Exothermic & Endothermic Reactions, Reaction Profiles & Bond Energy Calculations
Energy is conserved in all chemical reactions: an exothermic reaction transfers thermal energy to the surroundings so the temperature of the surroundings increases, whereas an endothermic reaction takes in thermal energy from the surroundings so the temperature decreases. At the molecular level, breaking chemical bonds in reactants is always an endothermic process (requires energy), and making new chemical bonds in products is always an exothermic process (releases energy). Examiners heavily test reaction profile diagrams, the polystyrene cup calorimetry required practical, and Higher Tier overall energy change calculations using bond energies.
- Exothermic Reactions
- Transfer energy from the reacting chemicals to the surroundings, causing the temperature of the surroundings (measured by a thermometer in the solution) to increase. Products have less chemical energy than the reactants, so the overall energy change (ΔH) is negative. Examples: combustion, oxidation, and acid-alkali neutralisation.
- Endothermic Reactions
- Take in energy from the surroundings into the reacting chemicals, causing the temperature of the surroundings to decrease. Products have more chemical energy than the reactants, so the overall energy change (ΔH) is positive. Examples: thermal decomposition, citric acid + sodium hydrogencarbonate, and sports injury cold packs.
- Reaction Profiles & Activation Energy (Ea)
- Diagrams showing the relative energies of reactants and products over the progress of a reaction. Activation energy (Ea) is the minimum amount of energy that colliding particles must have in order to react (the arrow from the reactant energy improve to the peak of the curve). The overall energy change is the vertical arrow from the reactant level to the product level.
Exothermic & Endothermic Reactions, Reaction Profiles & Bond Energy Calculations: Reaction Profile (ΔH & E_a) + Le Chatelier Yield Simulator
Toggle a catalyst on the enthalpy profile diagram or adjust temperature and pressure to see how dynamic equilibrium shifts.
Modelled on N₂(g) + 3H₂(g) ⇌ 2NH₃(g) (ΔH = −92 kJ/mol) (4 moles gas → 2 moles gas):
Higher pressure shifts equilibrium to the side with fewer gas moles (right, 4 → 2). Lower temperature favours the exothermic forward reaction.
Key Concepts & Exam Definitions (6)
- Exothermic Reactions
- Transfer energy from the reacting chemicals to the surroundings, causing the temperature of the surroundings (measured by a thermometer in the solution) to increase. Products have less chemical energy than the reactants, so the overall energy change (ΔH) is negative. Examples: combustion, oxidation, and acid-alkali neutralisation.
- Endothermic Reactions
- Take in energy from the surroundings into the reacting chemicals, causing the temperature of the surroundings to decrease. Products have more chemical energy than the reactants, so the overall energy change (ΔH) is positive. Examples: thermal decomposition, citric acid + sodium hydrogencarbonate, and sports injury cold packs.
- Reaction Profiles & Activation Energy (Ea)
- Diagrams showing the relative energies of reactants and products over the progress of a reaction. Activation energy (Ea) is the minimum amount of energy that colliding particles must have in order to react (the arrow from the reactant energy improve to the peak of the curve). The overall energy change is the vertical arrow from the reactant level to the product level.
- Bond Breaking vs Bond Making (Higher Tier)
- During a chemical reaction, energy must be supplied to break existing covalent bonds in the reactants (endothermic), and energy is released when new bonds in the products are formed (exothermic). In an exothermic reaction, the energy released by forming new bonds is greater than the energy needed to break existing bonds.
- Bond energy calculation formula (higher tier)
- Overall Energy Change = ∑(Bond Energies Broken in Reactants) - ∑(Bond Energies Made in Products). Count every individual covalent bond carefully by drawing out structural formulas first (e.g., CH₄ has 4 C-H bonds; 2H₂O has 4 O-H bonds; O₂ has 1 O=O double bond).
- Chemical Cells & Hydrogen Fuel Cells (Triple Science)
- Chemical cells produce a voltage from two metals of different reactivity dipped in an electrolyte; the bigger the difference in reactivity between the two metals, the greater the voltage produced. Non-rechargeable (alkaline) cells stop working when one reactant is used up; rechargeable cells use reversible reactions. Hydrogen fuel cells oxidise hydrogen electrochemically (2H₂ + O₂ -> 2H₂O) to produce electricity with water as the only product.
