Electron sharing, molecules, Lewis structures, bond types, the octet rule, and the properties of molecular compounds — the chemistry that underlies life, materials, and the atmosphere
The world is full of substances that do not conduct electricity, do not dissolve easily in water, and melt at temperatures far lower than the compounds you studied in Chapter 9. Air, water, sugar, alcohol, wax, gasoline, plastic, DNA — these substances behave very differently from table salt or magnesium oxide. The reason is that their atoms are not held together by opposite charges. Instead, their atoms share electrons.
Covalent bonding is the sharing of electrons between nonmetal atoms. When two atoms share electrons, both atoms gain stability without either one giving up an electron completely. The result is a molecule — a group of atoms held together by shared electron pairs and behaving as a single particle.
By the end of this chapter, you will understand how covalent bonds form and why sharing produces stability, how to draw Lewis structures for simple molecules, how single, double, and triple bonds differ in strength and length, what the octet rule is and when it has exceptions, and how the covalent nature of a compound determines its physical properties. You will also see how covalent bonding appears in living systems, materials science, medicine, and environmental chemistry.
Oxygen gas keeps us alive; nitrogen gas makes up most of the air we breathe; water is the solvent of life. All three are molecular compounds held together by covalent bonds. Yet nitrogen is almost completely unreactive, oxygen reacts readily with fuels and metals, and water dissolves almost everything. Why do covalently bonded molecules behave so differently from each other? By the end of this chapter, you will have the tools — Lewis structures, bond types, lone pairs, and the octet rule — to begin answering that question.
In Chapter 9, you learned that ionic bonding occurs when a metal atom transfers one or more electrons to a nonmetal atom. But what happens when two nonmetal atoms meet? Neither atom has a strong tendency to give up electrons. Both want to reach a stable electron configuration. Neither will simply hand its electrons to the other.
The answer is sharing. When two nonmetal atoms approach each other closely, their electron clouds begin to overlap. If both atoms have unpaired valence electrons, those electrons can be shared between the two nuclei simultaneously. Each shared pair of electrons is attracted to both nuclei at once — and this mutual attraction holds the atoms together.
A covalent bond is a chemical bond formed by the sharing of one or more pairs of electrons between two atoms. The word comes from co- (together) and valent (referring to valence electrons). Covalent bonds are partnerships between valence electrons, not transfers.
Consider two hydrogen atoms. Each has one electron and needs two to be stable. When they approach, their electrons begin to occupy the space around both nuclei. Both electrons are now attracted to both protons. The atoms have found a configuration that lowers the total energy of the system. This is the essence of covalent bonding: a shared pair of electrons attracts two nuclei simultaneously, pulling them toward each other and lowering energy.
When atoms are held together by covalent bonds, they form a molecule — the smallest unit of a covalently bonded substance that retains the chemical identity of that substance. Water, carbon dioxide, oxygen gas, and glucose are all molecules.
| Property | Ionic Compounds | Molecular Compounds |
|---|---|---|
| What electrons do | Transferred from metal to nonmetal | Shared between nonmetal atoms |
| Particles formed | Cations and anions in a lattice | Discrete molecules |
| Typical elements | Metal + nonmetal | Nonmetal + nonmetal |
| Melting point | Generally high | Generally low to moderate |
| Conductivity (liquid) | Conducts | Usually does not conduct |
| Solubility in water | Often soluble | Variable; depends on polarity |
Ionic bonding involves complete electron transfer. Covalent bonding involves electron sharing. This fundamental difference produces entirely different particle types and entirely different physical properties.
A molecule is a discrete group of atoms held together by covalent bonds, with a definite composition and a definite structure. This stands in contrast to ionic compounds, which exist as extended lattice networks without discrete units. A sodium chloride crystal is not a single molecule; it is a repeating grid of ions. But a water molecule is truly a unit of three specific atoms in a specific arrangement.
