Chapter 3: Thermodynamics Flashcards
Thermodynamics definition
Thermodynamics is the study of the flow of energy in the universe, as that flow relates to
work, heat, entropy, and the different forms of energy. Classical thermodynamics concerns
itself only with observations that can be made at the macroscopic level, such as
measurements of temperature, pressure, volume, and work.
Zeroth Law of Thermodynamics
The zeroth law of thermodynamics is based on a simple observation: when one object is in
thermal equilibrium with another object, say a cup of warm tea and a metal stirring stick, and
the second object is in thermal equilibrium with a third object, such as your hand, then the
first and third object are also in thermal equilibrium. As such, when brought into thermal
contact, no net heat will flow between these objects. Note that thermal contact does not
necessarily imply physical contact, as objects can be in thermal contact across space.
Kelvin (K)
Absolute Zero: Theoretical temp, no energy transfer is 0 Kelvins
Freezing point of water is 273 K
Boiling point of water is 373 K
F= 1.8 C + 32
C= K- 273, C= (F - 32) /2
K= C+ 273
Thermal expansion (included mnemonic)
Length, volume, solubility, and even the conductivity of matter change as a function of temperature.
When the temperature of an object changes, its length changes a lot (αLΔT).
Thermal Expansion ΔL = αLΔT
where ΔL is the change in length,
α is the coefficient of linear expansion,
L is the original length,
and ΔT is the change in temperature.
The coefficient of linear expansion is a constant that characterizes how a specific material’s length changes as the temperature changes.
This usually has units of K–1, although it may sometimes be quoted as °C–1. This difference is inconsequential because the unit size for the Kelvin and Celsius scales is the same.
example:
A metal rod of length 2 m has a coefficient of linear expansion of 10−6 K−1. It is cooled from 1080°C to 80°C. What is the final length of the rod?
ΔL = αLΔT = (10−6 K)(2 m)(80 K − 1080 K) = −2 × 10−3 m
The negative sign represents a decrease in length. The original length was 2 m; therefore, the final length is 2 − (2 × 10−3) = 1.998 m
Volumetric Thermal expansion (expansion of liquids)
ΔV = βVΔT
where ΔV is the change in volume,
β is the coefficient of volumetric expansion,
V is the original volume,
and ΔT is the change in temperature.
The coefficient of volumetric expansion is a constant that characterizes how a specific material’s volume changes as the temperature changes. Its value is equal to three times the coefficient of linear expansion for the same material (β = 3α).
Closed, Isolated, and Open thermodynamic systems
A system is the portion of the universe that we are interested in observing or manipulating. The rest of the universe is considered the surroundings.
Isolated systems are not capable of exchanging energy or matter with their surroundings. As a result, the total change in internal energy must be zero. Isolated systems are very rare,
although they can be approximated. A bomb calorimeter attempts to insulate a reaction from the surroundings to prevent energy transfer, and the entire universe can be considered an isolated system because there are no surroundings.
Closed systems are capable of exchanging energy, but not matter, with the surroundings. The classic experiments involving gases in vessels with movable pistons are examples of closed systems. For thermodynamic purposes, most of what will be encountered on Test Day will be a closed system or will approximate a closed system
Open systems can exchange both matter and energy with the environment. In an open system, not only does the matter carry energy, but more energy may be transferred in the form of heat or work. A boiling pot of water, human beings, and uncontained combustion reactions are all examples of open systems.
State functions
State functions are thermodynamic properties that are a function of only the current equilibrium state of a system. In other words, state functions are defined by the fact that they are independent of the path taken to get to a particular equilibrium state.
The state functions include pressure (P), density (ρ), temperature (T), volume (V), enthalpy (H), internal energy (U), Gibbs free energy (G), and entropy (S).
On the other hand, process functions, such as work and heat, describe the path taken to get to from one state to another.
First law of thermodynamics
Essentially, the first law of thermodynamics states that the change in the total internal energy of a system is equal to the amount of energy transferred in the form of heat to the system, minus the amount of energy transferred from the system in the form of work. The internal energy of a system can be increased by adding heat, doing work on the system, or some combination of both processes. The change in internal energy is calculated from the equation
ΔU = Q − W
where ΔU is the change in the system’s internal energy,
Q is the energy transferred into the system as heat,
and W is the work done by the system.
Heat and the second law of thermodynamics
second law of thermodynamics: objects in thermal contact and not in thermal equilibrium will exchange heat energy such that the object with a higher temperature will give off heat energy to the object with a lower temperature until both objects have the same temperature at thermal equilibrium.
Heat, then, is defined as the process by
which a quantity of energy is transferred between two objects as a result of a difference in temperature.
The SI unit for heat is the joule (J), 1 Cal ≡ 103 cal = 4184 J = 3.97 BTU
One calorie (little c) is the amount of heat required to raise 1 g of water one degree Celsius. One Calorie (big C) is the amount of heat required to raise 1 kg of water 1 degree Celsius, equal to 1000 calories.
Heat Transfer
Conduction is the direct transfer of energy from molecule to molecule through molecular
collisions. As this definition would suggest, there must be direct physical contact between the objects.
Convection is the transfer of heat by the physical motion of a fluid over a material. Because convection involves flow, only liquids and gases can transfer heat by this means. In
convection, if the fluid has a higher temperature, it will transfer energy to the material
Radiation is the transfer of energy by electromagnetic waves. Unlike conduction and convection, radiation can transfer energy through a vacuum. Radiation is the method by
which the Sun is able to warm the Earth, and a microwave warms food.
Specific heat contains mnemonic
The specific heat (c) of a substance is
defined as the amount of heat energy required to raise one gram of a substance by one degree Celsius or one unit kelvin. For example, the specific heat of liquid water is one calorie per gram per unit kelvin or 4.184 J/g x K.
q = mcΔT looks a lot like “q equals MCAT.”
m is the mass,
c is the specific heat of the substance,
and ΔT is the change in temperature (in Celsius or kelvins)
Heat transformation
When heat energy is added to or removed from a system that is experiencing a phase change, the amount of heat that is added or removed cannot be calculated with the equation q = mcΔT because there is no temperature change during a phase change.
Instead, the following equation is used:
q = mL
where q is the amount of heat gained or lost from the material,
m is the mass of the
substance,
and L is the heat of transformation or latent heat of the substance. ( Heat of fusion, heat of vaporization given)
Phase changes
Solid to liquid: fusion or melting
Liquid to solid: freezing or solidification AT MELTING POINT= HEAT OF FUSION
Liquid to gas: boiling, evaporation, or vaporization AT BOILING POINT= HEAT OF VAPORIZATION
Gas to liquid: condensation
Solid to gas: sublimation
Gas to solid: deposition
Thermodynamic processes
These processes are:
isothermal (constant temperature, and therefore no change in internal energy),
adiabatic (no heat exchange),
isovolumetric (no change in volume, and therefore no work accomplished;
also called isochoric).
Isobaric processes are those that occur at a constant pressure, and are of less focus on the MCAT.
Entropy
Entropy is the measure of the spontaneous dispersal of energy at a specific temperature: how much energy is spread out, or how widely spread out energy becomes in a process.
ΔS= Qrev/ T
where ΔS is the change in entropy,
Qrev is the heat that is gained or lost in a reversible process,
and T is the temperature in kelvin.
The units of entropy are usually J/mol x K.
When energy is distributed into a system at a given temperature, its entropy increases.
When energy is distributed out of a system at a given temperature, its entropy decreases.
