Covalent Bonding

Introduction

Covalent bonding is a way in which atoms become connected by sharing one or more pairs of electrons. It is especially common between nonmetal atoms. The shared electrons are attracted to the nuclei of both bonded atoms, which lowers the energy of the system and helps hold the atoms together.

This course is designed for Grades 9–10. You will learn how covalent bonds form, how to represent them with Lewis structures and dot-and-cross diagrams, how single, double, and triple bonds differ, and how bonding helps explain the properties of molecular substances and giant covalent structures.

By the end of the course, you should be able to explain covalent bonding using valence electrons, draw simple molecular structures, distinguish polar and nonpolar bonds, connect structure to physical properties, and apply the ideas to substances such as water, carbon dioxide, diamond, and graphite.

Foundations: Valence Electrons and Stability

Valence Electrons

Valence electrons are the electrons in the outermost occupied electron shell of an atom. They are the electrons most directly involved in chemical bonding. For many main-group elements, the group position in the periodic table helps you predict the number of valence electrons.

A useful school-level model is the octet rule: many atoms tend to form bonds that give them access to eight electrons in their valence shell. Hydrogen is an important exception because its first shell is full with two electrons. The octet rule is a helpful pattern rather than an absolute law; some molecules do not follow it exactly.

Why Sharing Can Hold Atoms Together

In a covalent bond, a shared pair of electrons occupies a region between two nuclei. Each negatively charged electron is attracted to both positively charged nuclei. At the same time, nucleus-nucleus and electron-electron repulsions still exist. A stable bond length occurs where the balance of attractions and repulsions gives a lower-energy arrangement than the separated atoms.

Forming a stable covalent bond usually releases energy, while breaking that bond requires energy. A stronger bond generally requires more energy to break. This idea becomes useful when you compare single, double, and triple bonds and when you study chemical reactions.

OpenStax Chemistry 2e: Covalent Bonding provides an open textbook explanation of electron sharing, bond formation, and bond polarity.

How Covalent Bonds Form

Shared Electron Pairs

A covalent bond is often represented by a line between two atoms. One line stands for one shared pair of electrons. In a dot diagram, the same bond can be shown as two electron dots between the atomic symbols.

Hydrogen provides the simplest example. Each hydrogen atom has one valence electron. When two hydrogen atoms share their electrons, each hydrogen can be considered to have access to two electrons, filling its first shell.

Chlorine provides another useful example. Each chlorine atom has seven valence electrons. Two chlorine atoms can share one pair, giving each atom an octet in the simple model.

Single, Double, and Triple Bonds

A single bond contains one shared electron pair. A double bond contains two shared electron pairs. A triple bond contains three shared electron pairs.

Examples include H–H as a single bond, O=O as a double bond, and N≡N as a triple bond. For the same pair of atoms, higher bond order is usually associated with a shorter and stronger bond, although exact bond strengths depend on the atoms and molecular environment.

The nitrogen molecule is especially stable because its two nitrogen atoms are connected by a triple bond. The large amount of energy needed to break this bond is one reason atmospheric nitrogen is relatively unreactive under ordinary conditions.

Representing Covalent Molecules

Lewis Structures

A Lewis structure uses element symbols, lines, and dots to represent valence electrons. A line between two atoms represents a bonding pair. Pairs of valence electrons not used in bonding are called lone pairs.

A practical method for many simple molecules is to count the available valence electrons, choose a reasonable arrangement of atoms, place shared pairs between bonded atoms, and then distribute the remaining electrons so that outer shells are complete where the model applies.

For methane, CH4, carbon forms four single covalent bonds to four hydrogen atoms. Each hydrogen reaches a duet, and carbon is surrounded by eight shared electrons.

Dot-and-Cross Diagrams

A dot-and-cross diagram uses different symbols for electrons that originally came from different atoms. The symbols help you track the contribution from each atom, but once the bond forms the electrons in the shared pair are not permanently assigned to one atom or the other.

