# Groupwork 6 Enthalpy

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*Work in groups on these problems. You should try to answer the questions without referring to your textbook. If you get stuck, try asking another group for help.*

In general chemistry, you were introduced to the concept of enthalpy, *H*. In physical chemistry, we see that enthalpy arises from the situation where we want to calculate the energy changes in chemical reactions in conditions where the pressure is constant. This groupwork activity intends to remind you of basic concepts about enthalpy from general chemistry and explore calculations when we are not at standard temperature and pressure (25˚C and 1 atm).

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Given that \(H=U+PV\), what is the general expression for the change in enthalpy when the pressure is constant?

The heat required to melt ice at 0˚C and 1 atm is \(q_p=\Delta H=6.01 kJ/mol\) while the internal energy for the same process is \(\Delta U\approx6.01 kJ/mol\). The heat required to boil water at 100˚C and 1 atm is \(q_p=\Delta H=40.7 kJ/mol\) while \(\Delta U=37.6 kJ/mol\). Why are \(\Delta H\) and \(\Delta U\)the same for melting ice but different for boiling water?

Hess's Law tells us that enthalpy changes for chemical reactions are additive. Thus, if we know the enthalpies for several reactions, we can add them up to get the enthalpy of another reaction. We also know that \(\Delta_rH(forward)=-\Delta_r(reverse)\) and that enthalpies depend on the amount of material used and produced. Thus \(\Delta_rH(1mole)=\frac{1}{2}\Delta_rH(2mol)\).

Find the enthalpy of the reaction

\(PCl_3(l)+Cl_2(g)\rightarrow PCl_5(s)\)

Given

\(2P(s)+3Cl_2(g)\rightarrow 2PCl_3(l)\) \(\Delta_rH=-640kJ\)

and

\(2P(s)+5Cl_2(g)\rightarrow 2PCl_5(s)\) \(\Delta_rH=-887kJ\)

The standard reaction enthalpy, \(\Delta_rH^{\circ}\), is defined as enthalpy change associated with one mole of a specified reagent when all reactants and products are in their standard states. The standard state for a gas is considered at 1 bar at the temperature of interest.

The standard molar enthalpy of formation, \(\Delta_fH^{\circ}\) for a substance is defined as the enthalpy change associated with forming one mole of a specified reagent from its constituent elements in their standard states at the given temperature. We usually think of this at standard temperature and pressure (25˚C and 1 atm or bar).

Given \(\Delta_fH^{\circ}(CCl_4(l))=-135.44kJ/mol\) and \(\Delta_fH^{\circ}(CCl_4(g))=-102.9kJ/mol\) at STP, what is the heat of vaporization for CCl_{4}?

Recommended: Practice problems 19-35 through 19-41 in the textbook to recall how to use Hess's law to calculate enthalpy changes