Exothermic and endothermic reactions Revision Notes

In an exothermic reaction, the total energy released when new bonds form in the products is greater than the energy required to break bonds in the reactants. This results in a net release of energy, usually in the form of heat, to the surroundings.

Because energy is conserved, the excess energy leaves the reacting system and transfers to the environment, increasing its thermal energy.

This transfer of thermal energy raises the temperature of the surroundings, which can be detected by a thermometer placed near the reaction.

• Energy released by bond formation > energy absorbed breaking bonds
• Net energy released as heat to surroundings
• Surrounding temperature increases due to energy transfer

When an exothermic reaction occurs, the temperature of the immediate environment rises. For example, mixing water with concentrated sulfuric acid releases heat, causing the mixture and container to become warm.

To measure this temperature change, a thermometer or temperature sensor can be placed in or near the reaction mixture.

The temperature rise is evidence that thermal energy is being released from the reaction to the surroundings.

• Temperature rise indicates heat release
• Use thermometer to measure temperature change
• Temperature increase confirms exothermic nature

Combustion reactions, such as burning methane gas, release heat energy and cause the surrounding air to warm.

Respiration in cells is a biological example where glucose reacts with oxygen, releasing energy used by the organism and heat to the environment.

Other examples include neutralisation reactions between acids and bases, which also release heat and raise the temperature of the solution.

• Combustion releases heat to surroundings
• Respiration produces energy and heat
• Neutralisation reactions are exothermic

When investigating exothermic reactions in the laboratory, it is important to use appropriate apparatus such as insulated containers to reduce heat loss to the wider environment and obtain accurate temperature readings.

Repeated trials improve reliability of temperature measurements.

Safety precautions include handling hot materials carefully and using protective equipment when working with reactive chemicals.

• Use insulated containers to measure temperature accurately
• Repeat experiments for consistent results
• Follow safety protocols when handling heat and chemicals

An endothermic reaction is a chemical process that requires energy to proceed, specifically absorbing thermal energy from its surroundings. This means that during the reaction, energy is taken in from the environment to break chemical bonds or cause other changes within the reacting substances.

Because energy is absorbed, the surroundings lose heat, which causes a measurable drop in temperature around the reaction site. This transfer of energy is essential to distinguish endothermic reactions from exothermic ones, where energy is released instead.

• Energy is absorbed from the surroundings during the reaction.
• The system’s internal energy increases as bonds break or new substances form.
• Surroundings experience a temperature decrease due to heat loss.

The temperature of the surroundings falls because thermal energy moves into the reacting system to supply the energy needed for the reaction. Since heat energy is leaving the surroundings, their temperature decreases according to the principles of heat transfer.

For example, if a reaction mixture is in a beaker at room temperature, the beaker and the air around it will cool down as the reaction proceeds. This cooling effect is a direct consequence of the endothermic nature of the reaction.

• Thermal energy flows from surroundings to the reacting system.
• Heat loss from surroundings causes temperature to drop.
• The magnitude of temperature change depends on the amount of energy absorbed and the heat capacity of the surroundings.

Photosynthesis is a classic example where plants absorb energy from sunlight to convert carbon dioxide and water into glucose and oxygen. This process requires energy input, making it endothermic.

Another example is the thermal decomposition of calcium carbonate (CaCO₃) into calcium oxide (CaO) and carbon dioxide (CO₂). Heat must be supplied for this reaction to occur, so it absorbs thermal energy from its environment.

In the laboratory, dissolving ammonium chloride (NH₄Cl) in water causes the solution to cool down, demonstrating an endothermic process as the salt absorbs heat from the water.

• Photosynthesis absorbs solar energy.
• Thermal decomposition reactions require heat input.
• Dissolving certain salts in water can absorb heat, lowering temperature.

To observe the temperature drop in an endothermic reaction, a thermometer can be placed in or near the reaction mixture. Careful measurement before, during, and after the reaction allows quantification of the temperature change.

