Thermodynamics Flashcards
What is heat?
Heat is energy transferred as a result of temperature differences.
What is work?
Work is done as a result of motion against an opposing force.
What are heat and work both methods of?
Heat and work are both methods of transferring energy to the system.
What is an isolated system?
A system not connected thermally or mechanically to the surroundings.
What is internal energy?
Internal energy is the capacity to do work (U).
State the molecular origins of internal energy (3).
- Kinetic energy from moving molecules.
- Potential energy possessed by molecules.
- Attractive and repulsive forces exerted by molecules on one another, related to electrostatic force between two charged particles.
Molecules possess potential energy as a result of…(3).
- Attractions between nuclei and electrons.
- Chemical bonds.
- Intermolecular attractions.
What is internal energy (U)?
The sum of all kinetic and potential energy contributions.
When will internal energy change?
Internal energy will change if its volume is changed, as this changes the average separation between molecules.
What is the relationship between work and internal energy?
- If work is done BY a system, its internal energy DECREASES.
- If work is done ON a system, its internal energy INCREASES.
Explain the meaning of the +/- signs of q and w.
w: POSITIVE means that work is done ON the system.
NEGATIVE means that work is done BY the system.
q: POSITIVE means that heat is supplied TO the system.
NEGATIVE means that heat is released FROM the system.
Give 3 ways of stating the first law of thermodynamics.
- The internal energy of an isolated system is constant ∆U = 0.
- Energy can neither be created nor destroyed, only interconverted between forms.
- The total amount of energy in the universe is constant.
What is expansion work? Give an example.
Expansion work involves a change in the volume of the system i.e. a reaction with gases.
What is non-expansion work? Give an example.
Non-expansion work doesn’t involve a change in volume i.e. chemical reaction done in a battery.
What are the consequences of the first law of thermodynamics? (2).
- Energy is conserved, heat + work are equivalent forms of energy.
- Chemical reactions at constant volume have no expansion work, therefore w = 0 and heat is the ONLY contribution to ∆U. Therefore, ∆U = qv.
What is meant by the term ‘state functions’?
Pressure, volume and temperature depend only on the state of the system and are independent of how the change was brought about.
What is meant by the term ‘path functions’?
Work and heat depend on the path between states and depend on how the change was brought about.
What is an exothermic reaction?
They give out energy/heat, the reacting system loses energy.
∆H < 0, heat is released from the system at constant pressure.
What is an endothermic reaction?
They take in energy/heat, reacting system gains energy.
∆H > 0, heat is absorbed by the system at constant pressure.
What is meant by the term molar enthalpy of melting?
The enthalpy change which occurs when 1 mol of a SOLID melts to form a liquid.
What is meant by the term molar enthalpy of vaporisation?
The enthalpy change which occurs when 1 mol of a LIQUID boils to form a gas.
State Hess’s Law.
∆H = ∆H(products) - ∆H(reactants)
∆H = ∆U + p∆V
What is meant by the term standard enthalpy of formation?
The standard reaction enthalpy for the formation of a compound from its elements in their most stable form.
How do we calculate the standard reaction enthalpy?
∆rH = Zn∆fH - Zn∆fH
Z: sum of
n: number of moles of each reactant/product
What is bond enthalpy?
Bond enthalpy is equal to the standard enthalpy change for the dissociation of a molecule in the gas phase in atoms.
XY (g) –> X (g) + Y (g)
What are the features of a spontaneous reaction? (3)
- Once started will continue without any intervention.
- Will move towards their equilibrium state without any external influence.
- Can be exothermic, endothermic or neither.
What is meant by the term entropy?
A measure of the amount of disorder.
Entropy increases for a spontaneous process, making ∆S positive.
State the equation for ∆S.
∆S = Sf - Sin
f: final
in: initial
∆S: JK-1
State the second law of thermodynamics. (2)
- The entropy of an isolated system increases during any spontaneous process.
- Spontaneous processes are those that increase the total entropy of the universe.
Explain what the values of ∆S tell us.
∆S(total) > 0, the reaction is spontaneous.
∆S(total) < 0, the reaction is non-spontaneous.
∆S(total) = 0, the reaction is at equilibrium.