When you draw a dot-and-cross diagram, check three things: the total number of valence electrons must be correct, shared pairs must lie between bonded atoms, and lone pairs must be shown where needed.

A Worked Example: Water

A water molecule, H2O, contains two O–H single bonds. Oxygen has six valence electrons, and each hydrogen has one. Oxygen shares one electron pair with each hydrogen. Two lone pairs remain on the oxygen atom.

The electron pairs around oxygen do not all repel one another equally. The molecule is bent rather than linear. At this level, you can connect this shape to the presence of two bonding pairs and two lone pairs around oxygen.

Electronegativity and Bond Polarity

Unequal Sharing

Electronegativity describes how strongly an atom attracts shared electrons in a bond. When two identical atoms bond, such as H2 or Cl2, their electronegativities are equal and the electrons are shared equally. This is a nonpolar covalent bond.

When atoms with different electronegativities form a covalent bond, the electrons may be shared unequally. The more electronegative atom attracts more electron density and develops a partial negative charge, while the other atom develops a partial positive charge. This is a polar covalent bond.

In water, oxygen attracts the bonding electrons more strongly than hydrogen. Each O–H bond is therefore polar.

Bond Polarity and Molecular Polarity Are Not the Same

A molecule can contain polar bonds without being polar overall. Molecular shape matters because bond dipoles have directions.

Carbon dioxide, CO2, contains two polar C=O bonds. The molecule is linear and symmetrical, so the two bond dipoles point in opposite directions and cancel. Water is bent, so its O–H bond dipoles do not cancel. Water therefore has a net molecular dipole.

This distinction is an important example of scientific reasoning: knowing the parts is not enough; you must also consider how the parts are arranged in space.

OpenStax Chemistry: Atoms First gives an open-access explanation of electronegativity and polar covalent bonding.

Molecular Substances and Their Properties

Strong Bonds Inside Molecules, Weaker Attractions Between Molecules

A common misconception is that molecular substances have low melting or boiling points because their covalent bonds are weak. The covalent bonds inside each molecule can be strong. During melting or boiling of a simple molecular substance, the main change is usually the separation of molecules from one another, not the breaking of the covalent bonds inside each molecule.

The attractions between separate molecules are called intermolecular forces. In many small molecular substances these forces are much weaker than covalent bonds. That is why many simple molecular substances have lower melting and boiling points than giant covalent materials.

Boiling water, for example, separates water molecules so they can move freely as a gas. It does not normally break the O–H covalent bonds inside each water molecule.

Electrical Conductivity

Most simple molecular substances do not conduct electricity well because they do not contain freely moving ions or a large supply of delocalized electrons. This is a broad trend, not a rule that should be applied without checking the structure of a particular substance.

Giant Covalent Structures

Diamond

Some substances do not consist of separate small molecules. In a giant covalent structure, also called a network covalent structure, very large numbers of atoms are joined by covalent bonds in a continuous network.

Diamond is an allotrope of carbon. Each carbon atom is covalently bonded to four other carbon atoms in a three-dimensional tetrahedral network.

Because changing diamond into a fluid or gas requires disrupting many strong covalent bonds in the network, diamond remains solid to extremely high temperatures; at ordinary pressure it tends to sublime rather than melt. Its rigid three-dimensional bonding also helps explain its extreme hardness. Diamond does not conduct electricity well because its valence electrons are localized in bonds and are not free to move through the structure.

Graphite

Graphite is another allotrope of carbon, but its structure is different. Each carbon atom is covalently bonded to three others in flat sheets of hexagonal rings. The sheets are held together much more weakly than the atoms within a sheet.

The weak attraction between layers allows them to slide over one another, helping explain why graphite is soft and useful in pencils and as a lubricant. Each carbon also contributes to delocalized electrons that can move along the layers, so graphite can conduct electricity.

Diamond and graphite show why structure determines properties. Both are made only of carbon atoms, but different bonding arrangements produce very different hardness, electrical conductivity, and uses.

