Simple Structure: Small Molecules Held by Covalent Bonds
The Atomic Quest for Stability
Everything you see, touch, and are made of is composed of atoms. But atoms rarely like to be alone. They are constantly seeking to become more stable, much like a chair is more stable with four legs than with three. For many atoms, this stability comes from having a full outer shell of electrons, a state often called a "noble gas configuration."[1]
Atoms can achieve this stable state in different ways. One way is to completely transfer electrons from one atom to another, forming ions that are held together by strong electrical attractions, called ionic bonds. The other primary way, which is the focus of our topic, is for atoms to share electrons. This sharing creates a powerful link called a covalent bond. When two or more atoms connect using covalent bonds, they form a molecule.
How Covalent Bonds Form: A Shared Partnership
Imagine two hydrogen atoms, each with a single electron. Individually, they are unstable. But if they move close enough, their electron orbitals can overlap. This overlap allows the two atoms to share their two electrons. Now, from the perspective of each hydrogen atom, it has access to two electrons, which fills its outer shell, making it as stable as the noble gas helium. This shared pair of electrons is the covalent bond that holds the $H_2$ molecule together.
Atoms can share more than one pair of electrons. A single bond involves one shared pair ($2$ electrons). A double bond involves two shared pairs ($4$ electrons), and a triple bond involves three shared pairs ($6$ electrons). For example, an oxygen molecule ($O_2$) is held together by a double bond, while a nitrogen molecule ($N_2$) is held together by a very strong triple bond.
| Bond Type | Shared Electron Pairs | Notation | Example Molecule |
|---|---|---|---|
| Single Bond | 1 | — or $:$ | $H_2$, $H_2O$, $CH_4$ |
| Double Bond | 2 | = or $::$ | $O_2$, $CO_2$ |
| Triple Bond | 3 | ≡ or $:::$ | $N_2$ |
Representing Small Covalent Molecules
Scientists use different models to represent molecules, each with its own advantages. Understanding these representations helps us visualize the otherwise invisible world of atoms and bonds.
Molecular Formula: This is the simplest representation. It shows the type and number of atoms in a molecule. For example, $H_2O$ tells us a water molecule has 2 hydrogen atoms and 1 oxygen atom. However, it doesn't show how they are connected.
Lewis Dot Structure: This model shows the arrangement of valence electrons[2] around the atoms and how they are shared to form bonds. Dots represent electrons, and lines represent bonds. For water ($H_2O$), the Lewis structure clearly shows the two O—H single bonds and the two lone pairs[3] of electrons on the oxygen atom.
Structural Formula: This is a simplified drawing that shows how atoms are connected. It uses lines to represent covalent bonds. A single line is a single bond, a double line is a double bond, etc. The structural formula for methane ($CH_4$) shows a central carbon atom with four single bonds, each connected to a hydrogen atom.
Properties of Simple Covalent Substances
The small size and specific bonding of these molecules give them a unique set of physical properties.
Low Melting and Boiling Points: The covalent bonds within a molecule are very strong. However, the forces between individual molecules, called intermolecular forces, are relatively weak. It doesn't take much energy to overcome these weak forces and separate the molecules from each other, which is why many small covalent molecules are gases or liquids at room temperature. For instance, oxygen ($O_2$) is a gas, and water ($H_2O$) is a liquid.
Poor Electrical Conductivity: Since they do not contain free electrons or charged ions, substances made of small covalent molecules do not conduct electricity, whether they are in the solid, liquid, or gaseous state. Pure water is a very poor conductor.
Solubility: The solubility of a covalent molecule depends on its polarity.[4] A general rule is "like dissolves like." Polar covalent molecules, like sugar, tend to dissolve in polar solvents like water. Non-polar covalent molecules, like oil or oxygen gas, do not dissolve well in water but may dissolve in non-polar solvents.
Molecules in Action: From Air to Soda
Let's look at some concrete examples of small covalent molecules that play crucial roles in our everyday lives.
