Thermochemistry

Thermochemistry Introduction

  • Thermochemistry is a branch of chemistry that studies the relationships between chemical reactions and energy changes.

The Nature of Energy

  • Chemistry vs. Energy

    • Chemistry focuses on matter, but it's crucial to understand that energy also influences matter.

  • Definition of Energy

    • Energy is defined as anything that has the capacity to do work.

  • Work

    • Work is described as a force acting over a distance.

  • Heat

    • Heat is the flow of energy caused by a difference in temperature.

    • Energy can be exchanged between objects through contact (e.g., collisions).

Energy, Heat, and Work

  • Energy

    • An object can possess energy as a quantity or as a collection of objects.

  • Exchange of Energy

    • Heat and work are the two main methods through which energy can be exchanged between objects.

Classification of Energy

  • Kinetic Energy

    • This is the energy of motion and represents energy being transferred.

  • Thermal Energy

    • Associated with temperature, thermal energy is a specific type of kinetic energy.

  • Potential Energy

    • Defined as energy stored in an object, this energy is associated with the object's composition and position.

    • Also known as chemical energy when referring to energy stored in the structure of a compound.

Manifestations of Energy

  • Different Forms of Energy

    • Kinetic Energy: Due to motion.

    • Potential Energy: Due to position or composition.

    • Thermal Energy: Associated with temperature.

    • Chemical Energy: Associated with positions of electrons and nuclei.

Units of Energy

  • Kinetic Energy Calculation

    • Kinetic energy (KE) is directly proportional to mass (m) and velocity (v):
      KE = rac{1}{2} m v^2

  • Joule Definition

    • One joule (J) is the energy needed to move a 1-kilogram mass at a speed of 1 m/s.

    • A joule (J) can also be defined as the energy to move a 1-kilogram mass a distance of 1 meter.

  • Calorie and Kilocalorie

    • A calorie (cal) is the amount of energy required to raise the temperature of one gram of water by 1 °C:
      1 ext{ cal} = 4.184 ext{ J}

    • A kilocalorie (kcal) is equivalent to 1000 cal and is often referred to as a food Calorie.

  • Kilowatt-hour

    • 1 kilowatt-hour (kWh) = 3.60 imes 10^6 ext{ J}

  • All conversion factors discussed are exact.

Studying Energy: The Universe Division

  • System

    • Defined as the material or process being studied regarding energy changes.

  • Surroundings

    • Everything else that the system can exchange energy with.

  • The focus is on the energy exchange between the system and surroundings.

The First Law of Thermodynamics: Law of Conservation of Energy

  • Thermodynamics Overview

    • Thermodynamics is the study of energy and its transformations.

  • First Law of Thermodynamics

    • Also known as the law of conservation of energy; it holds that the total amount of energy in the universe is constant.

    • It prohibits the design of perpetual motion machines that produce energy without an energy source.

Conservation of Energy

  • Energy cannot be created or destroyed; it can only be transformed.

  • The total energy present at the beginning of a process must equal the total energy present at the end.

Comparing Energies During Transfer

  • The amount of energy gained or lost by the system must equal the amount of energy lost or gained by the surroundings, demonstrating conservation of energy.

Energy Flow and Conservation of Energy

  • The principle of conservation of energy requires that the sum of energy changes in the system and surroundings equals zero.

Internal Energy

  • Definition of Internal Energy

    • Internal energy is the total kinetic and potential energy of all particles within a system:
      U = U{kinetic} + U{potential}

  • The change in internal energy of a system relies solely on the initial and final energy amounts, not the path taken during the change.

  • State Function

    • A state function provides a value that depends only on the initial and final conditions, disregarding the process used.

Example of State Function

  • Mountain Example

    • When climbing to the top of a mountain, regardless of the trail (long and winding vs. short and steep), the elevation difference remains constant. The elevation change is an example of a state function, depending only on the difference in elevation from base to peak.

Energy Flow in Chemical Reactions

  • Internal Energy Comparison

    • The total internal energy in 1 mole of carbon solid and 1 mole of another compound is greater than that of 1 mole of a different compound at consistent temperature and pressure.

  • Energy Release during a Reaction

    • Specific reactions may involve a net release of energy to the surroundings.

  • Energy Absorption

    • Following certain reactions, energy may be absorbed from the surroundings into the process.

Energy Flow Dynamics

  • When energy flows out of a system, it is considered negative, while energy flowing into the surroundings is positive.

  • Conversely, when energy flows into a system, it is positive, and energy flowing out of the surroundings is negative.

Summarizing Energy Flow

  • If reactants possess a higher internal energy than products, energy flow is negative, and vice versa.

  • Energy dynamics can be summarized by their associated signs based on energy flows.

Definitions of Energy, Heat, and Work

  • Heat (q)

    • Represents thermal energy exchange. - If heat is added to the system, it's positive. - If heat is lost from the system, it's negative.

  • Work (w)

    • Includes work done on the system (positive) or work done by the system (negative).

  • Change in Internal Energy (ΔE)

    • Positive if energy flows into the system, negative if it exits. - Shrinkage or expansion in energy flow leads to either exothermic or endothermic classifications.

