Thermochemistry

Introduction to Energy and Thermodynamics

Thermochemistry is a branch of thermodynamics that focuses on the study of energy changes, particularly heat, during chemical reactions. Energy is defined as the ability to do work or transfer heat, and it is a central concept in chemistry.

• Energy: The capacity to do work or transfer heat.

• Thermodynamics: The study of energy and its transformations.

• Thermochemistry: The study of chemical reactions and the energy changes involving heat.

The Nature of Chemical Energy

Chemical energy is primarily stored as potential energy in the form of electrostatic interactions between charged particles. The most important form is electrostatic potential energy (Eel).

• Electrostatic Potential Energy Formula:

• Unit of Energy: Joule (J), where

Attraction Between Ions

Electrostatic attraction occurs between oppositely charged ions. Energy is released when chemical bonds are formed (Eel < 0), and energy is consumed when bonds are broken (Eel > 0).

• Bond Formation: Releases energy (exothermic).

• Bond Breaking: Consumes energy (endothermic).

First Law of Thermodynamics

Law and Energy Conversion

The first law of thermodynamics states that energy can be converted from one form to another, but it cannot be created or destroyed. This principle is fundamental to understanding energy changes in chemical reactions.

• Energy Conservation: Total energy remains constant.

• Examples: Chemical energy converted to heat, sunlight converted to chemical energy in plants.

System and Surroundings

In thermochemistry, the universe is divided into the system (the part under study) and the surroundings (everything else).

• System: The portion singled out for study (e.g., reactants in a reaction).

• Surroundings: Everything outside the system.

Types of Systems

Systems are classified based on their ability to exchange energy and matter with their surroundings:

• Open System: Exchanges both energy and matter.

• Closed System: Exchanges energy but not matter.

• Isolated System: Exchanges neither energy nor matter.

Internal Energy

Definition and State Function

The internal energy (E) of a system is the sum of all kinetic and potential energies of its components. The change in internal energy (ΔE) is a state function, depending only on the initial and final states, not the path taken.

• Change in Internal Energy:

• State Function: Depends only on initial and final states.

Thermodynamic Quantities

Thermodynamic quantities have three parts: a number, a unit, and a sign. The sign indicates whether the system gains or loses energy.

• Positive ΔE: System gains energy from surroundings.

• Negative ΔE: System loses energy to surroundings.

Relating ΔE to Heat and Work

Energy exchange between system and surroundings occurs as heat (q) or work (w):

• Formula:

Sign Conventions for q, w, and ΔE

Endothermic and Exothermic Processes

Heat exchange between system and surroundings can be classified as:

• Endothermic: System absorbs heat (temperature drops).

• Exothermic: System releases heat (temperature rises).

State Functions

Definition and Path Independence

State functions depend only on the current state of the system, not on how it was reached. Internal energy is a state function, while heat and work are not.

• State Function: ΔE, enthalpy (H).

• Path Function: q, w (depend on the process).

Enthalpy

Definition and Calculation

Enthalpy (H) is a thermodynamic quantity that accounts for heat flow at constant pressure. It is defined as:

• Formula:

• Change in Enthalpy:

• At constant pressure, (heat gained or lost).

Enthalpy Change and Heat

• Endothermic Process: ΔH is positive.

• Exothermic Process: ΔH is negative.

Pressure–Volume Work

Mechanical work associated with a change in gas volume is called pressure–volume work:

• Formula:

• Work is negative when done by the system.

Enthalpies of Reaction

Heat of Reaction

The change in enthalpy for a reaction (ΔHrxn) is the enthalpy of products minus reactants:

• Formula:

Enthalpy Guidelines

• Extensive Property: Depends on the amount (moles) of reactant.

• Reverse Reaction: ΔH is equal in magnitude, opposite in sign.

• State Dependence: ΔH depends on the states (phases) of reactants and products.

Calorimetry

Measurement of Heat Flow

Calorimetry is used to measure heat flow in chemical reactions. The instrument used is a calorimeter.

• Heat Capacity: Energy required to raise temperature by 1 K (1°C).

• Specific Heat: Heat capacity per gram.

• Molar Heat Capacity: Heat capacity per mole.

• Constant-Pressure Calorimetry: Used for reactions in aqueous solution.

• Specific Heat of Water: 4.184 J/g·K.

• Calculation:

Bomb Calorimetry (Constant Volume)

• Used for reactions in a sealed vessel.

• Measures change in internal energy (ΔE).

• For most reactions, ΔE ≈ ΔH.

• Calculation:

Hess’s Law

Calculation of Enthalpy Changes

Hess’s law states that if a reaction is carried out in a series of steps, the overall ΔH is the sum of the enthalpy changes for the individual steps. This is possible because enthalpy is a state function.

• Enthalpy of Formation (ΔHf): Enthalpy change for forming a compound from its elements in their standard states.

• Standard Enthalpy of Formation (ΔHf°): Measured at 25°C, 1.00 atm.

• Calculation:

• ΔHf° for an element in its standard state is zero.

Bond Enthalpies

Definition and Application

Bond enthalpy is the energy required to break one mole of a particular bond in a gaseous substance. It is always positive, and the greater the bond enthalpy, the stronger the bond.

• Energy Released: When a bond forms between gaseous fragments.

• Calculation:

Bond Enthalpies and Enthalpies of Reaction

Bond energies can be used to estimate whether a reaction is endothermic or exothermic by comparing the energies required to break bonds and the energies released when new bonds are formed.

• Example: For the reaction CH4 + Cl2 → CH3Cl + HCl, calculate ΔH using bond enthalpies.