Oxygen Gas ($O_2$): The air we breathe is about 21% oxygen gas. Each oxygen molecule consists of two oxygen atoms held together by a strong double bond. Our bodies rely on these molecules for cellular respiration, the process that releases energy from food.
Water ($H_2O$): This is perhaps the most vital covalent molecule on Earth. Its simple structure of two hydrogen atoms bonded to one oxygen atom creates a "bent" shape that gives water its unique properties, such as high surface tension and the fact that ice is less dense than liquid water. Without these properties, life as we know it would not exist.
Carbon Dioxide ($CO_2$): A linear molecule with a central carbon atom double-bonded to two oxygen atoms ($O=C=O$). Plants use $CO_2$ and water during photosynthesis to create food. It is also the gas that gives carbonated drinks their fizz. In its solid form, known as dry ice, it sublimes directly from a solid to a gas.
Glucose ($C_6H_{12}O_6$): This is a larger but still discrete covalent molecule that serves as a primary source of energy for living organisms. The sugar you put in your tea or that fuels your brain is made of glucose molecules, each held together by a network of covalent bonds between carbon, hydrogen, and oxygen atoms.
| Molecule | Formula | State at Room Temp | Role in Daily Life |
|---|---|---|---|
| Oxygen | $O_2$ | Gas | Respiration, combustion |
| Water | $H_2O$ | Liquid | Solvent, essential for life |
| Carbon Dioxide | $CO_2$ | Gas | Photosynthesis, carbonation |
| Ammonia | $NH_3$ | Gas | Fertilizers, cleaning products |
| Methane | $CH_4$ | Gas | Natural gas, fuel |
Common Mistakes and Important Questions
Are covalent bonds strong or weak?
This is a common point of confusion. The covalent bonds inside a molecule are very strong. It takes a lot of energy to break the bond between two atoms in an $O_2$ or $H_2O$ molecule. However, the forces between different molecules (intermolecular forces) are weak. This is why it's easy to boil water (separate the $H_2O$ molecules from each other) but very difficult to split water into hydrogen and oxygen gas (breaking the O—H covalent bonds inside the molecules).
Why don't covalent compounds conduct electricity?
Electrical conductivity requires the movement of charged particles, such as free electrons (in metals) or free-moving ions (in ionic compounds that are dissolved or melted). Small covalent molecules are electrically neutral and do not have charged particles that are free to move. Therefore, they do not conduct electricity under any normal conditions.
What is the difference between a molecule and a compound?
All compounds are molecules, but not all molecules are compounds. A molecule is any structure held together by covalent bonds; it can be made of one type of atom (like $O_2$ or $N_2$) or different types of atoms. A compound is a substance made of two or more different types of atoms chemically bonded together. So, $O_2$ is a molecule but not a compound, while $H_2O$ is both a molecule and a compound.
Conclusion
The world of small covalent molecules is a fascinating and essential one. From the simple sharing of electrons between atoms emerges a vast array of substances with defined properties. Understanding the nature of the covalent bond—strong, directional, and responsible for holding atoms together within a molecule—allows us to explain why water is a liquid, why oxygen is a gas, and why sugar dissolves in tea. These simple structures are the foundational units of more complex biological and synthetic materials, making them a cornerstone of chemistry and a key to understanding the world at a molecular level.
Footnote
[1] Noble Gas Configuration: The electron configuration of noble gases (like Helium, Neon, Argon), which have a full outer shell of electrons. This configuration is very stable, and other atoms tend to gain, lose, or share electrons to achieve a similar arrangement.
[2] Valence Electrons: The electrons in the outermost shell, or energy level, of an atom. These are the electrons involved in forming chemical bonds.
[3] Lone Pair: A pair of valence electrons that is not shared between atoms and is therefore not involved in covalent bonding. Also known as a non-bonding pair.
[4] Polarity: A property of a molecule that has a slightly positive end and a slightly negative end due to unequal sharing of electrons in its covalent bonds. Water is a polar molecule.