Thermochemistry

THERMOCHENISTRY

Index 1.0 Introduction | | | | | | | | 1 | 2.0 Enthalpy Changes | | | | | | | 2,3 | 2.1 The Standard Conditions For Calculating Enthalpy Changes | | | | 3.0 Hess’s Law | | | | | | | | 4,5 | 3.1 The Applications of Hess’s Law | | | | | | 4.0 Standard Molar Enthalpy Change of Formation, ΔHof | | | | 6,7 | 4.1 The Stability of A Compound | | | | | | | 4.2 Using ΔHof Values To Predict The Relative Stability Of A Compound | | | 5.0 Standard Molar Enthalpy Change of Combustion, ΔHoc | | | 8,9 5.1 Standard Enthalpy Change of Combustion and Molecular Structure | | | 6.0 Standard Molar Enthalpy Change of Neutralisation, ΔHonuet | | | 10,11 | 6.1 The Standard Enthalpy Change of Neutralisation For Strong Acids | | | | and Strong Alkalis | | | | | | | | 6.2 The Standard Enthalpy Change of Neutralisation of Weak Acids and | | | | Weak Alkalis | | | | | | | | 7.0 Standard Molar Enthalpy Change of Atomisation, ΔHoat | | | 12 | 8.0 Electron Affinity | | | | | | | 13 | 9.0 Lattice Energy | | | | | | | 14 | 9.1 The Magnitude of Lattice Energy | | | | | | 10.0 The Born-Haber Cycle | | | | | | | 15 | 11.0 References | | | | | | | | 16 |

1.0 Introduction
Thermochemistry is the study of the energy and heat associated with chemical reactions and/or physical transformations. A reaction may release or absorb energy, and a phase change may do the same, such as in melting and boiling. Thermochemistry focuses on these energy changes, particularly on the system 's energy exchange with its surroundings. Thermochemistry is useful in predicting reactant and product quantities throughout the course of a given reaction. It is also used to predict whether a reaction is spontaneous or non-spontaneous, favorable or unfavorable. Endothermic reactions absorb heat. Exothermic reactions release heat. Thermochemistry coelesces the concepts of thermodynamics with the concept of energy in the form of chemical bonds. The subject commonly includes the enthalpy changes of reaction, formation, combustion, solution, neutralisation, atomisation, and lattice energy.

2.0 Enthalpy Changes
Enthalpy change is the name given to the amount of heat evolved or absorbed in a reaction carried out at constant pressure. It is given the symbol ΔH, read as "delta H". Enthalpy (H) is measured in the units of kJ. Hence, for the reaction
A + B → C + D, the enthalpy change is given by ΔH. ΔH = The total enthalpy of the products – The total enthalpy of the reactants = H2 - H1 where H2 = ( HC + HD ) and H1 = ( HA + HB ).
If the reaction is exothermic, heat is given out to the surroundings, the enthalpy change for the reaction, ΔH has a negative value because the heat content of the products is less than the heat content of the reactants. If the reaction is endothermic, heat is absorbed from the surroundings, the enthalpy change for the reaction, ΔH has a positive value because the heat content of the products is higher than the heat content of the reactants.

2.1 The Standard Conditions for the Calculating Enthalpy Changes
The numerical value of the enthalpy change is influenced by five factors. The first factor is the temperature of the experiment being carried out. The second factor is the physical states ( solid, liquid, gas ) of the reactants. The third factor is the allotropic forms of the reactants. The forth factor is the pressure of gaseous reactants and the last factor is the concentration of the reactants.
Under these conditions, the reactants and the products are said to be in their standard states. For example, the standard state of oxygen is a gas whereas the standard state of water is a liquid. Although the enthalpy change values are commonly quoted at standard conditions, measurements can be made at other temperatures.
The enthalpy change measured at standard conditions is described as the standard enthalpy changes of reaction or standard hear of reaction and given the symbol . The units of it are kJ mol-1.

