Chapter 6 Thermochemistry
Thermochemistry is the study of the energy evolved or absorbed in chemical reactions and any physical transformations, such as boiling and melting. Thermochemistry, generally, is concerned with the energy exchange accompanying transformations, such as mixing, chemical reactions, and including calculations of such quantities as the heat capacity, heat of formation, and enthalpy.
Energy is the capacity to do work. All forms of energy are capable of doing work, but not all forms of work pertain to chemistry (take physics). Work is subsequently a directed energy change from a process.
Types of energy:
-Radiant energy: solar energy; comes from the sun. It is the primary energy source of Earth, and allows life to exist.
-Thermal energy: the energy associated with random motion of atoms and molecules.
-Chemical energy: stored within structural units of chemical substances.
-Potential energy: energy available by virtue of an object's position.
Law of Conservation of Energy. The total quantity of energy in the universe is constant. This is the single most important law of our universe.
Energy Changes in Chemical Reactions
Three systems represented by water in a flask: (a) is an open system, and allows the exchange of both energy and mass with it's surroundings; (b) closed system which allows the exchange of energy but not mass; (c) an isolated system, which allows neither energy nor mass to be exchanged.
Heat - the transfer of thermal energy between two bodies that are different temperature.
Thermochemistry - the study of heat change in chemical reactions
System - specific part if the universe that is of interest to us.
Surroundings - the rest of the universe outside the system
Open system - can exchange mass and energy with the surroundings
Closed system - allows transfer of energy but not mass
Isolated system - does not allow transfer of either energy or mass
Exothermic - process that gives of heat or transfer energy to the Surroundings
Endothermic - heat has to be supplied to the system by the surroundings
This is an endothermic process
This is an exothermic process
Laws of Thermodynamics
In thermodynamics, there are four laws that describe the specifics for the transport of heat and work in thermodynamic processes. These laws can be applied to systems about which one knows nothing other than the balance of energy and matter transfer.
Zeroth Law - defines thermal equilibrium as a state of a system where there is no temperature change over time.
First Law of Thermodynamics - energy can be converted from one for to another, but cannot be created or destroyed
Second Law of Thermodynamics - entropy is always increasing
Third Law of Thermodynamics - impossible to cool an object to absolute zero
Sign Conventions for Work and Heat
Process
Sign
Work done by the system on the surroundings
-
Work done on the system by the surroundings
+
Heat absorbed by the system from the surroundings (endothermic process)
+
Heat absorbed by the surroundings from the system (exothermic process)
-
Work and Heat
the piston moves a distance h, against a pressure P and does work on the surroundings.
We know from physics that:
w = Fd
where w is work, F is the force acting, and d is distance.
The pressure is defined as P=F/A (the pressure P equals the force F applied over an area A), and when substituted into the previously stated work tells us that the work done on by the gas on the surroundings is:
w = - P * (delta V)
where P is pressure and, delta V is the change in volume.
The internal energy E of a system can be defined as the sum of the kinetic and potential energies of all the particles in the system. the internal energy can be changed by both work and heat. That is,
Delta E = q + w
where E is internal energy, q is heat, and w is work (previously defined).
Sample Question:
A gas is compressed from an initial volume of 5.75 L to a final volume of 1.23 L by an external pressure of 1.00 atm. During the compression the gas releases 128 J of heat.
In thermodynamics, the enthalpy (denoted as H, or specific enthalpy denoted as h) is a thermodynamic property of a fluid. It can be used to calculate the heat transfer during process taking place in a closed thermodynamic system under constant pressure. Enthalpy H is an arbitrary concept but the enthalpy change ΔH is more useful because it is equal to the change in the internal energy of the system, plus the work that the system has done on its surroundings. The enthalpy of a system H, can be defined as
H = E + PV
where E is the internal energy of the system, P is the pressure of the system, and V is the volume of the system.
change in enthalpy is very similar to the heat of reaction, especially in a constant pressure system.
Enthalpy is a common state function, since internal energy pressure and volume are all state functions. It is the heat at constant pressure.
Delta H = qp
Calorimetry
A calorimeter is a device that is used to determine the heat associated with a chemical reaction. Calorimetry is the science of measuring heat, and is based on observing changes in temperature due to heat absorption and dispersion. The heat capacity C of substance is defined as:
If the heat capacity is given per gram of a substance, it is called the specific heat capacity. The units are
Specific Heats of Some Common Substanaces
Substance
Specific Heat (J/gC)
Al
.9
Au
.129
C (graphite)
.72
C (diamond)
.502
Cu
.385
Fe
.444
Hg
.139
H2O
4.184
C2H5OH (ethanol)
2.46
The measurement of heat changes is performed using calorimetry, usually an enclosed chamber within which the change to be examined occurs. The temperature of the chamber is monitored using a thermometer. The energy released by the reaction is equal to the energy absorbed by the solution. This is equal, qualitatively, to the specific heat capacity x mass of solution x increase in temperature. More simply:
E = s x m x (Delta T)
Sample Question:
.35 grams of water are placed in a calorimeter. The water temperature drops from 34.79 C down to 30.01 C when .81 grams of a substance is placed into the calorimeter with it. The substance has a specific heat of .63 J/g-C. What was the initial temperature of the substance?
