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Energy is the driving force of changes. All changes are caused by energy, and the cause or energy can be in many forms: light, heat, work, electrical, mechanical (energy stored in a spring), chemical, etc. The changes are phenomena caused by energy, but more importantly, forms of energy inter-convert into one another during the changes.
Unlike some of these changes, forms of energy convert into one another always at a fixed rate. Exactly 4.184 J of mechanical work converts into 1.00 cal measured as heat, and vice versa.
Energy can neither be destroyed, nor created; it converts from one form (for example heat) to another form (say mechanical work) at a fixed rate. This is the fundamental principle of conservation of energy.
To indicate changes, we use d to represent the delta (D) commonly used in text books, because using the delta will cause the loading of this page into your computer very slow.
The internal energy E accounts for energy and work transferred to a system. This concept is another form of the principle of conservation of energy. The change in internal energy dE of a closed system increases by the amount of energy input to the system. Such input can be in the form of heat (q) or (mechanical or any other form) of work (w). Usually, it is formulated as
When energy is transferred from system A to system B, what happens to the energy? Does energy disappear? The internal energy as defined above shows that energy does not disappear, because the internal energy loss of system A equals to the internal energy gain of system B. The change in internal energy, dE, is negative for system A, but positive for system B. Energy is neither destroyed, nor created.
On the other hand, if double amounts of reactants (i.e. 4 moles of H2 and 2 moles of O2) are used, twice amount of energy. (2*(-571.7) =) -1043.4 kJ is released.
If dH is positive, at least that much energy must be supplied to carry out the endothermic reaction.
Because temperature, pressure, and concentration of reactants and products affect the amount of measured energy, the scientific community has agree upon a temperature of 273 K and 1 atm as the standard temperature and and pressure (STP). However, standard enthalpies are often given for data collected at 298 K.
The most stable state at the standard condition is the standard state. The enthalpy of an element at its standard state is assigned 0 for reference.
For example, at 1 atm, graphite is the most stable state of carbon. The standard enthalpy of combustion of carbon is the energy released (-394 kJ) when 1 mole of graphite reacts with oxygen in the reaction,
As another example, when 1.0 mole Zn reacts with sufficient amount of HCl solution (1.0 M), 150 kJ is released. Thus, we write standard energy of reaction for Zn as,
Similarly, a few more examples are given below. The enthalpies can be both positive and negative values.
The total enthalpy change in a reaction is the same whether the reaction occurs in one or several steps.
However, one should recognize that the enthalpy change is related to the amounts of reactants and products in the equation.
A simple application of Hess's law is to give the standard enthalpy of decomposition of CO2 from its standard enthalpy of formation,
Note that we change the sign of dHo if the reaction is reversed.
Enthalpy of reaction (dH) is measured when the reaction is carried out at constant pressure. When a bomb calorimeter is used, the volume does not change. The amount of energy measured is the internal energy dE. To convert dE into dH, we use the defined relationship,
The changes in pressure and volume (P V) work can be evaluated by the application of ideal gas law,
where dn is the total number of moles of gas of product Sn(products) substracting the total number of moles of gas of reactants, Sn(reactants).
The enthalpy of a reaction can be evaluated from the standard enthalpies of formation of all products and reactants.
The entropy (dS) is another important thermodynamic data, which is often listed together with dHof. The entropies are: P4O10, 229 J/(K mol); H2, 70 J/(K mol); H3PO4, 158 J/(K mol). and the entropy of the reaction is thus,
Don't worry. General chemistry students are not expected to know entropy and Gibb's free energy yet. A summary of thermochemistry has been given. Thermochemistry deals with the energy aspect of chemical reactions. A more complete study of energy is Thermodynamics. This link is divided into 7 units: Energy, Enthalpy, Hess's Law, Enthalpy of Formation, Entropy, Gibbs Free Energy and Conclusion. You may find its simple approach interesting. Some of the aspect has been included in the discussion of the above example.