Thermochemistry, Chemistry tutorial


The modern society mainly depends on energy for its existence. Any symptoms of an energy shortage - gasoline shortages, rolling blackouts of the electrical power, or big increases in the cost of natural gas are sufficient to shake people's confidence and roil the markets. Energy is vastly a chemical topic. Almost all of the energy on which we depend is derived from the chemical reactions, like the combustion of fossil fuels, the chemical reactions taking place in batteries, or the formation of biomass via photosynthesis. Think for an instant regarding some of the chemical methods that we encounter in the course of a usual day: We eat foods to produce the energy required to maintain our biological functions. We burn fossil fuels (like petroleum, coal, natural gas etc) to produce much of the energy that powers our offices and homes and that moves us from place to place via train, automobile or plane. We listen to tunes on battery-powered by MP3 players.

The relationship between chemical change and energy exhibits in different manners. Chemical reactions comprising foods and fuels discharge energy. By contrary, the splitting of water into hydrogen and oxygen needs an input of electrical energy. Likewise, the chemical method we call photosynthesis, which takes place in plant leaves, transforms one form of energy, radiant energy from the Sun, to the chemical energy. Chemical methods can do more than simply produce heat; they can do work, like turning an automobile starter, powering a drill, and so forth. What we get from all this is that chemical change usually comprises energy. If we are to correctly understand chemistry, we must as well comprehend the energy changes which accompany the chemical change.

The study of energy and its transformations is termed as thermodynamics (Greek meaning: thermo means heat and dynamics means power). This area of study start all through the Industrial Revolution as the relationships among work, heat and energy content of fuels were studied in an effort to make best use of the performance of steam engines. Nowadays thermodynamics is very significant in all streams of science and engineering. This feature of thermodynamics is termed as Thermochemistry.

Concept of Thermochemistry:

Thermochemistry is basically the study of the heat liberated or absorbed as an outcome of the chemical reactions. This is a branch of thermodynamics and is used via a broad range of engineers and scientists. For illustration, biochemists use Thermochemistry to comprehend bioenergetics, while chemical engineers apply Thermochemistry to design manufacturing plants. Chemical reactions comprise the conversion of a set of substances collectively termed to as 'reactants' to a set of substances collectively termed to as the 'products'. In the given balanced chemical reaction the reactants are gaseous methane, CH4 (g), and gaseous molecular oxygen, O2 (g), and the products are gaseous carbon-dioxide, CO2 (g), and liquid water H2O (l):

CH4 (g) + 2O2 (g) → CO2 (g) + 2 H2O (l)

The reactions in which a fuel combines by oxygen to produce water and carbon-dioxide are termed as combustion reactions. As natural gas comprises primarily of methane, it is expected that reaction will discharge heat. Reactions which discharge heat are known as exothermic reactions, and reactions which absorb heat are known as endothermic reactions.

The heat related with a chemical reaction based on the temperature and pressure at which the reaction is taken out. All the Thermochemical data presented here are for reactions taken out beneath standard conditions, which are a temperature of 298 K (24.85°C) and an applied pressure of one bar. The quantity of heat liberated in a reaction based on the amount of material undergoing the reaction. The chemical formulas which appear in a reaction each represent 1 mole of material; for illustration - the symbol CH4 stands for 1 mole of methane having a mass of 16 grams (0.56 ounces), and the 2O2 (g) states us that the 2 moles of oxygen are needed. Thermochemistry as well based on the physical state of the reactants and products. For illustration, the heat discharged in the equation above is 890 kilojoules (kJ); if, though, water in the gas phase is formed, H2O (g), the heat liberated is just 802 kJ. Reversing the reaction that discharges heat, yields a reaction wherein heat should be supplied for the reaction to take place. The given reaction absorbs 890 kJ.

CO2 (g) + 2H2O (l) → CH4 (g) + 2O2 (g)


The Energy is a potential to do work, like accelerating an object (that is, kinetic energy), lifting things up (that is, potential energy), generate electric power (that is, electric energy), increasing temperature of a system (that is, heat) and producing sound (that is, waves of energy). In physics, work is generally stated as force times the distance in the direction of the force, and such a definition states the very fundamental idea of energy.

Energy is the driving force of changes. All the changes are caused by energy, and the cause or energy can be in numerous forms: heat, light, work, electrical, mechanical (that is, energy stored in a spring), chemical and so on. The changes are phenomena caused via energy, however more significantly, forms of energy inter-convert to one another throughout the changes.

Dissimilar to some of these changes, forms of energy convert into one other always at a fixed rate. Precisely 4.184 J of mechanical work transforms into 1.00 cal measured as heat and vice-versa.

Conservation of Energy:

Energy can neither be created nor destroyed; it converts from one form (for illustration heat) to the other form (state mechanical work) at a fixed rate. This is the basic principle of conservation of energy.

To point out changes, we employed to represent the delta (D) generally employed in text books.

The internal energy 'E' accounts for energy and work transferred to the system. This concept is the other form of the principle of conservation of energy. The change in internal energy 'dE' of a closed system rises by the amount of energy input to the system. These input can be in the form of heat (q) or (mechanical or any other form) of work (w). Generally, it is formulated as:

dE = q + w

The heat transferred out and work done by the system is assigned negative signs via convention. This relationship is generally termed to as the first law of thermodynamics.