3.0 Hess’s Law
Hess’s law of constant heat summation states that the heat liberated or absorbed during a chemical reaction is independent of the route by which the chemical changes occurs, provided the initial and final conditions are the same.

Hess 's law allows the enthalpy change (ΔH) for a reaction to be calculated even when it cannot be measured directly. This is accomplished by performing basic algebraic operations based on the chemical equation of reactions using previously determined values for the enthalpies of formation.
Addition of chemical equations may lead to a net equation. If enthalpy change is known for each equation, the result will be the enthalpy change for the net equation. If the net enthalpy change is negative (ΔHnet < 0), the reaction is exothermic and is more likely to be spontaneous; positive ΔH values correspond to endothermic reactions. Entropy also plays an important role in determining spontaneity, as some reactions with a negative enthalpy change are nevertheless spontaneous.
Hess 's Law states that enthalpy changes are additive. Thus the ΔH for a single reaction can be calculated from the difference between the heat of formation of the products and the heat of formation of the reactants:

where the o superscript indicates standard state values.

3.1 Application of Hess’s Law
Hess’s law is very important, especially for determining the values of for the reactions that are difficult to carry out directly in a calorimeter. For the compound that can be synthesised through the reaction between the constituent elements, the enthalpies of formation of the compound can be determined by measuring the heat energy liberated or absorbed when the elements combine.
However, the enthalpy of formation for many compounds cannot be determined directly. For example, it is not possible to determine the enthalpy of formation of ethanol in the laboratory because it is not possible to synthesise ethanol from its constituent elements, carbon, hydrogen and oxygen. In such cases, Hess’s Law is required to calculate the enthalpy of formation from the enthalpy of combustion. Similarly, the lattice energy of an ionic compound is determined from the enthalpy change of atomisation energy, electron affinity and enthalpy change of formation by applying Hess’s Law.

4.0 Standard Molar Enthalpy Change of Formation, ΔHof
The standard enthalpy change of formation of a compound is the enthalpy change ( heat liberated or absorbed ) when one mole of the pure compound is formed from its elements under standard conditions. The standard enthalpy change of formation is given the symbol, ΔHof.
Most compounds have negative standard enthalpies of formation. They are called exothermic compounds. A few compounds have positive standard enthalpies of formation. They are described as endorthermic compounds.
By definition, the standard molar enthalpy change of formation of an element in its standard state is always zero. This is because no heat change is involved when an element is formed from itself.

4.1 The Stability of A Compound
There are two types of chemical stability and energetic stability. The standard enthalpy change of formation (ΔHof ) is only related to the energetic stability. Energetic stability is also known as the thermodynamic stability. In the other words, the value of ΔHof does not show the rate of reaction, which is indicated by the kinetic stability.
For example, diamond is energetically unstable but kinetically stable. At ordinary temperatures and pressures, diamond is not stable enerdetically compared with its allotrope, graphite.
C(diamond) → C(graphite) ΔHof = -2 kJ/mol
The negative value of ΔHof shows that diamond tends to change into graphite. However, the rate of change from diamond to graphite is extremely slow at room temperature. Hence, diamond is kinetically stable.

4.2 Using ΔHof Values to Predict the Relative Stability of A Compound
The standard enthalpy change of formation of a compound represents the energy transferred to or absorbed from the surroundings when chemical bonds are formed in the compound. ΔHof is therefore a measure of the energetic stability of the compound relative to its constituent elements. The more negative the value, the more energetically stable the compound.
For example, exothermic compounds with very negative values of ΔHof are very stable and do not decompose to their constituent elements readily. Conversely, endothermic compounds with very positive ΔHof values are unstable and tend to decompose (and may even explode) under normal conditions or on heating gently.
The enthalpy change of formation of a compound is a useful guide to its energetic stability and reactivity. In general endothermic compounds are unstable and highly reactive. When heated, many of them decompose readily into simple substances with the release of energy. It should be noted that exothermic compounds which are stable with respect to decomposition to its elements may be very reactive reagent.
It is not useful to compare ΔHof of compounds that do not contain the same elements, for example, carbon dioxide (CO2) with silicon (IV) chloride (SiCl4). In contrast, the comparison between the halogen hydrides is very useful because it shows the effect of the increasing size of the halogen atom on the energetic stability of the hydride.
For the halogens, the atomic size increases in the order F < Cl < Br < I
The smaller the atomic size of the halogen, the more stable the halogen hydride. Hence, hydrogen iodide, an endothermic compound, decomposes at low temperatures to its elements but hydrogen fluoride, a very exothermic compound, is very stable towards heat. The longer the covalent bond length, H-X, the weaker the covalent bond.

