Lecture_11__5.1-5.4__Annotated

Chapter 5: Energy and the First Law of Thermodynamics

Date: Tuesday, February 11

Learning Goals for Today’s Lecture

• 5.1.1: Relation of energy changes during bond making/breaking to electrostatic potential energy between charged particles.

• 5.2.1: Understanding the first law of thermodynamics in relation to energy changes in chemical processes.

• 5.2.2: Connection between internal energy changes, heat transfer (into/out of the system), and work (done on/by the system).

• 5.2.3: Differentiation of state functions from non-state functions.

• 5.3.1: Connection between enthalpy and internal energy (E), pressure (P), and volume (V) of a system.

• 5.3.2: Definition and calculation of pressure-volume (P-V) work under constant pressure.

• 5.3.3: Interpretation of the sign of ΔH regarding endothermic and exothermic processes.

• 5.4.1: Interpretation of the enthalpy of reaction (ΔH) using the enthalpies of products and reactants.

• 5.4.2: Understanding how ΔH changes when the balanced reaction equation is multiplied or reversed.

Energy

• Definition: The ability to do work or transfer heat.

• Unit: Joules (J).

• Focus: Thermodynamics—study of energy and its transformations, particularly thermochemistry, which examines chemical reactions and energy changes related to heat.

5.1 The Nature of Chemical Energy

• Electrostatic Potential Energy: Important form of potential energy in charged particles expressed as ( E = \frac{1}{2} k \frac{Q_1 Q_2}{d} ) where:

  • Q = charge,

  • d = distance between charges.

• Unit of Energy: Joules (1 J = 2 kg m²/s²).

Attraction Between Ions

• Nature: Electrostatic attraction occurs between oppositely charged ions.

  • Energy Dynamics:

    • Forming chemical bonds: Energy released (E_el < 0).

    • Breaking chemical bonds: Energy consumed (E_el > 0).

    • Lower energy indicates stability in a bond.

5.2 First Law of Thermodynamics

• Law: Energy can be converted from one form to another, but it is neither created nor destroyed.

  • Animation examples include converting chemical energy to heat for heating homes and plants converting sunlight into chemical energy.

System and Surroundings

• System: The part of the universe chosen for study (e.g., chemical reactions).

• Surroundings: Everything else, including equipment and environment beyond the system.

Types of Systems

1. Open System: Exchanges heat and mass with surroundings.

2. Closed System: Exchanges only heat with surroundings.

3. Isolated System: Does not exchange heat or mass with surroundings.

Internal Energy

• Definition: Change in internal energy (ΔE) is the final energy minus initial energy (ΔE = E_final - E_initial).

• State Function: Depends only on initial and final states, not on the path taken.

Thermodynamic Quantities

Components

1. Number

2. Unit

3. Sign

   • Positive ΔE: System gains energy.

   • Negative ΔE: System loses energy.

Internal Energy Concepts

• Internal energy (E): Sum of all kinetic and potential energies of all system components.

• Observations:

  • If ΔE < 0: Energy released to surroundings.

  • If ΔE > 0: Energy absorbed from surroundings.

Relating ΔE to Heat and Work

• Energy exchange with surroundings occurs via heat (q) and work (w):

  • ( ΔE = q + w )

Sign Conventions

Exchange of Heat (Endothermic and Exothermic)

• Endothermic Process: Heat absorbed from surroundings; temperature drops.

• Exothermic Process: Heat released to surroundings; commonly observed in explosions.

Pressure-Volume Work

Work Done by Gases

• The mechanical work linked with gas volume change is relevant in chemical reactions.

  • Expressed as ( w = -PΔV )

  • Work is negative as it reflects work done by the system.

Enthalpy (H)

• Defined for processes at constant pressure as ( H = E + PV ) where:

  • ΔH is used to account for heat flow during reactions at constant pressure.

Enthalpy Change and Heat Direction

• Endothermic Process: ΔH is positive; heat absorbed.

• Exothermic Process: ΔH is negative; heat released.

Enthalpy of Reaction

• Defined as ΔH = H_products - H_reactants.

• Heat of Reaction (ΔH_rxn): Enthalpy associated with reactions.

Enthalpy Guidelines

1. Extensive Property: Depends on amount of substance (moles).

2. Reaction Sign Reversal: ΔH changes sign in reverse reaction.

3. State Dependence: Enthalpy changes depend on physical states of reactants/products.