The molecular formula of a compound tells you the exact number of each type of atom in one molecule. Examples: H2O (two hydrogen, one oxygen), CO2 (one carbon, two oxygen), NH3 (one nitrogen, three hydrogen), CH4 (one carbon, four hydrogen), C6H12O6 (glucose). Molecular formulas tell you composition but not structure.
Several elements exist naturally as two-atom molecules. These are called diatomic molecules. When writing chemical equations, always write these elements with the subscript 2 — writing a bare symbol like "O" represents a single atom, not the naturally occurring form.
| Element | Symbol | State at Room Temperature |
|---|---|---|
| Hydrogen | H₂ | Gas |
| Nitrogen | N₂ | Gas |
| Oxygen | O₂ | Gas |
| Fluorine | F₂ | Gas |
| Chlorine | Cl₂ | Gas |
| Bromine | Br₂ | Liquid |
| Iodine | I₂ | Solid |
In a chemical equation, writing H instead of H₂ or O instead of O₂ is chemically incorrect. It represents a single, unstable atom rather than the stable naturally occurring form. Always write the subscript 2 for these seven elements when they appear as pure substances.
A molecular formula gives the actual number of atoms in one molecule. An empirical formula gives the simplest whole-number ratio of atoms. For glucose (C6H12O6), the empirical formula is CH2O. For hydrogen peroxide (H2O2), the empirical formula is HO. For water, the two are the same. You will work with empirical formulas when analyzing experimental data.
When two atoms form a covalent bond, one pair of electrons is shared between them. Each atom typically contributes one electron to the shared pair. The shared pair then belongs to both atoms simultaneously.
It is the valence electrons — those in an atom's outermost energy level — that participate in covalent bonding. Inner electrons remain close to their own nucleus and do not participate in bonding.
Consider hydrogen fluoride (HF). Hydrogen has one valence electron; fluorine has seven. They share one electron pair. Hydrogen now counts two electrons — matching helium. Fluorine now counts eight electrons — matching neon. Both atoms achieve stability through sharing.
In a Lewis structure, a shared pair of electrons between two atoms is written as a line (a bond) between those atoms. Each line represents two electrons being shared. Electrons that are not shared remain on their own atom as lone pairs — shown as pairs of dots.
H — H Hydrogen molecule: one shared pair (two electrons)
H — F: Hydrogen fluoride: one shared pair; three lone pairs on fluorine
Electrons not used in bonding are called lone pairs (or nonbonding pairs). They are not passive extras — lone pairs affect molecular shape, polarity, and reactivity. The lone pair on nitrogen in ammonia is what makes ammonia a base. The lone pairs on oxygen in water are responsible for water's bent shape and high boiling point.
Atoms can share more than one pair of electrons. The number of shared pairs determines whether the bond is single, double, or triple — and has a direct effect on bond strength and bond length.
| Bond Type | Shared Pairs | Electrons Shared | Relative Strength | Relative Length | Symbol |
|---|---|---|---|---|---|
| Single | 1 | 2 | Weakest | Longest | — |
| Double | 2 | 4 | Intermediate | Intermediate | = |
| Triple | 3 | 6 | Strongest | Shortest | ≡ |
Single bonds are found in H2, HCl, H2O, and CH4 (methane). Double bonds appear in O2 (O=O), CO2 (O=C=O), and formaldehyde. Triple bonds are found in N2 (N≡N) and CO (C≡O). Nitrogen's triple bond is famously strong — this is why N2 is so unreactive despite making up 78% of the atmosphere.
As the number of shared electron pairs increases between the same two types of atoms: Bond strength increases. More pairs pull the nuclei together more strongly, requiring more energy to break. Bond length decreases. More pairs draw the atoms closer together. These comparisons are valid only between bonds of the same element pair.
Bond strength comparisons (single vs. double vs. triple) are valid only when comparing bonds between the same pair of elements. A C=O double bond is stronger than a C—O single bond. But do not compare, say, a C—C bond to an N=N bond without additional data — different element pairs have different inherent bond strengths.
