Energy Changes
Objectives:
- Determine the types of heat changes (∆H) in physical and chemical processes.
- Interpret graphical representations of heat changes.
- Relate the physical state of a substance to its degree of orderliness.
- Determine the conditions for spontaneity of a reaction.
- Relate ∆H°, ∆S° and ∆G° as the driving forces for chemical reactions.
- Solve simple problems based on the relationship: ∆G° = ∆H° – T∆S°.
Key Concepts and Formulas
Energy Changes (∆H): The enthalpy change (∆H) measures the heat absorbed or released during a process. A positive ∆H (+∆H) indicates an endothermic reaction (heat absorbed), while a negative ∆H (–∆H) indicates an exothermic reaction (heat released).
Example: When Na dissolves in water, the process is exothermic (–∆H), releasing heat.
Endothermic Reaction: A reaction that absorbs heat from its surroundings (∆H > 0).
Example: Dissolution of ammonium chloride (NH₄Cl) in water is endothermic, making the solution feel cold.
Exothermic Reaction: A reaction that releases heat (∆H < 0).
Example: Combustion of potassium (K) in water releases a large amount of heat (exothermic).
Graphical Representations of Heat Changes (Heating & Cooling Curves)
Energy changes in substances are often shown on graphs to help us understand how heat affects temperature during physical and chemical processes. Below is a detailed explanation of heating and cooling curves with examples.
Heating and Cooling Curves: These are graphs that show how the temperature of a substance changes as heat is added or removed. They help explain what happens during physical state changes (like melting and boiling) and when a substance is heated or cooled without changing state.
Example: When heating ice, it first melts into water at 0°C without a rise in temperature. Once fully melted, heating continues to raise the temperature of the water. At 100°C, water boils to steam, again without temperature rise during the phase change.
How to Read a Heating Curve
When heat is applied to a substance (e.g., water), it shows distinct stages:
- Stage 1: Heating of the solid – temperature increases as heat is added.
- Stage 2: Melting – the solid turns to liquid at constant temperature (energy breaks bonds).
- Stage 3: Heating of the liquid – temperature of the liquid rises as more heat is added.
- Stage 4: Boiling – liquid changes to gas at constant temperature (energy used to overcome intermolecular forces).
- Stage 5: Heating of the gas – temperature of gas rises as heat is added.
Graphical Representation
The diagram below illustrates a typical heating curve of water:
Figure: A typical heating curve for water showing phase transitions and temperature changes during heating.
Important Observations:
- Flat sections of the curve (plateaus) represent **phase changes** (melting and boiling). During these phases, temperature stays constant even though heat is being added.
- Sloped sections represent temperature changes of solid, liquid, or gas states without a phase change.
- The formula used during sloped segments (without phase change) is: q = m × c × ∆T, where:
- q = heat absorbed/released
- m = mass of the substance
- c = specific heat capacity
- ∆T = change in temperature
Practical Example of Heating Curve:
Example: If you heat ice at –10°C until it becomes steam, you will first raise the temperature of the ice to 0°C, then melt it (constant temperature), then heat water to 100°C, then boil it (constant temperature), and finally heat steam above 100°C.
Summary:
- During **phase changes**, temperature remains constant as energy is used to break or form bonds.
- During heating or cooling within a single phase, temperature changes as energy increases or decreases the kinetic energy of particles.
Entropy and Order-Disorder Phenomena
Entropy (S) measures the degree of disorder or randomness in a system. An increase in entropy indicates more disorder.
Entropy (S): A measure of the randomness or disorder within a system. Higher entropy means more disorder.
Example: When gases mix, the system becomes more disordered, increasing entropy.
Order-Disorder Phenomenon: The physical state of a substance is related to its entropy. Solids (high order) have low entropy, whereas gases (high disorder) have high entropy.
Example: Dissolving a salt in water increases entropy because the orderly crystalline lattice breaks down into randomly dispersed ions.
Spontaneity of Reactions
The spontaneity of a reaction is determined by the Gibbs free energy change (∆G). The relationship is given by:
Gibbs Free Energy: ∆G° = ∆H° – T∆S°
Explanation: For a reaction to be spontaneous at constant temperature and pressure, ∆G must be negative. If ∆G = 0, the reaction is at equilibrium; if ∆G is positive, the reaction is non-spontaneous.
Worked Example: If a reaction has ∆H° = –50 kJ and ∆S° = –100 J/(mol·K) at 298 K, then convert ∆S° to kJ: –0.1 kJ/(mol·K) and compute ∆G° = –50 – (298 × –0.1) = –50 + 29.8 = –20.2 kJ. Since ∆G° is negative, the reaction is spontaneous.
Criterion for Equilibrium: A reaction at equilibrium has ∆G° = 0.
Example: For a reaction where the forward and reverse processes occur at the same rate, ∆G° = 0, indicating dynamic equilibrium.
Comments
Post a Comment
We’d love to hear from you! Share your thoughts or questions below. Please keep comments positive and meaningful, Comments are welcome — we moderate for spam and civility; please be respectful.