Covalent Bonding in Everyday Materials

Covalent bonding is central to many materials around you. Water contains polar covalent bonds. Oxygen and nitrogen in air are covalent molecules. Carbon dioxide contains double bonds. Fuels and many plastics contain covalently bonded carbon frameworks. Proteins, carbohydrates, fats, and DNA depend on extensive networks of covalent bonds.

Understanding bonding helps you move from a chemical formula to an explanation. Instead of memorizing that a substance has a particular property, you can ask: What particles are present? How are the atoms bonded? Are there separate molecules or a continuous network? Are there mobile charged particles or delocalized electrons? What forces must be overcome during a physical change?

Common Misconceptions

Misconception 1: Covalent bonding means electrons are transferred. In covalent bonding, electrons are shared. Electron transfer is the basic school-level model for ionic bonding.

Misconception 2: A shared pair always lies exactly halfway between the atoms. In a polar covalent bond, electron density is pulled toward the more electronegative atom.

Misconception 3: All covalent substances have low melting points. Many simple molecular substances do, but giant covalent structures such as diamond have very high melting points.

Misconception 4: Boiling breaks covalent bonds. Boiling a simple molecular substance generally overcomes intermolecular attractions rather than breaking bonds inside each molecule.

Misconception 5: Polar bonds always make a polar molecule. Molecular shape can cause bond dipoles to cancel, as in carbon dioxide.

Interactive Tasks

Quiz: Test Your Knowledge

What happens to electrons in a covalent bond? (They are shared between atoms) (!They are transferred completely to one atom) (!They disappear from the valence shell) (!They are converted into protons)

Which pair of elements is most likely to form a covalent bond in the simple school-level model? (Two nonmetals) (!Two metals) (!A metal and a nonmetal only) (!Two noble gases only)

How many electron pairs are shared in a single covalent bond? (One pair) (!Two pairs) (!Three pairs) (!Four pairs)

Which molecule contains a double covalent bond between identical atoms? (Oxygen) (!Hydrogen) (!Chlorine) (!Methane)

Which molecule contains a triple covalent bond? (Nitrogen) (!Water) (!Methane) (!Hydrogen chloride)

What does electronegativity describe? (An atom's attraction for bonding electrons) (!The number of neutrons in an atom) (!The mass of one mole of a substance) (!The speed of electrons in a wire)

What makes a covalent bond polar? (Unequal sharing of bonding electrons) (!Complete transfer of all valence electrons) (!Equal sharing between identical atoms) (!The absence of any electron pair)

Why do many simple molecular substances have relatively low boiling points? (Intermolecular attractions are relatively weak) (!Covalent bonds inside molecules are always weak) (!Their atoms contain no electrons) (!Their molecules are held together by metallic bonds)

Why does diamond remain solid to extremely high temperatures? (Many strong covalent bonds form a giant network) (!Diamond contains freely moving ions) (!Its molecules are held by weak forces only) (!Each carbon atom is bonded to only one neighbor)

Why can graphite conduct electricity? (It has delocalized electrons that can move along layers) (!It contains liquid metal between the layers) (!Its carbon atoms become positive ions) (!It has no covalent bonds)

Memory Game

Covalent bond	Shared pair of electrons attracted by two nuclei
Lone pair	Pair of valence electrons not used in bonding
Electronegativity	Tendency of an atom to attract bonding electrons
Double bond	Bond made from two shared electron pairs
Diamond	Three-dimensional network in which each carbon bonds to four others
Graphite	Layered carbon structure with mobile delocalized electrons

Drag and Drop

Match the correct terms.	Topic
Single bond	One shared electron pair
Triple bond	Three shared electron pairs
Polar covalent bond	Unequal sharing of bonding electrons
Simple molecular substance	Separate molecules with intermolecular attractions
Giant covalent structure	Continuous network of covalently bonded atoms

...