A Lewis structure (also called a Lewis dot structure or electron dot structure) is a diagram showing the arrangement of atoms in a molecule, all bonding pairs of electrons (as lines between atoms), and all lone pairs on each atom (as pairs of dots). Lewis structures help you count electrons, check the octet rule, and predict molecular geometry.
Step 1 — Count valence electrons. Add up valence electrons from every atom. This is your total electron budget.
Step 2 — Identify the central atom. Usually the atom that appears once in the formula, or the least electronegative atom (not hydrogen, which is always terminal).
Step 3 — Connect atoms with single bonds. Each single bond uses 2 electrons from your budget.
Step 4 — Complete octets on outer atoms first. Place lone pairs on surrounding atoms (2 electrons for hydrogen, 8 for others). Subtract used electrons from your budget.
Step 5 — Place remaining electrons on the central atom as lone pairs.
Step 6 — Check the central atom's electron count. If it has fewer than 8, convert lone pairs on adjacent atoms into double or triple bonds until the central atom achieves an octet.
Step 7 — Verify the total. Count all electrons in your diagram. This must equal the total from Step 1.
Total electrons: 4 (bonds) + 4 (lone pairs) = 8 ✓Total electrons: 4 bonds × 2 e⁻ = 8; 4 lone pairs × 2 e⁻ = 8; Total = 16 ✓Total electrons: 6 (bonds) + 2 (lone pair) = 8 ✓Bonding pairs are pairs of electrons shared between two atoms — they create and sustain the covalent bond. Lone pairs (nonbonding pairs) belong to only one atom and do not participate in bonding. Both occupy space around an atom, and both exert repulsive forces on nearby electron pairs. Lone pairs, however, repel more strongly than bonding pairs — a fact that determines the shape of molecules.
| Molecule | Bonding Pairs on Central Atom | Lone Pairs on Central Atom | Shape | Bond Angle |
|---|---|---|---|---|
CH₄ (methane) | 4 | 0 | Tetrahedral | 109.5° |
NH₃ (ammonia) | 3 | 1 | Trigonal pyramidal | ~107° |
H₂O (water) | 2 | 2 | Bent | ~104.5° |
Each additional lone pair on the central atom pushes the bonding pairs closer together, compressing the bond angle. This is why water has a bent shape rather than a straight line — the two lone pairs on oxygen force the hydrogen atoms together. You will study molecular geometry in full in Chapter 12.
Lone pairs give molecules the ability to donate electrons to other species. Nitrogen's lone pair in ammonia allows it to act as a base by accepting a hydrogen ion. Oxygen's two lone pairs in water allow water to act as both a base and a hydrogen bond donor. These chemical roles are explored in Chapter 23 (acids and bases) and Chapter 28 (biological chemistry).
Atoms tend to bond in ways that give each atom eight electrons in its outer energy level. This configuration matches the electron arrangement of the noble gases, which are chemically stable. Hydrogen is the primary exception — it follows the duet rule, achieving stability with just two electrons (the helium configuration).
| Element | Valence Electrons | Electrons Needed | Bonds Typically Formed |
|---|---|---|---|
| H | 1 | 1 (duet rule) | 1 |
| C | 4 | 4 | 4 |
| N | 5 | 3 | 3 |
| O | 6 | 2 | 2 |
| F | 7 | 1 | 1 |
| Cl | 7 | 1 | 1 |
The octet rule is a useful guideline, not an absolute law. Several important exceptions exist at the high school level.
| Exception Type | Examples | Explanation |
|---|---|---|
| Hydrogen and helium | H in any compound; He (no bonds) | Stability with 2 electrons (duet rule) |
| Electron-deficient (boron) | BF₃, BCl₃ | Boron forms 3 bonds, leaving only 6 electrons; can accept additional pairs |
| Expanded octet (Period 3+) | SF₆ (12 e⁻ on S), PCl₅ (10 e⁻ on P) | Access to d orbitals allows more than 8 electrons around the central atom |
| Odd-electron molecules | NO (nitrogen monoxide) | Odd number of valence electrons; at least one atom cannot complete an octet; called free radicals — very reactive |
At the high school level, you will work primarily with molecules where the octet rule holds. When you encounter a Period 3 or higher central atom (S, P, Cl, etc.) or boron, be aware that exceptions may apply. If your Lewis structure requires a Period 3 central atom to have more than 8 electrons to bond correctly with all surrounding atoms, this is expected — not an error.
