of molecules based on how bonding and lone pairs arrange themselves around a central atom. The shape then determines molecular polarity.
Polarity, in turn, governs intermolecular forces (dipole-dipole, hydrogen bonding, London dispersion), which explain boiling points, solubility and states of matter. So Lewis structures are the first domino in a chain that runs through most of introductory chemistry.
The ionic/covalent distinction you practised here also connects to electronegativity differences and the bonding spectrum, which you will meet when studying polar covalent bonds.
Lewis structure, Lewis dot diagram, electron dot diagram, electron dot structure, dot and cross diagram, valence electrons, covalent bonding, ionic bonding, metallic bonding, octet rule, duet rule, bonding pair, lone pair, non-bonding pair, single bond, double bond, triple bond, bond order, bond length, bond strength, molecular formula, polyatomic ion, resonance structure, F₂, CO₂, H₂S, CH₄, H₂O, H₂O₂, OH⁻, NO₃⁻, CaCO₃, VSEPR, molecular geometry, Chemistry for WA, Pearson Chemistry, Worksheet 4.1