Thermodynamics and Thermochemistry for the DAT
/Learn key DAT concepts about thermodynamics and thermochemistry, plus practice questions and answers
Learn key DAT concepts about thermodynamics and thermochemistry
Table of Contents
Part 1: Introduction to thermodynamics and thermochemistry
Part 2: Principles of thermochemistry
a) Laws of thermodynamics
b) Endothermic and exothermic reactions
c) Spontaneous and nonspontaneous reactions
d) Gibbs free energy
Part 3: Systems and processes
a) Types of systems
b) Types of processes
Part 4: Heat transfer
a) Heat transfer
b) Hess’s law
Part 5: High-yield terms
Part 6: Questions and answers
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Part 1: Introduction to thermodynamics and thermochemistry
Thermodynamics and thermochemistry are fundamental branches of physical chemistry that govern the behavior of energy and matter in chemical systems. Thermodynamics deals with the study of energy transformations and the relationships between heat, work, and the properties of systems. It provides insights into the direction and extent of chemical reactions, as well as the efficiency of energy conversion processes. Thermochemistry, a subset of thermodynamics, focuses specifically on the heat changes associated with chemical reactions. It involves the measurement and calculation of enthalpy changes, heat capacities, and other thermodynamic properties to understand the heat flow accompanying chemical reactions and to predict reaction outcomes. Pay attention to the high-yield terms in bold, and prep for the DAT with exam-style practice questions and answers.
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Part 2: Principles of thermochemistry
a) Law of thermodynamics
There are three fundamental laws of thermodynamics to know for the DAT. In the realm of thermochemistry, a system encompasses the molecules, bonds, and atoms involved in a chemical reaction. A system’s environment is anything not participating in the reaction, such as a solvent, ambient air, or the room where the reaction occurs.
The zeroth law of thermodynamics asserts that if two thermodynamic systems are in thermal equilibrium with a third system, then all three systems are in thermal equilibrium with each other. In simpler terms, if systems A and B are in thermal equilibrium, and systems B and C are in thermal equilibrium, then systems A, B, and C are all in thermal equilibrium. This concept introduces temperature as a measure of heat, signifying that even if the three systems exchange energy, there would be no net heat exchange and all systems would have the same temperature.
Temperature and heat are distinct quantities. Temperature represents the average kinetic energy of a substance and is measured in Celsius (C), Kelvin (K), or Fahrenheit (F). Heat is the transfer of energy resulting from a temperature difference between substances and is measured in joules (J). Importantly, heat transfer is necessary for a change in temperature. Furthermore, heat is a process function, while temperature is a state function. A state function yields a property whose value remains constant regardless of the path taken to reach the final state. State functions are in contrast to process functions, where the property's value changes based on preceding steps.
The first law of thermodynamics, often called the law of energy conservation, states that energy is conserved. Energy cannot be created or destroyed but can only be transferred between objects and systems. For a closed system—one that can exchange energy but not matter—the first law of thermodynamics can be expressed as:
ΔU = Q-W
ΔU = change in system’s internal energyQ = heat added to the system
W = work done by the system
Thus, the internal energy of a system can be transferred into heat loss or gain, or into forms of work.
The second law of thermodynamics states that systems tend toward increasing entropy. This law introduces the concept of entropy as a measure of disorder. To better understand entropy, consider the relative entropies of gas, liquid, and solids. Gaseous molecules possess the highest entropy, followed by liquid, while solids have the lowest entropy.
Entropy can be described by the following equation:
ΔS = k * ln(W)
ΔS = change in entropyk = Boltzmann’s constant
ln(W) = natural logarithm of W
In this context, W does not signify work, but instead denotes the total count of potential microstates available to the system. A microstate encompasses every conceivable arrangement of particles' orientations within the system. Consequently, as the system's particle count increases, the number of microstates also increases, leading to a rise in system entropy. The Boltzmann constant, denoted by k or kB, is a fundamental constant employed in various thermodynamic formulas. While you don’t need to memorize this value, recognizing that it is a defined constant proves beneficial in understanding thermodynamic principles.
b) Endothermic and exothermic reactions
Enthalpy and entropy are fundamental concepts in thermodynamics, essential for understanding the energy changes that occur during chemical reactions. Enthalpy (denoted as H) represents the total heat content of a system, encompassing both internal energy and the energy required to maintain constant pressure. It quantifies the heat absorbed or released by a system at constant pressure, with exothermic reactions releasing heat (negative ΔH) and endothermic reactions absorbing heat (positive ΔH). Enthalpy changes are crucial indicators of a reaction's heat flow and can be determined experimentally through calorimetry, where the heat absorbed or released by a reaction is measured.
