Liquids, Solids, Thermodynamics, and Phase Transitions

Comprehensive flashcards covering key concepts from a Chemistry lecture on liquids, solids, thermodynamics, and phase transitions. These cards are designed to help students review and prepare for exams.

Solid

Definite shape and volume, fixed position, vibrate due to energy, closely packed.

Liquid

Definite volume, variable shape, short-ranged ordering, ability to move around, packed together.

Gas

Variable shape and volume, random ordering, complete freedom of motion, large amounts of space, can expand/contract.

Volume

Amount of 3D space occupied.

Temperature

Measure of possessed kinetic (thermal) energy.

Kinetic Molecular Theory (KMT)

State of matter depends on the amount of kinetic energy and the strength of attractions between particles.

More Kinetic Energy (KE)

More motion of particles and more freedom possible.

More motion

Higher temperature

Ideal Gas

Particles have complete freedom of motion.

Liquids (KMT)

Limited translational freedom; some particles have enough KE for rotational/vibrational freedom, some do not.

Solids (KMT)

No translational or rotational freedom, only vibrational.

Intermolecular Forces (IMFs)

Attractive forces between opposite charges (Formal/Informal).

Ion-Ion

Cation (+) and anion (-).

Dipole-Dipole

Polar molecules with partial charges (δ+ and δ-).

London Dispersion Forces

Temporary dipole of a nonpolar molecule due to fluctuations in electron distribution.

Hydrogen Bonding

Special type of dipole-dipole interaction involving hydrogen bonded to a very electronegative atom (O, N, F).

Phase Changes (Increasing Energy)

Melting/Fusion (s -> l), Sublimation (s -> g), Boiling/Vaporization/Evaporation (l -> g).

Phase Changes (Decreasing Energy)

Solidification/Freezing (l -> s), Condensation (g -> l), Deposition (g -> s).

Vaporization

Occurs at the surface of a liquid; molecules with enough KE overcome IMFs.

Vapor Pressure

Amount of molecules in gas phase above its liquid, dependent on molecule and temperature.

Condensation

Some gas molecules lose energy due to collisions and are recaptured by the liquid.

Open Container (Vaporization/Condensation)

Vapor escapes and spreads out, rate of vaporization > rate of condensation; net loss of liquid.

Closed Container (Vaporization/Condensation)

Vapor cannot spread out; at equilibrium, rate of vaporization = rate of condensation.

Volatile Liquids

Liquids that evaporate easily.

Nonvolatile Liquids

Hard to evaporate.

Boiling Point

The temperature at which a liquid's vapor pressure equals the external pressure.

Enthalpy

Thermal energy; enthalpy of vaporization.

Sublimation

Solid to gas; no liquid stage.

Deposition

Gas to solid.

Specific Heat

Related to heat; resistance to change in temperature; molecular and phase-specific.

Phase Diagram

Plot of pressure vs. temperature; experimentally determined.

Triple Point

Pressure and temperature where all 3 states are in equilibrium.

Critical Point

Pressure and temperature above which a substance exists as a 'Supercritical Fluid'.

Thermodynamics

Study of energy and its relationship to heat and work.

Zeroth Law of Thermodynamics

If system A is in equilibrium with system B, and system B is in equilibrium with system C, then A is also in equilibrium with C.

First Law of Thermodynamics

Energy cannot be created nor destroyed, only changed from one form to another.

Calorimetry

Experimental application of 1st law.

Bomb Calorimetry

Holds volume constant; ΔV = 0; ΔU = q = mcΔT.

Enthalpy

H = U + PV; enthalpy = heat at constant pressure; ΔH = q.

Standard State

State of a material at a defined set of conditions (pure gas/solid/liquid at 1 atm, specified temperature, most stable form).

Standard Enthalpy of Change

Enthalpy change when all reactants and products are in standard states.

Standard Enthalpy of Formation

Enthalpy change for a reaction forming 1 mol of a pure compound from its constituent elements in their standard states.

Hess' Law

Any reaction can be written as the sum of formation reactions or their reverses; ΔHrxn = ΣnΔHf(products) - ΣnΔHf(reactants).

Reversible Process

A process that stays in equilibrium the entire time.

Irreversible Process

A process that occurs because of some change (immediate, sudden, etc.).

Spontaneity

Whether a process will occur naturally under a given set of conditions.

Spontaneous Process

Occurs without input of energy.

Nonspontaneous Process

Requires input of energy/change in conditions.

Clausing

Discovered amount of 'reversible heat' (qrev) was related to Temperature (T) and some other quantity, Entropy (S).

Boltzmann

Statistical model of Entropy (probability).

Microstates

Specific configurations of all possible locations and/or energies of atoms/molecules in a system.

Entropy

Disorder, Chaos.

Second Law of Thermodynamics

Entropy of an isolated system must always increase; any isolated system evolves towards equilibrium where entropy is maximized.

Third Law of Thermodynamics

Entropy of a system approaches a constant value as Temperature approaches absolute zero (0 Kelvins).

State Functions

Only final and initial states are considered.

Dissolution

Energy dispersion; less contained particles, more (w, Г).

Gibb's Free Energy

Difficult to determine spontaneity because highly temperature/condition dependent; requires entropy of system and surroundings.

Gibb's Free Energy equation

G = H - TS; energy available for process to occur.

Equilibrium (AG = 0)

Point where a process goes from non-spontaneous to spontaneous.

Calculating Standard Free Energy

Use ΔH° and ΔS° (and T) to find ΔG°.