Exothermic and endothermic reactions Revision Notes

An exothermic reaction is a chemical reaction that releases thermal energy to its surroundings. This release occurs because the total energy needed to break the bonds in the reactants is less than the energy released when new bonds form in the products.

During the reaction, energy is transferred from the reacting system to the environment, usually as heat. This transfer raises the temperature of the surroundings, which can be detected using a thermometer.

The energy change in an exothermic reaction is negative, indicating that energy leaves the system. This energy change is often called the enthalpy change (ΔH), with ΔH < 0 for exothermic processes.

• Energy released when bonds form is greater than energy absorbed to break bonds.
• Thermal energy flows from the system to the surroundings.
• Temperature of surroundings increases due to heat release.

Chemical bonds store energy. Breaking bonds requires energy input (endothermic step), while forming bonds releases energy (exothermic step). In an exothermic reaction, the energy released by bond formation exceeds the energy required to break bonds.

This net release of energy appears as heat, increasing the kinetic energy of surrounding particles, which raises the temperature of the environment.

For example, in the combustion of methane (CH₄), carbon–hydrogen bonds break and new carbon–oxygen and oxygen–hydrogen bonds form in carbon dioxide and water, releasing heat:

• Breaking bonds absorbs energy.
• Forming bonds releases energy.
• Net energy release causes temperature rise in surroundings.

The combustion of methane is a common exothermic reaction:

• CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) + energy (heat)

To identify an exothermic reaction experimentally, measure the temperature of the surroundings before and during the reaction using a thermometer.

An increase in temperature indicates that heat is being released to the environment.

Common practical examples include: combustion reactions, neutralisation reactions between acids and alkalis, and respiration in living organisms (though respiration is a biological process).

• Use a thermometer to record initial and final temperatures.
• A temperature rise confirms an exothermic process.
• Repeat measurements improve reliability.

When asked to state or identify an exothermic reaction, focus on the release of thermal energy and the consequent temperature rise in the surroundings.

Use correct terminology such as 'thermal energy', 'temperature increase', and 'energy released to the environment'.

Include balanced chemical equations where relevant to show the reaction and energy changes clearly.

• State that exothermic reactions release heat to surroundings.
• Mention temperature rise as evidence of heat release.
• Provide balanced equations to support explanations.

An endothermic reaction is a chemical process that absorbs thermal energy from its surroundings. During such a reaction, energy is taken in to break bonds in the reactants, which requires more energy than is released when new bonds form in the products.

Because energy is absorbed, the surroundings lose heat, resulting in a decrease in their temperature. This energy transfer is essential to the reaction proceeding.

The enthalpy change (ΔH) for an endothermic reaction is positive, indicating that the system gains energy from the environment.

• Endothermic reactions absorb heat energy.
• Energy is required to break bonds in reactants.
• Overall energy absorbed exceeds energy released.
• Surroundings lose heat, causing temperature drop.
• ΔH (enthalpy change) is positive.

A key indicator of an endothermic reaction is a measurable drop in temperature of the surroundings. For example, if a reaction mixture is in a beaker, the temperature of the beaker or solution will decrease during the reaction.

This temperature change can be measured using a thermometer placed in the reaction mixture or in contact with the reaction vessel.

The temperature drop confirms that thermal energy is absorbed from the surroundings into the reaction system.

• Use a thermometer to measure temperature before and during the reaction.
• A decrease in temperature indicates an endothermic process.
• Temperature drop results from heat absorption by the reaction.

Common examples of endothermic reactions include: the dissolution of ammonium nitrate in water and photosynthesis in plants.

In the dissolution of ammonium nitrate, solid ammonium nitrate absorbs heat from water as it dissolves, causing the solution temperature to fall.

Photosynthesis absorbs energy from sunlight to convert carbon dioxide and water into glucose and oxygen, storing energy in chemical bonds.

• Dissolution of ammonium nitrate in water: NH4NO3(s) → NH4+(aq) + NO3−(aq) (absorbs heat)
• Photosynthesis: 6CO2(g) + 6H2O(l) + energy → C6H12O6(aq) + 6O2(g)
• Both processes absorb energy, causing surroundings to cool.

The enthalpy change (ΔH) quantifies the heat energy absorbed or released during a reaction at constant pressure.

For endothermic reactions, ΔH is positive because energy is absorbed from the surroundings.

This contrasts with exothermic reactions, where ΔH is negative due to energy release to the surroundings.

• ΔH > 0 for endothermic reactions.
• Energy absorbed breaks bonds in reactants.
• Energy absorbed exceeds energy released in bond formation.
• Positive ΔH corresponds to temperature decrease in surroundings.

