First+Law+of+Thermodynamics-+enthalpy;+heats+of+formation,+reaction,+fusion,+and+vaporization;+Hess's+Law;+calorimetry

=__First Law of Thermodynamics__= The First Law of Thermodynamics states that energy is conserved. In other words, energy is never created nor destroyed, but only transferred. In this transfer of energy between a system and its surroundings, the energy lost by the system must be the same amount of energy that is gained by the surroundings and vise versa. The positive or negative transfer of energy is described using the terms **endothermic**, when the system gains energy from its surroundings and heat is absorbed, and **exothermic**, when the system loses energy to its surroundings and therefore heat is lost. Some students may be familiar with reaction coordinates like the ones below, which show exothermic (A) and endothermic (B) reactions in terms of heat and time.

=__Enthalpy__= **Enthalpy**, H, accounts for the heat flow in processes occurring at constant pressure when no other forms of work are performed other than expanding or compressing gasses. This is a change in quantity of energy of heat content. However, because H cannot be measured on its own, chemists must instead measure the change in enthalpy, ∆H, which is equal to the enthalpy of the products minus the enthalpy of the reactants (∆H = Hproducts – Hreactants). Also, it is important to note that:


 * 1) as can be seen in the reaction coordinates above, the ∆H of an exothermic reaction is negative, while the ∆H of an endothermic reaction is positive.
 * 2) enthalpy is a state function, meaning that its value depends only on the present state of the system, not on the path the system took to reach that state.
 * 3) enthalpy is an extensive property, meaning that the value of ∆H is directly proportional to the amount of reactants consumed in the reaction. Therefore, if the amount of reactants is doubled, then the value of ∆H will likewise double. This is important to know when applying Hess' Law
 * 4) though the quantity of heat lost or gained varies, the //standard state// of 25°C and 1 atmosphere are used.

Below is an example of an enthalpy diagram which shows the formation of carbon dioxide and water from methane and oxygen:  =__Heat of Reaction__= Also known as enthalpy of reaction, ∆Hrxn, the **heat of reaction** is the change in enthalpy, or heat flow, as a reaction occurs. If the ∆Hrxn is positive, the reaction is endothermic, meaning that the system absorbs energy from its environment as the reaction occurs. Conversely, if the ∆Hrxn is negative, the reaction is exothermic, meaning that the system releases energy into its environment as the reaction occurs. Two important things to note about ∆Hrxn are:

=__Calorimetry__= **Calorimetry** is the measurement of heat flow, which is often used to calculate the calorie content of food by burning that food and measuring the energy that is released. The container that insulates the system so that heat flow can be meaured is called a **calorimeter**, which is insulated to prevent heat from escaping from the system and into the surroundings. Thus, when a calorimeter insulates the system perfectly (and remains under constant pressure), the energy gained by the system is the same amount of energy produced from the chemical reaction taking place. This is expressed in the equation //q//soln = -//q//rxn (the heat gained by the solution is the same in magnitude but opposite in sign to the heat produced by the reaction). However, not all substances absorb heat at the same rate or magnitude. For this reason, an object's temperature change when it absorbs a certain amount of energy is determined by its **heat capacity**, C, which is the amount of energy required to raise the temperature of a substance 1K. The larger a substance's value of C is, the more energy is required to raise its temperature. The heat capacity of one mole of a substance is called the **molar heat capacity**, Cmolar, and the heat capacity of one gram of a substance is called the **specific** **heat capacity**, //s//, which is calculated using the equation //q//=//sm// ∆T, where //q// is the amount of heat transferred, //m// is the mass of the substance, and ∆T is the temperature change of the substance. For instance, the specific heat of water is 4.184J because 209J are required to increase the temperature of 50.0g of water by 1.00K. =__Hess' Law__= Hess' Law states that "if a reaction is carried out in a series of steps, the enthalpy change, ∆H, for the overall reaction will be equal to the sum of the enthalpy changes for the individual steps." In this way, only a few experimental measurements are necessary to caclulate the ∆H for many different reactions. In the following example, the ∆H of a combustion step and that of a condensation step are used to calculate the ∆H for the overall reaction:  Also, it is important to note that whatever is done to the reaction must also be done to its ∆H because of the fact that enthalpy is an extensive property. For example, if the equation for the combustion of CH4, as shown above, were doubled, the ∆H would also have to be doubled. =__Heat of Formation__= **Heat of formation**, also known as enthalpy of formation or ∆Hf, is the amount of energy required to form a compound from its constituent elements. The **standard enthalpy of formation**, ∆H°f, is the amount of energy required to produce one mole of a compound from that compound's elements, when all substances are in their standard states of 1 atm and 298K. For example, ½ mole of H2(//g//) reacts with 1/2 mole of Cl2(//g//) to form one mole of HCl(//g//), which has a ∆H°f of -92.30kJ. Similarly, the **heat of vaporization** is the ∆H of a reaction in which a liquid is converted to a gas, and the **heat of fusion** is the ∆H of a reaction in which a solid is converted into a liquid. Many chemistry textbooks and reference books contain appendixes of these values, and some common standard enthalpies of formation can be found [|here].
 * 1) when a reaction is reversed, the sign of ∆Hrxn is also reversed because the enthalpy change for a reaction is equal in magnitude but opposite in sign to the enthalpy change of the reverse reaction.
 * 2) the ∆Hrxn is directly proportional to the reaction, so whatever is done to the reaction (like halving the reaction's coefficients) must also be done to the value of ∆Hrxn.
 * 3) the enthalpy change of a reaction depends on the state of the reactant and products. For example, if the product of the combustion of methane is H2O(//g//), then the ∆Hrxn would be -802kJ, while if the product is H2O(//l//), then the ∆Hrxn would be -890kJ.