How Chemical Reactions Work
The short answer
Chemical reactions fundamentally involve the rearrangement of electrons between atoms, forming new chemical bonds and breaking old ones, while the atoms themselves remain intact. For instance, burning wood combines carbon with oxygen to form carbon dioxide and ash, releasing energy.
This lesson explains the fundamental principles governing how atoms interact and transform, from their basic structure to the energy changes that drive chemical processes.
By BasicsTutor · Created with AI and fact-checked · Updated Oct 25, 2026 · How we make lessons
All matter is made of atoms, and atoms are mostly empty space.
Matter, whether solid, liquid, or gas, is composed of tiny, discrete units called atoms. Crucially, these atoms are not solid spheres but are mostly empty space, with a dense, positively charged nucleus at their center surrounded by a cloud of much lighter, negatively charged electrons. The scale is immense: if an atom were the size of a football stadium, its nucleus would be a marble at the center, and its electrons like tiny flies buzzing far away. This understanding emerged from experiments like Ernest Rutherford's gold foil experiment in 1911. When he fired positively charged alpha particles at a thin gold sheet, most passed straight through, suggesting atoms were largely empty. However, a few were deflected at large angles, indicating a tiny, dense, positive core – the nucleus. This observation directly disproved the earlier "plum pudding" model, which proposed that atoms were a diffuse, uniform blob of positive charge with electrons embedded throughout. Most people's intuition, based on our experience with macroscopic objects, is that atoms are solid, miniature billiard balls that physically "touch" when they interact. This leads to the misconception that solid objects are solid because their atoms are packed tightly and touch each other directly. Because atoms are mostly empty space, the "solidity" of matter and the way atoms interact chemically isn't about physical contact in the way we usually imagine. Instead, it's governed by the electromagnetic forces between the electrons of one atom and the nucleus of another, and between their respective electron clouds. These forces create repulsions and attractions that dictate how close atoms can get and how they bond, without ever truly "touching" in the mechanical sense. This means all physical contact, from hitting a wall to holding a cup, is essentially a result of these forces of repulsion between electron clouds.
Imagine two strong magnets pushing against each other. You can feel a definite "force" and "resistance" when you try to bring them together, even though they aren't physically touching. The forces between atoms work similarly: the electron clouds (negative poles) repel each other, while the nuclei (positive poles) attract the electrons. The overall interaction is a balance of these forces, creating a sense of solidity without direct contact.
- Atoms are the fundamental building blocks of all matter.
- Atoms are mostly empty space, not solid spheres.
- Interactions between atoms are primarily governed by electromagnetic forces, not physical contact.
Chemical reactions involve the rearrangement of electrons, not the destruction or creation of atoms.
Because atoms are mostly empty space, their interactions are dominated by their electron clouds. When a chemical reaction occurs, the atoms themselves do not change their identity or disappear. Instead, the process involves the breaking of existing chemical bonds and the formation of new ones. This bond making and breaking is fundamentally about the redistribution and sharing of valence electrons (the outermost electrons) between atoms. The nuclei of the atoms remain untouched, maintaining the element's identity. This principle is a cornerstone of modern chemistry and is encapsulated by the Law of Conservation of Mass, famously established by Antoine Lavoisier. His experiments showed that in any closed system, the total mass of the reactants before a chemical change equals the total mass of the products after the change. This can only be true if the number and types of atoms remain constant, merely changing their partners. The understanding of atomic structure, with a stable nucleus and mobile electrons, further clarifies that chemical reactions are electron-level events, distinct from nuclear reactions that do change atomic identity. A common misconception is that chemicals are "destroyed" or "created" during a reaction, or that one element can spontaneously transform into another, like lead into gold (a goal of medieval alchemy). While new substances with different properties are formed, the underlying atoms are merely reorganized; the number of carbon atoms, for example, remains the same before and after a reaction, they just end up in different molecules. Because atoms are conserved, chemists can precisely predict the amounts of reactants needed and products formed using stoichiometry. This foundational truth allows for the synthesis of new materials, the design of efficient industrial processes, and the understanding of biological pathways, all based on the predictable rearrangement of existing atoms through electron interactions. Every synthetic drug, every plastic, and every metabolic process in your body relies on this precise electron choreography.
