CHEM1010 W12.1

Energy in Chemical Reactions

Electrostatic Potential Energy

• Like Charges (e.g., two positive ions):
  • When far apart, there's minimal interaction.
  • Bringing them together requires work due to repulsive forces.
  • The closer they get, the more work is needed, increasing electrostatic potential energy.
  • Releasing them converts potential energy into kinetic energy (movement).
• Opposite Charges (e.g., positive and negative ions):
  • Attract each other.
  • They accelerate towards each other, decreasing potential energy.
  • Separating them requires work, increasing potential energy.

Thermodynamics and the First Law

• Thermodynamics: Deals with work, potential energy, and kinetic energy relationships.
• First Law of Thermodynamics (Conservation of Energy):
  • Energy is neither created nor destroyed.
  • It's converted from one form to another (e.g., work to potential energy, potential energy to kinetic energy).

System and Surroundings

• System:
  • The chemical species reacting (reactants and products).
  • Can be a liquid or a gas.
• Surroundings:
  • The reaction vessel or container.
  • The atmosphere around the reaction.
• Universe:
  • Combination of the system and the surroundings.
  • Defining the system and surroundings helps track matter and energy.

Types of Systems

• Open System:
  • Both energy and matter can be transferred between the system and surroundings.
  • Example: Open mug of hot coffee.
    • System: The coffee liquid.
    • Surroundings: Mug, saucer, table, air.
    • Matter transfer: Evaporation of coffee (steam).
    • Energy transfer: Coffee heating the mug and surroundings.
• Closed System:
  • Energy can be transferred, but matter cannot.
  • Example: Closed cup of coffee with a lid.
    • Matter is contained.
    • Energy transfer: Coffee cooling down and releasing heat to the surroundings, glass gets hot.
• Isolated System:
  • Neither energy nor matter can be transferred.
  • Example: Coffee in a sealed, insulated mug.
    • Approximation: True isolated systems are hard to achieve.

Internal Energy (E)

• Represents the total kinetic and potential energies within a system.
• We primarily focus on the change in internal energy ($\Delta E$).
• $\Delta E = E<em>{\text{final}} - E</em>{\text{initial}}$, where:
  • $\Delta E$ is the change in internal energy.
  • $E_{\text{final}}$ is the final internal energy.
  • $E_{\text{initial}}$ is the initial internal energy.

Changes in Internal Energy

• If \Delta E < 0 (negative):
  • The system has lost energy to the surroundings.
• If \Delta E > 0 (positive):
  • The system has gained energy from the surroundings.

Combustion of Hydrogen and Oxygen

• $2H<em>2 + O</em>2 \rightarrow 2H_2O$
• Initial internal energy is higher than the final internal energy.
• \Delta E < 0
• Energy is released to the surroundings as light, heat, and sound.
• Kinetics describes the rate of reaction, while thermodynamics describes the start and end energies.

Endothermic and Exothermic Reactions (Heat Flow/Enthalpy)

• Enthalpy (H): Heat flow at constant pressure.
• Focus on the change in enthalpy ($\Delta H$).
• Endothermic Reactions:
  • \Delta H > 0. The system gains heat from the surroundings.
  • Feels cold because the system is drawing energy from the surroundings.
  • Key Distinction: Heat (energy flow) vs. Temperature (degrees).
• Exothermic Reactions:
  • \Delta H < 0. The system loses heat to the surroundings.
  • Often characterized by an increase in temperature in the surroundings.
  • Heat flows from the system to the surroundings.

State Functions

• Internal energy (E) and enthalpy (H) are state functions.
• We only care about the initial and final states, not the path taken.

Heat Flow and Enthalpy of Reaction

• Enthalpy of Reaction ($\Delta H_{\text{reaction}}$) or Heat of Reaction

• The amount of heat flow is proportional to the amount of reagents.

  • Example: Combustion of Methane ($CH_4$)
  • $CH<em>4 + 2O</em>2 \rightarrow CO<em>2 + 2H</em>2O$
  • One mole of methane releases approximately 890 kJ of heat to the surroundings.
  • Heat is transferred from the system to the surroundings (enthalpy is reduced).
  • Two moles of methane release 1780 kJ.
  • Ten moles of methane release 8900 kJ.
  • $\Delta H_{\text{reaction}}$ is often expressed in kJ/mol.

• Reverse Reaction:

  • The reverse reaction has the same magnitude of $\Delta H$ but with the opposite sign.
  • Example: $CO<em>2 + 2H</em>2O \rightarrow CH<em>4 + 2O</em>2$ has $\Delta H = +890$ kJ/mol.

Phase and Heat Flow

• The phase (solid, liquid, gas) of reactants and products affects $\Delta H$.
• Phase changes (e.g., evaporation) involve changes in enthalpy.

Calorimetry

• A process to determine the amount of heat being transferred.
• Measures changes in temperature to quantify heat flow.

Example: Formation of Liquid vs. Gaseous Water

• $2H<em>2(g) + O</em>2(g) \rightarrow 2H_2O(l) \quad \Delta H = -890 \text{ kJ}$
• $2H<em>2(g) + O</em>2(g) \rightarrow 2H_2O(g) \quad \Delta H = -802 \text{ kJ}$
• The difference in $\Delta H$ accounts for the condensation of water vapor to liquid water.
• The overall enthalpy change is the sum of the individual steps.

Hess's Law

• The total enthalpy change for a reaction is the sum of the enthalpy changes for each step, regardless of the number of steps.

• It doesn't matter how you get there, only the initial and final states.

• Example:

  • $2H<em>2(g) + O</em>2(g) \rightarrow 2H_2O(g)$

  • $2H<em>2O(g) \rightarrow 2H</em>2O(l)$

  • $\Delta H<em>{\text{total}} = \Delta H</em>1 + \Delta H_2$

Standard Enthalpy of Formation

• The enthalpy change when a compound is formed from its elements in their standard states.
• Using Hess’s law, it's possible to calculate the enthalpy of a reaction using standard enthalpies of formation.
• Reaction: $C<em>3H</em>8 + 5O<em>2 \rightarrow 3CO</em>2 + 4H_2O$
• We can look up standard enthalpies of formation for $C<em>3H</em>8$, $CO<em>2$, and $H</em>2O$.

Applying Hess's Law with Standard Enthalpies of Formation

• Imagine breaking reactants into elemental blocks and then reassembling them into products.
• The overall $\Delta H$ is the sum of the enthalpy changes for breaking down reactants and forming products.
• $\Delta H<em>{\text{reaction}} = \sum \Delta H</em>{\text{f, products}} - \sum \Delta H<em>{\text{f, reactants}}$, where $\Delta H</em>{\text{f}}$ is the standard enthalpy of formation
• Diagrammatical Representation:
  • Reactants (e.g., propane and oxygen) at a high energy level.
  • Products (e.g., $CO_2$ and water) at a lower energy level.
  • Breaking reactants into elements and then forming products allows us to account for energy changes (bookkeeping).
• The combustion reaction process involves taking all starting materials and accounting for the change in enthalpy from elemental forms and formation.
• Does not describe reaction's mechanism, only start and end states.

Next Steps

• The next lecture will consider entropy to determine whether reactions will spontaneously occur.