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
6.1.A.1 Temperature Changes Indicate Energy Changes:.
1. Thermodynamics Basics:
Thermodynamics is the science of energy transfer and change. It is regulated by four basic laws:
i. Zeroth Law:
Two systems, if they are in thermal equilibrium with a third system, will be in thermal equilibrium with one another, which allows the notion of temperature.
ii. First Law (Energy Conservation):
Energy cannot be created or destroyed, only transferred or converted. The change in internal energy of a system is the heat added minus the work done by the system.
iii. Second Law (Entropy):
The entropy of an isolated system will always increase over time, i.e., energy always spreads. This is the reason why processes are irreversible and no process can be 100% efficient.
iv. Third Law:
As temperature decreases to absolute zero, the entropy of a system approaches a minimum (typically zero for ideal crystals), i.e., absolute zero cannot be achieved.
v. Energy Conservation:
Energy is always conserved in any process. Energy can be converted from one form to another (e.g., work to heat), and this is captured in the First Law of Thermodynamics.
2. Heat vs. Temperature:
| Property | Heat | Temperature |
|---|---|---|
| Definition | Energy transferred due to a temperature difference. | Measure of the average kinetic energy of particles. |
| Unit of Measurement | Joules (J) or calories. | Celsius (°C), Fahrenheit (°F), Kelvin (K). |
| What it Measures | The total energy being transferred. | The degree of hotness or coldness of an object. |
| Nature | Energy in transit. | A scalar quantity representing thermal state. |
| Effect | Causes a change in temperature or state. | Indicates how hot or cold a substance is. |
| Example | Heat flowing from a stove to water. | The temperature of the water (e.g., 70°C). |
3. Specific Heat Capacity:
Specific Heat Capacity (c) is the amount of heat required to raise the temperature of a unit mass of a substance by 1°C (or 1 K). It is a property that defines the amount of heat that a material will take for a given change in temperature.
Formula:
Where:
– ( Q ) = Heat energy absorbed or released (Joules)
– ( m ) = Mass of substance (kg)
– ( c ) = Specific heat capacity (J/kg·°C or J/kg·K)
– ( Delta T ) = Temperature change (°C or K)
i. Why Materials Respond to Heat and Changes in Temperature:
a. Materials with High Specific Heat Capacity:
– Require more heat to change their temperature.
– Examples: Water, concrete.
– Effect: These materials heat up and cool down slowly, and thus are used for temperature control (e.g., water as a coolant or in thermal storage).
b. Materials with Low Specific Heat Capacity:
– Require less heat to change their temperature.
– Examples: Metals like iron, aluminum.
– Effect: Such materials easily heat up and cool rapidly, and hence are suitable for applications with rapid temperature changes (e.g., cooking pots).
4. Heat Transfer Mechanisms:
| Mechanism | Medium | How Heat is Transferred | Example |
|---|---|---|---|
| Conduction | Solids | Heat transferred through direct contact between particles | Metal spoon in hot water, heating a pan |
| Convection | Liquids & Gases | Heat transferred by the movement of fluid (rise of hot, sinking of cold) | Boiling water, air circulation in a room |
| Radiation | No medium (can occur in a vacuum) | Heat transferred through electromagnetic waves | Heat from the Sun, warmth from a fire |
6.1.A.2 Energy Changes in a System: Endothermic & Exothermic Processes:
1. Endothermic and Exothermic Processes:
Endothermic Processes (Heat Absorbing):
– Take heat from surroundings, lowering temperature.
– Examples: Melting, evaporation, sublimation, photosynthesis.
Exothermic Processes (Heat Releasing):
– Release heat to surroundings, raising temperature.
– Examples: Freezing, condensation, burning, respiration.
Phase Changes:
– Endothermic: Melting, evaporation (heat is absorbed).
– Exothermic: Freezing, condensation (heat is released).
2. Phase Changes:
i. Melting (Melting or Fusion):
– Energy absorbed: It takes heat to break bonds among particles of a solid and melt it to turn into liquid.
– Illustration: Melt water from ice.
ii. Freezing (Solidification):
– Energy liberated: Upon cooling, a liquid is given when its particles release energy and coalesce, relinquishing heat to the surroundings.
– Example: Water turning into ice.
iii. Boiling (Vaporization):
– Energy absorbed: Energy is required to break the intermolecular forces between molecules of a liquid to transform them into a gas.
– Example: Boiling water into steam.
In brief:
– Melting & Boiling: Endothermic (heat absorbed).
– Freezing: Exothermic (heat released).
3. Heat Transfer:
i. Conduction:
– Direct contact transfer: Heat moves through a material from a hot area to a cold area by direct particle-to-particle collision.
– Example: A metal spoon heating up in a boiling pot.
ii. Convection:
– Transfer by fluid motion: Heat is transferred through fluids (gases or liquids) by hot particles rising and cold particles falling, creating a circulation.
– Example: Warm air rising in a room, or water boiling.
iii. Radiation:
– Transfer by electromagnetic waves: Heat is transferred through space or in the air through infrared radiation without a medium.
– Example: Heat from the Sun on Earth or campfire heat.
4. Conservation of Energy:
The Law of Conservation of Energy informs us that energy cannot be created or destroyed, but may be transformed from one form to another.
– Energy Transformation: Energy may be transformed from one form to another, for example, from kinetic to potential energy, or from chemical to thermal energy, but the total amount of energy in a closed system remains always constant.
– Example: In a moving car, chemical energy in the fuel is transformed into kinetic energy (motion) and thermal energy (heat).
– Transfer of Energy: Energy is also transferred from one system to another, e.g., when heat is transferred from a hot body to a cold body.
In brief: Energy is conserved — it is transformed but the amount remains the same.
