Kaplan MCAT General Chemistry Review

Because systems can be in different equilibrium states at different temperatures and pressures, a set of standard conditions has been defined for measuring the enthalpy, entropy, and Gibbs free energy changes of a reaction. The standard conditions are defined as 25°C (298 K) and 1 atm. Don’t confuse standard conditions with standard temperature and pressure (STP), for which the temperature is 0°C (273 K) and 1 atm. This common MCAT Test Day mistake is easily made but also easily avoided. You’ll use standard conditions for thermodynamic problems of enthalpy or free energy, but you’ll use STP for ideal gas calculations.

MCAT Expertise

On the MCAT, be sure that you do not confuse standard conditions in thermodynamics with standard temperature and pressure (STP) in gas law calculations (see Chapter 7).

Under standard conditions, the most stable form of a substance is called the standard state of that substance. You should recognize the standard states for some elements and compounds commonly encountered on the test. For example, H₂ (g), H₂O (l), NaCl (s), O₂ (g), and C (s) (graphite) are the most stable forms of these substances under standard conditions. Recognizing whether or not a substance is in its standard state is important for thermochemical calculations, such as heats of reactions and, in particular, the heat of formation. The changes in enthalpy, entropy, and free energy that occur when a reaction takes place under standard conditions are called the standard enthalpy, standard entropy, and standard free energy changes, respectively, and are symbolized by H°, S°, and G°.

The rest of this chapter will focus on the following state functions: enthalpy, entropy (energy dispersal), and free energy, with some concluding remarks on spontaneity.

Heat

Before we can examine the first of the four state functions that are the focus of this chapter, we must address the topic of heat, which is a source of some confusion for students. Perhaps the greatest barrier to a proper understanding of heat is the semantic conflation of the terms heat and temperature. Many people use these terms interchangeably in everyday conversation. We might say, for example, that a midsummer afternoon was unbearably hot or that the temperature exceeded 100°F. Both convey the sense that the day was very, very warm, but what makes sense in everyday conversation needs to be clarified for a proper understanding of thermochemical principles. Temperature (T) is related to the average kinetic energy of the particles of the substance whose temperature is being measured. Temperature is the way that we scale how hot or cold something is (but not necessarily how hot or cold something feels to us: For reasons that will become clear in our discussion of specific heat, blacktop will feel hotter to our bare feet than a wooden boardwalk would, even if they are at the same temperature). We are familiar with a few temperature scales: Fahrenheit, Celsius, and Kelvin. The average kinetic energy of the particles in a substance is related to the thermal energy of the substance, but because we must also include consideration of how much substance is present to calculate total thermal energy content, the most we can say about temperature is that when a substance’s thermal energy increases, its temperature increases also. Nevertheless, we cannot say that something that is hot necessarily has greater thermal energy (in absolute terms) than a substance that is cold. For example, we might determine that a large amount of cool water has greater total heat content than a very small amount of very hot water.

Key Concept

Remember that heat and temperature are different. Heat is a specific form of energy that can enter or leave a system, while temperature is a measure of the average kinetic energy of the particles in a system.

Heat (Q) is the transfer of energy from one substance to another as a result of their difference in temperature. In fact, the zeroth law of thermodynamics implies that objects are in thermal equilibrium only when their temperatures are equal. Heat is therefore a process function, not a state function: We can quantify how much thermal energy is transferred between two or more objects as a result of their difference in temperatures by measuring the heat transferred.

The first law of thermodynamics states that the change in the total internal energy (U) of a system is equal to the amount of heat (thermal energy) transferred (Q) to the system minus the amount of work (W) (another form of energy transfer by the application of force through displacement) done by the system. This can be expressed mathematically as follows:

U = Q - W