Calculating ΔE and Identifying Reactions

  • Example Scenarios

    • Balloon heating with 900 J of heat and doing 422 J of work.

    • A 50 g sample of water cools from 30°C to 15°C, losing 3140 J of heat.

    • A reaction releases 8.65 kJ of heat without any work done.

Heat Exchange Principles

  • Defined as the thermal energy transfer between a system and its surroundings, primarily caused by thermal differences in temperature.

  • Heat flows from the higher-temperature matter to the lower-temperature matter until thermal equilibrium is reached.

Thermal Energy Transfer

  • Heat flow dynamics indicate that when two objects with different temperatures are in contact, heat transfers from hot to cold, resulting in equal energy exchanges.

Measuring Heat Capacity


  • The relationship between heat absorbed and temperature rise is expressed through heat capacity (C).

    • C's units are Joules per kelvin (J/K).


  • Specific Heat Capacity

    • Specific heat capacity (Cᵢ) is the energy required to raise the temperature of one gram of a substance by one degree Celsius.

    • Molar heat capacity is the energy needed to raise the temperature of one mole of a substance.


  • Specific Heat Example Table

    Substance

    Specific Heat Capacity (J/g·°C)


    Lead

    0.128


    Gold

    0.128


    Silver

    0.235


    Copper

    0.385


    Iron

    0.449


    Aluminum

    0.903


    Water

    4.18

    Quantitative Heat Energy Relationships

    • Heat capacity depends on two factors: mass and specific heat of the material.

    • This relationship allows for the calculation of heat absorbed by knowing mass, specific heat, and temperature change.

    Energy Transfer in Chemical Reactions

    • When energy is exchanged with surroundings (e.g., metal submerged in water), changes depend on mass and specific heat capacities of both substances.

    Specific Heat Calculation Example

    • To calculate specific heat of lead with a heating scenario involving specific energy values and mass changes.

    Pressure-Volume Work

    • PV Work Definition

      • Defined as work performed by volume changes against external pressure.

    • Volume expansion leads to work being done on the surroundings.

    Measuring Energy Changes: Calorimetry

    • Calorimetry Overview

      • Technique used to measure thermal energy exchanged between a reaction and its surroundings.

    • Bomb Calorimeter

      • A sealed insulated container used for combustion reactions at constant volume.

    Enthalpy (H)

    • Defined as the sum of internal energy and the product of pressure and volume of a system. - H = U + PV

    • Enthalpy change (ΔH) relates to heat transfer during processes maintained at constant pressure.

    Understanding Endothermic and Exothermic Reactions

    • Endothermic Reactions

      • Characterized by a positive ΔH, indicating heat absorption by the system. - They feel cold to the touch.

    • Exothermic Reactions

      • Signified by a negative ΔH, representing heat release by the system. - They feel warm to the touch.

    Molecular Perspectives of Energy Changes in Reactions

    • Exothermic Reactions

      • The surrounding temperature rises as potential energy of reactants is converted into heat. - The products possess less energy than reactants.

    • Endothermic Reactions

      • The surrounding temperature drops as heat is absorbed. - The products have more potential energy than reactants to accommodate energy gain.

    Enthalpy of Reaction Overview

    • Enthalpy changes are extensive properties; thus, larger quantities of reactants yield larger ΔH values.

    • Calculations often require considering the number of moles in a balanced reaction.

    Calorimetry at Constant Pressure

    • Typically executed in aqueous solutions, calorimetry ensures constant pressure due to interaction with the atmosphere. - Numerous experimental setups utilize nested foam cups.

    Specific Heat and Reaction Temperature Calculations

    • Examples provided showcase calculations of specific heat and thermal dynamics in calorimetry scenarios.

    Relationships Involving ΔH

    • Reactions multiplied by factors will also multiply ΔH by the same factor. - Reversing reactions changes ΔH's sign. - Overall heat changes in a reaction follow the sum of individual steps according to Hess's law.

    Standard Conditions and Enthalpy of Formation

    • Defined standard states of materials assist in establishing comparative enthalpy values. - The standard enthalpy change (ΔH°) describes changes under standard states. - The formation enthalpy (ΔHf°) refers to energy change to form 1 mole of a pure compound from its constituents in standard states.

    Example Values for Standard Enthalpies of Formation

    Substance

    ΔHf° (kJ/mol)

    Br(g)

    111.9

    Br2(l)

    0

    H2(g)

    0

    O2(g)

    0

    H2O(g)

    -241.8

    H2O(l)

    -285.8

    C(s, graphite)

    0

    Calculating Standard Enthalpy Changes

    • Changes in overall reactions can be assessed based on formation reactions.

    • The formula for ΔHrxn is given by summing the ΔH values of components across their formation states.

    Examples of Reaction Calculations

    • The section details how to calculate ΔH for specified reactions using prior data and sum methodologies.

    Environmental Considerations of Energy Use

    • In the US, average annual energy usage exceeds 100,000 kWh per person, predominantly from fossil fuels (coal, natural gas, and petroleum).

    • Demonstrates the reliance on combustible materials derived from ancient biological matter, affecting sustainability and ecological balances.