When 5.0 g of ammonium nitrate dissolves in 50 cm³ of water, the temperature drops from 20.0 °C to 15.0 °C.

Calculate the heat absorbed by the solution, assuming the solution has the same density and specific heat capacity as water (density = 1.0 g/cm³, specific heat capacity = 4.18 J/g°C).

Step 1: Calculate mass of solution: mass = volume × density = 50 cm³ × 1.0 g/cm³ = 50 g. Step 2: Calculate temperature change: ΔT = 15.0 °C − 20.0 °C = −5.0 °C (temperature decreases). Step 3: Calculate heat absorbed (q): q = mass × specific heat capacity × ΔT = 50 g × 4.18 J/g°C × (−5.0 °C) = −1045 J. The negative sign indicates heat is lost by the surroundings; the reaction absorbs this heat, so the reaction absorbs +1045 J.

• Mass of solution = 50 g
• Temperature change = −5.0 °C
• Heat absorbed by reaction = +1045 J
• Heat lost by surroundings = −1045 J

Reaction pathway diagrams graphically represent the energy changes that occur during a chemical reaction as it progresses from reactants to products. The vertical axis shows the potential energy of the system, while the horizontal axis represents the progress of the reaction over time.

These diagrams help visualise the energy required to start a reaction and the overall energy change, allowing distinction between exothermic and endothermic reactions. Key features include the energy levels of reactants and products, the activation energy, and the enthalpy change (ΔH).

• Vertical axis: potential energy (kJ/mol).
• Horizontal axis: reaction progress.
• Shows energy barrier and overall energy change.

Activation energy is the minimum energy that reacting particles must have to successfully collide and form products. It appears as a peak in the reaction pathway diagram, representing the energy barrier to the reaction.

Even exothermic reactions require activation energy to start. The height of this peak above the reactants’ energy level is the activation energy.

The activation energy influences the reaction rate: higher Ea means fewer particles have enough energy to react at a given temperature, so the reaction is slower.

• Ea is the energy difference between reactants and the highest point on the diagram.
• Ea must be overcome for the reaction to proceed.
• Ea affects reaction speed but not the overall energy change.

In exothermic reactions, the products have lower energy than the reactants, indicating that energy is released to the surroundings, usually as heat.

The enthalpy change (ΔH) is negative because the system loses energy.

The diagram shows reactants at a higher energy level, a peak representing activation energy, then products at a lower energy level than reactants. The difference in height between reactants and products is the magnitude of ΔH (negative).

• ΔH < 0 (negative).
• Energy released to surroundings.
• Products lower in energy than reactants.
• Example: Combustion of methane: CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔH ≈ -890 kJ/mol.

In endothermic reactions, the products have higher energy than the reactants, meaning energy is absorbed from the surroundings.

The enthalpy change (ΔH) is positive because the system gains energy.

The diagram shows reactants at a lower energy level, a peak for activation energy, then products at a higher energy level than reactants. The difference in height between products and reactants is the magnitude of ΔH (positive).

• ΔH > 0 (positive).
• Energy absorbed from surroundings.
• Products higher in energy than reactants.
• Example: Thermal decomposition of calcium carbonate: CaCO3(s) → CaO(s) + CO2(g) ΔH ≈ +178 kJ/mol.

Breaking chemical bonds requires energy input and is therefore endothermic.

Forming chemical bonds releases energy and is exothermic.

The overall enthalpy change of a reaction depends on the balance between energy absorbed to break bonds and energy released when new bonds form: ΔH = total bond energy absorbed − total bond energy released.

• Bond breaking: energy absorbed (endothermic).
• Bond formation: energy released (exothermic).
• Overall ΔH depends on bond energies.

Consider a reaction pathway diagram where reactants start at 200 kJ/mol, the activation energy peak is at 350 kJ/mol, and products are at 100 kJ/mol.

Calculate the activation energy and enthalpy change.

Activation energy, Ea = peak energy − reactants energy = 350 − 200 = 150 kJ/mol (energy barrier). Enthalpy change, ΔH = products energy − reactants energy = 100 − 200 = −100 kJ/mol (exothermic).

• Ea = 150 kJ/mol (energy needed to start reaction).
• ΔH = −100 kJ/mol (energy released overall).
• Reaction is exothermic because ΔH is negative.

Enthalpy change, symbolised as ΔH, is the amount of thermal energy absorbed or released by a chemical system during a reaction at constant pressure. It reflects the difference in enthalpy between products and reactants.

Mathematically, ΔH = H_products − H_reactants, where H represents the enthalpy (heat content) of a substance.

