Energy Changes in Reactions: Exothermic or Endothermic
The Core Concepts: Energy, Bonds, and Heat
All substances contain stored energy within the chemical bonds that hold their atoms together. During a reaction, bonds in the reactants are broken, which requires an input of energy. New bonds are then formed to make the products, and this process releases energy. Whether the overall reaction feels hot or cold depends on the balance between these two steps.
If more energy is released from forming new bonds than is used to break old ones, the reaction is exothermic. The excess energy is released as heat, so $\Delta H$ is negative. Conversely, if breaking bonds requires more energy than forming new ones releases, the reaction is endothermic. The reaction draws heat from its surroundings, so $\Delta H$ is positive. This can be visualized in energy diagrams.
Visualizing Energy Changes: Reaction Profiles
A reaction profile (or energy diagram) is a graph that shows how the energy of the system changes from reactants to products. It clearly illustrates the difference between exothermic and endothermic processes.
| Feature | Exothermic Reaction | Endothermic Reaction |
|---|---|---|
| Energy Change ($\Delta H$) | $\Delta H < 0$ (Negative) | $\Delta H > 0$ (Positive) |
| Heat Flow | Releases heat to the surroundings | Absorbs heat from the surroundings |
| Product Energy vs. Reactant Energy | Products have lower energy | Products have higher energy |
| Diagram Shape | Products line is below reactants | Products line is above reactants |
| Example | Combustion of wood: $C + O_2 \rightarrow CO_2 + heat$ | Photosynthesis: $6CO_2 + 6H_2O + sunlight \rightarrow C_6H_{12}O_6 + 6O_2$ |
From Everyday Life to Industry: Real-World Applications
The principles of exothermic and endothermic reactions are not confined to the laboratory; they are active all around us, in our bodies, homes, and the technologies we use.
Exothermic Applications: These reactions are harnessed for heat and power. The combustion of fossil fuels in car engines and power plants is exothermic. Hand warmers use the rapid oxidation of iron to generate heat. Our bodies rely on exothermic reactions during cellular respiration to break down food and produce warmth and energy: $C_6H_{12}O_6 + 6O_2 \rightarrow 6CO_2 + 6H_2O + energy (ATP)$. Even the setting of concrete involves an exothermic hardening process.
Endothermic Applications: These reactions are crucial for cooling and energy storage. Instant cold packs contain a pouch of water and a solid like ammonium nitrate. When the pouch is broken, the solid dissolves in an endothermic process, absorbing heat from the surroundings and causing the pack to become cold. Similarly, baking bread involves endothermic steps where heat is absorbed to cause rising and browning. Plants use the endothermic process of photosynthesis to store solar energy as chemical energy in sugars.
Measuring and Calculating Energy Changes
Scientists can measure the energy change in a reaction using a calorimeter. This device insulates the reaction to prevent heat loss, allowing for precise measurement of temperature change. The basic calculation involves the equation:
Where: $q$ is heat energy (in Joules), $m$ is mass (in grams), $c$ is specific heat capacity (J/g°C), and $\Delta T$ is temperature change (°C).
For a reaction in solution, if the temperature increases ($\Delta T > 0$), $q$ is positive, meaning the surroundings gained heat, so the reaction is exothermic and $\Delta H$ is negative. If the temperature decreases, the opposite is true. This measured heat change, when divided by the number of moles reacted, gives the enthalpy change per mole ($\Delta H$ in kJ/mol).
Important Questions
Q1: Is ice melting an endothermic or exothermic process?
Ice melting is an endothermic process. To change from solid ice to liquid water, heat energy must be absorbed from the surroundings to break the hydrogen bonds between water molecules. This is why an ice cube cools your drink—it pulls heat energy into itself to melt.
Q2: Can a reaction be both exothermic and endothermic?
No, a single chemical reaction, as a whole, is classified as either net exothermic or net endothermic based on its overall $\Delta H$. However, every reaction involves both endothermic steps (bond breaking) and exothermic steps (bond forming). The classification depends on which of these two opposing energy changes is greater.
Q3: Why is the sign of $\Delta H$ negative for exothermic reactions?
The sign convention comes from the system's point of view. In an exothermic reaction, the system (the chemicals) loses energy to the surroundings as heat. In thermodynamics, a loss of energy from the system is given a negative sign. Since $\Delta H = H_{final} - H_{initial}$, and the final enthalpy ($H_{products}$) is lower than the initial enthalpy ($H_{reactants}$), the result is a negative number.
Conclusion
Footnote
1. Enthalpy ($H$): A thermodynamic property of a system, equivalent to the total heat content. It is the sum of internal energy plus the product of pressure and volume. In simple terms for chemistry, it represents the total energy stored in chemical bonds.
2. $\Delta H$ (Delta H): The symbol for change in enthalpy. It is the difference in heat content between the products and reactants of a chemical reaction.
3. Calorimeter: An apparatus used to measure the amount of heat involved in a chemical reaction or physical change.
4. Specific Heat Capacity ($c$): The amount of heat required to raise the temperature of 1 gram of a substance by 1 degree Celsius.
5. ATP (Adenosine Triphosphate): The primary energy-carrying molecule found in all living organisms' cells. It captures chemical energy obtained from the breakdown of food molecules and releases it to fuel other cellular processes.