Answer:
0.35 x 4.18 ( 30.01 - 34.79) + 0.81 x 0.63 ( 30.01 - T)=0
- 6.99 + 15.3 - 0.510 T = 0
8.31 = 0.510 T
T = 16.3 °C
Standard Enthalpy
Enthalpy is a state function, as such the change in enthalpy is going from some initial state to some final state, is independent of the pathway. This means that in going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes one step or a series of steps. This principle is Hess's Law.
The total enthalpy of a system cannot be measured directly; the enthalpy change of a system can be however. Enthalpy change is defined by the following equation:
ΔH = H final − H initial
where
ΔH is the enthalpy change, H final is the final enthalpy of the system, measured in joules. In a chemical reaction, H final is the enthalpy of the products. H initial is the initial enthalpy of the system, measured in joules. In a chemical reaction, H initial is the enthalpy of the reactants.
For an endothermic reaction, the system's change in enthalpy is equal to the energy absorbed in the reaction, including the energy lost by the system and gained from compression from its surroundings. A relatively easy way to determine whether or not a reaction is exothermic or endothermic is to determine the sign of ΔH. If ΔH is positive, the reaction is endothermic, heat is absorbed by the system due to the products of the reaction having a greater enthalpy than the reactants. On the other hand if ΔH is negative, the reaction is exothermic, that is the overall decrease in enthalpy is achieved by the generation of heat.
Hess's Law
The process works by treating each chemical equation as a mathematical equation, adding and subtracting multiples of each to cancel each variable (chemical). The yield arrow acts as an equals sign, and different states of elements count as different variables.
Laboratory
Coffee Cup Calorimetry
For this lab you will need:
two or more styrofoam cups
a lid with holes for a thermometer and stirrer
a thermometer
a stirrer
hydrochloric acid
magnesium strips (approximately .1g)
First, nestle your styrofoam cups into each other. The more cups, the more isolated the system will be from it's surroundings. Fill your calorimeter with 100 ml of .1M hydrochloric acid. Set your thermometer into the acid, and allow it to become stable. Drop a magnesium strip into the acid and seal the cup. Watch the temperature rise, while stirring, and wait for it to cap. Record temperature change. Try to calculate the kcal/mole for the reaction. It should be around .5.
Clean out your cup using standard laboratory cleanup procedures.
References
All images and information was drawn from myself or the following sources:
Chang, Chemistry, 8th Edition
Zumdahl, 4th Edition
Thermochemistry is the study of the energy evolved or absorbed in chemical reactions and any physical transformations, such as boiling and melting. Thermochemistry, generally, is concerned with the energy exchange accompanying transformations, such as mixing, chemical reactions, and including calculations of such quantities as the heat capacity, heat of formation, and enthalpy.
Table of Contents
Energy
Energy is the capacity to do work. All forms of energy are capable of doing work, but not all forms of work pertain to chemistry (take physics). Work is subsequently a directed energy change from a process.Types of energy:
-Radiant energy: solar energy; comes from the sun. It is the primary energy source of Earth, and allows life to exist.
-Thermal energy: the energy associated with random motion of atoms and molecules.
-Chemical energy: stored within structural units of chemical substances.
-Potential energy: energy available by virtue of an object's position.
Law of Conservation of Energy. The total quantity of energy in the universe is constant. This is the single most important law of our universe.
Energy Changes in Chemical Reactions
Heat - the transfer of thermal energy between two bodies that are different temperature.
Thermochemistry - the study of heat change in chemical reactions
System - specific part if the universe that is of interest to us.
Surroundings - the rest of the universe outside the system
Open system - can exchange mass and energy with the surroundings
Closed system - allows transfer of energy but not mass
Isolated system - does not allow transfer of either energy or mass
Exothermic - process that gives of heat or transfer energy to the Surroundings
Endothermic - heat has to be supplied to the system by the surroundings
Laws of Thermodynamics
In thermodynamics, there are four laws that describe the specifics for the transport of heat and work in thermodynamic processes. These laws can be applied to systems about which one knows nothing other than the balance of energy and matter transfer.
Zeroth Law - defines thermal equilibrium as a state of a system where there is no temperature change over time.
