Difficulty: Introductory to Intermediate | Prerequisites: Part 1 of these notes (ionic bonding, valence electrons, periodic table groups).
Covalent bonding is the other major type of chemical bond, and it governs the chemistry of most organic molecules, gases, and biological systems. Where ionic bonding transfers electrons, covalent bonding shares them. This part covers how to draw Lewis structures, predict molecular formulas for covalent compounds, determine molecular shape using VSEPR theory, and distinguish ionic from covalent compounds. Molecular geometry matters because the shape of a molecule determines its polarity, reactivity, and biological function.
Covalent bonds form when two nonmetal atoms share electrons rather than transferring them. You can draw Lewis structures to show how electrons are shared (bonding pairs) and where unshared electrons sit (lone pairs), then use VSEPR theory to predict the three-dimensional shape of the molecule based on how electron groups repel each other around the central atom.
Covalent bond
A chemical bond formed when two atoms share one or more pairs of electrons. This typically occurs between two nonmetals. Think of it as two atoms holding onto the same electrons rather than one giving them away.
Lewis structure (electron dot structure)
A diagram that shows the bonding between atoms and the lone pairs of electrons in a molecule. Dots represent electrons, and lines represent shared pairs (bonds).
Bonding pair
A pair of electrons shared between two atoms in a covalent bond. Each single line in a Lewis structure represents one bonding pair.
Lone pair (non-bonding pair)
A pair of valence electrons that belong to one atom and are not shared with another atom. Lone pairs still take up space around the atom and affect molecular shape.
Single bond
One shared pair of electrons between two atoms (one line in a Lewis structure). Example: Cl-Cl in Cl2.
Double bond
Two shared pairs of electrons between two atoms (two lines). Example: O=O in O2.
Triple bond
Three shared pairs of electrons between two atoms (three lines). Example: the N-to-N bond in N2.
Diatomic molecule
A molecule made of exactly two atoms. Seven elements exist naturally as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2. The mnemonic is "HOFBrINCl" (or "Have No Fear Of Ice Cold Beer").
VSEPR theory (Valence Shell Electron Pair Repulsion)
A model that predicts molecular geometry based on the idea that electron groups (bonds and lone pairs) around a central atom repel each other and arrange themselves as far apart as possible. In simple terms, electron clouds push each other away, and the shape of the molecule is whatever arrangement minimises that repulsion.
Molecular geometry
The three-dimensional shape of a molecule, determined by the positions of its atoms (not its lone pairs). Common shapes include linear, bent, trigonal planar, trigonal pyramidal, and tetrahedral.
Linear
A molecular shape where all atoms are arranged in a straight line, with a bond angle of 180 degrees. All diatomic molecules (e.g. Cl2, N2, ClF) are linear.
Bent (angular)
A molecular shape where two bonds and one or two lone pairs on the central atom create a V-shape. Bond angle is less than the ideal (around 104.5 to 120 degrees depending on the number of lone pairs). Example: SF2, H2O.
Trigonal pyramidal
A molecular shape where three bonds and one lone pair on the central atom create a pyramid shape (like a tripod). Bond angle is about 107 degrees. Example: NCl3, NH3.
Tetrahedral
A molecular shape where four bonds around a central atom point toward the corners of a tetrahedron, with bond angles of 109.5 degrees. Example: CH4.
Electronegativity
A measure of how strongly an atom attracts shared electrons in a covalent bond. The greater the electronegativity difference between two atoms, the more polar the bond. When the difference is very large (roughly above 1.7), the bond is considered ionic rather than covalent.
Covalent bonds form between nonmetal atoms that both need to gain electrons to reach a full octet.
Rather than transferring electrons, the atoms share one or more pairs of electrons so that both atoms count those shared electrons toward their octet.
The number of covalent bonds an atom forms typically equals the number of electrons it needs to complete its octet:
Group 17 (7 valence electrons, need 1 more): forms 1 bond (e.g. F, Cl, Br, I)
Group 16 (6 valence electrons, need 2 more): forms 2 bonds (e.g. O, S)
Group 15 (5 valence electrons, need 3 more): forms 3 bonds (e.g. N, P)
Group 14 (4 valence electrons, need 4 more): forms 4 bonds (e.g. C)
Count the total number of valence electrons from all atoms in the molecule.
Draw a skeleton structure with the central atom (usually the least electronegative) in the middle, connected to surrounding atoms by single bonds.
