Integrated SequencePhysics Chemistry Organic Biology

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Special points of emphasis

Work, Energy, and Power


Intermolecular Forces

Heat and Temperature

The Ideal Gas and Kinetic Theory

The First Law of Thermodynamics


The States of Matter

Heating curves are typically used to show the process of phase change. Heating curves represent the relationship between heat flow and temperature. The heat which flows into the material resolves itself as some combination of internal energy increase and work.

When heat is added to a solid below its melting point, the temperature begins to rise. Rising temperature means that the average kinetic energy of the particles is increasing (although with greater range of vibration along lines of intermolecular force, the electrostatic potential energy component is also increasing).

If the solid is heated at its melting point, the heating curve shows that the temperature remains constant until the solid has melted. The melting process requires energy. The enthalpy change of melting is called the heat of fusion, and does not represent an increase in the kinetic energy of the particles, hence no temperature increase. The heat flow is allowing the particles to escape from the rigid intermolecular binding of the solid state into the less tight arrangement of the liquid. When mutually attracting charged particles (the polarities which lead to intermolecular force) are moved apart from one another, the electrostatic potential energy of the system is increased. The heat flow at the melting point increases the internal energy. In other words, at the melting temperature, it is not the kinetic energy of the particles but the electrostatic potential energy between the particles that increases.

After all of the solid has melted, heating the liquid raises its temperature until the boiling point is reached. Head added to the liquid at its boiling point is absorbed as the heat of vaporization as the liquid boils at constant temperature. Similar to the heat of fusion, the internal energy increase brought about by the heat of vaporization is not increasing the average kinetic energy of the particles during the process of vaporization but the electrostatic potential energy increases as the particles are separated each from each other. Remember, though, that the heat flow is not only increasing the internal energy in vaporization. The phase change from liquid to gas is accompanied by a large change in the volume of the system, so as the heat flows in, it must also perform pressure-volume work.

After all the liquid has been vaporized, and the system is entirely gaseous, the addition of more heat then raises the temperature of the gas. At that point, the internal energy increase corresponding to heat flow is all dedicated to increasing particle kinetic energy.


The States of Matter

Chemical Thermodynamics and the Equilibrium State

At the boiling point, the liquid and gas phases are in equilibrium. What does that mean? Let's take a moment to go into a bit of depth.

The equilibrium balance of liquid and gas occurs when there is an equal probability for the distribution of matter and energy for the system and surroundings, an equal probability for heat to flow in to form the vapor or heat to flow out and result in condensation. This is another way of saying that the free energy of the liquid and gaseous phasese are equal. Under the precise thermodynamic conditions of pressure and temperature at the boiling point neither state is favored. At the equilibrium point, there is an even chance that what occurs here is being undone over there. If energy flows this way, it is just as likely to flow the other way. For the boiling liquid, the randomness of gas molecules makes the formation of gas likely to happen, but there is also likelihood in heat dissipating out into the surroundings when gas condenses to form liquid. At the boiling point, the conditions of temperature and pressure make each direction equally probable.

But if you increase the pressure, you increase the free energy of the gas relative to the liquid, and so condensation occurs. The change occurs in the direction that leads to a free energy decrease. What happened when you increased the pressure, actually, is that you increased the stakes of the pressure-volume work in condensation. This means that the compression will correspond to a larger negative enthalpy change.

Similarly, increase the temperature and you will increase the free energy of the liquid relative to the gas, and so the vapor is favored at the higher temperature. At the higher temperature, you decresed the the entropy stakes of heat flow (remember, the entropy change due to heat flow, ΔS = ΔH/T). At higher temperature, equilibrium shifts more toward the endothermic direction.

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