Thermochemistry: Energy and Chemical Reactions

Chapter 5: Thermochemistry

• This covers the study of energy and its transformations, especially the relationship between chemical reactions and energy changes involving heat.

Section 1: The Nature of Chemical Energy

• Energy is defined as the capacity to do work or produce heat.
• Thermodynamics is the study of energy and its transformations.
• Thermochemistry is a type of thermodynamics that studies the relationships between chemical reactions and energy changes involving heat.
• Radiant energy (solar energy) and thermal energy (energy associated with random motion of molecules) are types of energy
• Potential energy represents energy available by virtue of an object's position and kinetic energy represents energy of motion
• All forms of energy can be converted from one form to another.
• The Law of Conservation of Energy states energy cannot be created or destroyed but only converted

Section 2: First Law of Thermodynamics

• Most chemical reactions absorb or produce energy, primarily in the form of heat.
• Heat is the transfer of thermal energy between two bodies at different temperatures.
• The first law of thermodynamics states that energy cannot be created or destroyed but can be converted into different forms.
• Analyzing energy changes requires defining the system, which is the specific part of the universe of interest.
• The surroundings are the rest of the universe outside the system.
• Systems can be open, closed, or isolated.
• An exothermic process releases heat and transfers thermal energy to the surroundings
• An endothermic process requires heat to be applied to the system.

Section 3: Enthalpy

• Internal energy of a system is the sum of kinetic and potential energies of its components.
• "ΔE = q + w" represents the change in a system's internal energy where "q" is heat and "w" is work
• "#" and a + or - sign indicate the direction of energy flow
• ΔΕ is negative if exothermic because energy leaves and work is done on the surroundings; positive if endothermic because energy enters and work is done on the system.
• Enthalpy (H) helps measure the flow of heat for a system
• ΔH represents change in enthalpy as in ΔH = H(products) - H(reactants)
• If products have more enthalpy, heat will be absorbed (ΔH is positive, and the reaction is endothermic);
• If reactants have more enthalpy, heat is released (ΔH is negative, and the reaction is exothermic).

Section 4: Enthalpies of Reaction

• The enthalpy change accompanying a reaction is called the enthalpy of reaction or the heat of reaction
• The reaction 2H2(g) + O2(g) → 2H2O(g) has a △H of -483.6kJ
• A negative enthalpy change indicates the reaction is exothermic
• Enthalpy is an extensive property, contingent on the state of reactants and products.
• Enthalpy change for a reaction equals the magnitude but is opposite in sign to the reverse reaction.

Section 5: Calorimetry

• Calorimetry is the science of measuring heat by observing temperature changes when a system absorbs or discharges energy as heat.
• A calorimeter is an instrument used experimentally to measure heat change.
• Specific heat of a substance is the amount of heat required to raise the temperature of one gram of the substance by 1 degree C; units are J/g°C.
• Molar heat capacity measures heat capacity of one mole of a substance.
• Calorimeters often have two cups nested inside each other
• The inner cup holds the solution where rxn occurs
• Temperatures are measured initially and after the reaction starts
• If the solution heats up, heat is released by the reaction (exothermic); if it cools down, heat is absorbed by the reaction (endothermic).

Section 6: Hess's Law

• During reaction, if temperature rises to 31.9C from 25C, then the enthalpy can he calculated using ΔH
• During calorimetry problems, the enthalpy of the solution is calculated using "q = (mass)(specific heat)(ΔTemperature)"
• According to Hess's Law; for a given set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.
• If a reactions is reversed, the sign of AH is reversed and AH is directly proportional to the reactants.
• If the coefficients in a balance equation are multipled by an integer, the ΔΗ is also multiplied by the same integer.

Section 7: Enthalpies of Formation

• At standard states, substances are at 1 atm and often 298K (rather than 273K as in STP).
• A degree sign (°) indicates a standard state.
• Charts provide ΔH° values for elements and compounds for calculating the standard enthalpy of reaction (ΔH°rxn).
• Hess's Law can be used to calculate standard enthalpy of reaction. Additionally, reactions also provide standard enthalpy of reaction

Section 8: Bond Enthalpies

• Energy changes from chemical reactions which involve heat are closely linked to changes in the formation and breaking of chemical bonds.
• Bond enthalpy represents the enthalpy change for breaking a particular bond in one mole of a gaseous substance.

Section 9: Foods and Fuels

• Chemical reactions that produce heat are combustion reactions
• The amount of energy released when one gram of any substance is combusted is the fuel value, which fuel values can be measured with calorimeters.
• The bodies energy comes from carbohydrates and fat, then broken down for muscles, temp control, and tissue repair
• Any unused excess is stored as fat.
• Fossil fuels like natural gas, petroleum, and coal are commonly used as fuels due to their bonds breaking and releasing heat, often for electricity.

Explore thermochemistry, the study of energy and its transformations in relation to chemical reactions. Learn about the nature of chemical energy, including potential and kinetic energy. Understand the First Law of Thermodynamics and the Law of Conservation of Energy.