ESSENTIAL Flashcards
(31 cards)
State the Law of Conservation of Energy.
Energy may be converted from one form to another, but it can neither be created nor destroyed (ΔE = q + w).
Classify the combustion of methane in terms of heat flow.
It is an exothermic process because heat is released from the system to the surroundings.
Provide one example of an endothermic reaction mentioned in the slide deck.
The formation of nitric oxide from nitrogen and oxygen gas is endothermic, requiring heat absorption.
What sign convention is used for q in exothermic versus endothermic reactions?
q < 0 for exothermic (heat leaves the system); q > 0 for endothermic (heat enters the system).
Relate bond strength to potential-energy differences between reactants and products in an exothermic reaction.
• In an exothermic reaction, energy is released to the surroundings, usually as heat.
• Reactants have weaker bonds and higher potential energy.
• Products have stronger bonds and lower potential energy.
• More energy is released when new bonds form than is absorbed to break the original bonds.
• This results in a net loss of potential energy from the system.
• The difference in energy is released as heat.
Explain the energy balance between bond breaking and bond forming in endothermic processes.
Energy is absorbed to break strong bonds in the reactants
New bonds formed in products are weaker, releasing less energy
More energy in than out → net energy gain
Overall increase in potential energy
How does the concept of enthalpy change (ΔH) capture these potential-energy differences?
ΔH equals the heat exchanged at constant pressure and reflects the net difference between bond-breaking (endothermic) and bond-forming (exothermic) energies.
Express enthalpy in terms of internal energy, pressure, and volume.
H = E + P V.
Why is enthalpy especially convenient for constant-pressure chemistry?
Under constant P, the heat absorbed or released equals ΔH directly, so no separate P ΔV work term is needed.
Give the thermodynamic definition of entropy.
Entropy (s) is a state function that quantifies the degree of disorder or energy dispersal in a system. Statistically, it is defined as:; S = kb ln W.
Can entropy change (ΔS) be negative? Under what circumstance?
Yes; ΔS < 0 when a system becomes more ordered and the number of accessible microstates decreases (e.g., gas → liquid). This is allowed as long as the surroundings gain enough entropy to satisfy the Second Law.
What qualitative ordering of phase entropies is always valid?
S₍solid₎ < S₍liquid₎ ≪ S₍gas₎.
State the Second Law succinctly.
In any spontaneous process the entropy of the universe increases (ΔS_univ > 0).
Why can energy be conserved while entropy increases?
The First Law tracks quantity of energy, whereas the Second Law tracks its quality or dispersal; energy can stay constant while its disorder (entropy) rises.
Give an everyday example illustrating the Second Law.
Hot coffee cooling as heat disperses into the room, raising ΔS_univ > 0.
State the Third Law.
The entropy of a perfect crystalline substance approaches zero as T → 0 K.
Explain why residual entropy can exist even at 0 K.
Molecular orientational disorder (molecules rotated/twisted in a different way) leaves multiple microstates, giving S > 0 despite 0 K.
How does the Third Law permit absolute entropies to be tabulated?
With S(0 K) defined as zero, integrating Cᵖ/T from 0 K upwards yields absolute S values.
Why does entropy rise continuously with temperature even in a single phase?
Increasing thermal agitation populates additional vibrational and rotational microstates.
Define Gibbs free energy.
Represents the maximum amount of usable energy available to do work in a system at constant T and P. Tells us whether a process is spontaneous or not. G = H − T S.
Explain why ΔG = 0 characterizes equilibrium.
At that point neither the forward nor reverse process is spontaneous; the system’s free energy is at its minimum.
Describe how ΔG connects enthalpy and entropy contributions.
ΔG = ΔH − T ΔS shows enthalpy drives at low T, entropy at high T; their balance sets spontaneity.
List the six fundamental phase transitions in order of increasing entropy change.
Deposition < Condensation < Freezing < Melting < Vaporisation < Sublimation.
How does phase favorability change with temperature in terms of Gibbs free energy?
A phase becomes stable when its ΔG relative to other phases is lowest; ΔG shifts sign as T ΔS grows larger than ΔH.