Repeated trials improve reliability, and insulating the reaction vessel reduces heat exchange with the wider environment, making the temperature change clearer.

Safety precautions include handling chemicals carefully and avoiding direct contact with hot or cold substances produced during the reaction.

• Use a thermometer to measure temperature changes.
• Repeat measurements to ensure accuracy.
• Insulate the reaction vessel to minimise external heat loss or gain.
• Follow safety protocols when handling chemicals and equipment.

In an energy profile diagram for an endothermic reaction, the products have higher energy than the reactants, indicating energy absorption.

The difference in energy between products and reactants corresponds to the amount of thermal energy absorbed from the surroundings.

Activation energy is the initial energy needed to start the reaction, but the overall energy change is positive for endothermic reactions, reflecting net energy absorption.

• Products have higher energy than reactants in endothermic reactions.
• Overall energy change (ΔH) is positive.
• Activation energy must be overcome to initiate the reaction.

Reaction pathway diagrams graphically represent the energy changes that occur during a chemical reaction. They plot the potential energy of the system against the progress of the reaction, from reactants to products. These diagrams help visualise how energy is absorbed or released and the energy barrier that must be overcome for the reaction to proceed. Understanding these diagrams is essential for analysing whether a reaction is exothermic or endothermic.

• Energy is plotted on the vertical axis, reaction progress on the horizontal axis.
• Reactants start at an initial energy level; products finish at a different energy level.
• The highest point on the curve represents the activation energy.

In an exothermic reaction, the products have lower energy than the reactants. This means that energy is released to the surroundings, usually as heat, causing the temperature of the surroundings to rise. The difference in energy between reactants and products is negative, indicating a net release of energy.

The activation energy is the energy needed to start the reaction by breaking bonds in the reactants. After this energy barrier is overcome, new bonds form in the products, releasing more energy than was absorbed.

• Products energy < Reactants energy.
• Overall energy change (ΔH) is negative.
• Activation energy is shown as the peak above reactants' energy level.
• Example: Combustion of methane (CH₄ + 2O₂ → CO₂ + 2H₂O) releases energy.

In an endothermic reaction, the products have higher energy than the reactants, indicating that energy is absorbed from the surroundings. This absorption causes the temperature of the surroundings to decrease.

The activation energy is still required to initiate the reaction by breaking bonds. However, the energy absorbed to break bonds is greater than the energy released when new bonds form, resulting in a net energy gain by the system.

• Products energy > Reactants energy.
• Overall energy change (ΔH) is positive.
• Activation energy is shown as the peak above reactants' energy level.
• Example: Photosynthesis (6CO₂ + 6H₂O + energy → C₆H₁₂O₆ + 6O₂) requires energy input.

Activation energy is the minimum energy required to start a chemical reaction. It corresponds to the energy needed to break bonds in the reactants so that new bonds can form to create products.

On a reaction pathway diagram, activation energy is represented by the height difference between the reactants' energy level and the peak of the curve. Both exothermic and endothermic reactions require activation energy, but the overall energy change differs.

A catalyst lowers the activation energy, making the reaction proceed faster without changing the overall energy change.

• Activation energy is always positive and must be overcome for reaction to proceed.
• It is the energy difference between reactants and the highest point on the diagram.
• Catalysts reduce activation energy but do not affect ΔH.

Breaking chemical bonds requires energy input (endothermic process), while forming bonds releases energy (exothermic process). The overall energy change of a reaction depends on the balance between these two processes.

In exothermic reactions, the energy released by bond formation exceeds the energy needed to break bonds, resulting in a net release of energy.

In endothermic reactions, more energy is required to break bonds than is released by forming new bonds, so energy is absorbed from the surroundings.

• Bond breaking = energy absorbed (endothermic).
• Bond forming = energy released (exothermic).
• Net energy change = energy absorbed - energy released.

Activation energy is the smallest amount of energy that reacting particles need to collide with in order to start a chemical reaction. Even if particles collide, the reaction will not occur unless they have at least this minimum energy. This energy is required to break bonds in the reactants and allow new bonds to form in the products.

