HENRY'S LAW
I did the first one im not sure how to do the 2 and the 3rd question.
1. What is the Henry's law constant for
CO2 at 20{C}?
C=kP
K=c/p
= 3.70*10^-2/1.00atm= 0.037
2. What pressure is required to achieve
a CO2 concentration of 6.00×10−2 at
20{C}?
Given: C (mol/L)= 6.00×10−2
T(C)= 20C
3. At 1 atm, how many moles of CO2 are
released by raising the temperature
of 1 liter of water from 20{C} to 25
{C}?
Given: P(atm)= 1.00
k( mol/L.atm)= 3.40×10−2
T= 25C
You don't give all of the information but I assume you looked up the k for CO2 correctly for k at 20 degrees C.
For #2, won't that be just
C=kP
P = k/c . Substitute for k and c and calculate P? Right? or did I miss something?
For #3.
You know C in #2. You can convert this to moles since you know the molarity. M = moles/L and you know the volume is 1 L.
Now you have a new k for a new temperature of 25 degrees C. Using 1 atm P, calculate a new C and a new moles. Then subtract moles in #2 from moles in #3 for the difference. That is the amount CO2 released at the elevated T. Check my work.
thank you i got the answers
2. To find the pressure required to achieve a CO2 concentration of 6.00×10^(-2) at 20°C, you can use Henry's law equation: C = kP.
Given:
C = 6.00×10^(-2)
T = 20°C (which is 293.15 K)
k = 0.037 (from the previous question)
Rearrange the equation to solve for P: P = C/k
Substitute the given values into the equation:
P = (6.00×10^(-2)) / 0.037
Calculating the value, you will get:
P ≈ 1.62
Therefore, the pressure required to achieve a CO2 concentration of 6.00×10^(-2) at 20°C is approximately 1.62 atm.
3. To find the number of moles of CO2 released by raising the temperature of 1 liter of water from 20°C to 25°C at 1 atm, you can use Henry's law equation: C = kP.
Given:
P = 1.00 atm
k = 3.40×10^(-2)
T1 = 20°C (which is 293.15 K)
T2 = 25°C (which is 298.15 K)
V = 1 liter
First, use the equation to find the CO2 concentration at the initial temperature:
C1 = kP = (3.40×10^(-2))(1.00) = 3.40×10^(-2) mol/L
Next, use the combined gas law to calculate the final CO2 concentration at the new temperature:
C2 = C1 × (T2/T1) = (3.40×10^(-2))(298.15/293.15)
Now, multiply the final concentration by the volume to calculate the number of moles of CO2 released:
moles = C2 × V
Substitute the values and calculate:
moles ≈ (3.40×10^(-2))(298.15/293.15)(1)
The final result will give you the number of moles of CO2 released.
To solve the second and third questions, we need to use Henry's law equation: C = kP, where C is the concentration of CO2 in mol/L, k is the Henry's law constant, and P is the partial pressure of CO2 in atm.
Let's start with the second question:
2. What pressure is required to achieve a CO2 concentration of 6.00×10^−2 at 20°C?
Given: C (mol/L) = 6.00×10^−2
T (°C) = 20°C
We know that the Henry's law equation is C = kP. Rearranging the equation, we have P = C/k.
Substituting the given values, we have P = 6.00×10^−2 / k.
To find the pressure, we need to know the value of k for CO2 at 20°C. Unfortunately, the value of k is not provided in the question. You may need to refer to a table or obtain the value from a reliable source. Once you have the value of k, substitute it into the equation to find the pressure (P).
Now, let's move on to the third question:
3. At 1 atm, how many moles of CO2 are released by raising the temperature of 1 liter of water from 20°C to 25°C?
Given: P (atm) = 1.00
k (mol/L.atm) = 3.40×10^−2
T = 25°C
In this question, we need to calculate the number of moles of CO2 released due to the increase in temperature. We can use the same Henry's law equation: C = kP.
Given that the initial concentration is 0 (since there is no CO2), we can rewrite the equation as: ΔC = kΔP, where ΔC is the change in concentration and ΔP is the change in partial pressure.
Since the pressure is constant at 1 atm, ΔP = 0. Subtracting the initial concentration (0) from the final concentration, we have ΔC = C - 0 = C.
Therefore, ΔC = C = kΔP = k * 0 = 0.
This means that no CO2 will be released by raising the temperature of 1 liter of water from 20°C to 25°C when the initial pressure is 1 atm.