What is the total change in entropy when 10.0 g of ice melts at 25.0°C? DHfus = 6.01 kJ/mol
The answer is 1.03 J/K but idk how to find it.
Well, let's start by breaking the ice... I mean, the problem down.
To find the total change in entropy, we need to consider both the heat transferred and the temperature change.
First, we need to convert the mass of ice into moles. We can do this by dividing the mass by the molar mass of water. If I'm not mistaken, the molar mass of water is about 18.015 g/mol.
So, 10.0 g of ice is approximately 10.0 g / 18.015 g/mol ≈ 0.555 mol.
Next, we want to find the heat transferred, which is given as ΔHfus (the enthalpy of fusion), equal to 6.01 kJ/mol. But hold your horses, we need to convert that to J/mol because entropy is in joules per kelvin (J/K).
6.01 kJ/mol = 6.01 x 10^3 J/mol.
Now, let's plug in the values into the formula:
ΔS = ΔHfus / T.
ΔS = (6.01 x 10^3 J/mol) / (25.0 + 273.15 K)
Now, let's calculate this... drumroll, please.
ΔS ≈ 1.03 J/K.
And there you have it, the total change in entropy when 10.0 g of ice melts at 25.0°C is approximately 1.03 J/K. Keep cool!
To find the total change in entropy when ice melts, you can use the equation:
ΔS = q / T
where ΔS represents the change in entropy, q is the heat transferred, and T is the temperature.
First, we need to calculate the heat transferred during the fusion of the ice. The heat transferred (q) can be determined using the equation:
q = n * ΔH
where n is the number of moles and ΔH is the enthalpy of fusion.
To find the number of moles, we can use the molar mass of water, which is approximately 18.015 g/mol.
Number of moles = mass (in grams) / molar mass
Number of moles = 10.0 g / 18.015 g/mol
Now we can calculate the heat transferred:
q = (10.0 g / 18.015 g/mol) * 6.01 kJ/mol
Next, we need to convert the heat transferred from kJ to J:
q = (10.0 g / 18.015 g/mol) * (6.01 kJ/mol * 1000 J/kJ)
Now that we have the heat transferred (q), we can calculate the change in entropy:
ΔS = q / T
However, in this case, we are not given the final temperature at which the ice melts. So we cannot directly calculate the change in entropy using the given values.
Hence, without knowing the final temperature, it is not possible to find the exact value of the total change in entropy.
To find the total change in entropy, we can use the following equation:
ΔS = ΔH_fus / T
Where:
ΔS is the change in entropy
ΔH_fus is the molar enthalpy of fusion (given as 6.01 kJ/mol)
T is the temperature in Kelvin (to support SI units and absolute temperature scale)
First, we need to convert the molar enthalpy of fusion from kJ/mol to J/mol:
6.01 kJ/mol * (1000 J / 1 kJ) = 6010 J/mol
Next, we need to determine the number of moles of ice being melted. To do this, we will use the molar mass of water (18.015 g/mol):
10.0 g * (1 mol / 18.015 g) ≈ 0.555 mol
Now, we convert the temperature from Celsius to Kelvin:
25.0°C + 273.15 K = 298.15 K
Finally, we substitute the values into the equation to calculate ΔS:
ΔS = (6010 J/mol) / (0.555 mol) / (298.15 K) ≈ 1.03 J/K
The total change in entropy when 10.0 g of ice melts at 25.0°C is approximately 1.03 J/K.
H2O(s) ==> H2O(l)
dG = dHfus - TdS(fus)
dG = 0 since the system 9s at equilibrium
dH fusion you know from the tables. I think is is about 334 J/g.
T is 273.15, solve for dS.