Use the experimental values to calculate the enthelpy change in the system.

• Frank665
In summary, the conversation discusses using experimental values to calculate the enthalpy change in a system, specifically involving the quantities of reactants and the initial and final temperatures. It is assumed that the specific heat of water is used to find the heat exchanged, as this is typically the case at constant pressure. The speaker also mentions not being able to provide computational assistance as this is a homework assignment.
Frank665
Use the experimental values to calculate the enthalpy change in the system.

Observation Experiment
Quantity of reactant 1 100.0 mL of 1.00 M KOH(aq)
Quantity of reactant 2 100.0 mL of 1.00 M HBr(aq)
Initial Temperature (oC) 20.0
Final Temperature (oC) 22.5

Frank665 said:
Use the experimental values to calculate the enthalpy change in the system.

Observation Experiment
Quantity of reactant 1 100.0 mL of 1.00 M KOH(aq)
Quantity of reactant 2 100.0 mL of 1.00 M HBr(aq)
Initial Temperature (oC) 20.0
Final Temperature (oC) 22.5

If your teacher didn't give you the (constant pressure, I assume) specific heats, it should mean he assumes you use the one of water, so you can easily find the heat exchanged (at constant pressure it's = enthalpy variation).
Of course I won't make the computations for you, this is an homework.

Enthalpy change is a thermodynamic property that measures the amount of heat energy released or absorbed during a chemical reaction. It can be calculated using the experimental values of the quantities of reactants and the change in temperature.

In this experiment, 100.0 mL of 1.00 M KOH(aq) and 100.0 mL of 1.00 M HBr(aq) were mixed, resulting in a temperature increase from 20.0 oC to 22.5 oC. To calculate the enthalpy change, we can use the formula:

ΔH = -q = -mCΔT

Where:
ΔH = enthalpy change (in J)
q = heat exchanged (in J)
m = mass of solution (in kg)
C = specific heat capacity of solution (in J/(g*K))
ΔT = change in temperature (in K)

First, we need to calculate the mass of the solution. This can be done by converting the volumes of the reactants to masses using their respective molar masses (39.1 g/mol for KOH and 80.9 g/mol for HBr) and the concentration (1.00 M):

Mass of KOH = (100.0 mL) * (1.00 mol/L) * (39.1 g/mol) = 3.91 g
Mass of HBr = (100.0 mL) * (1.00 mol/L) * (80.9 g/mol) = 8.09 g

Total mass of solution = 3.91 g + 8.09 g = 12.00 g

Next, we need to calculate the specific heat capacity of the solution. This can be done by taking the weighted average of the specific heat capacities of KOH and HBr:

C = [(3.91 g) * (4.18 J/(g*K)) + (8.09 g) * (4.18 J/(g*K))] / (12.00 g) = 4.18 J/(g*K)

Finally, we can plug these values into the formula to calculate the enthalpy change:

ΔH = -(12.00 g) * (4.18 J/(g*K)) * (22.5 oC - 20.0 oC) = - 100.08 J

Therefore, the enthalpy change in this system is

1. What is enthalpy change?

Enthalpy change is the amount of energy transferred in a chemical or physical process, usually measured in joules (J) or kilojoules (kJ).

2. How do you calculate enthalpy change?

To calculate enthalpy change, you need to use the equation: ΔH = Σ(Hproducts) - Σ(Hreactants), where Σ represents the sum of the enthalpies of the products and reactants.

3. What are experimental values?

Experimental values are the values obtained from conducting an experiment or performing a measurement in a controlled setting. These values may differ from theoretical values due to experimental error.

4. Why is it important to use experimental values in calculating enthalpy change?

Using experimental values allows for a more accurate calculation of enthalpy change, as these values are obtained through actual measurements and reflect the real conditions of the system. Theoretical values are based on ideal conditions and may not accurately represent the system.

5. Are there any limitations to using experimental values in calculating enthalpy change?

Yes, there are limitations to using experimental values in calculating enthalpy change. These include experimental error, uncertainties in measurements, and the assumption that the system is in a closed system and there is no energy lost to the surroundings.

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