For example, when hydrogen gas reacts with oxygen gas to form water, the hydrogen and oxygen molecules must collide with enough energy to break their existing bonds before new bonds in water molecules can form.

• Activation energy is an energy barrier to reaction.
• Only collisions with energy equal to or greater than activation energy lead to reaction.

Activation energy determines whether a reaction will proceed when particles collide. If the activation energy is high, fewer particles have enough energy to react, so the reaction is slower. If the activation energy is low, more particles can react upon collision, so the reaction is faster.

This explains why some reactions happen quickly at room temperature, while others require heating or a catalyst to proceed at a noticeable rate.

• Higher activation energy means slower reaction rate.
• Lower activation energy means faster reaction rate.
• Catalysts work by lowering activation energy.

A reaction pathway diagram shows the energy changes during a reaction. The activation energy is represented by the energy difference between the reactants and the highest point on the curve, called the transition state or activated complex.

This peak represents the energy barrier that reactants must overcome. The difference between the reactants’ energy and the products’ energy shows whether the reaction is exothermic or endothermic, but the activation energy is always the energy needed to start the reaction.

• Activation energy = energy at peak − energy of reactants.
• It is the energy barrier to reaction progress.
• Visible as the highest point on the reaction energy curve.

Consider the combustion of methane gas. Methane and oxygen molecules must collide with enough energy to break bonds before new bonds in carbon dioxide and water form. The activation energy for this reaction is relatively high, so a spark or flame is needed to start it.

In contrast, the reaction between sodium and water is very fast because the activation energy is low; the reaction starts immediately on contact.

• Methane combustion requires ignition energy (high activation energy).
• Sodium reacts instantly with water (low activation energy).

Activation energy is a fundamental concept explaining why reactions do not occur spontaneously upon collision of particles. It sets the minimum energy threshold for a reaction to begin and influences the reaction rate.

Understanding activation energy helps in controlling reactions, such as using catalysts to lower it and speed up industrial processes.

• Activation energy is the minimum energy for reaction initiation.
• It affects reaction speed by controlling how many collisions are effective.
• It is represented on reaction energy diagrams as the energy barrier.
• Catalysts reduce activation energy, increasing reaction rates.

Reaction pathway diagrams graphically represent the energy changes that occur during a chemical reaction. They plot the potential energy of the system against the progress of the reaction, from reactants to products. These diagrams help visualise how energy is absorbed or released and the energy barrier that must be overcome for the reaction to proceed.

Key components of these diagrams include: the energy level of reactants, the energy level of products, the activation energy (the peak energy barrier), and the overall energy change of the reaction (difference between reactants and products).

• Reactants: substances present at the start of the reaction.
• Products: substances formed by the reaction.
• Activation energy (Ea): minimum energy required to start the reaction.
• Overall energy change (ΔH): difference in energy between reactants and products.

In an exothermic reaction, energy is released to the surroundings, usually as heat, causing the temperature of the surroundings to rise. On the diagram, the products have lower energy than the reactants, indicating a net release of energy.

The activation energy is shown as the energy difference between the reactants and the highest point on the curve (the transition state). The overall energy change is the difference between the reactants and products energy levels, represented as a downward arrow.

For example, the combustion of methane (CH₄ + 2O₂ → CO₂ + 2H₂O) is exothermic. The diagram shows reactants at a higher energy level than products, with a peak representing the activation energy needed to initiate combustion.

• Products energy level is lower than reactants.
• Energy is released to surroundings (negative ΔH).
• Activation energy is the energy barrier to start the reaction.
• Temperature of surroundings increases.

Endothermic reactions absorb energy from the surroundings, causing the temperature of the surroundings to decrease. On the diagram, the products have higher energy than the reactants, indicating a net absorption of energy.

The activation energy is again the energy difference between the reactants and the peak of the curve. The overall energy change is positive, shown as an upward arrow from reactants to products.

