Thermochemistry
All chemical reactions involve energy, and the study of energy and its transformations is known as thermodynamics
Energy is the capacity to do work or produce heat. It can't be created or destroyed and is a state function
Kinetic Energy is the energy of motion. KE =
mv2
Potential Energy is the energy of a position relative to other objects. Potential energy is stored energy that results from the
attractions and repulsions an object experiences in relation to other objects. Potential energy is usually converted into kinetic
energy and is energy at rest.
System and Surroundings
The portion singled out for study is called the system (object doing or receiving/reactants and products)
Everything else is regarded as the surroundings (equipment/containers)
Transferring Energy: Work and Heat
A force is any kind of push or pull exerted on an object
Energy used to cause an object to move against a force is called work
P =
W = (F)(D) W = -P V
The other way in which energy is transferred is as heat. Heat is the energy transferred from a hotter object to a colder one.
The First Law of Thermodynamics
One of the most important observations in any chemical reaction and transfer of energy is the fundamental idea that energy cannot
be created nor destroyed. Energy is conserved. Any energy that is lost by the system must be gained by the surroundings or vice
versa.
Internal Energy
Internal energy is the sum of all the kinetic and potential energies of all the components of the system.
?E = Efinal – Einitial
Relating ?E to Heat and Work
?E = q + w
Heat absorbed into system (+q)
Heat released by system (-q)
Work done on system (+w)
Work done by system (-w)
Endothermic and Exothermic Processes (Exothermic is favorable)
When a process occurs in which the system absorbs heat, we say that the process is endothermic. During an endothermic process,
heat flows into the system from its surroundings.
Ex: Melting of ice
A process that results in the evolution of heat is called exothermic. During an exothermic process, heat flows out of the system
and into its surroundings.
Ex: Combustion of gasoline
State Functions (dealing with present state)
q and w are not state functions, even though ?E is a state function.
Enthalpy
Enthalpy is the heat absorbed or released under constant pressure
Like internal energy, enthalpy is a state function.
The change in enthalpy, ?H, equals the heat, qp, gained or lost by the system when the process occurs under constant pressure
?H = Hfinal – Hinitial = qp
The subscript p indicates that pressure is constant
The sign on ?H indicates the direction of heat transfer during a process that occurs at constant pressure. A positive value of ?H
indicates that the system has gained heat from the surroundings. A negative value of ?H indicates that the system has released
heat to its surroundings.
Enthalpies of Reaction
The enthalpy change of a reaction is given by the equations
?H = H(products) – H(reactants)
The enthalpy change that accompanies a reaction is called the enthalpy of reaction or merely the heat of reaction and is sometimes
written as ?Hrxn
The enthalpy of a system is often viewed as a measure of how much heat is stored as potential energy in the system, or its heat
content.
Enthalpy is an extensive property, which means that the magnitude of ?H is directly proportional to the amount of reactant
consumed in the process.
Calorimetry
The temperature change experienced by an object when it absorbs a certain amount of energy is determined by its heat capacity.
We define heat capacity of an object as the amount of heat required to raise its temperature by 1 Kelvin, or 1 degree Celsius.
The heat capacity of one mol of a substance is called its molar heat capacity. The heat capacity of one gram of substance is called
its specific heat capacity or merely its specific heat.
q = (Cp)(m)(?T)
For example, 209 J is required to increase the temperature of 50.0 grams of water by 1.00 K. Thus, the specific heat of water is
209J = (Cp)(50.0g)(1.00K) = 4.18 J/g-K
Constant – Pressure Calorimetry
A coffee cup calorimeter is often used in general chemistry labs to illustrate the principles of calorimetry. Because the
calorimeter is not sealed, the reaction occurs under the essentially constant pressure of the atmosphere.
Andrew Rosen
qsoln = -qrxn
qsoln = (specific heat of solution)(grams of solution)(?T) = -qrxn
qp = E +P V
Bomb Calorimetry (Constant – Volume Calorimetry)
Combustion reactions, usually involving an organic compound, reacts with oxygen and is conveniently studied in a Bomb
Calorimeter.
To calculate the heat of combustion from the measured temperature increase in the bomb calorimeter, it is necessary to know the
heat capacity of the calorimeter, Ccal
This measurement can be achieved by combusting a sample that releases a known quantity of heat. For example, it is known that
combustion of exactly 1 gram of benzoic acid in a bomb calorimeter produces 26.38 kJ of heat. Suppose that 1.000 grams of
benzoic acid is combusted in a calorimeter and it causes a temperature increase of 4.857 degrees Celsius. The heat capacity of the
calorimeter is then given by
Ccal = 26.38 kJ / 4.857 degrees Celsius = 5.431 kJ / degree Celsius
Once we know the Ccal, we can measure temperature changes produced by other reactions, and from these we can calculate the
heat involved in the reaction, qrxn
qrxn = (-Ccal)(?T)
Hess’s Law of Heat Summation (Easier to go from high to low)
∑ ∑
Long list of equations. Use your imagination and manipulate them. Can use fractions.
When switching reactants and products flip the sign of H
When multiplying reactions by a constant also multiply the value of H by a constant
Add and subtract correct values of H when canceling reactions and producing the net equation
Combustion reactions are extremely exothermic (very favorable)
Enthalpies of Formation
The standard enthalpy of a reaction is defined as the enthalpy change when all the reactants and products are in their standard
states. We denote a standard enthalpy as ?Ho
, where the superscript o
indicates standard state conditions (1 atm and 298 K)
The standard enthalpy of formation of a compound, ?Hf
o
, is the change in enthalpy for the reaction that forms 1 mol of the
compound from its elements, with all substances in their standard states.
By definition, the standard enthalpy of formation of the most stable form of any element is zero because there is no formation
reaction needed when the element is already in its standard state. Thus, the values of ?Hf
o
for C(graphite), H2(g), O2(g), and the
standard states of other elements are zero by definition.
Foods and Fuels
The energy released when 1 gram of a material is combusted is often called its fuel value. Because fuel values represent the heat
released in a combustion, fuel values are positive numbers.
Foods
Most of the energy our bodies need comes from carbohydrates and fats. Carbohydrates are decomposed in the intestines into
glucose. Glucose is soluble in blood, and in the human body, it is known as blood sugar. It is transported by the blood to cells,
where it reacts with O2(g), H2O(l), and energy.
The average fuel value of carbohydrates is 17 kJ/g
Fats also produce CO2 and H2O in their metabolism and in their combustion in a bomb calorimeter. Fats are well suited to serve
as the body’s energy reserve for two reasons. They are insoluble in water, which permits their storage in the body; and the
produce more energy per gram then either proteins or carbohydrates.
The average fuel value of carbohydrates is 38 kJ/g
Proteins are used by the body mainly as building materials for organ walls, skin, hair, muscle, and so forth.
The average fuel value of proteins is 17 kJ/g
Fuels
During the complete combustion of fuels, carbon is converted to CO2 and hydrogen is converted to H2O, both of which have large negativde enthalpalies.