Enthalpy is a state function, meaning its change depends only on the initial and final states, not on the reaction pathway.

• ΔH is measured in kilojoules per mole (kJ/mol).
• Constant pressure conditions are typical in laboratory and atmospheric reactions.
• A negative ΔH indicates energy release; a positive ΔH indicates energy absorption.

In exothermic reactions, the system releases thermal energy to the surroundings, causing the temperature of the surroundings to rise.

Since energy leaves the system, the enthalpy change is negative (ΔH < 0).

Examples include combustion of fuels, respiration, and many neutralisation reactions.

• Example: Combustion of methane: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) ΔH = −890 kJ/mol
• Observation: Temperature of surroundings increases.
• Energy is given out as heat, light, or both.

Endothermic reactions absorb thermal energy from the surroundings, causing the temperature of the surroundings to decrease.

Because energy is taken into the system, the enthalpy change is positive (ΔH > 0).

Examples include photosynthesis and thermal decomposition reactions.

• Example: Thermal decomposition of calcium carbonate: CaCO₃(s) → CaO(s) + CO₂(g) ΔH = +178 kJ/mol
• Observation: Temperature of surroundings decreases.
• Energy must be supplied continuously for the reaction to proceed.

Chemical equations can include the enthalpy change to indicate whether the reaction is exothermic or endothermic.

The sign and magnitude of ΔH provide information about the energy change per mole of reaction as written.

Including state symbols clarifies the physical states of reactants and products, which affect enthalpy values.

• Example exothermic: 2H₂(g) + O₂(g) → 2H₂O(l) ΔH = −572 kJ/mol
• Example endothermic: N₂(g) + O₂(g) + energy → 2NO(g) ΔH = +180 kJ/mol
• State symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution.

Consider the reaction: H₂(g) + Cl₂(g) → 2HCl(g) ΔH = −184 kJ/mol.

Since ΔH is negative, the reaction is exothermic, releasing 184 kJ of energy per mole of H₂ reacted.

This energy release can be detected by a temperature rise in the surroundings during the reaction.

• If the reaction absorbed 184 kJ instead, ΔH would be +184 kJ, indicating an endothermic process.
• Always check the sign of ΔH to classify the reaction type.

Activation energy, symbolised as Ea, is the minimum amount of energy that reacting particles must possess during a collision for a chemical reaction to take place. It is the energy barrier that must be overcome to convert reactants into products.

Without sufficient energy to overcome this barrier, particles collide but do not react, resulting in no chemical change.

• Activation energy is measured in kilojoules per mole (kJ/mol).
• It is a threshold energy for successful collisions.

A reaction pathway diagram shows the energy changes during a reaction. The reactants start at a certain energy level, and to form products, the system must reach a higher energy state called the activated complex or transition state.

The difference in energy between the reactants and this activated complex is the activation energy (Ea). After reaching this peak, energy is released or absorbed as products form.

For example, in the reaction between hydrogen and iodine gases: H2(g) + I2(g) → 2HI(g), the activation energy corresponds to the energy needed to break the H–H and I–I bonds before new H–I bonds form.

• Activation energy is always positive because energy input is needed to start the reaction.
• The height of the activation energy barrier affects how fast the reaction proceeds.

Only collisions with energy equal to or greater than Ea result in a reaction. Increasing the temperature increases the proportion of particles with sufficient energy, thus increasing the reaction rate.

A higher activation energy means fewer particles have enough energy to react at a given temperature, so the reaction is slower.

Catalysts work by providing an alternative reaction pathway with a lower activation energy, increasing the rate without being consumed.

• Activation energy controls the speed of a reaction.
• Lower Ea means faster reaction; higher Ea means slower reaction.
• Temperature and catalysts influence the number of particles overcoming Ea.

Consider two reactions: Reaction A with Ea = 50 kJ/mol and Reaction B with Ea = 100 kJ/mol at the same temperature.

Since Reaction A has a lower activation energy, more particles have sufficient energy to react, so Reaction A proceeds faster than Reaction B.

This explains why some reactions occur rapidly at room temperature while others require heating.

• Activation energy is a key factor in determining reaction speed.
• Comparing Ea values helps predict relative reaction rates.

Reaction pathway diagrams graphically represent the energy changes that occur during a chemical reaction. The vertical axis shows the potential energy of the system, while the horizontal axis represents the progress of the reaction from reactants to products.

These diagrams help visualise the energy required to initiate a reaction and the overall energy change, providing insight into whether a reaction releases or absorbs energy. Key features include the energy levels of reactants and products, the activation energy (Ea), and the enthalpy change (ΔH).

