When water freezes, the phase change that occurs is exothermic (DH = -6.02 kJ). Based on the change in enthalpy, you would expect that water would always freeze. Use the concepts of entropy and free energy to explain why this phase change is favourable only below o degrees Celsius
I would look up delta S for ice and for liquid water and use that in DG = DH -TdeltsS to show that difference between 273 K and other temperatures.
Calculate the standard Gibbs free energy
NH3(g) + HCl(G) --> NH4Cl
Ive tried but I keep getting a different answer
How do I get to the correc tanswer which is -91.2 kJ?
What numbers are you using. I looked up numbers and came up with -89.7 kJ. Close.
Balanced equation
NH3(g) + HCl(G) --> NH4Cl
deltaH = (1 mol)(-314.4 kJ/mol) - (1 mol)(-45.9 kJ/mol) plus (1 mol)(-92.3 kJ/mol)
=-176.2 kJ
deltas = (1 mol)(94.6) - (1 mol)(192.78) plus (1 mol)(-186.90)
= 0.00588
DG = DH - TDS
=176.2 kJ / 0.00588
=29995
I mean -176.2 -(278)(o.oo588)
= -177.8
Standard Gibbs free energy is DGfo.
I looked up delta Gfo and found
NH4Cl(s) = -201.5 kJ/mol
NH3(g) =-16.5
HCl(g) = -95.3
Using those numbers I get =-89.7. Check to see what those numbers are in your text table. I think you are calculating G when you want Gof
is this on a website?
It may be but I picked these values out of a freshman college chemistry text.
To understand why the phase change of water from liquid to solid (freezing) is favorable only below 0 degrees Celsius, we need to consider both entropy and free energy.
Entropy (S) refers to the measure of randomness or disorder in a system. In general, substances tend to increase their entropy, meaning they become more disordered, when they transition from a solid to a liquid or a gas phase. This is because in the solid state, particles are more closely packed and have a fixed arrangement, while in the liquid and gas states, they have more freedom to move and occupy different positions.
When water freezes, its phase change from liquid to solid results in a decrease in entropy. This is because, in the solid state, water molecules become more ordered, forming a regular lattice structure. As a result, the overall randomness or disorder of the system decreases, reducing the entropy.
On the other hand, free energy (G) represents the energy available to do useful work. It takes into account both the enthalpy (H) and entropy (S) of a system, using the equation: G = H - T*S.
When determining whether a phase change is favorable, we need to consider the free energy change (∆G) of the transition. A negative ∆G indicates a favorable or spontaneous process, whereas a positive ∆G suggests an unfavorable or non-spontaneous process.
For water to freeze, the enthalpy change (∆H) is negative (-6.02 kJ) since the process is exothermic. However, at temperatures above 0 degrees Celsius, the positive entropy change (∆S) dominates the equation, leading to a positive ∆G. This means that even though the enthalpy change favors the freezing process, the increase in entropy counteracts it, making the phase change unfavorable and non-spontaneous.
Below 0 degrees Celsius, the positive entropy change becomes smaller as the water molecules arrange themselves more rigidly in the solid state due to reduced thermal motion. At this point, the decrease in entropy due to the formation of the solid lattice is more significant than the increase in entropy caused by the lower temperature. As a result, the overall ∆G becomes negative, making the freezing process favorable and spontaneous.
Therefore, the combination of enthalpy and entropy effects, as described by the change in free energy, explains why water freezes only below 0 degrees Celsius, despite the exothermic nature of the phase change.