5.0 Standard Molar Enthalpy Change of Combustion, ΔHoc
The standard enthalpy change of combustion (ΔHoc) of a substance is the standard enthalpy change when one mole of the pure substance is completely burned in excess oxygen under standard conditions. Since heat is always given out during the combustion of a substance, the value of ΔHoc is always negative. For example, the standard enthalpy change of combustion for methane is -890 kJmol-1. This means that when one mole of methane is burnt in excess oxygen, 890kL of heat is liberated. That is, CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l) ΔHoc = -890kLmol-1
Any device that is used to measure temperature changes resulting from the enthalpy changes due to chemical reactions or physical changes is called calorimeter. The standard molar enthalpy change of combustion of a substance can be determined by using (a) a simple calorimeter (b) a flame calorimeter (c) a bomb calorimeter

5.1 Standard Enthalpy Change of Combustion and Molecular Structure
When chemical changes take place, energy must be supplied to break the bonds between atoms in the reactants and energy is liberated when new bonds are formed between these atoms to produce the products. In the other words, bond breaking in an endothermic process while bond forming is an exothermic process. Therefore, the enthalpy change of reaction is the difference in energy between the bond-breaking process and the bond-forming process.
For example, the similar values of ΔHoc is expected in butane and methylpropane because both the compounds contain the same number of atoms and the same type of chemical bonds. This suggests that each chemical bond makes a specific contribution to the total enthalpy change.

Hydrocarbon | Butane | Methylpropane | Molecular Formula | C4H10 | C4H10 | Structure Formula | | | ΔHoc (kJmol-1) | -2877 | -2870 |
Table: The heats of combustion of butane and methylpropane

6.0 Standard Molar Enthalpy Change of Neutralisation, ΔHonuet
The standard molar enthalpy of neutralisation is the enthalpy change when one mole of H+ (aq) ions from an acid reacts with one mole of OH- (aq) ions from an alkali to form one mole of water molecules under standard conditions. The standard molar enthalpy change of neutralisation is often called enthalpy change of neutralisation. The value of ΔHonuet is always negative (exothermic).

6.1 The Standard Enthalpy Change of Neutralisation For Strong Acids and Strong Alkalis
The fundamental change common to all neutralisation reactions is the reaction between H+ (aq) ions from the acid and OH- (aq) ions from the alkali to form water molecules.
H+ (aq) + OH- (aq) → H2O (l)
Strong acids and strong alkalis dissociate almost completely in aqueous solution. This is confirmed by the fact that the standard molar enthalpy change of neutralisation of any strong acid and any strong base is almost constant, that is, -57.3 kJmol-1. For example,
NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l) ΔHo = -57.3 kJmol-1
½ Ca (OH)2 (aq) + HNO3 (aq) → ½ Ca (NO3)2 (aq) + H2O (l) ΔHo = -57.3 kJmol-1

6.2 The Standard Enthalpy Change of Neutralisation of Weak Acids and Weak Alkalis
If a weak acid or a weak alkali is used, the standard enthalpy change of neutralisation is different from -57.3 kJmol-1. For example,
HCN (aq) +KOH (aq) → KCN (aq) + H2O (l) ΔHo = -11.7 kJmol-1
CH3COOH (aq) + NaOH (aq) → CH3COONa (aq) + H2O (l) ΔHo = -55.6 kJmol-1

