Thermodynamics Flashcards
Topic 4, Lectures 19-22 - Tom Slater
Thermodynamics uses
Determines whether a reaction will take place
How much useful energy a reaction can generate
The ratio of reactants:products at equilibrium
Where it applies
Classical thermodynamics only deals with large systems and doesn’t apply at the level of individual molecules because it is dependent on the statistics of molecular distribution.
What can’t you determine with thermodynamics?
How fast reactions will occur/reach equilibrium
First law of thermodynamics
Energy cannot be created or destroyed, only converted between forms.
Zeroth law of thermodynamics
If two systems are in thermal equilibrium with a third system, they are also in equilibrium with each other. A=C, B=C, A=B. Temperature is a reliable way of measuring thermal balance
Thermal equilibrium
Where two thermal systems in contact stop exchanging heat, and have the same temperature.
Systems and surroundings
Chemical reactions represent the systems and anything external is the surroundings. Systems can be open,closed and isolated.
Open systems
Systems in which both matter and energy can be exchanged, e.g., a lidless pot of boiling water, the steam (matter) escapes and heat (energy) is released into the atmosphere.
Closed systems
Systems in which only energy can be exchanged, e.g., a reaction in a vessel with a lid, only heat (energy) is released but matter remains in the vessel.
Isolated systems
Systems in which neither matter nor energy are exchanged, e.g., a completely insulated vessel of boiling water. Though no system is perfectly isolated in real life.
Heat
Energy transfer that causes disorderly molecular motion. Heat transfer from system to surroundings causes random motion.
Work
Energy transfer that causes organised molecular motion. Work done causes motion of molecules in the same direction.
First law equation (closed system)
∆𝑈 = ∆𝑞 + ∆𝑤
The change in energy of a system is determined by the thermal energy added to the system plus the work done to a system.
Internal energy
The total kinetic and potential energy of the constituents of a system, rather than the whole system.
Internal energy of earth would account for the energy and motion of all things on earth but would not account for the energy associated with its motion around the sun.
Extensive variables
Depends on the amount of material (grams, moles, etc.)
Examples: heat, work, energy, mass, volume
Intensive variables
Independent of amount of material
Examples: pressure, temperature
Work done equation
∆𝑤 = − |𝐹 |∆𝑑
When distance is in the opposite direction to the force
The value of distance becomes negative
Expansion work
Ignoring the effects of friction and gravity, the expansion of gases can affect pressure and temperature. This has implications in areas such as pistons.
State functions
Properties of a system that depend only on the state of the system, as opposed to the path taken to reach that state.
Examples: internal energy, enthalpy, entropy, volume, pressure, gibbs free energy.
Process functions
Features of a system that depend on the path taken to transition between states. e.g. work done, heat. For example, the work done by expanding gas depends on whether pressure or temperature was constant during the expansion.
Maximising work
∆𝑤 = − 𝑉𝑖/𝑉𝑓 (𝑝𝑒𝑥𝑡𝑑𝑉)
Shows that most work is done when the external pressure is as close to the internal pressure as possible. Work done is also maximised when pressure is reduced gradually (infinitessimal changes.)
Work done—ideal gas at constant temperature
-𝑛𝑅𝑇(𝑙𝑛 𝑉𝑓/𝑉𝑖) calculates the maximum amount of work done
Expansion work equation
∆𝑤 = −𝑝𝑒𝑥𝑡∆𝑉
Heat capacity
Tells us the amount of energy required to raise the temperature of a material by a fixed amount.
C = dq/dT
dq = heat energy supplied
dT = change in temperature
When would expansion work done be 0?
When volume is constant, therefore ΔU=Δq+Δw becomes ΔU=Δq
Measuring internal energy change
Say a bomb calorimeter contained a sample of high-pressure O2 and heat released is absorbed by a surrounding water bath
Measure the temperature change of the bath after ignition, using an extrapolated cooling curve to allow for heat loss to the surroundings.
Heat obtained using HC of the whole system
Bomb calorimeter
A device used to measure heat changes at constant volume.
Heat capacity at constant volume
Tells us the change in internal energy needed to change the temperature by a certain amount. Cv = (Du/Dt)v (partial derivative as volume is constant)
Internal energy and heat capacity of ideal gases
Given that the Ek of one particle is 3/2KbT, and that Kb = R/avogadros, the Ek of n moles = 3/2nRT.
At constant volume = 3/2nR
Enthalpy
A quantity that is easier to measure in standard reactions than internal energy due to most chemical reactions being carried out at atmospheric pressure.
H = U+pV
Enthalpy change at constant pressure
At constant pressure, change in H = change in q, meaning enthalpy change is equal to energy supplied as heat.
Exothermic Δ𝐻 < 0
Endothermic Δ𝐻 > 0.
Standard states
The form of a substance at one bar of pressure and 298K
Enthalpy of non-standard conditions
calculated using H = 𝐻⦵ + corrections
Enthalpy of formation
Enthalpy change associated with forming a molecule from its constituent elements in standard states.
