SAT Chemistry Chemical Reactions and Thermochemistry - Thermochemistry
SAT Chemistry Chemical Reactions and Thermochemistry - ThermochemistryIn general, all chemical reactions either liberate or absorb heat. The origin of chemical energy lies in the position and motion of atoms, molecules, and subatomic particles. The total energy possessed by a molecule is the sum of all the forms of potential and kinetic energy associated with it.
The energy changes in a reaction are due, to a large extent, to the changes in potential energy that accompany the breaking of chemical bonds in reactants to form new bonds in products.
The molecule may also have rotational, vibrational, and translational energy, along with some nuclear energy sources. All these make up the total energy of molecules. In beginning chemistry, the greatest concern in reactions is the electronic energy involved in the making and breaking of chemical bonds.
Because it is virtually impossible to measure the total energy of molecules, the energy change is usually the experimental data that we deal with in reactions. This change in quantity of energy is known as the change in enthalpy (heat content) of the chemical system and is symbolized by ΔH.
CHANGES IN ENTHALPY
Changes in enthalpy for exothermic and endothermic reactions can be shown graphically, as in the examples below.
Notice that the ΔH for an endothermic reaction is positive, while that for an exothermic reaction is negative. It should be noted also that changes in enthalpy are always independent of the path taken to change a system from the initial state to the final state.
Because the quantity of heat absorbed or liberated during a reaction varies with the temperature, scientists have adopted 25°C and 1 atmosphere pressure as the standard state condition for reporting heat data. A superscript zero on ΔH (i.e., ΔH°) indicates that the corresponding process was carried out under standard conditions. The standard enthalpy of formation (ΔH°ƒ) of a compound is defined as the change in enthalpy that accompanies the formation of 1 mole of a compound from its elements with all substances in their standard states at 25°C. This value is called the molar heat of formation.
To calculate the enthalpy of a reaction, it is necessary to write an equation for the reaction. The standard enthalpy change, AH, for a given reaction is usually expressed in kilocalories and depends on how the equation is written. For example, the following equations express the reaction of hydrogen with oxygen in two ways:
H2(g) + 1/2 O2(g) → H2O(g) ΔH°ƒ = -241.8 kj
2H2(g) + O2(g) → 2H2O(g) ΔH°ƒ = -483.6 kj
Experimentally, ΔH°ƒ, for the formation of 1 mole of H2O(g) is -241.8 kj. Since the second equation represents the formation of 2 moles of H2O(g), the quantity is twice -241.8, or -483.6 kj. It is assumed that the initial and final states are measured at 25°C and 1 atmosphere, although the reaction occurs at a higher temperature.
How much heat is liberated when 40.0 grams of H2(g) reacts with excess 0(g)? The reaction equation is:
Notice that the physical state of each participant must be given since the phase changes involve energy changes. Combustion reactions produce a considerable amount of energy in the form of light and heat when a substance is combined with oxygen. The heat released by the complete combustion of 1 mole of a substance is called the heat of combustion of that substance. Heat of combustion is defined in terms of 1 mole of reactant, whereas the heat of formation is defined in terms of 1 mole of product. All substances are in their standard state. The general enthalpy notation, ΔH, applies to heats of reaction, but the addition of a subscript c, ΔHc, specifically indicates heat of combustion.
ADDITIVITY OF REACTION HEATS AND HESS’S LAW
Chemical equations and ΔH° values may be manipulated algebraically. Finding the ΔH for the formation of vapor from liquid water shows how this can be done.
Since we want the equation for H2O(1) -+ H2O(g), we can reverse the second equation. This changes the sign of ΔH.
The principle underlying the preceding calculations is known as Hess’s Law of Heat Summation. This principle states that, when a reaction can be expressed as the algebraic sum of two or more other reactions, the heat of the reaction is the algebraic sum of the heats of these reactions. This is based upon the First Law of Thermodynamics, which, simply stated, says that the total energy of the universe is constant and cannot be created or destroyed.
These laws allow calculations of ΔH’s that cannot be easily determined experimentally. An example is the determination of the ΔH°ƒ of CO from the ΔH°ƒ of CO2
Some commonly used standard heats of formation, designated as ΔH°ƒ, Eire listed in Table Ⓐ in the Tables for Reference section.
An alternative (and easier) method of calculating enthalpies is based on the concept that ΔH°reaction is equal to the difference between the total enthalpy of the reactants and that of the products. This can be expressed as follows:ΔH°reaction = ∑(sum of)ΔH°f (products) - ∑(sum of)ΔH°/(reactants)
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