First Law of Thermodynamics - energy can be converted from one for to another, but cannot be created or destroyed
Second Law of Thermodynamics - entropy is always increasing
Third Law of Thermodynamics - impossible to cool an object to absolute zero
Work and Heat
We know from physics that:
w = Fd
where w is work, F is the force acting, and d is distance.
The pressure is defined as P=F/A (the pressure P equals the force F applied over an area A), and when substituted into the previously stated work tells us that the work done on by the gas on the surroundings is:
w = - P * (delta V)
where P is pressure and, delta V is the change in volume.
The internal energy E of a system can be defined as the sum of the kinetic and potential energies of all the particles in the system. the internal energy can be changed by both work and heat. That is,
Delta E = q + w
where E is internal energy, q is heat, and w is work (previously defined).
Sample Question:
A gas is compressed from an initial volume of 5.75 L to a final volume of 1.23 L by an external pressure of 1.00 atm. During the compression the gas releases 128 J of heat.
Answer:
-1.00 atm (1.23L - 5.75L) = 4.52L*atm
4.52L*atm x 101.325 J/(L*atm) = 457.989J
457.989J - 128J = 330J
Enthalpy
In thermodynamics, the enthalpy (denoted as H, or specific enthalpy denoted as h) is a thermodynamic property of a fluid. It can be used to calculate the heat transfer during process taking place in a closed thermodynamic system under constant pressure. Enthalpy H is an arbitrary concept but the enthalpy change ΔH is more useful because it is equal to the change in the internal energy of the system, plus the work that the system has done on its surroundings. The enthalpy of a system H, can be defined as
H = E + PV
where E is the internal energy of the system, P is the pressure of the system, and V is the volume of the system.
change in enthalpy is very similar to the heat of reaction, especially in a constant pressure system.
Enthalpy is a common state function, since internal energy pressure and volume are all state functions. It is the heat at constant pressure.
Delta H = qp
Calorimetry
A calorimeter is a device that is used to determine the heat associated with a chemical reaction. Calorimetry is the science of measuring heat, and is based on observing changes in temperature due to heat absorption and dispersion. The heat capacity C of substance is defined as:
If the heat capacity is given per gram of a substance, it is called the specific heat capacity. The units are
E = s x m x (Delta T)
Sample Question:
.35 grams of water are placed in a calorimeter. The water temperature drops from 34.79 C down to 30.01 C when .81 grams of a substance is placed into the calorimeter with it. The substance has a specific heat of .63 J/g-C. What was the initial temperature of the substance?
Answer:
0.35 x 4.18 ( 30.01 - 34.79) + 0.81 x 0.63 ( 30.01 - T)=0
- 6.99 + 15.3 - 0.510 T = 0
8.31 = 0.510 T
T = 16.3 °C
Standard Enthalpy
Enthalpy is a state function, as such the change in enthalpy is going from some initial state to some final state, is independent of the pathway. This means that in going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes one step or a series of steps. This principle is Hess's Law.
The total enthalpy of a system cannot be measured directly; the enthalpy change of a system can be however. Enthalpy change is defined by the following equation:
ΔH = H final − H initial
where
ΔH is the enthalpy change, H final is the final enthalpy of the system, measured in joules. In a chemical reaction, H final is the enthalpy of the products. H initial is the initial enthalpy of the system, measured in joules. In a chemical reaction, H initial is the enthalpy of the reactants.
For an endothermic reaction, the system's change in enthalpy is equal to the energy absorbed in the reaction, including the energy lost by the system and gained from compression from its surroundings. A relatively easy way to determine whether or not a reaction is exothermic or endothermic is to determine the sign of ΔH. If ΔH is positive, the reaction is endothermic, heat is absorbed by the system due to the products of the reaction having a greater enthalpy than the reactants. On the other hand if ΔH is negative, the reaction is exothermic, that is the overall decrease in enthalpy is achieved by the generation of heat.
Hess's Law
The process works by treating each chemical equation as a mathematical equation, adding and subtracting multiples of each to cancel each variable (chemical). The yield arrow acts as an equals sign, and different states of elements count as different variables.
Laboratory
Coffee Cup Calorimetry
For this lab you will need:
First, nestle your styrofoam cups into each other. The more cups, the more isolated the system will be from it's surroundings. Fill your calorimeter with 100 ml of .1M hydrochloric acid. Set your thermometer into the acid, and allow it to become stable. Drop a magnesium strip into the acid and seal the cup. Watch the temperature rise, while stirring, and wait for it to cap. Record temperature change. Try to calculate the kcal/mole for the reaction. It should be around .5.
Clean out your cup using standard laboratory cleanup procedures.
References
All images and information was drawn from myself or the following sources:Chang, Chemistry, 8th Edition
Zumdahl, 4th Edition