Heats of Reactions:

The enthalpy of reaction, 'dH', is the energy [heat (q) and work (d (PV))] discharged in a reaction. The Thermochemical equation is generally represented in the form:

2 H2 (g, 1atm) + O2 (g) = 2 H2O (l),   dH = - 571.7 kJ

This equation signifies that whenever 2 moles of H2 gas react by 1 mole of O2 gas, 571.7 kJ of energy is liberated (lost to the surrounding). Therefore, if 1 mole of H2 and one half mole of O2 react, half of 571.7 kJ or 285.9 kJ of energy is liberated.

H2 (g, 1atm) + 0.5 O2 = H2O (l),   dH = - 285.9 kJ

On the other hand, if double amounts of reactants (that is, 4 moles of H2 and 2 moles of O2) are employed, double amount of energy. (2*(-571.7) =) -1043.4 kJ is liberated.

If dH is positive then at least that much energy should be supplied to carry out the endothermic reaction.

The Standard Enthalpy of Reaction:

For ease in application, 1 atm for gas and 1.0 M for solutions were taken as the 'standard conditions', and data collected at standard conditions were termed as the standard data like standard enthalpy of reaction and standard enthalpy of formation. Such values are condensed and summarized in hand-books for engineers and scientists in their applications.

As pressure, temperature and concentration of reactants and products influence the quantity of measured energy, the scientific community has agree on a temperature of 273 K and 1 atm as the standard temperature and pressure (STP). Though, standard enthalpies are frequently given for data collected at 298 K.

The most stable state at standard condition is the standard state. The enthalpy of an element at its standard state is assigned zero (0) for reference.

For illustration, at 1 atm, graphite is the most stable state of carbon. The standard enthalpy of combustion of carbon is the energy liberated (- 394 kJ) whenever 1 mole of graphite reacts by oxygen in the reaction,

C (graphite) + O2 = CO2,   dHo = - 394 kJ.

As the measurement is completed at the standard condition, a superscript 'o' is generally placed on the right side of H in most literature. Incidentally, the equation above is for the formation of CO2, and the enthalpy of reaction occurs to be the enthalpy of formation of CO2, designated as dHof = - 394 kJ, as evaluated in the next paragraph.

As the other illustration, whenever 1.0 mole Zn reacts by the adequate amount of HCl solution (1.0 M), 150 kJ is liberated. Therefore, we write standard energy of reaction for Zn as,

Zn + 2 HCl (aq) = H2 (g) + ZnCl2 (aq),    dHo = -150 kJ.

The standard Enthalpy of Formation:

The combination of elements at their standard states resultant in one compound is termed as a formation reaction. Whenever enthalpy of formation is computed at the standard condition, it is termed as the standard enthalpy of formation. The standard enthalpy of combustion of carbon illustrated earlier,

C (graphite) + O2 = CO2,       dHof = - 394 kJ.

is the formation of CO2 from elements at their standard states. Therefore, dHof of - 394 kJ is as well the standard enthalpy of the formation of CO2.

Likewise, a few more illustrations are described below. The enthalpies can be both positive and negative values.

1/2 O2 = O,                           dHof = 249.17 kJ (as well bond energy of O=O).

1/2 H2 = H,                            dHof = 217.96 kJ (as well bond energy of H-H).

H2 + O2 = H2O2 (aq),            dHof = -191.17 kJ

1/2 H2 + 1/2 Cl2 = HCl (g),    dHof = - 92.31 kJ.

1/2 H2 + 1/2 Br2 = HBr (g),    dHof = - 36.40 kJ.

Hess's Law:

Hess's law is the other interpretation of the principle of conservation of energy. As the changes in energy are independent of the path, they totally depend on the initial and final state of the system. Therefore, if it takes some steps to reach the final state from the initial state, the changes in energy are additative. The law defines:

The net enthalpy change in a reaction is similar whether the reaction takes place in one or several steps.

Though, one must recognize that the enthalpy change is associated to the amounts of reactants and products in the equation.

A simple application of Hess's law is to provide the standard enthalpy of decomposition of CO2 from its standard enthalpy of formation,

CO2 (g) = C (graphite) + O2 (g), dHo = 394 kJ.

It will be noted that we change the sign of dHo if the reaction is reversed.

Measurements of Energy Changes:

Different experimental methods have been designed to measure or compute the energy changes in a chemical reaction. It is essential to know the heat capacity of the system. Accurate measurements need carefully designed calorimeters.

The Enthalpy of reaction (dH) is measured if the reaction is taken out at constant pressure. If a bomb calorimeter is employed, the volume doesn't change. The amount of energy evaluated is the internal energy dE. To transform dE into dH, we make use of the defined relationship,

dH = dE + d (PV).

The changes in pressure and volume (P V) work can be computed by the application of ideal gas law,

d (PV) = dn R T,

Here, dn is the total number of moles of gas of product ∑n(products) subtracting the total number of moles of gas of reactants, ∑n(reactants).

dn = ∑n(products) - ∑n(reactants)

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