Chem - Chapter 9.5 study guide

Thermochemistry

Measuring ΔH for Chemical Reactions: Constant-Pressure Calorimetry

Thermochemistry involves the study of energy changes, particularly heat, during chemical reactions. Constant-pressure calorimetry is a common method for measuring the enthalpy change (ΔH) of reactions.

• Enthalpy (ΔH): Defined as the sum of a system's internal energy plus the product of its pressure and volume. It represents the heat content of a system at constant pressure.

• Measuring Enthalpy Changes: In a coffee-cup calorimeter, reactions are performed at constant pressure, and the heat exchanged is measured to determine ΔH.

• Constant-Volume vs. Constant-Pressure Calorimetry: Constant-volume calorimetry (e.g., bomb calorimeter) measures changes at fixed volume, while constant-pressure calorimetry measures changes at fixed pressure, which is typical for reactions in open containers.

• Relationship to Heat Flow: At constant pressure, the change in enthalpy for a reaction or process equals the heat flow into or out of the system.

• Endothermic and Exothermic Reactions:

  • Endothermic: ΔH > 0; the system absorbs heat from the surroundings.

  • Exothermic: ΔH < 0; the system releases heat to the surroundings.

  The sign of ΔH and the temperature change of a solution can be used to predict whether a reaction is endothermic or exothermic.

Equation:

$\Delta H = q_p$

where $q_p$ is the heat at constant pressure.

Hess's Law: Relationships Involving ΔHrxn

Hess's Law states that the total enthalpy change for a reaction is the same, no matter how many steps the reaction is carried out in. This allows calculation of enthalpy changes for complex reactions by combining known enthalpy changes for simpler reactions.

• Using Reaction Enthalpies: You can multiply a chemical equation by a factor, reverse a chemical equation, or sum a series of chemical equations to find the overall enthalpy change.

• Application of Hess's Law: Use Hess's Law to calculate the enthalpy change for a reaction from a series of steps.

Equation:

$\Delta H_{\text{overall}} = \sum \Delta H_{\text{steps}}$

Determining Enthalpies of Reaction from Standard Enthalpies of Formation

The standard enthalpy of formation ($\Delta H_f^\circ$) is the enthalpy change when one mole of a compound is formed from its elements in their standard states. This concept is essential for calculating the enthalpy changes of reactions using tabulated values.

• Standard State: The most stable physical form of an element or compound at 1 bar pressure and a specified temperature (usually 25°C).

• Common Naturally Occurring Diatomic Molecules: $O_2(g)$, $H_2(g)$, $N_2(g)$, $F_2(g)$, $Cl_2(g)$, $Br_2(l)$, $I_2(s)$

• Thermochemical Equations: Write balanced equations for the formation of compounds from their elements in standard states.

• Calculating Enthalpy of Reaction: Use the enthalpies of formation of products and reactants to determine the overall enthalpy change.

Equation:

$\Delta H_{\text{rxn}} = \sum n_p \Delta H_f^\circ(\text{products}) - \sum n_r \Delta H_f^\circ(\text{reactants})$

where $n_p$ and $n_r$ are the stoichiometric coefficients of products and reactants, respectively.

Example: Calculating ΔHrxn Using Standard Enthalpies of Formation

• Given the reaction: $C_2H_4(g) + 3 O_2(g) \rightarrow 2 CO_2(g) + 2 H_2O(l)$

• Find $\Delta H_{\text{rxn}}$ using tabulated $\Delta H_f^\circ$ values for $CO_2(g)$, $H_2O(l)$, and $C_2H_4(g)$.

Additional info: Standard enthalpies of formation for elements in their standard states are zero by definition.