Enthalpy of formation for elements
zero
Enthalpy change of vaporisation
The energy change associated with forming a gaseous molecule from its liquid form
Enthalpy of combustion
The energy change associated with the burning of one mole of a species under standard conditions
Enthalpy of solution
Hess’s law
H is a state function, so the enthalpy of a given reaction can be obtained as the sum of standard enthalpies for a sequence of reactions that lead from the same reactants to the same products.
Calculating enthalpy change using Hess’s law
reactants -> standard-state-elements -> products
Δ𝐻⦵= Δ𝑓𝐻⦵[𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠] − Δ𝑓𝐻⦵[𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠]
Born-Haber cycles
A closed path of transformations starting and ending at the same point. Used generally when the enthalpy change or one pathway has an unknown value.
Volumetric enthalpy change
Enthalpy per unit volume
Gravimetric enthalpy change
Enthalpy per unit mass
Second law of thermodynamics
The entropy of an isolated system increases in the course of a spontaneous change ie: the entropy of an isolated system can never decrease. Heat flows naturally from hot to cold, unless work is done to reverse it.
Third law of thermodynamics
As temperature reaches absolute 0
(0K, -273C) the entropy of a perfect crystal also approaches 0. It is impossible to cool an object to absolute 0 though.
Entropy
A measure of how disordered a system is. Increases from solid to gas, increases with temperature.
Thermodynamic definition of entropy
dS(sys) = dq(reversible)/T
Change in entropy at constant T = integral of above equation = Δ𝑞𝑟𝑒𝑣/𝑇
When can entropy of a system decrease?
When the entropy of the surroundings increase by a greater amount.
Direction of heat flow
Hot to cold as the amounts of heat gained by separate objects are equal and opposite, and the magnitude of the entropy change for the hot object is less than that for the cold. Entropy increases as a result of direction.
Going from molar heat capacity to SI units
Multiply MCp by the number of moles of the species.
Difference between system and surroundings entropy
System is dependent on the reversibility whereas the surroundings are not. If there were no difference, they would always be at equilibrium.
Reversibility
Corresponds to performing infinitesimal changes under equilibrium conditions, ensuring maximum work is being done, thus minimising heat exchange.
Reversible reaction at constant temperature
dS = nRln(Vf/Vi) for system
Entropy change of surroundings would be the negative of the above equation under reversible conditions, otherwise greater.
Gibbs free energy equation
Δ𝑆𝑢𝑛𝑖𝑣𝑒𝑟𝑠𝑒 = Δ𝑆𝑠𝑦𝑠𝑡𝑒𝑚 + Δ𝑆𝑠𝑢𝑟𝑟𝑜𝑢𝑛𝑑𝑖𝑛𝑔𝑠
= Δ𝑆𝑠𝑦𝑠𝑡𝑒𝑚 − Δ𝐻𝑠𝑦𝑠/𝑇
(at constant pressure)
G = H - TS
Gibbs free energy values
When G = 0, change is reversible (at equilibrium)
Else, G < 0
How does volume affect G?
An increase in volume leads to a decrease in G. Due to
ΔG = -𝑇𝑛 𝑅 ln 𝑉1/𝑉0
How is G affected by temperature?
Always decreases as temperature increases.
How is G affected by pressure?
Always increases as pressure increases.
Mixing
Mixing substances increases the entropy of a system. However, the entropy of the surroundings may decrease due to the energy required to mix the substances.
Gibbs energy of mixed systems
Need to know how molar Gibbs energy changes (chemical potential)
With one substance:
𝜇 = 𝐺/𝑛 = 𝐺𝑚
Chemical potential with phase change
𝑑𝐺 = 𝜇𝐵𝑑𝑛 − 𝜇𝐴𝑑𝑛 = 𝜇𝐵 − 𝜇𝐴 𝑑𝑛
At equilibrium, mu is the same for all phases.
Mixture of ideal gases
𝐺 = 𝑛𝐴𝜇𝐴 + 𝑛𝐵𝜇𝐵
This works for ideal gases, as there is no energy of interaction. If there are inter-component interactions, muA changes with the composition of the mixture.
Gibbs-Duhem euqation
Sum𝐽 𝑛𝐽𝑑𝜇𝐽 = 0
Gibbs energy change on mixing
∆𝑚𝑖𝑥𝐺 = 𝐺𝑓𝑖𝑛𝑎𝑙 − 𝐺𝑖𝑛𝑖𝑡𝑖𝑎𝑙 = 𝑛𝐴𝑅𝑇𝑙𝑛(𝑝𝐴/𝑝) + 𝑛𝐵𝑅𝑇𝑙𝑛(𝑝𝐵/𝑝)
Sample quantities
The total amount of substance in a given system. Dependent on the size or amount of a system. E.g. mass, volume, total energy.
Molar quantities
Derived by dividing the sample quantity by the number of moles. Expressed as per mole of a substance. E.g. molar mass, molar volume and molar energy