Distribute remaining electrons as lone pairs to satisfy the octet rule on each outer atom first, then on the central atom.
If the central atom does not have a full octet, convert lone pairs from outer atoms into double or triple bonds.
All diatomic molecules are linear (two atoms can only arrange themselves in a straight line).
Cl2: each chlorine has 7 valence electrons, shares 1 pair. Single bond, with 3 lone pairs on each Cl. Shape: linear.
S2: each sulphur has 6 valence electrons, shares 2 pairs. Double bond, with 2 lone pairs on each S. Shape: linear.
N2: each nitrogen has 5 valence electrons, shares 3 pairs. Triple bond, with 1 lone pair on each N. Shape: linear.
VSEPR counts the total electron groups (bonds + lone pairs) around the central atom, then determines the shape based on where the atoms (not lone pairs) end up.
2 electron groups, 0 lone pairs: linear (180 degrees). Example: ClF, BeCl2.
3 electron groups, 0 lone pairs: trigonal planar (120 degrees). Example: BF3.
3 electron groups, 1 lone pair: bent (less than 120 degrees). Example: SO2.
4 electron groups, 0 lone pairs: tetrahedral (109.5 degrees). Example: CH4.
4 electron groups, 1 lone pair: trigonal pyramidal (about 107 degrees). Example: NCl3, NH3.
4 electron groups, 2 lone pairs: bent (about 104.5 degrees). Example: SF2, H2O.
Key point: lone pairs take up more space than bonding pairs, which compresses the bond angles.
SF2 (sulphur difluoride): S has 2 bonds and 2 lone pairs, giving 4 electron groups. Shape: bent.
NCl3 (nitrogen trichloride): N has 3 bonds and 1 lone pair, giving 4 electron groups. Shape: trigonal pyramidal.
ClF (chlorine monofluoride): two atoms, one bond. Shape: linear.
Covalent compounds between two nonmetals use Greek prefixes to indicate the number of each atom:
1 = mono (often omitted for the first element), 2 = di, 3 = tri, 4 = tetra, 5 = penta, 6 = hexa
The second element's name ends in "-ide."
Examples: SF2 = sulphur difluoride, NCl3 = nitrogen trichloride, ClF = chlorine monofluoride, N2O4 = dinitrogen tetroxide.
The quick rule: if the compound contains a metal and a nonmetal, it is ionic. If it contains only nonmetals, it is covalent.
Ionic examples: strontium phosphide (Sr3P2), cesium nitride (Cs3N). Both contain a metal bonded to a nonmetal.
Covalent examples: beryllium bromide (BeBr2), phosphorus trifluoride (PF3), sulphur diiodide (SI2), hydrogen sulphide (H2S). These are nonmetal-nonmetal compounds.
Note: beryllium is technically a metal, but BeBr2 has significant covalent character because beryllium is small with a high charge density. This is a known edge case. For exam purposes, the metal/nonmetal rule is the expected approach unless your professor notes otherwise.
Electron Groups | Lone Pairs | Molecular Shape | Bond Angle | Example |
|---|---|---|---|---|
2 | 0 | Linear | 180 degrees | BeCl2, CO2 |
3 | 0 | Trigonal planar | 120 degrees | BF3 |
3 | 1 | Bent | Less than 120 degrees | SO2 |
4 | 0 | Tetrahedral | 109.5 degrees | CH4 |
4 | 1 | Trigonal pyramidal | About 107 degrees | NH3, NCl3 |
4 | 2 | Bent | About 104.5 degrees | H2O, SF2 |
Number | Prefix |
|---|---|
1 | mono- |
2 | di- |
3 | tri- |
4 | tetra- |
5 | penta- |
6 | hexa- |
7 | hepta- |
8 | octa- |
"Mono-" is dropped from the first element's name (we say "carbon dioxide," not "monocarbon dioxide"), but is kept on the second element ("carbon monoxide").
Molecular geometry determines how molecules interact with each other and with biological systems. Water's bent shape makes it polar, which is why it dissolves salts and sustains life. Drug molecules work by fitting into receptor proteins like a key in a lock, and that fit depends entirely on the molecule's three-dimensional shape. The tetrahedral geometry of carbon is the reason organic chemistry (and therefore all of biochemistry) is built on carbon backbones.
