AP Chemistry 6.1 Endothermic and Exothermic Processes Study Notes - New Syllabus Effective fall 2024
AP Chemistry 6.1 Endothermic and Exothermic Processes Study Notes.- New syllabus
AP Chemistry 6.1 Endothermic and Exothermic Processes Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Explain the relationship between experimental observations and energy changes associated with a chemical or physical transformation.
Key Concepts:
- Energy of Phase Changes
- Exothermic & Endothermic Reactions
- Energy Diagrams
- Thermal Energy & Molecular Collisions
Temperature Changes and Energy Changes in a System
Temperature changes in a system indicate that energy changes have occurred within that system. When energy is transferred into or out of a system, its temperature changes accordingly.
- Temperature is a measure of the average kinetic energy of particles in a substance.
- An increase in temperature corresponds to an increase in the average kinetic energy of particles (energy absorbed).
- A decrease in temperature corresponds to a decrease in kinetic energy (energy released).
- Temperature change (\( \mathrm{\Delta T} \)) serves as a macroscopic indicator of microscopic energy transfer.
Key Idea: A change in temperature always signifies that energy has been transferred between the system and its surroundings, either as heat or work.
Example
When 100 g of water is heated from \( \mathrm{25^{\circ}C} \) to \( \mathrm{50^{\circ}C} \), what does this temperature increase indicate about energy transfer?
▶️ Answer / Explanation
- The temperature increase indicates that the system (water) absorbed energy from its surroundings.
- This absorbed energy increased the average kinetic energy of the water molecules.
- Thus, the process is endothermic with respect to the water sample.
Endothermic and Exothermic Processes
Energy changes in a system can be described as endothermic or exothermic processes, including heating and cooling, phase changes, or chemical transformations. These classifications depend on the direction of energy flow between the system and the surroundings.
Key Properties:
- Endothermic process: The system absorbs energy from the surroundings.
- Exothermic process: The system releases energy to the surroundings.
- For an endothermic reaction: \( \mathrm{q > 0} \) (heat flows into the system).
- For an exothermic reaction: \( \mathrm{q < 0} \) (heat flows out of the system).
- Temperature increase in the surroundings → exothermic; temperature decrease → endothermic.
Whether a process is endothermic or exothermic depends on the sign of heat flow (q) and the direction of energy transfer between the system and surroundings.
General Thermochemical Representation:
Endothermic: \( \mathrm{A + energy \rightarrow B} \)
Exothermic: \( \mathrm{C \rightarrow D + energy} \)
Example
Classify each process and indicate the direction of heat flow:
- (a) \( \mathrm{H_2O(s) \rightarrow H_2O(l)} \)
- (b) \( \mathrm{H_2O(l) \rightarrow H_2O(s)} \)
- (c) \( \mathrm{CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O + \text{energy}} \)
▶️ Answer / Explanation
- (a) Endothermic — energy is absorbed to overcome intermolecular hydrogen bonds during melting.
- (b) Exothermic — energy is released as liquid water molecules form a solid crystalline lattice.
- (c) Exothermic — combustion releases large amounts of energy as heat and light; heat flows from system → surroundings.
- In all cases, the sign of q indicates the energy direction:
- Endothermic: \( \mathrm{q > 0} \)
- Exothermic: \( \mathrm{q < 0} \)
Energy Flow in Chemical Reactions (System and Surroundings)
When a chemical reaction occurs, the energy of the system either decreases (in an exothermic reaction), increases (in an endothermic reaction), or remains the same. Energy transfer between the system and surroundings can occur as heat transfer (q) or work (w).
Key Properties:
For any energy exchange between system and surroundings:
\( \mathrm{\Delta E_{system} = q + w} \)
If the system releases energy:
\( \mathrm{q_{system} < 0} \) and \( \mathrm{q_{surroundings} > 0} \) → Exothermic reaction.
If the system absorbs energy:
\( \mathrm{q_{system} > 0} \) and \( \mathrm{q_{surroundings} < 0} \) → Endothermic reaction.
Energy is conserved in all processes:
\( \mathrm{q_{system} = -q_{surroundings}} \)
Temperature changes in the surroundings are an indicator of the direction of energy flow.
The energy gained or lost by a system during a reaction is equal in magnitude and opposite in sign to the energy lost or gained by its surroundings.
General Relationships:
Exothermic reaction: \( \mathrm{System \rightarrow Surroundings} \)
Endothermic reaction: \( \mathrm{Surroundings \rightarrow System} \)
Example
When 50.0 g of \( \mathrm{NaOH} \) dissolves in water, the temperature of the solution increases from \( \mathrm{25.0^{\circ}C} \) to \( \mathrm{45.0^{\circ}C} \). Describe the energy flow and classify the process.
▶️ Answer / Explanation
- The increase in temperature indicates that energy was released into the surroundings (the water solution).
- The system (dissolving NaOH) lost energy: \( \mathrm{q_{system} < 0} \).
- The surroundings gained energy: \( \mathrm{q_{surroundings} > 0} \).
- Thus, this dissolution is an exothermic process.
- Energy conservation is satisfied: \( \mathrm{q_{system} = -q_{surroundings}} \).
Energy Changes in Solution Formation
The formation of a solution may be an exothermic or endothermic process, depending on the relative strengths of the intermolecular/interparticle interactions before and after the dissolution process.
Key Properties:
When a solute dissolves in a solvent, three energy changes occur:
- Solute–solute separation (breaking solute particles apart) — endothermic.
- Solvent–solvent separation (creating space for solute particles) — endothermic.
- Solute–solvent interaction (formation of new interactions) — exothermic.
The overall enthalpy change of solution formation is given by:
\( \mathrm{\Delta H_{solution} = \Delta H_{solute} + \Delta H_{solvent} + \Delta H_{mix}} \)
- If \( \mathrm{\Delta H_{solution} < 0} \): the process is exothermic — energy is released when solute and solvent interact strongly.
- If \( \mathrm{\Delta H_{solution} > 0} \): the process is endothermic — more energy is required to separate particles than is released upon mixing.
- The magnitude of these terms determines whether the dissolution feels warm or cool to touch.
Key Idea: Solution formation involves a balance between the energy required to separate solute and solvent particles and the energy released when new solute–solvent interactions form.
Thermochemical Representation:
Endothermic dissolution: \( \mathrm{\text{solute(s)} + \text{solvent(l)} + \text{energy} \rightarrow \text{solution(aq)}} \)
Exothermic dissolution: \( \mathrm{\text{solute(s)} + \text{solvent(l)} \rightarrow \text{solution(aq)} + \text{energy}} \)
Example
When ammonium nitrate (\( \mathrm{NH_4NO_3} \)) dissolves in water, the temperature of the solution decreases. Explain why the process is endothermic and describe the energy changes that occur.
▶️ Answer / Explanation
- Energy is absorbed from the surroundings to break the ionic bonds in \( \mathrm{NH_4NO_3(s)} \) and hydrogen bonds in water molecules — both are endothermic steps.
- New ion–dipole interactions form between \( \mathrm{NH_4^+} \), \( \mathrm{NO_3^-} \), and water molecules — an exothermic process.
- However, the energy absorbed in separation exceeds the energy released upon mixing: \( \mathrm{|\Delta H_{solute} + \Delta H_{solvent}| > |\Delta H_{mix}|} \).
- Thus, the overall \( \mathrm{\Delta H_{solution} > 0} \), meaning the process is endothermic.
- The temperature of the solution drops because energy is drawn from the surroundings to complete dissolution.