Thermochemistry
CHY113 OpenStax
Chapter 5: Thermochemistry
Learning Objectives
5.1 Energy Basics
• Define energy, distinguish types of energy, and describe the nature of energy changes that accompany chemical and physical changes
• Distinguish the related properties of heat, thermal energy, and temperature
• Define and distinguish specific heat and heat capacity, and describe the physical implications of both
• Perform calculations involving heat, specific heat, and temperature change
5.2 Calorimetry
• Explain the technique of calorimetry
• Calculate and interpret heat and related properties using typical calorimetry data
5.3 Enthalpy
• State the first law of thermodynamics
• Define enthalpy and explain its classification as a state function
• Write and balance thermochemical equations
• Calculate enthalpy changes for various chemical reactions
• Explain Hess’s law and use it to compute reaction enthalpies
Resources
5.1 Energy Basics
• Define energy, distinguish types of energy, and describe the nature of energy changes that accompany chemical and physical changes
• Distinguish the related properties of heat, thermal energy, and temperature
• Define and distinguish specific heat and heat capacity, and describe the physical implications of both
• Perform calculations involving heat, specific heat, and temperature change
Thermochemistry: The study of the heat absorbed or released during chemical and physical changes.
Energy: The capacity to supply heat or do work.
Law of conservation of energy: During a chemical or physical change, energy can be neither created nor destroyed,
although its form can change.
Two main types of energy:
Potential energy: the energy an object has because of its relative position, composition, or condition
Kinetic energy: the energy that an object possesses because of its motion
Thermal energy is a type of kinetic energy (KE) associated with the random motion of atoms and molecules.
Temperature is a quantitative measure of “hot” or “cold.”
Fast moving molecules -> High thermal energy -> “Hot”
Slow moving molecules -> Low thermal energy -> “Cold”
Heat (q) is the transfer of thermal energy between two bodies at different temperatures.
Heat flow (a redundant term, but one commonly used) increases the thermal energy of one body and decreases the thermal energy of the other.
When two substances are placed in contact, thermal energy will always flow from the high temperature substance to the low temperature substance.
Heat flow will continue until both substances are at the same temperature.
A change that releases heat is called an exothermic process.
A change that absorbs heat is an endothermic process.
The units of heat are the calorie, cal (historical), or the joule, J (the SI unit), where 1 cal = 4.184 J.
One calorie is the amount of heat necessary to raise 1 gram of water 1 K (or 1 ˚C since the Kelvin degree and the Centigrade degree have the same quantity of temperature).
The amount of temperature increase (or decrease) of a material can be quantified by either it's specific heat, c or s
(in units of J/g˚C) - an intensive property, or by the object's heat capacity, C (in units of J/˚C) - an extensive property.
The equations for heat transfer are: q = mc∆T where m is mass in grams, c is specific heat in J/g˚C, and
∆T = Tfinal -Tinitial, or q = C∆T where C is the heat capacity in J/˚C and ∆T = Tfinal -Tinitial.
Heat (flow) can be measured by a change in temperature of an amount of material. This process is called calorimetry.
So if 10. J of heat were added to 1.0 grams of water and 10. J of heat were added to 1.0 gram of gold, the gold's final temperature would be much hotter.
Add 10. J heat to 1.0 gram of water: 10. J = (1.0g)(4.184 J/g˚C)(∆T) ; ∆T = 10. J/(1.0g)(4.184 J/g˚C) = +2.4 ˚C
Add 10. J heat to 1.0 gram of gold: 10. J = (1.0g)(0.129 J/g˚C)(∆T) ; ∆T = 10. J/(1.0g)(0.129 J/g˚C) = +76 ˚C
5.2 Calorimetry
• Explain the technique of calorimetry
• Calculate and interpret heat and related properties using typical calorimetry data
Calorimetry is used to measure the amount of heat transferred to or from a substance.
In thermodynamics, we define the universe = system + surroundings.
For calorimetry, the reaction is the system and the jacket of water is the surroundings
In this cartoon of a bomb calorimeter in action, either (a) an exothermic process occurs and heat, q is negative, indicating that thermal energy is transferred from the system to its surroundings, or (b) an endothermic process occurs and heat, q, is positive, indicating that thermal energy is transferred from the surroundings to the system.
You can make a crude calorimeter very easily...
Figure 5.12 A Coffee Cup Calorimeter
A calorimetry experiment: Consider a hot piece of metal dropped into the cool water of a calorimeter.
When the hot metal is placed in the water within a calorimeter, heat will transfer from the metal (M) to the water (W).
The temperature of metal will decrease.
The temperature of water will increase.
Eventually, thermal equilibrium will be reached and both objects will have the same temperature.
Since this is done in a calorimeter, the heat exchange is only between metal and water.
The net change in heat is zero. (Conservation of Energy)
Rearranging shows that the heat gained by the metal is equal to the heat lost by the water.
The heat of both substances is equal in magnitude but opposite in sign.