The properties of ionic compounds arise from strong electrostatic attraction between oppositely charged ions in a crystal lattice. The properties of molecular compounds arise from a different set of forces — the attractions between neutral molecules, called intermolecular forces. These forces are typically much weaker than ionic bonds, and this difference produces all of the characteristic properties of molecular compounds.
Most molecular compounds have relatively low melting and boiling points. Water melts at 0°C. Methane melts at –183°C. Rubbing alcohol boils at only 83°C. This occurs because separating molecules requires overcoming only the weak intermolecular forces between them — not the covalent bonds within them. During melting and boiling, covalent bonds are not broken. The intermolecular attractions are. Ionic compounds, by contrast, require much higher temperatures to disrupt their strong lattice forces.
Ionic compounds dissolve in water to produce ions that carry electric current. Molecular compounds consist of neutral molecules. When they dissolve or melt, they do not typically produce ions, so they do not conduct electricity. Pure water is a poor conductor because it contains only neutral H2O molecules. Note that some molecular compounds — particularly acids — do produce ions when dissolved in water through a chemical process called ionization. HCl gas dissolves in water and releases H+ and Cl– ions, making hydrochloric acid conductive. But this is a chemical reaction, not simple dissolution of pre-existing ions.
The solubility of molecular compounds varies widely and depends largely on polarity. Polar molecular compounds tend to dissolve well in water, because water is also polar — sugar, ethanol, and acetic acid dissolve readily. Nonpolar molecular compounds tend to be insoluble in water but dissolve in nonpolar solvents — oil, wax, and most hydrocarbons do not dissolve in water. The guiding principle is "like dissolves like" — polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes. You will study polarity in Chapter 12.
| State | Examples | Explanation |
|---|---|---|
| Gas | H₂, O₂, N₂, CO₂, CH₄, NH₃ | Very weak intermolecular forces; thermal energy easily separates molecules |
| Liquid | H₂O, ethanol, acetone, acetic acid | Moderate intermolecular forces; molecules held loosely together |
| Solid | Glucose, sucrose, wax, many plastics | Larger molecules; stronger intermolecular forces outweigh thermal energy |
Virtually every molecule in a living cell is held together by covalent bonds. Proteins, carbohydrates, lipids, and nucleic acids — the four major classes of biological molecules — are all molecular compounds built on covalent frameworks. Covalent bonding is not merely one type of bonding among many; it is the chemistry of life itself.
Carbon's four valence electrons allow it to form four covalent bonds, including bonds to other carbon atoms. This makes carbon capable of forming chains, rings, branched structures, and complex three-dimensional frameworks of extraordinary variety. No other element rivals carbon's structural diversity. Oxygen's double bonds in sugars and fats store chemical energy. Nitrogen's lone pair in amino acids allows proteins to form their characteristic shapes. Hydrogen bonding — a special attraction involving hydrogen covalently bonded to N, O, or F — holds DNA's two strands together and gives proteins their secondary structure.
Polymers are long-chain molecules made of repeating units called monomers, connected by covalent bonds. Plastics, nylon, rubber, polyester, and many fibers are polymers. The covalent bonds in polymers give them strength, flexibility, and durability.
Diamond is an extreme covalent solid — every carbon atom is bonded to four neighbors in a tetrahedral arrangement extending throughout the entire crystal. This network of covalent bonds makes diamond the hardest natural material known. Graphite — another form of pure carbon — consists of layers of hexagonally bonded carbon atoms. The covalent bonds within each layer are strong; the intermolecular forces between layers are weak, which is why graphite is soft enough to write with.