Imagine building different structures with a set of LEGO bricks. You can break apart a 'house' (reactants) and reassemble the same bricks into a 'car' (products). The individual bricks (atoms) remain unchanged, but their arrangement and the overall structure (molecule) are entirely different. The way the bricks connect (chemical bonds) is what changes.
- Chemical reactions are about rearranging atoms, not changing their identity.
- The Law of Conservation of Mass underpins all chemical changes.
- Electron movement and sharing are central to bond breaking and forming.
Chemical reactions occur when atoms or molecules collide with sufficient energy and correct orientation.
For chemical bonds to break and new ones to form, the reactant atoms or molecules must physically collide with each other. However, not all collisions are effective; for a reaction to proceed, the colliding particles must possess a minimum amount of energy, known as the activation energy, and they must collide in a specific orientation that allows the relevant parts of the molecules to interact. This is because bonds are specific: a carbon atom needs to meet an oxygen atom at the right 'angle' and with enough force to allow electrons to rearrange. This principle is explained by Collision Theory. At the molecular level, particles are constantly in motion, colliding with each other due to their kinetic energy. The kinetic energy of these particles increases with temperature. If the collision energy is below the activation energy barrier, the particles simply bounce off each other without reacting. The activation energy represents the energy required to distort bonds or overcome electron cloud repulsion to reach a transitional state where rearrangement can occur. The correct orientation ensures that the atoms involved in the bond changes are positioned to do so, like fitting a key into a lock. Many people assume that if two reactive substances are mixed, they will immediately and completely react. This overlooks the necessity of collisions with specific energy and alignment. For instance, natural gas and oxygen are both highly reactive but don't react at room temperature unless given a spark (which provides the activation energy), or if heated to their ignition temperature. Without this initial energy input, collisions are too weak to overcome the existing bonds. This principle explains why factors like temperature, concentration, and surface area affect reaction rates. Increasing temperature provides more particles with sufficient kinetic energy to overcome the activation barrier. Increasing concentration or surface area increases the frequency of collisions. Catalysts work by providing an alternative reaction pathway with a lower activation energy, allowing more collisions to be effective at a given temperature and speeding up the reaction without being consumed themselves, which is crucial in industrial processes and biological systems.
Imagine trying to open a specific lock with a key. You need to insert the key (collision) into the keyhole (correct orientation), and then you need to turn it with enough force (sufficient energy) to unlock it. Just bumping the key against the lock, or inserting it backwards, won't work, even if you have the right key.
- Collisions between reactant particles are necessary for a reaction.
- Collisions must have enough energy (activation energy) to be effective.
- Reacting particles must collide in the correct spatial orientation.
Energy is always conserved, but chemical reactions involve a redistribution of energy, often as heat or light.
Because atoms are rearranged rather than created or destroyed, and bonds represent stored potential energy, the total energy of a closed system remains constant during a chemical reaction. Energy is neither created nor destroyed, but it transforms from one form to another. In chemical reactions, this often means chemical potential energy (stored in the electrons within bonds) is converted into kinetic energy (heat), electromagnetic energy (light), or work, or vice-versa. Reactions that release energy to the surroundings are called exothermic, while those that absorb energy are endothermic. This is a direct application of the First Law of Thermodynamics, also known as the Law of Conservation of Energy. Chemical bonds themselves represent stored potential energy. Forming stable bonds generally releases energy (because the products are at a lower potential energy state), while breaking bonds requires an input of energy. When a reaction proceeds, the net difference between the energy required to break old bonds and the energy released by forming new ones determines whether the reaction overall releases or absorbs energy. A common misunderstanding is that exothermic reactions "create" heat or light. Instead, they convert stored chemical potential energy into these observable forms, similar to how a stretched rubber band releases stored energy as kinetic energy when let go. Similarly, endothermic reactions don't "destroy" cold; they absorb energy from their surroundings (often as heat), making the surroundings feel colder as that energy is used to form higher-energy bonds. Understanding energy changes is crucial for controlling and utilizing chemical reactions. We use exothermic reactions to generate power (combustion of fuels in engines, power plants), provide warmth (heating pads), or initiate processes (explosives). Endothermic reactions are used for cooling (instant cold packs) or to store energy for later use (photosynthesis stores solar energy as chemical bonds in glucose). The energy balance dictates whether a reaction will be a source or sink of energy for its environment.