Crossword Puzzle

Valence	Which type of electron is found in an atom's outer occupied shell and is important in bonding?
Molecule	What neutral group of covalently bonded atoms can exist as a separate particle?
Polarity	What property describes an uneven distribution of electrical charge in a bond or molecule?
Diamond	Which carbon allotrope has a rigid three-dimensional covalent network?
Graphite	Which carbon allotrope has layers and can conduct electricity?
Electronegativity	What term means an atom's tendency to attract bonding electrons?

LearningApps

Cloze Text

Open-Ended Tasks

Easy

Lewis structure: Draw Lewis structures for H2, Cl2, H2O, and CH4, then circle every shared electron pair.
Molecular model: Build simple models of H2O, CO2, and CH4 using safe classroom materials, photograph them, and label the bonds and approximate shapes.
Covalent bond: Create a one-page illustrated explanation for a younger learner showing the difference between sharing electrons and transferring electrons.
Water molecule: Make a short narrated slide or poster explaining why water contains polar covalent bonds and identify its lone pairs.

Standard

Bond polarity: Use an electronegativity table to compare H-H, H-Cl, O-H, and C-O bonds, then explain which are more polar and why.
Carbon dioxide: Produce a diagram or short video showing why CO2 can contain polar bonds while the molecule has no overall dipole.
Intermolecular force: Design a demonstration or analogy that distinguishes covalent bonds inside molecules from attractions between separate molecules, and explain the limits of your analogy.
Allotrope: Compare diamond and graphite in a two-column evidence chart linking carbon bonding arrangement to hardness, conductivity, and uses.

Advanced

Bond energy: Investigate reliable bond-energy data for selected single, double, and triple bonds and write a short analysis of the relationship between bond order, strength, and bond length.
Molecular geometry: Create 3D models of several small molecules and use electron-pair repulsion ideas to explain why their shapes differ.
Materials science: Choose a covalent material such as silicon dioxide, a polymer, graphene, or silicon carbide and produce a research poster connecting microscopic bonding to macroscopic properties.
Science communication: Interview a chemistry teacher, laboratory technician, engineer, pharmacist, or materials scientist about where covalent bonding matters in their work, then produce an edited article or video that checks the claims against trustworthy sources.

Learning Assessment

Structure and properties: Given unfamiliar descriptions of two substances, decide which is more likely to be simple molecular or giant covalent and justify your choice using melting point and conductivity evidence.
Lewis structure: Draw and check Lewis structures for several molecules, then explain how your electron count confirms that no valence electrons have been lost from the diagram.
Bond polarity: Compare two covalent bonds using electronegativity data and explain how the difference affects the distribution of electron density.
Molecular polarity: Use both bond polarity and molecular shape to predict whether H2O and CO2 have net dipoles, and defend your reasoning.
Carbon allotrope: Explain how the bonding structures of diamond and graphite lead to different hardness and electrical conductivity.
Chemical change: Analyze a heating or boiling scenario and distinguish whether the process mainly overcomes intermolecular attractions or breaks covalent bonds.

Evidence of Learning

Knowledge: You can define covalent bonding, identify valence electrons, distinguish bonding pairs from lone pairs, and explain single, double, and triple bonds.

Skills: You can draw and check simple Lewis structures, interpret dot-and-cross diagrams, use electronegativity to reason about bond polarity, and use molecular shape to discuss overall polarity.

Products: Strong evidence may include accurate molecular models, annotated diagrams, a comparison of diamond and graphite, a short investigation, or a science communication product that uses trustworthy sources.

Reasoning: You can connect microscopic structure to macroscopic properties instead of relying only on memorized facts.

Transfer: You can apply covalent-bonding ideas to unfamiliar molecules and materials, recognize the limits of simple models such as the octet rule, and ask what evidence would test your explanation.

OERs on the Topic

The English Wikipedia article below gives a broad overview of covalent bonding and links to related concepts.

You can also review OpenStax Chemistry 2e: Covalent Bonding, an openly licensed textbook section that explains shared electron pairs, bond formation, electronegativity, and polarity.