Quick-Fire Flashcards (6)
Click any card to check the answer, then mark it as Known to track your progress.
In an exothermic reaction, how does the energy of the products compare to the energy of the reactants?
Define activation energy.
Explain in terms of bond breaking and bond making why a reaction is endothermic overall (3 marks).
State the formula for calculating overall energy change from bond energies.
In the temperature change required practical (adding NaOH to HCl), why does the temperature increase at first and then decrease as more NaOH is added?
Give two advantages and one disadvantage of hydrogen fuel cells compared to rechargeable lithium-ion batteries in electric cars (Triple).
Exam-Style Questions & Mark Schemes (4)
Write or think through your answer first, then click Check Mark Scheme to see the exact points examiners award marks for.
- Q1
Hydrogen reacts with chlorine to form hydrogen chloride gas according to the equation: H-H + Cl-Cl -> 2 H-Cl. Given the bond energies: H-H = 436 kJ/mol, Cl-Cl = 242 kJ/mol, and H-Cl = 431 kJ/mol, calculate the overall energy change for the reaction and explain whether the reaction is exothermic or endothermic. [4 marks]
Click “Check Mark Scheme” to compare your answer, 1. Total energy taken in to break bonds in reactants = 1 x (H-H) + 1 x (Cl-Cl) = 436 + 242 = 678 kJ/mol [1 mark]. 2. Total energy released when making bonds in products = 2 x (H-Cl) = 2 x 431 = 862 kJ/mol [1 mark]. 3. Overall energy change = Bonds broken: Bonds made = 678 - 862 = -184 kJ/mol [1 mark]. 4. The reaction is exothermic because the overall energy change is negative (more energy is released forming the 2 H-Cl bonds than is taken in breaking the H-H and Cl-Cl bonds) [1 mark].
- Q2
Methane burns in oxygen: CH₄ + 2O₂ -> CO₂ + 2H₂O. Using bond energies: C-H = 413 kJ/mol, O=O = 498 kJ/mol, C=O = 805 kJ/mol, and O-H = 464 kJ/mol, calculate the overall energy change for the complete combustion of 1 mole of methane. (Higher Tier) [4 marks]
Click “Check Mark Scheme” to compare your answer, 1. Bonds broken in reactants: 4 x (C-H) + 2 x (O=O) = (4 x 413) + (2 x 498) = 1,652 + 996 = 2,648 kJ/mol [1 mark]. 2. Bonds made in products: 2 x (C=O in CO₂) + 4 x (O-H in 2H₂O) = (2 x 805) + (4 x 464) = 1,610 + 1,856 = 3,466 kJ/mol [1 mark]. 3. Overall energy change = Bonds broken: Bonds made = 2,648 - 3,466 [1 mark]. 4. Overall energy change = -818 kJ/mol [1 mark].
- Q3
Describe how a student could investigate how the mass of zinc powder added to 25 cm³ of copper(II) sulfate solution affects the maximum temperature rise of the reaction. Include how to minimise errors. [5 marks]
Click “Check Mark Scheme” to compare your answer, 1. Place a polystyrene cup inside a glass beaker (for stability and thermal insulation) and use a measuring cylinder to add 25 cm³ of copper(II) sulfate solution [1 mark]. 2. Place a lid with a hole on the cup, insert a digital thermometer (reading to 0.1°C), and record the initial temperature of the solution [1 mark]. 3. Weigh out 0.2 g of zinc powder on an electronic balance, add it to the solution, replace the lid, and stir continuously [1 mark]. 4. Record the highest (maximum) temperature reached and calculate the temperature rise (T_max - T_initial) [1 mark]. 5. Repeat the experiment with fresh 25 cm³ samples of copper(II) sulfate of the same concentration using 0.4 g, 0.6 g, 0.8 g, and 1.0 g of zinc powder [1 mark].