The standard enthalpy change for a weak acid and a strong acid is lower than -57.3 kJmol-1. This is because the weak acid does not dissociate completely. Therefore when the acid reacts, some of the heat energy liberated during neutralisation is absorbed in the dissociation acid

7.0 Standard Molar Enthalpy Change of Atomisation, ΔHoat
The standard enthalpy change of atomisation of an element is the heat energy absorbed when one mole of gaseous atoms are formed from its element under standard conditions. For example,
Na (s) → Na (g) ΔHoat = +108 kJmol-1
½ H2 (g) → H (g) ΔHoat = +216.0 kJmol-1
For liquid, the enthalpy change of atomisation includes the enthalpy change of vaporisation, that is, the heat energy required to change a liquid to a vapour. For solids, the enthalpy change of atomisation includes the enthalpy change of fusion (that is, the heat energy required to change a solid to a liquid) and the enthalpy change of vaporisation. In some solid elements, the enthalpy change of atomisation is equal to the enthalpy change of sublimation of the solid. The enthalpy changes of atomisation of the noble gases (Ne, Ar etc.) are zero.

8.0 The Electron Affinity
The electron affinity is also known as the electron-gain enthalpy. Anions are formed by atoms accepting electrons. The electron affinity of an element is the enthalpy change when a mole of gaseous atoms accepts a mole of electrons to form a mole of uninegative charged ions. For example,
Cl (g) + e- → Cl- ΔHo = -364 kJmol-1
The electron affinity of an element refers to atoms and ions in their gaseous states. The value of the first electron affinity is always negative, that is, heat energy is liberated when an anion is formed.
When an atom has already accepted an electron, the atom becomes a negatively charged ion, and this negative charge will repel the addition of a second electron. As a result, energy must be absorbed to overcome the repulsive forces between the two negatively charged particles. Therefore, the second and successive electron sffinities of any element are always positive.
Electron affinity is not the same as electronegativity. The electron affinity is ameasure of the ability of atoms in the gaseous phase of an element to accept electrons to form gaseous negative ions, while electronegativity is a measure of the ability of an atom to attract pairs of electrons in the covalent bond of a molecule to itself (without forming an ion).

9.0 Lattice Energy
The lattice energy is also known as the lattice enthalpy. The forces of attraction that join together ions of opposite charges in a crystal lattice are electrostatic in nature. The lattice energy is defined as the enthalpy change when one mole of crystalline substance is formed from its gaseous ions.
The value of lattice energy is always negative because the formation of ionic bonds is an exothermic reaction. Conversely, heat energy must be supplied (that is, endothermic) to break the strong bonds between ions in the crystal.

9.1 The Magnitude of Lattice Energy
The magnitude of lattice energy depends on two main factors, the charge on the ions and the inter-ionic distance. (1) The charge on the ions
The greater the charges on the ions, the greater the attraction between them and the greater will be the value of the lattice energy. (2) The distance between the ions
The smaller the distance between the ions, the greater the attraction between them and the greater will be the value of the lattice energy.

10.0 The Born-Haber Cycle
The Born-Haber cycle can be defined as a cycle of reactions used for calculating the lattice energy of ionic crystalline solids.
This involved the process: (1) Atomisation of the elements: enthalpy of atomisation (2) Formation of the cation: ionisation energy (3) Formation of the anion: electron affinity (4) Formation of the ionic crystal: lattice energy
For example,

The Born-Haber cycle for sodium cloride

11.0 References
1. Tan Y. T., “Physical Chemistry for STPM,” Penerbitan Fajar Bakti SDN.BHD., 2004
2. http://chemistry.bd.psu.edu/jircitano/BH.html
3. http://en.wikipedia.org/wiki/Thermochemistry

References: 1. Tan Y. T., “Physical Chemistry for STPM,” Penerbitan Fajar Bakti SDN.BHD., 2004 2. http://chemistry.bd.psu.edu/jircitano/BH.html 3. http://en.wikipedia.org/wiki/Thermochemistry