Thermodynamics Basics Quiz

Thermodynamics

• Thermodynamics studies how energy is transformed from one form to another and transferred from one place to another.
• A system is the region of interest (e.g., a flask of gas, a reaction mixture).
• The surroundings are everything else in the universe surrounding the system.

Systems, States, and Energy

• A system is the material or process being studied.
• The surroundings comprise everything else in the universe.
• Systems can be open (exchange mass and energy), closed (exchange energy only), or isolated (no exchange).

Work and Energy

• Work is the process of achieving motion against an opposing force.
• Work = force x distance moved.
• The unit for work is the joule (J).
• Energy is the capacity to do work (and ultimately, lift a weight).
• A system can perform expansion work (volume change) or nonexpansion work (e.g., a battery).

Work and Energy (internal energy)

• Internal energy (U) is the total energy stored in a system.
• ΔU = Ufinal - Uinitial.
• AU represents the change in internal energy.
• Energy transfer to a system as work increases internal energy (w is +); energy leaving the system as work decreases internal energy (w is -).
• The energy transferred as work (w) is denoted by w.
• When a system expands, w = -P ext ΔV, where P ext is the external pressure.
• If a system expands, w is negative.

Heat

• Heat is the exchange of thermal energy between a system and its surroundings due to a temperature difference.
• Heat flows from higher temperature to lower temperature.
• Thermal equilibrium is reached when both objects have the same temperature.
• The energy transferred as heat is denoted by q.
• Heat capacity (C) is a measure of how resistant a system is to changes in temperature.

Heat: Units

• The SI unit of energy is the joule (J).
• Calorie (cal) is also used, with 1 cal = 4.184 J exactly.
• Nutritional calorie (Cal) equals 1000 calories (kcal).

Heat exchange

• Heat transfer into a system increases internal energy (q is positive); heat transfer out of a system decreases internal energy (q is negative).
• An exothermic process releases heat, and an endothermic process absorbs heat.

Adiabatic and Diathermic Walls

• Adiabatic walls prevent energy transfer as heat.
• Diathermic walls allow energy transfer as heat.

Heat capacity

• Heat capacity is the amount of heat required to raise the temperature of a substance by a degree.
• Specific heat capacity is the amount of heat required to raise the temperature of one gram of a substance by 1 °C.
• Molar heat capacity is the amount of heat required to raise the temperature of one mole of a substance by 1 °C.

Specific Heat Capacity and Molar Heat Capacity

• Specific heat capacity (C): amount of heat to raise 1 g of a substance by 1°C.
• Units: J/g°C.
• Molar heat capacity (C m ): amount of heat to raise 1 mol of a substance by 1°C.
• Units: J/mol°C

Enthalpy

• Enthalpy (H) is a state function used for constant pressure processes.
• H = U + PV
• At constant pressure, the change in enthalpy (ΔH) is equal to the heat transferred (q).

The First Law

• The first law of thermodynamics states that the total energy of an isolated system remains constant.
• ΔU = q + w.

State Functions

• State functions depend only on the current state of a system, not on how that state was reached.
• Internal energy (U) and enthalpy (H) are state functions.

Work done by a system

• The work done by a system is not a state function because it depends on the path taken.

Heat is not a state function

• Heat is not a state function because it depends on the path taken.

Taking different paths

• The change in internal energy (∆U) for an isothermal reversible expansion of an ideal gas is zero.

Enthalpy of Physical Changes

• The enthalpy change for a phase change (e.g., vaporization, fusion) is equal to the heat supplied at constant pressure.
• Enthalpy of vaporization (ΔH vap ) and enthalpy of fusion (ΔH fus ) are positive values for processes that absorb heat.
• Enthalpy of condensation and freezing are negative values for processes that release heat.

Heating Curves

• A heating curve is a graph of temperature versus heat added.
• During phase changes, temperature remains constant.

The Heat Capacity of Materials

• Specific heat capacity measures how much heat energy raises the temperature of a substance.
• Materials with high heat capacity require more energy to raise temperature.

Heat of reaction

• Enthalpy of reaction is the heat absorbed or released in a reaction.

Enthalpy of Reaction

• The enthalpy change in a chemical reaction is an extensive property, meaning it depends on the amount of reactants.

Thermochemical Equations

• Stoichiometric coefficients in thermochemical equations indicate the number of moles reacting to give the indicated enthalpy change.

Measuring ΔH Calorimetry at Constant Pressure

• Calorimetry is used to physically measure the heat released or absorbed in a reaction.
• Constant pressure calorimetry (often a coffee-cup calorimeter) is used for reactions in aqueous solution.

Measuring heat

• Heat transfer is measured with a calorimeter, recording the temperature change.

The Measurement of Heat

• Heat transfer in calorimetry is proportional to the change in temperature of the calorimeter.

Bomb Calorimetry

• Bomb calorimetry measures heat transfer at constant volume.
• Internal energy (ΔU) change during a reaction is measured by measuring the temperature change.

Hess's Law

• Hess's Law states that the enthalpy change for an overall reaction is the sum of the enthalpy changes for the individual steps.

Standard Enthalpies of Formation

• By using the standard enthalpy of formation, we can calculate enthalpy change of any reaction.

Standard Enthalpy of Formation

• The standard enthalpy of formation, ∆H°f, of a substance is the standard enthalpy change per mole of formula units for the formation of that substance from its elements in their most stable forms.

Calculating Standard Enthalpy Change using Standard Enthalpies of Formation

• The change in standard enthalpy for a reaction is equal to the sum of the standard enthalpy of formations of the products minus the sum of the standard enthalpy of formations of the reactants.

The Born Haber Cycle

• The Born-Haber cycle is a method to calculate the lattice energy of an ionic substance.
• Hess's Law is used to add up the heats of other processes.

The Variation of Reaction Enthalpy with Temperature

• The reaction enthalpy depends on temperature and might change sign if the temperature changes.
• For exothermic and endothermic reactions

Entropy and Disorder:

• Entropy is a measure of disorder and randomness.
• Entropy of an isolated system increases in the course of any spontaneous change.
• Entropy changes accompanying changes in physical state of solids, liquids, and gases.
• The third law of thermodynamics states that the entropy of a perfect crystal at absolute zero (0 K) is zero

Standard Molar Entropies

• Standard molar entropies are tabulated for substances.
• Standard molar entropies help to calculate the entropy change of a reaction.

Global Changes in Entropy, Spontaneity

• Calculating the total entropy change of the universe, considering both changes in the system and surroundings, establishes reaction spontaneity.
• Spontaneity depends on both enthalpy and entropy changes.

Gibbs Free Energy

• Gibbs free energy (G) is a state function to determine spontaneity and direction of a reaction at constant temperature and pressure.

• AG = ΔH - TΔS.