Endothermic and Exothermic Reactions - Chemistry

1. Fundamental Concepts

Energy Change Definition：Chemical reactions are classified by the net energy exchange with the surroundings. The total energy of reactants and products follows the law of conservation of energy.
Endothermic Reactions：Reactions that absorb thermal energy from the surroundings. The system gains energy, and the temperature of the external environment usually decreases.
Exothermic Reactions：Reactions that release thermal energy to the surroundings. The system loses energy, and the temperature of the external environment usually increases.
Enthalpy Change(ΔH)：A core thermodynamic quantity in US high school chemistry to quantify reaction energy.
- Endothermic: (positive)
- Exothermic: (negative)

Bond Energy Basis：Breaking chemical bonds requires absorbing energy; forming chemical bonds releases energy. The sign of ΔH is determined by the difference between bond-breaking energy and bond-forming energy.
- Endothermic: Energy absorbed for bond breaking > Energy released for bond formation.
- Exothermic: Energy released for bond formation > Energy absorbed for bond breaking.
Energy Profile Diagram：A standard analytical tool in US chemistry exams. The vertical axis represents enthalpy (H), horizontal axis represents reaction progress.
- Endothermic: Enthalpy of products > Enthalpy of reactants.
- Exothermic: Enthalpy of products < Enthalpy of reactants.
Common Misconception：The spontaneity of a reaction is not determined solely by . Spontaneity depends on Gibbs free energy (), which is covered in advanced high school and college chemistry.
Phase Changes：Phase transitions are physical processes classified by energy change. Melting, evaporation, sublimation are endothermic; freezing, condensation, deposition are exothermic.

Easy

1. Exothermic：Combustion of methane

$$\ce{CH4(g) + 2O2(g) -> CO2(g) + 2H2O(l)}\quad \Delta H < 0$$

2. Endothermic：Melting of ice into liquid water

$$\ce{H2O(s) -> H2O(l)}\quad \Delta H > 0$$

Medium
1. Exothermic：Neutralization reaction between strong acid and strong base

$$\ce{HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l)}\quad \Delta H < 0$$

2. Endothermic：Decomposition of calcium carbonate (heating required)

$\ce{CaCO3(s) ->[\Delta] CaO(s) + CO2(g)}$

Hard
1. Exothermic：Formation reaction of ammonia (Haber process, industrial application)

$$\ce{N2(g) + 3H2(g) <=> 2NH3(g)}\quad \Delta H < 0$$

2. Endothermic：Dissolution of ammonium nitrate in water (common in cold packs)

$$\ce{NH4NO3(s) -> NH4+(aq) + NO3-(aq)}\quad \Delta H > 0$$

Identify Reaction Type by ：Directly judge using the sign of enthalpy change given in the problem. Positive is endothermic, negative is exothermic.
Judge by Phenomenon：Use experimental phenomena for judgment. Temperature rise of the system indicates exothermic; temperature drop indicates endothermic. Continuous heating required for reaction usually suggests endothermic.
Bond Energy Calculation：Apply the formula (Bond energy of reactant bonds broken) (Bond energy of product bonds formed). Substitute the bond energy values provided in the question to calculate and determine the reaction type.
Energy Profile Analysis：Locate the enthalpy values of reactants and products on the energy diagram. Compare their magnitudes to confirm if the reaction is endothermic or exothermic, and identify activation energy and enthalpy change.
Phase Change & Practical Application Judgment：Classify common phase changes and daily chemical applications (cold packs, hot packs, combustion, etc.) by memorizing typical examples to quickly answer multiple-choice and short-answer questions.