State the equation for ∆S when heat (qrev) is added reversibly.
∆S = qrev/T
How do we achieve reversible heat flow?
The temperature of the system needs to be at thermal equilibrium with the surroundings.
State the equation for ∆S for a phase transition.
∆S = ∆H/T
Outline the features of a phase transition (3).
- Temperature of a substance remains constant while a phase transition occurs.
- Heat is transferred reversibly.
- If the transition takes place at constant pressure, the heat supplied = enthalpy change ∆H.
State the Boltzmann formula and what is calculates.
S = kB lnW
kB: boltzmann constant (1.381 x 10-23 JK-1)
W: number of ways of arranging molecules and their energies.
The Boltzmann formula calculates absolute S values.
State the third law of thermodynamics.
The entropy of a perfect crystal is S=0 at T=0K.
Explain why we cannot have a perfect crystal.
Residual entropy arises.
No thermal motion.
Entropy arises from the number of possible orientations a molecule can take up.
State the equation for Gibbs free energy and what it measures.
∆G = ∆H - T∆S
A measure of the total entropy in the system and the surrounding.
What is the significance of a negative ∆G value?
Implies the reaction occurs spontaneously, favoured by a positive ∆S and a negative ∆H.
Draw the phase diagram and annotate.
See notes.
- Triple point: all three phase boundaries intersect and all three phases coexist. This point is different for every molecule.
- Critical point: above this, the molecule becomes a supercritical fluid.
State the Clausius-Clapeyron equation and explain.
ln (p2/p1) = -(∆vapH)/R (1/T2 - 1/T1)
This equation gives the form of the vapour-liquid boundary, where ∆G = 0.
State the term dynamic equilibrium.
When the rate of the forward reaction(s) equals that of the reverse reaction, with no net change in composition (∆G = 0).
How do thermodynamics and kinetics relate to equilibrium?
Thermodynamics: calculates the position of equilibrium and whether products or reactants will eventually dominate.
Kinetics: used to assess how quickly equilibrium is achieved.
Outline the relationship between equilibrium and ∆G.
-∆rG: more reactants than equilibrium mixture.
+∆rG: more products than equilibrium mixture.
∆rG = 0: at equilibrium.
When ∆G = 0, ∆rG = -RTlnK
What is meant by the symbol K?
K: equilibrium constant which is equal to the reaction quotient for the equilibrium composition.
Outline the relationship between K and ∆G.
-∆rG, K>1: PRODUCTS are favoured
+∆rG, K<1: REACTANTS are favoured
State le Chatelier’s principle.
When a change is made to a system in dynamic equilibrium, the equilibrium responds so as to minimise the effect of change.
State how the Haber process is arranged to give the maximum yield possible (2).
- Ammonia gas product is continuously removed by dissolving it in water to cause more products to form.
- Equilibrium responds to a pressure increase by decreasing the total amount of gas, this favours the products in the Haber process.
State the equation for the Haber process.
N2(g) + 3H2(g) –> 2NH3(g)
State the relationship between equilibrium and pressure (2).
- The value of the equilibrium constant does NOT depend on the overall pressure.
- The composition of the equilibrium mixture DOES depend on the overall pressure.
State the relationship between equilibrium and temperature (2).
- For ENDOthermic reactions, increasing temp (T) favours products as heat is absorbed.
- For EXOthermic reactions, decreasing temp (T) favours products as heat is released.
State the temperature dependence of equilibrium, van’t Hoff, equation.
ln (K2/K1) = -(∆rH)/R (1/T2 - 1/T1)
How does this relate to the graph of a straight line?
ln K vs. 1/T is a straight line if ∆rH is constant over the required temperature range.
The slope: -∆rH/R
How would this graph (ln K vs. 1/T) look for an EXOthermic reaction?
Positive slope.
K decreases as T increases.
Favours reactants.
How would this graph (ln K vs. 1/T) look for an ENDOthermic reaction?
Negative slope.
K increases as T increases.
Favours products.
How does a catalyst alter a reaction?
Increases the rate at which the reaction mixture reaches equilibrium, doesn’t alter the composition.
Lowers the temperature required for a significant ROR.
