Q: A student mixes 50.0 mL of 1.0 M HCl with 50.0 mL of 1.0 M NaOH in a calorimeter. The temperature rises from $$22.0\ ^\circ\mathrm{C} to$$ $$28.5\ ^\circ\mathrm{C}. $$Assuming the density of the solution is $$1.00\ \mathrm{g/mL}$$ and the specific heat is $$4.18\ \mathrm{J/(g\cdot{}^{\circ}C)}$$, what is the heat absorbed by the solution ($$q_{\mathrm{water}}) in kJ$$?

$$q_{\mathrm{water}} = 100.0\ \mathrm{g} \times 4.18\ \mathrm{J/(g\cdot{}^{\circ}C)} \times (28.5 - 22.0)\ ^\circ\mathrm{C} = 100.0 \times 4.18 \times 6.5 = 2717\ \mathrm{J} = 2.72\ \mathrm{kJ}$$

Q: A student dissolves $$2.00\ \mathrm{g} of $$NaOH(s) in $$100.0\ \mathrm{mL} of $$water. The temperature increases by $$5.0\ ^\circ\mathrm{C}. $$What is the enthalpy change per mole of NaOH ($$\Delta H) $$for this process? (Specific heat $$= 4.18\ \mathrm{J/(g\cdot{}^{\circ}C)}$$, molar mass of NaOH $$= 40.0\ \mathrm{g/mol}$$, assume density $$= 1.00\ \mathrm{g/mL})$$

$$q = 100.0\ \mathrm{g} \times 4.18\ \mathrm{J/(g\cdot{}^{\circ}C)} \times 5.0\ ^\circ\mathrm{C} = 2090\ \mathrm{J}$$; $$\mathrm{mol\ NaOH} = 2.00/40.0 = 0.0500$$; $$\Delta H = -2090/0.0500 = -41,800\ \mathrm{J/mol} = -41.8\ \mathrm{kJ/mol}$$

Question 1

Which of the following best describes the main purpose of using a Styrofoam cup in a calorimetry experiment?

Accepted Answer

To minimize heat loss to the surroundings, acting as an insulated container.

Question 2

A student mixes 50.0 mL of 1.0 M HCl with 50.0 mL of 1.0 M NaOH in a calorimeter. The temperature rises from $$22.0\ ^\circ\mathrm{C} to$$ $$28.5\ ^\circ\mathrm{C}. $$Assuming the density of the solution is $$1.00\ \mathrm{g/mL}$$ and the specific heat is $$4.18\ \mathrm{J/(g\cdot{}^{\circ}C)}$$, what is the heat absorbed by the solution ($$q_{\mathrm{water}}) in kJ$$?

Accepted Answer

$$q_{\mathrm{water}} = 100.0\ \mathrm{g} 	imes 4.18\ \mathrm{J/(g\cdot{}^{\circ}C)} 	imes (28.5 - 22.0)\ ^\circ\mathrm{C} = 100.0 	imes 4.18 	imes 6.5 = 2717\ \mathrm{J} = 2.72\ \mathrm{kJ}$$

Question 3

In the calorimetry experiment, which assumption is made about the density of the solution?

Accepted Answer

The density of the solution is assumed to be the same as the density of water, $$1.00\ \mathrm{g/mL}.$$

Question 4

Which of the following equations correctly represents Hess’s Law for the combination of the following reactions?

1. $$\mathrm{NaOH}(s) \rightarrow \mathrm{NaOH}(aq)\ \Delta H_1$
2. $$\mathrm{NaOH}(s) + \mathrm{HCl}(aq) \rightarrow \mathrm{NaCl}(aq) + \mathrm{H_2O}(l)\ \Delta $$H_2

To $$obtain:
3. $$\mathrm{NaOH}(aq) + \mathrm{HCl}(aq) \rightarrow \mathrm{NaCl}(aq) + \mathrm{H_2O}(l)\ \Delta H_3$

Accepted Answer

$$\Delta H_3$ = \Delta H_2$ - \Delta H_1$

Question 5

A student dissolves $$2.00\ \mathrm{g} of $$NaOH(s) in $$100.0\ \mathrm{mL} of $$water. The temperature increases by $$5.0\ ^\circ\mathrm{C}. $$What is the enthalpy change per mole of NaOH ($$\Delta H) $$for this process? (Specific heat $$= 4.18\ \mathrm{J/(g\cdot{}^{\circ}C)}$$, molar mass of NaOH $$= 40.0\ \mathrm{g/mol}$$, assume density $$= 1.00\ \mathrm{g/mL})$$

Accepted Answer

$$q = 100.0\ \mathrm{g} 	imes 4.18\ \mathrm{J/(g\cdot{}^{\circ}C)} 	imes 5.0\ ^\circ\mathrm{C} = 2090\ \mathrm{J}$$; $$\mathrm{mol\ NaOH} = 2.00/40.0 = 0.0500$$; $$\Delta H = -2090/0.0500 = -41,800\ \mathrm{J/mol} = -41.8\ \mathrm{kJ/mol}$$

Question 6

Why is the heat capacity of the calorimeter ($$q_{\mathrm{calorimeter}}) $$often neglected in introductory calorimetry experiments?