5.3 Enthalpy
• State the first law of thermodynamics
• Define enthalpy and explain its classification as a state function
• Write and balance thermochemical equations
• Calculate enthalpy changes for various chemical reactions
• Explain Hess’s law and use it to compute reaction enthalpies
Thermochemistry is a branch of chemical thermodynamics, the science that deals with the relationships between heat, work, and other forms of energy.
Substances act as reservoirs of energy, meaning that energy can be added to them or removed from them.
Energy is stored in a substance when the kinetic energy of its atoms or molecules is raised.
Energy is also stored in the chemical bonds in a substance.
The total of all possible kinds of energy present in a substance is called the internal energy (U), sometimes symbolized as E.
The relationship between the change in internal energy (∆U), heat (q), and work (w) is summarized in the
First Law of Thermodynamics:
Energy can be transported into a system, resulting in an increase in internal energy.
The system absorbs heat from the surroundings: +q.
Or the surroundings do work on the system: +w.
Energy can also be transferred out of a system, resulting in a decrease in internal energy.
The system releases heat to the surroundings: –q.
Or the system does work on the surroundings: –w.
A type of work called expansion work (or pressure-volume work) occurs when:
A system pushes back the surroundings against a restraining pressure.
The surroundings compress the system.
Mathematically we define pressure-volume work as
Internal energy is an example of a state function (or state variable).
The value of a state function depends only on the state that a system is in, and not on how that state is reached.
Altitude or elevation is also a state function.
In an internal combustion engine of a car, when gasoline is burned, heat flows out of the engine but the engine also does compression work at the pistons to move the car. The change in internal energy therefore has a heat flow component and a pressure-volume work component. The birth of thermodynamics was the study of how to obtain the most work possible from a reaction.
Now we need to define a new state function that will allow quantification of energy changes during reactions that incorporates both internal energy and work.
Enthalpy (H) is defined as the sum of a system’s internal energy (U) and the mathematical product of its pressure (P) and volume (V):
Enthalpy is also a state function.
Enthalpy values for specific substances cannot be measured directly; only enthalpy changes (∆H) for chemical or physical processes can be determined.
Commonly, chemical and physical processes occur at constant pressure.
At constant pressure, the enthalpy change (ΔH) is:
Chemists use thermochemical equations to represent the changes in both matter and energy.
In a thermochemical equation, the enthalpy change (ΔH) of a reaction is written to the right of the balanced equation.
For example the combustion of hydrogen gas is written:
H2(g) + ½O2(g) ⟶ H2O(l) ΔH = −286 kJ/mol
which indicates that 286 kJ of heat are produced from combustion of 1 mole of hydrogen gas.
If the direction of a chemical equation is reversed, the arithmetic sign of ΔH is also reversed.
ΔH < 0 indicates an exothermic reaction.
ΔH > 0 indicates an endothermic reaction.
Enthalpy is an extensive quantity - burn 1 mole, get 286 kJ of heat, burn 2 moles and get 572 kJ of heat.
The ΔH of a reaction depends on the physical state of the reactants and products.
Enthalpy changes are typically tabulated for reactions in which both the reactants and products are at the same conditions.
A standard state is a commonly accepted set of conditions used as a reference point for the determination of properties under other different conditions.
The IUPAC standard state conditions and enthalpy comments:
Pressure of 1 bar (or 1 atm; 1 bar = 0.987 atm)
Solutions at 1 M concentration
We will include a superscript “o” in the enthalpy change symbol to designate standard state conditions.
The usual (but not technically standard) temperature is 298.15 K.
The symbol, ∆H˚, is used to indicate an enthalpy change for a process occurring under these conditions.
The enthalpy changes for many types of chemical and physical processes are available in the reference literature.
Recall ∆H is an extensive property—it is proportional to the amount of substances involved.
Often ∆H is normalized to a per-mole basis.
Standard enthalpy of combustion (ΔH°c) is the enthalpy change when exactly 1 mole of a substance burns (combines vigorously with oxygen) under standard state conditions. This is sometimes called the “heat of combustion.”
Standard enthalpy of formation (ΔH°f) is an enthalpy change for a reaction in which exactly 1 mole of a pure substance is formed from its constituent free elements in their most stable states under standard state conditions. By definition, the standard enthalpy of formation of an element in its most stable form is equal to zero under standard conditions.
Hess's Law
There are two ways to determine the amount of heat involved in a chemical change:
Measure it experimentally.
Calculate it from other experimentally determined enthalpy changes.
Some reactions are difficult, if not impossible, to investigate and make accurate measurements experimentally.
For these reactions, the amount of heat involved must be calculated, usually using Hess’s law.
Hess’s law: If a process can be written as the sum of several stepwise processes, the enthalpy change of the total process equals the sum of the enthalpy changes of the various steps.
Hess’s law is valid because enthalpy is a state function.
Enthalpy changes depend only on where a chemical process starts and ends, not on the path it takes from start to finish.
The formation of CO2(g) from its elements [C(graphite) + O2(g)] can be thought of as occurring in two steps, which sum to the overall reaction, as described by Hess’s law. The horizontal lines represent enthalpies. For an exothermic process, the products are at lower enthalpy than are the reactants.
reference: https://chemdictionary.org/hesss-law/