6.1.A.3 Energy Changes in Chemical Reactions: Exothermic & Endothermic Processes:
1. Exothermic & Endothermic Reactions:
i. Exothermic Reactions (Energy Released)
– Energy is released to the environment, usually in the form of heat or light.
– The products possess less energy than the reactants.
– Example: Combustion (like burning gasoline or wood).
ii. Endothermic Reactions (Energy Absorbed)
– Energy is absorbed from the environment, usually in the form of heat.
– The products possess more energy than the reactants.
– Example: Photosynthesis (plants absorbing sunlight to make food).
In short:
– Exothermic: Energy released.
– Endothermic: Energy absorbed.
2. Energy Diagrams:
Energy diagrams show the way in which a system’s energy changes over the course of a reaction. They plot the energy of the system against reaction progress.
i. Exothermic Reaction:
– Energy decreases: Products have lower energy than reactants, and energy is released.
– Diagram: Graph shows a decrease in energy as the reaction takes place.
– Example: Combustion (burning fuel).
ii. Endothermic Reaction:
– Energy increases: Products have more energy than reactants, and energy is absorbed.
– Diagram: Graph shows a rise in energy as reaction takes place.
– Example: Photosynthesis.
iii. Key Points:
– Activation Energy: The top point on the diagram represents the energy that needs to be put in to start reaction (activation energy).
– Exothermic: Energy released, graph drops.
– Endothermic: Energy absorbed, graph rises.
6.1.A.4 Solution Formation: Exothermic vs. Endothermic Processes:
1. Exothermic & Endothermic Processes:
Dissolution Energy Changes:
– Exothermic Dissolution: Energy released when the solute dissolves, since solute-solvent forces are stronger than solute-solute forces.
– Example: Dissolving NaOH in water.
– Endothermic Dissolution: Energy absorbed in dissolving the solute, as it takes more energy to dissociate solute molecules than is released by solute-solvent intermolecular forces.
– Example: Dissolving ammonium nitrate in water.
2. Intermolecular Interactions:
i. Before Dissolution:
– Solute-Solute Interactions: Intermolecular forces between the particles of the solute. For example, in a solid solute like salt, the ions are held by ionic forces or molecular forces.
– Solvent-Solvent Interactions: Forces between solvent molecules, for example, hydrogen bonds in water or van der Waals forces in nonpolar solvents.
ii. During Dissolution:
– The solute particles must overcome solute-solute interactions (e.g., to rupture ionic bonds or molecular forces).
– The solvent molecules must also overcome solvent-solvent interactions so that space can be available for the solute.
iii. After Dissolution:
– Solute-Solvent Interactions: Once the solute dissolves, there are new interactions between the solvent particles and the solute, e.g., hydrogen bonding in water or dipole-dipole interactions in polar solvents.
– These new interactions keep the solute trapped in the solvent, creating a homogeneous solution.
Summary:
– Before: Solute and solvent particles are held together by their own intermolecular forces.
– After: New solute-solvent interactions are formed, creating a stable solution.
OLD Content :
Endothermic and Exothermic Processes
- Energy: the capacity to do work or to produce heat
- J = SI unit for energy
- kJ = 103 J
- J = SI unit for energy
- Work: force acting over a distance.
- W = – P Δ V
- Potential Energy: due to position or composition (stored energy) → can be converted to work
- Ex: water behind a dam, saturated fat (energy stored in bonds), attractive and repulsive forces
- Kinetic Energy: energy due to motion of the object
- M = mass of object in kg; v = velocity of object in m/sec
- State function: value that depends on the state of the substance, not how that state was reached
- Ex: internal energy, pressure, volume, density, enthalpy
- Note: work and heat are not state functions
- Ex: internal energy, pressure, volume, density, enthalpy
System and Surroundings
- System: that on which we focus attention
- The solute, solution reactants and products, reaction
- Surroundings: everything else in the universe
- Ourselves (ie hands), a thermometer, reaction vessel, the solvent are part of the surroundings
- When we measure temp changes during a chemical reaction, we are measuring the surroundings
- Universe = System + Surroundings
Exo and Endothermic
- Heat exchange/temp change accompanies:
- Healing or cooling a substance
- Phase changes
- Dissolving solutes
- Chemical reactions
- Exothermic: heat (free energy) is released by the system (to the surroundings)
- Endothermic: Heat is absorbed by the system (from the surroundings)
Exothermic Processes
- Cooling an object
- Phase changes:
- Freezing; condensation; deposition
- Some chemical & dissolution reactions (feel hot) → temp of surroundings is increasing
- Electron affinity:
- Solute and Solvent Interactions: in exothermic reactions, the solute and solvent particles are more strongly attracted to each other than they are too themselves
- In an exothermic process, some kind of bonds or attractive force is forming!
- “Free to form”
- Justify: The energy required to break bonds (positive) is less than the energy released when the bonds are formed (negative) → process is energetically favorable → ΔH⁰ = negative
Endothermic Processes
- Heating an object
- Phase changes:
- Melting; vaporization; sublimation
- Some chemical and dissolution reactions (feels cold) → temp of surroundings is decreasing
- In dissolving the change in energy is equal to the sum of the energies required to separate solute particles from one another and solvent particles from one another minus the energy released when attractions between solute particles and solvent particles form
→ are breaking the forces of attraction between sodium and its valence electron- Solute and Solvent Interactions: in endothermic reactions, the solute and solvent particles are more strongly attracted to themselves (stronger interactions) than they are too each other.
- In an exothermic process, some kind of bond or attractive force is breaking
- “Takes (energy) to break”
- Justify: The energy required to break bonds is more than the energy released when the bonds are formed → ΔH⁰ = positive
- If temp graph is shown: because temp decreased, energy flows into the process
Why is dissolving an ionic compound an ambiguous change (can be classified as both physical and chemical)