- Q4
Write the overall equation and the two electrode half-equations for a hydrogen-oxygen fuel cell operating in an alkaline/acidic electrolyte, and evaluate the use of hydrogen fuel cells vs petrol engines in cars. (Triple Science) [5 marks]
Click “Check Mark Scheme” to compare your answer, 1. Overall equation: 2H₂(g) + O₂(g) -> 2H₂O(l) [1 mark]. 2. Anode (oxidation) half-equation: H₂ -> 2H⁺ + 2e⁻ (or 2H₂ + 4OH⁻ -> 4H₂O + 4e⁻) [1 mark]. 3. Cathode (reduction) half-equation: O₂ + 4H⁺ + 4e⁻ -> 2H₂O (or O₂ + 2H₂O + 4e⁻ -> 4OH⁻) [1 mark]. 4. Advantage of fuel cell: Produces only water at the point of use (zero carbon dioxide, carbon monoxide, or nitrogen oxide emissions), and has fewer moving parts so it is more energy-efficient than a combustion engine [1 mark]. 5. Disadvantage of fuel cell: Hydrogen is an explosive gas requiring heavy, high-pressure tanks, refuelling infrastructure is limited, and producing hydrogen by electrolysis or steam reforming often still burns fossil fuels [1 mark].
Common Exam Questions (4)
Remember 'BENDO-MEXO': Which is endothermic and which is exothermic in bond breaking vs bond making?
Use the acronym BENDO-MEXO: Breaking bonds is ENDOthermic (energy must be taken in to pull bonded atoms apart), and Making bonds is EXOthermic (energy is released when new bonds form). A reaction is overall exothermic if more energy is released when making new bonds than was taken in to break the old bonds.
How do you calculate the overall energy change (ΔH) of a reaction from bond energies? (Higher Tier)
Use the formula: Overall Energy Change = Total Energy of Bonds Broken (Reactants): Total Energy of Bonds Made (Products). Be careful: this is Reactants MINUS Products. If your final answer is negative (e.g., -184 kJ/mol), the reaction is exothermic; if your final answer is positive (e.g., +92 kJ/mol), the reaction is endothermic.
Why do you place a polystyrene cup inside a glass beaker with a lid in the temperature change required practical?
Expanded polystyrene is an excellent thermal insulator (contains trapped air pockets), which minimises unwanted thermal energy transfer (heat loss or gain) between the reaction mixture and the room. Placing a lid on top reduces heat loss by evaporation and convection, and standing the light polystyrene cup inside a heavy glass beaker prevents it from tipping over.
What are everyday examples of exothermic and endothermic reactions?
Exothermic reactions: combustion (burning fuels), neutralisation (acid + alkali), many oxidation reactions, respiration, and self-heating cans / hand warmers. Endothermic reactions: thermal decomposition (e.g., heating calcium carbonate), the reaction of citric acid with sodium hydrogencarbonate (sherbet), photosynthesis, and instant sports injury cold packs (dissolving ammonium nitrate in water).
Need more questions on Exothermic & Endothermic Reactions, Reaction Profiles & Bond Energy Calculations, or want to organise your own notes?
Type a specific exam question you are stuck on, pick one of the prompts below, or paste your class notes to build a custom set of flashcards and practice questions on this page.
Related study guides
Rates of Reaction & Catalysts
The rate of a chemical reaction measures how fast reactants are converted into products. Under collision theory, chemical reactions only occur when particles collide with energy greater than or equal to the activation energy; changing temperature, concentration, gas pressure, surface area, or adding a catalyst alters collision frequency or energy.
Reversible Reactions, Dynamic Equilibrium & Le Chatelier's Principle
In a reversible reaction, denoted by the symbol ⇌, the products of the reaction can react together to re-form the original reactants; if the forward reaction is exothermic, the reverse reaction is always endothermic by the exact same amount of energy. When a reversible reaction occurs in a closed system, it reaches dynamic equilibrium when the rate of the forward reaction equals the rate of the reverse reaction and the concentrations of reactants and products remain constant. For Higher Tier students, examiners heavily test Le Chatelier's Principle, requiring predictions of how changing temperature, pressure, or concentration shifts the position of equilibrium and alters equilibrium yield.
Making Soluble Salts, Neutralisation & Strong vs Weak Acids
Acids produce hydrogen ions (H⁺) in aqueous solution, whereas alkalis are soluble bases that produce hydroxide ions (OH⁻); in any neutralisation reaction, these ions combine according to the ionic equation H⁺(aq) + OH⁻(aq) -> H₂O(l). Pure, dry crystals of a soluble salt are prepared either by reacting an acid with an excess of an insoluble base (such as copper oxide) or by performing an accurate titration between an acid and a soluble alkali (such as sodium hydroxide). Examiners regularly test the 6-mark required practical method for making copper sulfate crystals, titration mole calculations, and the Higher Tier distinction between strong/weak acids and concentrated/dilute acids.