• Vertical axis: Potential energy (arbitrary units).
• Horizontal axis: Reaction progress from reactants to products.
• Curve shows energy changes during bond breaking and bond forming.

In exothermic reactions, energy is released to the surroundings, so the products have lower energy than the reactants. This results in a negative enthalpy change (ΔH < 0).

The diagram starts with reactants at a higher energy level, rises to a peak representing the activation energy, then falls to a lower energy level for the products.

Activation energy (Ea) is the energy difference between the reactants and the highest point on the curve (the transition state). The enthalpy change (ΔH) is the difference between the reactants and products energy levels.

• ΔH is negative (products lower than reactants).
• Ea is the energy barrier to start the reaction.
• Energy is released as bonds form in products.

Endothermic reactions absorb energy from the surroundings, so the products have higher energy than the reactants, giving a positive enthalpy change (ΔH > 0).

The diagram begins with reactants at a lower energy level, rises to the activation energy peak, and ends with products at a higher energy level.

Activation energy (Ea) remains the energy difference between reactants and the peak, while ΔH is the energy difference between reactants and products.

• ΔH is positive (products higher than reactants).
• Ea is the minimum energy needed to initiate the reaction.
• Energy is absorbed to break bonds and form higher-energy products.

To construct a reaction pathway diagram, first draw the vertical energy axis and horizontal reaction progress axis. Label the reactants and products energy levels according to given data.

Mark the activation energy (Ea) as the energy difference from reactants to the peak of the curve. Label the enthalpy change (ΔH) as the difference between reactants and products energy levels, indicating whether it is positive or negative.

Use a smooth curve starting at reactants energy, rising to the peak (transition state), then descending to products energy. Include arrows and labels for clarity.

• Label reactants and products energy levels.
• Mark activation energy (Ea) from reactants to peak.
• Indicate enthalpy change (ΔH) between reactants and products.
• Use clear, consistent units and arrows.

Given: Reactants energy = 100 kJ/mol, Products energy = 40 kJ/mol, Activation energy (Ea) = 60 kJ/mol.

Step 1: Draw vertical axis labelled 'Energy (kJ/mol)' and horizontal axis labelled 'Reaction progress'.

Step 2: Mark reactants at 100 kJ/mol and products at 40 kJ/mol on the vertical axis, with reactants on the left and products on the right horizontally. Since products are lower, ΔH = 40 - 100 = -60 kJ/mol (exothermic). Label ΔH accordingly with a downward arrow from reactants to products energy level, marked -60 kJ/mol or ΔH = -60 kJ/mol. Step 3: Calculate the peak energy (transition state) as reactants energy + Ea = 100 + 60 = 160 kJ/mol. Mark this peak above reactants energy. Label the activation energy (Ea) as the energy difference from reactants (100 kJ/mol) to peak (160 kJ/mol), i.e. 60 kJ/mol, with an upward arrow and label Ea = 60 kJ/mol. Step 4: Draw a smooth curve starting at reactants (100), rising to peak (160), then descending to products (40). Step 5: Label all parts clearly: reactants, products, Ea, ΔH, and the transition state (peak).

• Calculate ΔH = products energy - reactants energy.
• Calculate peak energy = reactants energy + Ea.
• Draw curve and label energies and changes.
• Identify reaction as exothermic if ΔH < 0.

Chemical bonds hold atoms together in molecules and compounds. To break these bonds, energy must be supplied to overcome the attractive forces between atoms. This process requires energy input and is therefore endothermic.

Conversely, when new bonds form between atoms, energy is released as the atoms achieve a more stable arrangement. This release of energy makes bond formation an exothermic process.

The energy required to break a bond is called the bond dissociation energy or bond energy, usually expressed in kilojoules per mole (kJ/mol). Similarly, the energy released when a bond forms is equal in magnitude to the bond energy but opposite in sign.

• Breaking bonds: energy absorbed (endothermic).
• Forming bonds: energy released (exothermic).

The enthalpy change (ΔH) of a chemical reaction is the net energy change resulting from breaking bonds in the reactants and forming bonds in the products.

Mathematically, this can be expressed as: ΔH = Total energy absorbed in breaking bonds − Total energy released in forming bonds.

If more energy is absorbed breaking bonds than is released forming bonds, ΔH is positive and the reaction is endothermic overall. If more energy is released forming bonds than is absorbed breaking bonds, ΔH is negative and the reaction is exothermic overall.

• ΔH = Σ (bond energies of bonds broken) − Σ (bond energies of bonds formed)
• Positive ΔH: endothermic reaction.
• Negative ΔH: exothermic reaction.