Entropy (denoted as S), as previously stated, is a measure of the degree of disorder or randomness within a system. It quantifies the number of microstates available to a system and reflects the system's tendency to move towards a more disordered state. Entropy increases with the dispersal of energy and matter, leading to greater randomness in the system. In chemical reactions, entropy changes (ΔS) play a vital role in determining a reaction's spontaneity, with spontaneous processes characterized by an increase in entropy (positive ΔS).
c) Spontaneous and nonspontaneous reactions
The relationship between enthalpy, entropy, and temperature is described by the Gibbs free energy equation. This equation helps predict whether a reaction will proceed spontaneously based on changes in enthalpy and entropy.
The Gibbs free energy equation is given by:
ΔG = ΔH - TΔS
ΔG is the change in Gibbs free energyT is the temperature of the system
ΔH is the change in enthalpy
ΔS is the change in entropy
It's essential to recognize that thermodynamic properties, such as Gibbs free energy, enthalpy, and entropy, are typically presented in relation to a reference value or a previous measurement. These properties are expressed as changes relative to this baseline value, allowing for comparisons and analysis.
The value of ΔG serves as a crucial determinant in assessing the spontaneity of a process—whether it's spontaneous, nonspontaneous, or in dynamic equilibrium. Exergonic reactions have a negative ΔG. This means the change in free energy is negative, making these reactions spontaneous. Conversely, endergonic reactions have a positive ΔG and are nonspontaneous.
Temperature alterations (T) have a significant impact on reaction spontaneity. An increase in temperature can make a nonspontaneous reaction spontaneous, particularly if the reaction is temperature-dependent. The sign of a reaction's free energy change may also hinge on temperature if both ΔH and ΔS of the reaction are either positive or negative. Since temperature is always measured in Kelvin for the Gibbs free energy equation, it's inherently positive (T > 0). These considerations are vital for understanding and predicting reaction behavior under various conditions.
If ΔH > 0 and ΔS < 0, then ΔG > 0. The reaction will be endergonic and will be non-spontaneous.
If ΔH < 0 and ΔS > 0, then ΔG < 0. The reaction will be exergonic and spontaneous.
However, if the signs of ΔH and ΔS are the same (both negative or both positive), the spontaneity or non-spontaneity of the reaction will depend on the magnitude of T. ΔG could be negative, equal to zero, or positive.
| ΔG | ΔH | ΔS | T |
|---|---|---|---|
Certain chemical reactions exhibit high favorability, resulting in an extremely negative value of ΔG (ΔG << 0). For example, the hydrolysis of ATP to produce ADP and inorganic phosphate (Pi) is highly favorable. The energy liberated during this reaction can be utilized to drive other chemical processes. In glycolysis, for instance, ATP hydrolysis is coupled with a nonspontaneous reaction to propel it forward.
Certain chemical reactions exhibit high favorability, resulting in an extremely negative value of ΔG (ΔG << data-preserve-html-node="true" 0). For example, the hydrolysis of ATP to produce ADP and inorganic phosphate (Pi) is highly favorable. The energy liberated during this reaction can be utilized to drive other chemical processes. In glycolysis, for instance, ATP hydrolysis is coupled with a nonspontaneous reaction to propel it forward.
d) Gibbs free energy
Frequently, you might come across two seemingly similar values that are actually quite distinct: ΔG and ΔG°. ΔG° signifies the change in Gibbs free energy under standard conditions. The symbol ° denotes that the given value is applicable under standard conditions, which typically entail a temperature of 298K, pressure of 1 atm, and reactant concentrations of 1 M. This value can be calculated as:
ΔG° = -RTln(K)
K is the equilibrium constant for the reactionR is the ideal gas constant 8.314 J mol-1 K-1
T is the temperature of the reaction
ΔG denotes the alteration in free energy under nonstandard conditions, implying any temperature besides 298 K, pressure other than 1 atm, and reactant concentrations differing from 1 M. If the non-standard conditions are specified, ΔG can be computed using the standard change in Gibbs free energy, ΔG°. The computation for the change in Gibbs free energy under non-standard conditions is as follows:
ΔG = ΔG° + RTln(Q) = -RT ln(K/Q)
Q is the reaction quotientK is the equilibrium constant
R is the ideal gas constant 8.314 J mol-1K-1
T is the temperature of the reaction
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