Thermodynamics and Phase Diagrams Flashcards

Comprehensive practice flashcards covering Thermodynamics, Phase Equilibrium, and Metallurgy Principles based on the Learning Unit 1 transcript.

Define the Zeroth Law of Thermodynamics.

It states that if two systems are separately in thermal equilibrium with a third system, then the first two systems are in thermal equilibrium with each other.

According to the First Law of Thermodynamics, what is the relationship between internal energy, heat, and work?

The relationship is given by the equation $\Delta U = Q - W$, stating that energy can be converted between forms but cannot be created or destroyed.

What does the Second Law of Thermodynamics state regarding entropy in an isolated system?

The state of entropy of the entire universe, as an isolated system, will always increase in any natural and spontaneous process.

Define the Third Law of Thermodynamics in terms of absolute zero.

The value of entropy ($s$) of a completely pure crystalline structure is zero at absolute zero ($0\,kelvin$) temperature.

How is the change in entropy ($\Delta S$) defined in relation to heat transfer ($Q$) and temperature ($T$)?

$$\Delta S = \frac{\Delta Q}{T}$$

Rank the three states of matter (solid, liquid, gas) from lowest to highest entropy.

$$\text{Solid} < \text{Liquid} < \text{Gas}$$

Differentiate between stable and metastable equilibrium states.

Stable equilibrium exists when the object is in its lowest energy condition ($dG = 0$), while metastable equilibrium exists when additional energy ($\Delta G$) must be introduced before the object can reach true stability.

What is the criteria for unstable equilibrium?

Unstable equilibrium exists when no additional energy is needed before reaching metastability or stability.

State the fundamental equation for Gibbs free energy ($G$).

$G = H - TS$, where $H$ is enthalpy, $T$ is absolute temperature, and $S$ is entropy.

Define Enthalpy ($H$) in terms of internal energy ($E$), pressure ($P$), and volume ($V$).

$$H = E + PV$$

Why is Enthalpy ($H$) considered approximately equal to Internal Energy ($E$) in condensed phases?

In solids and liquids, the $PV$ term is usually very small, making $H \thickapprox E$.

What thermodynamic criterion must be met for any phase transformation to occur?

$\Delta G = G_2 - G_1 < 0$ (a decrease in Gibbs free energy).

What is the mathematical definition of equilibrium in a system at constant temperature and pressure?

$$dG = 0$$

Give a real-world example of stable and metastable equilibrium states according to the transcript.

Graphite (stable) and diamond (metastable) at room temperature and pressure.

What is the Gibbs’ phase rule for a single-component system?

$$F = C - P + 2$$

How is specific heat ($C_p$) at constant pressure defined?

$$C_p = \frac{dH}{dT}$$

What is the reference level for measuring Enthalpy ($H$) for a pure element?

$H = 0$ for a pure element in its most stable state at $298\,K$ ($25^{\circ}C$).

At the equilibrium melting temperature ($T_m$), what is the relationship between the Gibbs free energy of the solid and liquid phases?

$G_{solid} = G_{liquid}$ and $\Delta G = 0$.

Define the Clapeyron equation as described in the text.

A fundamental thermodynamic formula describing the relationship between pressure, temperature, and volume during a phase transition.

What is the driving force ($\Delta G$) for solidification when a liquid is undercooled by $\Delta T$ below $T_m$?

\Delta G \thickapprox \frac{L \times \text{\Delta} T}{T_m}, where $L$ is the latent heat of fusion.

What is the condensed Gibbs’ phase rule for most binary systems where pressure is constant?

$$F + P = C + 1$$

Define an ideal binary solution.

A solution where atoms are distributed randomly, the interchange of atoms causes no change in potential energy, $\Delta H_{mix} = 0$, and it obeys Raoult’s Law.

What does Chemical potential ($\mu$) measure?

It is a measure of the ‘escaping tendency’ of molecules.

According to the regular solution model, what does an interaction parameter $\Omega < 0$ signify?

Exothermic behavior where atoms prefer unlike neighbors (stronger A–B bonds), leading to infinite mutual solubility.

Differentiate between short-range order (SRO) and long-range order (LRO).

SRO involves a preference for A–B bonds over a local random distribution ($\Omega < 0$), while LRO involves atom sites becoming non-equivalent (A-sites and B-sites) at specific atom ratios, forming a superlattice.

What defines an Intermediate phase?

A solid solution that has a crystal structure different from either of the pure components A and B.

What occurs to LRO structures as temperature increases toward a critical point?

The degree of order decreases until, above some critical temperature, long-range order is lost completely.

What are Laves phases and what governs their formation?

They are phases based on $MgCu_2$, $MgZn_2$, and $MgNi_2$ where relative atomic size (ratio of $1.1\text{--}1.6$) allows atoms to fill space most efficiently.

What defines 'electron phases' such as $\alpha$ and $\beta$ brasses?

Phases where stability is determined by the relative valency and the number of valency electrons per unit cell.

In which type of intermediate compound does electronegativity play a dominant role?

Normal valency compounds, such as $Mg_2Sn$, where ionic bonds are formed.