Silicon dioxide (SiO2), the major component of sand and quartz, is a covalent network solid with a three-dimensional network of covalent bonds — giving it high hardness and a very high melting point.
Every pharmaceutical drug is a molecular compound. Its three-dimensional shape — determined entirely by its covalent bonds — controls how it interacts with proteins and receptors in the body. Anesthetics, antibiotics, vitamins, and hormones all depend on the precise geometry of their covalent frameworks.
The key molecules of Earth's atmosphere — N2, O2, CO2, H2O, O3 — are all molecular compounds. Ozone formation and destruction, the greenhouse effect, and acid rain all involve covalent reactions. Photosynthesis and respiration are covalent transformations: carbon dioxide and water become glucose and oxygen, then glucose and oxygen become carbon dioxide and water again — driving most of life on Earth.
In this chapter, you have learned what covalent bonds are, how to draw Lewis structures, and how bond type and lone pairs affect molecular properties. Chapter 12 will extend this further, explaining how lone pairs and bonding pairs determine the three-dimensional shape of molecules, why some molecules are polar and others are not, and how polarity and molecular shape determine the intermolecular forces that govern boiling points, solubility, and biological function.
Answer each question in one to three sentences before consulting notes or the chapter text.
Draw the Lewis structure for each molecule. Show all bonding pairs and lone pairs. Verify your total electron count matches the valence electron total.
Identify whether each molecule contains single bonds only, at least one double bond, or at least one triple bond. Then state the total number of covalent bonds in the molecule.
For each atom in the molecule described, state whether the octet rule is satisfied. If not, identify the exception and explain it.
Each statement or structure below contains an error. Identify the error and write the correct version.
What is wrong, and why does it matter? What should be written instead?
What is wrong with this structure? What should the correct Lewis structure show?
Identify the error and provide a counterexample of a covalent compound that does not dissolve in water.
What is the error in this reasoning? How should it be corrected?
What is wrong? How many electrons should the correct structure show? What type of bond should connect the nitrogen atoms?
Provide one specific example showing why this claim is false.
What is wrong with this reasoning? What is the correct formula and how many bonds does oxygen actually form?
Is this exactly true? What is a more accurate and complete statement about double vs. single bond strength?
Most covalent compounds do not conduct electricity, but acids — which are molecular compounds — ionize in water to produce ions. Hydrochloric acid, sulfuric acid, and nitric acid are all covalent compounds that produce conductive solutions. Conductivity depends on whether ions are present in solution, not strictly on whether the original substance was ionic or molecular.
Lone pairs are chemically essential. They determine molecular shape, participate in hydrogen bonding, allow molecules to act as acids or bases, and enable reactions with other molecules. The lone pair on nitrogen in ammonia is what makes ammonia a base. The lone pairs on oxygen in water determine its bent shape and high boiling point.
Breaking bonds requires energy input — it is endothermic. Forming bonds releases energy — it is exothermic. A reaction is exothermic overall when more energy is released forming product bonds than is absorbed breaking reactant bonds. This idea is central to thermochemistry in Chapter 20.
The octet rule is a guideline that works well for Period 2 elements but breaks down for hydrogen (duet rule), boron (often 6 electrons), and Period 3 elements (expanded octets). Recognizing when the octet rule applies and when exceptions are expected is part of chemical reasoning.
An atom is the fundamental unit of an element. A molecule is two or more atoms bonded together. Oxygen as an element exists as O₂ molecules, not as individual O atoms under normal conditions. A lone O atom would be an extremely reactive species very different from the stable oxygen gas you breathe.
Chapter 10 gave you the tools to understand and draw covalent molecules. But chemistry requires precise language as well as precise structures. Chapter 11 will teach the complete system for naming chemical compounds and writing chemical formulas from names. You will learn how to name ionic compounds — including those with transition metals and polyatomic ions — and how to name binary molecular compounds, acids, and common compounds by their classical names. You will practice converting between names and formulas in both directions. Nomenclature is the shared language of chemistry — universal, precise, and built directly on the bonding concepts you have now mastered in Chapters 9 and 10.