Students often confuse electron geometry with molecular geometry. Electron geometry counts all electron groups (bonds and lone pairs). Molecular geometry describes only where the atoms sit. A molecule with 4 electron groups but 2 lone pairs has tetrahedral electron geometry but bent molecular geometry.
All diatomic molecules are linear. Students sometimes assign more complex shapes (like trigonal bipyramidal or bent) to diatomic molecules such as Cl2 or N2. Two atoms can only form a straight line.
Students frequently forget to count lone pairs when predicting shape. SF2 has 2 bonds, which might suggest linear, but S also has 2 lone pairs, making 4 electron groups total and giving a bent shape.
The ionic vs covalent distinction trips students up with borderline elements like beryllium and hydrogen. The general rule (metal + nonmetal = ionic, nonmetal + nonmetal = covalent) works for exams, but real bonding is a spectrum.
Drawing Lewis structures and predicting molecular shape from them is one of the most commonly tested skills in general chemistry.
Expect questions that give you a molecular formula and ask for the Lewis structure, the number of bonding pairs and lone pairs, and the molecular geometry.
The VSEPR shape table (electron groups, lone pairs, resulting shape, bond angle) is worth memorising. It appears on exams in both direct and applied forms.
Naming covalent compounds with Greek prefixes is tested separately from ionic naming. Know which naming system applies to which type of compound.
Classifying compounds as ionic or covalent based on their constituent elements is a standard exam question.
True or False: A molecule made of two atoms can have a bent shape.
Answer: False. All diatomic molecules are linear.
Fill in the blank: Nitrogen (Group 15) typically forms ______ covalent bond(s).
Answer: 3
True or False: Lone pairs on the central atom do not affect molecular geometry.
Answer: False. Lone pairs repel bonding pairs and change the molecular shape.
Fill in the blank: A molecule with 4 electron groups and 1 lone pair on the central atom has a ______ shape.
Answer: Trigonal pyramidal
True or False: A compound made of two nonmetals is covalent.
Answer: True.
Q: Draw the Lewis structure of H2S and predict its molecular geometry.
A: Sulphur is the central atom with 6 valence electrons. Each hydrogen contributes 1, giving 8 total. S forms 2 single bonds to H and has 2 lone pairs. Four electron groups with 2 lone pairs gives a bent shape (about 104.5 degrees).
Q: Why is NCl3 trigonal pyramidal rather than trigonal planar?
A: Nitrogen has 5 valence electrons. It forms 3 bonds to 3 chlorine atoms and retains 1 lone pair. The lone pair occupies the fourth electron group position, pushing the three Cl atoms down into a pyramidal arrangement. Trigonal planar would require 0 lone pairs on the central atom.
Q: Classify each compound as ionic or covalent: (a) PF3, (b) Cs3N, (c) BeBr2, (d) H2S.
A: (a) PF3, covalent (both nonmetals). (b) Cs3N, ionic (Cs is a metal, N is a nonmetal). (c) BeBr2, classified as covalent in this lab context (Be has significant covalent character). (d) H2S, covalent (both nonmetals).
Q: N2 contains a triple bond. How many lone pairs does each nitrogen have in N2?
A: Each nitrogen has 5 valence electrons. Three are used in the triple bond (3 bonding pairs shared), leaving 2 electrons = 1 lone pair on each nitrogen.
Q: What is the molecular geometry of a molecule with 4 electron groups and 2 lone pairs on the central atom? Give an example from the lab.
A: Bent. SF2 is the example: sulphur has 2 bonds to fluorine and 2 lone pairs.
Molecular geometry connects directly to polarity and intermolecular forces, which determine boiling points, solubility, and physical properties of substances. The distinction between ionic and covalent bonding feeds into understanding solubility rules, electrical conductivity of solutions, and electrochemistry. Lewis structures are also the foundation for resonance structures and formal charge, which come up later in the course.
Covalent bonding, Lewis dot structures, electron dot diagrams, VSEPR theory, molecular geometry, molecular shape, linear molecule, bent molecule, trigonal pyramidal, tetrahedral, lone pairs, bonding pairs, single bond, double bond, triple bond, diatomic molecules, HOFBrINCl, naming covalent compounds, Greek prefixes, ionic vs covalent, metal nonmetal rule, electronegativity, polar covalent, nonpolar covalent, bond angle, electron group geometry, general chemistry lab 4, Purdue CHM, Lewis structure practice