An example is the thermal decomposition of calcium carbonate (CaCO₃ → CaO + CO₂), which requires heat input. The diagram shows reactants at a lower energy level than products, with a peak representing the activation energy needed to break bonds.

• Products energy level is higher than reactants.
• Energy is absorbed from surroundings (positive ΔH).
• Activation energy is the energy barrier to start the reaction.
• Temperature of surroundings decreases.

Activation energy (Ea) is the minimum energy that reacting particles must have to successfully collide and form products. It corresponds to the peak of the energy curve in the reaction pathway diagram.

Even in exothermic reactions, activation energy is required to break bonds in the reactants before new bonds form in the products. This energy barrier explains why some reactions do not occur spontaneously at room temperature.

Catalysts work by lowering the activation energy, allowing reactions to proceed faster without changing the overall energy change.

• Ea is the energy difference between reactants and the highest point on the curve.
• Ea must be overcome for the reaction to proceed.
• Catalysts reduce Ea, speeding up reactions.
• Ea is always positive regardless of reaction type.

When sketching reaction pathway diagrams, it is important to label all key features clearly: reactants, products, activation energy, and overall energy change.

The vertical axis represents potential energy, while the horizontal axis represents the progress of the reaction (not time).

Use arrows to indicate the direction of energy change: downward for exothermic and upward for endothermic reactions. Activation energy is shown as the energy difference from reactants to the peak of the curve, often labelled with 'Ea'. The overall energy change (ΔH) is the difference between reactants and products energy levels, labelled accordingly.

• Label reactants and products at correct energy levels.
• Show activation energy as the peak above reactants.
• Indicate overall energy change with an arrow between reactants and products.
• Use clear, consistent axis labels and units if applicable.

Chemical bonds hold atoms together in molecules and compounds. These bonds represent stored energy because atoms are more stable when bonded than when separate. To change the arrangement of atoms in a chemical reaction, bonds must be broken and new bonds formed.

The energy involved in these processes is crucial to understanding whether a reaction releases or absorbs energy overall. This energy change determines if a reaction is exothermic or endothermic.

• Chemical bonds store potential energy.
• Changing bonds involves energy transfer.

Breaking a chemical bond means overcoming the attractive forces holding atoms together. This requires an input of energy to separate the atoms.

Because energy must be absorbed from the surroundings to break bonds, bond breaking is an endothermic process.

For example, breaking the H–H bond in hydrogen gas (H₂) requires energy to separate the two hydrogen atoms.

• Energy input is needed to break bonds.
• Bond breaking absorbs energy from the environment.
• This process is endothermic.

When atoms come together to form a chemical bond, energy is released because the bonded state is more stable than the separate atoms.

This release of energy to the surroundings makes bond formation an exothermic process.

For example, when two hydrogen atoms form an H–H bond, energy is released as the bond forms.

• Energy is released when bonds form.
• Bond formation transfers energy to the environment.
• This process is exothermic.

A chemical reaction involves both breaking bonds in the reactants and forming bonds in the products.

The total energy change depends on the balance between energy absorbed to break bonds and energy released when new bonds form.

If more energy is released forming bonds than absorbed breaking bonds, the reaction is exothermic overall (net energy release). Conversely, if more energy is absorbed breaking bonds than released forming bonds, the reaction is endothermic overall (net energy absorption).

• Reactions involve bond breaking and bond formation.
• Net energy change = energy absorbed (breaking) – energy released (forming).
• Determines if reaction is exothermic or endothermic overall.

In methane combustion, C–H bonds in methane and O=O bonds in oxygen are broken (energy absorbed). New C=O bonds in carbon dioxide and O–H bonds in water are formed (energy released).

The energy released forming the new bonds is greater than the energy absorbed breaking the original bonds, so combustion is exothermic.

This explains why burning methane releases heat and light.

• Breaking C–H and O=O bonds requires energy.
• Forming C=O and O–H bonds releases more energy.
• Net result: exothermic reaction releasing heat.