Consider the reaction of hydrogen gas with chlorine gas to form hydrogen chloride:

H₂(g) + Cl₂(g) → 2HCl(g) (balanced equation) In this reaction:

Bonds broken: one H–H bond and one Cl–Cl bond (energy absorbed). Bonds formed: two H–Cl bonds (energy released).

• Breaking: H–H and Cl–Cl bonds (endothermic).
• Forming: 2 × H–Cl bonds (exothermic).

Given bond energies: H–H = 436 kJ/mol, Cl–Cl = 243 kJ/mol, H–Cl = 431 kJ/mol.

Calculate the enthalpy change for the reaction: H₂(g) + Cl₂(g) → 2HCl(g)

Step 1: Calculate total energy absorbed breaking bonds: = 436 kJ (H–H) + 243 kJ (Cl–Cl) = 679 kJ/mol Step 2: Calculate total energy released forming bonds: = 2 × 431 kJ (H–Cl) = 862 kJ/mol Step 3: Calculate ΔH: ΔH = Energy absorbed − Energy released = 679 kJ − 862 kJ = −183 kJ/mol The negative value indicates the reaction is exothermic, releasing 183 kJ per mole of reaction.

• ΔH = (436 + 243) − (2 × 431) = −183 kJ/mol
• Negative ΔH means energy is released overall.

Do not confuse bond breaking with bond formation; breaking always requires energy input, forming always releases energy.

Enthalpy change depends on the difference between total bond energies broken and formed, not just one or the other.

Bond energies are average values and can vary slightly depending on molecular environment, so calculated ΔH values are approximate.

• Breaking bonds is endothermic; forming bonds is exothermic.
• Calculate ΔH by subtracting energy released from energy absorbed.
• Negative ΔH indicates exothermic reaction; positive ΔH indicates endothermic reaction.

Chemical bonds hold atoms together in molecules. Each type of bond (single, double, triple) between specific atoms has an associated bond energy, which is the amount of energy required to break one mole of those bonds in the gaseous state.

Breaking bonds is an endothermic process because energy must be supplied to overcome the attractive forces holding the atoms together.

Conversely, forming bonds releases energy, making it an exothermic process. The energy released corresponds to the bond energy of the newly formed bonds.

• Bond energy is measured in kilojoules per mole (kJ/mol).
• Breaking bonds absorbs energy (positive energy change).
• Forming bonds releases energy (negative energy change).

The enthalpy change (ΔH) of a reaction can be estimated by considering the total energy required to break all the bonds in the reactants and the total energy released when new bonds form in the products.

The formula used is: \( \Delta H = \text{Total energy to break bonds} - \text{Total energy released forming bonds} \).

If ΔH is negative, the reaction is exothermic (energy released overall). If ΔH is positive, the reaction is endothermic (energy absorbed overall).

• Sum the bond energies of all bonds broken in the reactants.
• Sum the bond energies of all bonds formed in the products.
• Subtract the total bond energy of products from that of reactants to find ΔH.

Consider the combustion of methane: \( \mathrm{CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O} \).

Step 1: Identify bonds broken in reactants: 4 C–H bonds in methane, 2 O=O double bonds in oxygen.

Step 2: Identify bonds formed in products: 2 C=O double bonds in carbon dioxide, 4 O–H bonds in water (2 molecules × 2 O–H bonds each).

• Bond energies (approximate values): C–H = 412 kJ/mol, O=O = 498 kJ/mol, C=O (in CO₂) = 799 kJ/mol, O–H = 463 kJ/mol.
• Calculate total energy to break bonds: (4 × 412) + (2 × 498) = 1648 + 996 = 2644 kJ.
• Calculate total energy released forming bonds: (2 × 799) + (4 × 463) = 1598 + 1852 = 3450 kJ.
• Calculate ΔH: 2644 kJ − 3450 kJ = −806 kJ (exothermic).

A frequent mistake is to reverse the subtraction order in the formula, leading to incorrect sign for ΔH.

Another error is to omit some bonds or use incorrect bond energies, causing inaccurate results.

Always ensure bond energies correspond to the correct bonds and states (gaseous molecules).

• Use the formula ΔH = bonds broken − bonds formed, not the reverse.
• Include all bonds broken and formed in the calculation.
• Use consistent and accurate bond energy values from reliable data tables.

Write balanced chemical equations to identify all bonds involved clearly.

Show all steps: list bonds broken and formed with their energies, perform sums, and calculate ΔH with units.

State whether the reaction is exothermic or endothermic based on the sign of ΔH.

• Label all bond energies with units (kJ/mol).
• Show substitution into the formula explicitly.
• Include the final answer with correct units and sign.
• Explain the meaning of the sign of ΔH in context.