Unit 6: Thermodynamics

Using Hess's law: Given C + O₂ → CO₂, ΔH = -393.5 kJ and CO + ½O₂ → CO₂, ΔH = -283.0 kJ. Find ΔH (in kJ) for C + ½O₂ → CO.

For a reaction at constant temperature, which combination of ΔH and ΔS always results in a spontaneous reaction (ΔG < 0)?

Using Hess's law, given: C(s) + O₂(g) → CO₂(g), ΔH = -393.5 kJ H₂(g) + ½O₂(g) → H₂O(l), ΔH = -285.8 kJ C₂H₆(g) + 7/2 O₂(g) → 2CO₂(g) + 3H₂O(l), ΔH = -1560.7 kJ The enthalpy of formation of C₂H₆(g) is

When 50.0 mL of 1.00 M HCl at 25.0°C is mixed with 50.0 mL of 1.00 M NaOH at 25.0°C in a coffee-cup calorimeter, the temperature rises to 31.9°C. Assuming the density of the solution is 1.00 g/mL and the specific heat is 4.18 J/(g·°C), the enthalpy of neutralization per mole of water formed is

The bond energy of H-H is 436 kJ/mol, Cl-Cl is 242 kJ/mol, and H-Cl is 431 kJ/mol. The enthalpy change for H₂(g) + Cl₂(g) → 2HCl(g) is

The standard enthalpy of formation (ΔH°f) of O₂(g) is

For a system at constant pressure, which relationship correctly describes the enthalpy change?

If the reaction A → B has ΔH = +50 kJ, the reaction 2B → 2A has ΔH equal to

Hess's Law states that:

Bond energy can be used to estimate [math] because:

In calorimetry, [math] is used to calculate:

The standard enthalpy of formation ([math]) of an element in its standard state is:

An endothermic reaction:

The enthalpy change for a reaction can be calculated from standard enthalpies of formation using:

A coffee cup calorimeter measures reactions at constant:

Using the given enthalpies, what is the standard enthalpy of formation of CH₄(g)?

What is the enthalpy change of the dissolution per mole of NaOH?

Using bond energies, estimate the enthalpy change for this reaction.

Which process is endothermic?

If equal masses of each substance absorb the same amount of heat, which will show the largest temperature increase?