Ch._7-1
Chemistry: A Molecular Approach - Chapter 7 Thermochemistry
Nature of Energy
Energy: capacity to do work or produce heat.
Defined as: Energy = Work = Force × Distance.
Work (w): a force acting over a distance.
Heat (q): energy transferred between objects due to temperature difference.
Heat flows from a warmer object to a cooler object, and only changes can be detected in heat flow.
System and Surroundings
Chemical reactions and physical state changes involve:
Release of heat or absorption of heat.
System: part of the universe involved in the chemical reaction or process.
Surroundings: everything else in the universe.
Exothermic vs. Endothermic Reactions
Endothermic Reactions:
Heat flows into the system from surroundings.
Surroundings cool down; system warms up.
q is positive.
Exothermic Reactions:
Heat flows out of the system into surroundings.
System cools down; surroundings warm up.
q is negative.
Additional Details on Endothermic and Exothermic
Endothermic: involves melting, boiling, sublimation; heat absorbed.
Exothermic: involves freezing, condensation, deposition; heat released.
Units of Energy
Common Units:
Calorie: quantity of heat needed to raise the temperature of 1 g of water by 1 °C.
Joule (J): SI unit of heat and energy.
Relationship: 1 cal = 4.184 J; 1 Cal (kilocalorie) = 1000 cal = 4184 J.
Conservation of Energy
Law of Conservation of Energy: energy cannot be created or destroyed, only transformed.
Total energy remains constant during energy transfers.
Example: Chemical energy converted to heat to warm homes; sunlight converted to chemical energy by plants.
First Law of Thermodynamics
Energy in the universe is constant; can’t create energy without a source.
Internal energy ( (ΔE)) of a system: sum of kinetic and potential energies of particles.
Change in internal energy: (ΔE = q + w)
State Functions
State Functions: do not depend on the path taken (e.g., temperature, pressure, volume).
Example: Bank balance is a state function regardless of how money was deposited or withdrawn.
Non-State Functions: heat and work depend on the path taken (e.g., work done to lift an object varies based on the method).
Energy Exchange
Heat and work are the only ways to transfer energy between a system and its surroundings.
Conventions for q and w:
q > 0: heat from surroundings to system.
q < 0: heat from system to surroundings.
w > 0: work done by surroundings on the system.
w < 0: work done by the system on surroundings.
Heat Capacity
Heat Capacity (C): the amount of heat needed to raise the temperature of an object.
Measured in J/°C or J/K.
Directly proportional to mass and specific heat of the material.
Specific Heat Capacity: heat required to raise 1 g of substance by 1 °C.
Heat Exchange Examples
Water has a high specific heat, moderating temperature changes.
Land heats and cools faster than water due to lower specific heat.
Relationship and Calculation of Heat Energy Absorbed
Example calculations for heat energy lost or gained using ( q = m \times C_s \times ΔT ).
Specific heat of aluminum (Al) is used for calculations showing heat energy transfer.
Enthalpy (
(H))
Enthalpy: heat absorbed or released at constant pressure.
Endothermic reactions: positive (ΔH); Exothermic reactions: negative (ΔH).
Hess's Law
States the change in enthalpy of reactions is independent of the route taken.
Total enthalpy change is the sum of individual steps in multistep reactions.
Standard Conditions for Enthalpy
Defined conditions for enthalpy change: pure gas at 1 atm, pure solids or liquids at stable form, and 1 M solutions.
Enthalpy of formation of elements in standard state is zero.
Calculating Standard Enthalpy Changes
ΔH ΔH °𝒓𝒙𝒏 = Σnp ΔH ºf (products) – Σnr ΔH ºf (reactants)
means sum.
n is the coefficient of the reaction.
Δ - A change in enthalpy
o - A degree signifies that it's a standard enthalpy change.
f - The f indicates that the substance is formed from its
°𝒓𝒙𝒏 = Σnp ΔH ºf (products) – Σnr ΔH ºf (reactants)
means sum
Example Problems
Several practical examples demonstrating calculation of enthalpy changes, combustion reactions, and application of Hess’s law.
Conclusion
Knowledge of thermochemistry is crucial for understanding energy transformations in chemical reactions.