Accepted Answer

Because the heat absorbed by the calorimeter is assumed to be negligible compared to the heat absorbed by the solution.

Question 7

In Reaction 2, $$\mathrm{NaOH}(s) + \mathrm{HCl}(aq) \rightarrow \mathrm{NaCl}(aq) + \mathrm{H_2O}(l)$$, a student uses $$2.00\ \mathrm{g}$$ NaOH and $$50.0\ \mathrm{mL} of$$ $$1.0\ \mathrm{M}$$ HCl. What is the limiting reactant?

Accepted Answer

Both reactants are present in stoichiometric amounts; neither is limiting.

Question 8

A student performs Reaction 5: $$\mathrm{NaOH}(s) + \mathrm{CH_3COOH}(aq) \rightarrow \mathrm{NaCH_3COO}(aq) + \mathrm{H_2O}(l). If $$the measured $$\Delta H_5$ is -55.0$$\ \mathrm{kJ/mol}$$ and $$\Delta $$H_4 ($$dissolution of NaOH) is $$-44.0\ \mathrm{kJ/mol}$$, what is the calculated $$\Delta H_6$ for $$\mathrm{NaOH}(aq) + \mathrm{CH_3COOH}(aq) \rightarrow \mathrm{NaCH_3COO}(aq) + \mathrm{H_2O}(l)$$ using Hess’s Law?

Accepted Answer

$$\Delta H_6$ = \Delta H_5$ - \Delta H_4$ = -55.0 - (-44.0) = -11.0\ \mathrm{kJ/mol}$$

Question 9

Which of the following is a valid reason for rinsing the temperature probe with deionized water and drying it before each measurement?

Accepted Answer

To prevent contamination between different solutions and ensure accurate temperature readings.

Question 10

If the measured $$\Delta H_3$ for $$\mathrm{NaOH}(aq) + \mathrm{HCl}(aq) \rightarrow \mathrm{NaCl}(aq) + \mathrm{H_2O}(l)$$ is -57.0$$\ \mathrm{kJ/mol}$$ and the calculated $$\Delta $$H_3$$ from Hess’s Law is $$-55.0\ \mathrm{kJ/mol}$$, what is the percent error?

Accepted Answer

$$\mathrm{Percent\ error} = \left|\frac{-57.0 - (-55.0)}{-55.0}ight| 	imes 100\% = \left|\frac{-2.0}{-55.0}ight| 	imes 100\% = 3.6\%$$

Question 11

Which of the following best explains why the enthalpy change ($$\Delta H) $$measured in this experiment is equal to the heat flow ($$q) $$for the reaction?

Accepted Answer

Because the reactions are performed at constant pressure, so $$\Delta H = q_p.$$

Question 12

A student accidentally uses $$60.0\ \mathrm{mL} of$$ $$1.0\ \mathrm{M}$$ HCl and $$40.0\ \mathrm{mL} of$$ $$1.0\ \mathrm{M}$$ NaOH$$\ (aq) in $$Reaction 3. Which reactant is limiting, and how many moles of water are formed?

Accepted Answer

NaOH is limiting; $$0.0400\ \mathrm{mol} of $$water are formed.

Question 13

Which of the following is NOT a valid assumption made in the calorimetry experiment described?

Accepted Answer

The calorimeter absorbs a significant amount of heat.

Question 14

In the context of this experiment, which of the following best describes the application of Hess’s Law?

Accepted Answer

It allows the calculation of the enthalpy change for a reaction by combining the enthalpy changes of two or more other reactions.

Question 15

A student is asked to calculate the enthalpy change per mole of NaOH for Reaction 2. The student uses $$2.00\ \mathrm{g}$$ NaOH, $$50.0\ \mathrm{mL}$$ water, $$50.0\ \mathrm{mL}$$ HCl, and observes a $$7.0\ ^\circ\mathrm{C}$$ temperature increase. What is the correct calculation for $$\Delta H in$$ $$\mathrm{kJ/mol}$$? (Specific heat $$= 4.18\ \mathrm{J/(g\cdot{}^{\circ}C)}$$, molar mass NaOH $$= 40.0\ \mathrm{g/mol})$$

Accepted Answer

Total mass $$= 100.0\ \mathrm{g}$$; $$q = 100.0 	imes 4.18 	imes 7.0 = 2926\ \mathrm{J}$$; $$\mathrm{mol\ NaOH} = 2.00/40.0 = 0.0500$$; $$\Delta H = -2926/0.0500 = -58,520\ \mathrm{J/mol} = -58.5\ \mathrm{kJ/mol}$$

Question 16

Which of the following real-world scenarios best illustrates the practical application of calorimetry as described in this experiment?

Accepted Answer

Determining the energy content of food by measuring the heat released during combustion.

Calorimetry and Hess’s Law: Measuring Enthalpy Changes in Chemical Reactions

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