Thermochemistry Equations & Formulas - Lecture Review & Practice Problems

This paragraph introduces the concept of internal energy and the equation ΔE = Q + W, where ΔE is the change in internal energy, Q is the heat energy, and W is the work done. It explains that Q is negative when heat leaves the system (exothermic process) and positive when heat enters the system (endothermic process). The paragraph also covers the measurement of heat in joules and the conversion factors between kilojoules, joules, and calories. Work is represented by W and is calculated using the formula W = -PΔV, where P is pressure, and ΔV is the change in volume. The sign of W depends on whether work is done on or by the system. Finally, it presents a practice problem involving the absorption of heat by a gas and the expansion of the gas to calculate the change in internal energy.

The second paragraph delves into the calculation of heat energy using the formula Q = mcΔT, where m is mass, c is specific heat capacity, and ΔT is the change in temperature. It uses water as an example, noting its specific heat capacity is 4.184 joules per gram per degree Celsius. An example calculation is provided for heating 50 grams of water from 25°C to 75°C, resulting in the requirement of 10,460 joules of energy. The paragraph also explains how to calculate heat energy during phase changes using the formula Q = mΔH or Q = nΔH, depending on the units of enthalpy change (ΔH). It further illustrates the calculation using the heat of fusion for water as an example, converting grams of ice to moles and then calculating the energy required to melt the ice into water.

This paragraph discusses thermal chemical equations, using the combustion of propane as an example. It outlines how to balance chemical equations by ensuring the number of atoms for each element is the same on both sides. The paragraph also explains how to calculate the heat energy released during a reaction using the molar ratio, converting grams to moles and then to kilojoules. An example is given where 64 grams of oxygen reacts, and the resulting kilojoules of heat energy released are calculated. The concept of endothermic and exothermic reactions is also explained, noting the placement of heat energy values on either side of the reaction arrow.

The fourth paragraph focuses on the enthalpy of formation and how it can be used to estimate the enthalpy of a reaction. Enthalpy of formation is defined as the heat absorbed or released when a compound is produced from its elements in their natural standard state. The paragraph provides an example using methane and oxygen to form carbon dioxide and water, with given heats of formation for CO2, H2O, and CH4. It demonstrates how to calculate the enthalpy of the reaction by summing the products' enthalpies and subtracting the reactants' enthalpies. The calculation results in an estimated enthalpy of -184 kilojoules per mole for the reaction.

The final paragraph introduces Hess's Law, which is used to estimate the enthalpy of a reaction by manipulating and adding two known reactions to match a third, unknown reaction. An example is given with three reactions, where the enthalpies for the first and second reactions are provided. By doubling the first reaction and reversing the second, the paragraph shows how to achieve the third reaction and add the enthalpies to find the enthalpy for the net reaction, which is calculated to be 800 kilojoules in the example. This method is a practical application of Hess's Law for estimating reaction enthalpies.