Think of a loaded spring. Compressing it stores potential energy. When released, this potential energy is converted into kinetic energy (motion) and often sound (vibration). The total energy in the system (spring + environment) remains constant, just changing forms. Chemical bonds are like these springs, storing potential energy that can be released or absorbed during reactions.
- Energy is conserved in chemical reactions, only changing forms.
- Exothermic reactions release chemical potential energy (e.g., as heat or light).
- Endothermic reactions absorb energy from their surroundings.
Reactions tend towards states of lower potential energy and higher entropy.
The spontaneity of a chemical reaction, meaning whether it will happen on its own without continuous external energy input, is not solely determined by whether it releases energy (exothermic). Instead, reactions tend to proceed in directions that lead to a decrease in the system's Gibbs free energy (ΔG). Gibbs free energy combines two fundamental driving forces: the tendency to move to a lower potential energy state (enthalpy, ΔH) and the tendency to move to a more disordered or dispersed state (entropy, ΔS). This principle is a consequence of the Second Law of Thermodynamics, which states that the total entropy (disorder or energy dispersal) of an isolated system can only increase over time, or remain constant in ideal cases. For any spontaneous process, the universe's entropy must increase. Chemical systems achieve this by either releasing energy to increase the entropy of the surroundings (exothermic reactions that become more disordered, like combustion) or by increasing their own internal disorder (entropy) so significantly that it outweighs any energy absorption (like a solid dissolving). The Gibbs free energy equation (ΔG = ΔH - TΔS, where T is the absolute temperature) mathematically expresses this balance. A common misconception is that all spontaneous reactions must be exothermic, or that reactions "want" to reach a state of perfect "stability" (implying low energy and high order). This ignores the crucial role of entropy. For example, ice melting at room temperature is an endothermic process (it absorbs heat from the surroundings) but is spontaneous because the increase in disorder (liquid water has higher entropy than solid ice) is significant enough to drive the process. This is a genuinely counterintuitive insight: the universe fundamentally prefers disorder. Many everyday processes, from the corrosion of metals to the dispersal of a fragrance, are driven by this universal tendency towards greater entropy. Understanding this balance between energy and disorder allows chemists to predict reaction spontaneity, design reactions that would otherwise seem impossible (by coupling them with highly favorable ones), and comprehend why seemingly "unfavorable" processes (like the biological synthesis of complex molecules) can occur by being linked to highly favorable ones (like ATP hydrolysis). It reframes chemical change not just as an energy exchange, but as a fundamental move towards the most probable, dispersed arrangements of matter and energy.
Imagine a carefully stacked deck of cards versus a shuffled deck. The stacked deck is low in entropy (ordered) but high in potential to become disordered. Given a small nudge, it will spontaneously fall and spread out, increasing its entropy. The shuffled deck is high in entropy (disordered) and requires significant effort to restore order. Similarly, chemical systems tend to spontaneously move from ordered, higher free energy states to more disordered, lower overall free energy states.
- Spontaneity is determined by Gibbs free energy, a balance of energy (enthalpy) and disorder (entropy).
- Reactions tend towards states of lower potential energy and higher entropy (more disorder).
- The universe fundamentally moves towards increased overall entropy for spontaneous processes.
Put it into practice
- Experiment with Cooking Temperatures: Next time you cook, observe how temperature changes reaction rates. For instance, notice the difference in browning (Maillard reaction) when sautéing vegetables at low vs. high heat. This demonstrates the impact of kinetic energy on collision effectiveness and activation energy.
- Read Product Labels for Active Ingredients: When buying cleaning supplies, medications, or even some foods, identify the 'active ingredients.' Understand that these chemicals work by specific reactions (e.g., acids dissolving scale, oxidizers breaking down stains, enzymes breaking down starches) rather than simply 'cleaning' or 'healing.'
- Observe Everyday Corrosion: Pay attention to how metals corrode (e.g., rust on iron tools, tarnish on silver jewelry). This is a slow, spontaneous chemical reaction driven by the tendency towards lower potential energy and higher entropy, often involving oxidation. Consider how protective coatings (like paint or galvanization) prevent these reactions by creating barriers that inhibit necessary collisions.
- Explore Fermentation or Composting: If you have a garden, start a compost pile, or try a simple fermentation project (like making sauerkraut or sourdough starter). These biological processes are complex chains of chemical reactions, often driven by microorganisms, that convert organic matter into simpler, more stable (and higher entropy) forms, driven by the principles of spontaneity.