Enthalpy and Entropy

Lattice enthalpy is the energy change required for the formation of one mole of an ionic compound to be formed from its gaseous ions under standard conditions. It is a measure of the strength of the ionic bonding in a giant ionic lattice. 

Lattice enthalpy can not be measured directly and must be calculated indirectly using an energy cycle - a Born-Haber cycle. 

There are several key enthalpy changes involved in the Born-haber cycle: 

Enthalpy change of formation - the enthalpy change that takes place when one mole of a compound is made from its elements under standard conditions.    Na + 1/2Cl2 --> NaCl 

Enthalpy change of atomisation - the enthalpy change that takes place for the formation of one mole of gaseous atoms from the element in its standard state under standard conditions. It is always an endothermic reaction because bonds are being broken.     Na(s) --> Na(g)

First Ionisation energy - the enthalpy change required to remove one electron from each atom in one mole of gaseous atoms to form one mole of mole of gaseous 1+ ions. This is endothermic becuse energy is required to overcome the attraction of ions to the electrons.                          Na(g) --> Na+(g) + e-

First eletron affinity - the enthalpy change required to add one electron to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions. It is the opposite of Ionisation energy. First electron affinities are exothermic because the electron is being attracted to the nucleus. Second electron affinites are always endothermic because the negative electron repels the electron away.  Cl(g) + e- --> Cl-

Standard enthalpy change of solution

 It is the enthalpy change that takes place when one mole of a solute dissolves fully in a solvent. If the solvent is water, then the ions will end up surrounded by water molecules as aqueous ions. This is because the delta + and - partial charges on water are attracted to the charged ions. Na+Cl- (s) + aq ---> Na+(aq) + Cl-(aq)  It can be exothermic or endothermic. 

Determining enthalpy change of sol^ experimentally: 

1) Weigh a sample of ionic compound                                                                                          

2) Fill a beaker with water and measure the temperature.                                                          

3) Add the ionic compound to the water, stir, and wait for the temperature to no longer change       and then record it.                                                                                                                          

4) From the data you have obtained, calculate q = mcAT, work out NOM of ionic compound and divide q value by NOM to get enthalpy change of solution. 

nb. for mass use the mass of total solution - (mass of water + mass of solute) 

When a solid ionic compund dissolves in water, two process take place: 

• Ionic lattice breaks up forming seperate gasseous ions (opposite to lattice enthalpy) 
• The gaseous ions interact with dipoles on water molecules to produce hydrated aqueous ions. This part of the process is called Enthalpy change of hydration.

Enthalpy change of hydration is the enthalpy change that takes place when gaseous ions are dissolved in water to form one mole of aquesous ions.

Cl- (g) + aq --> Cl-(aq) 

You can also use a type of energy cycle to work out unknown enthalpy changes. 

An example of an enthalpy cycle to show process of disolving an ionic product in water:

LATTICE ENTHALPY: the values of lattice enthalpies depend on ionic size and ionic charge.

Effect of ionic size: 

As ionic radius increases, attraction between the ions decreases. This means that it is easier to seperate ions. Lattice enthalpy is therefore less negative as ionic radius increases/ you go down a group because more energy is required to make an ionic lattice. Melting point decreases as ionic size increases as less energy is required to seperate ions. 

Effect of ionic charge: 

As ionic charge increases, attraction between ions also increases. This makes it harder to seperate ions. Lattice enthalpy therefore is more negative as ionic charge increases/across the beginning of a period because it is easier to maintain the ionic lattice. The melting point also increases as ionic charge increases as more energy is required to seperate ions. 

The magnitude of lattice enthalpy is a good indication of a melting point. The more negative (exothermic) a lattice enthalpy, the higher the melting point will be. 

HYDRATION: these are also affected by ionic charge and ionic size. 

ionic size: As the ionic radius increases, attraction between ions and the water molecules in solution decreases, due to greater distance from the nucleus. This means that hydration becomes less negative. It is harder for water molecules to surround the gaseous ions to make aqueous ions.

ionic charge: As the ionic charge increases, attraction between the ions and water molecules increases. This means that hydration becomes more negative. It is easier for the water molecules to surround and ineract with the ions and aqueous ions can be formed more easily. 

This allows you to predict solubility: 

• If the magnitude of hydration enthalpies is greater than that of the lattice enthalpies, then the overall enthalpy change is exothermic = will disslove. 
• If the magnitude of hydration is less than that of lattice enthalpies, then the overall enthalp change is endothermic = will not dissolve.

However this does not provide the whole picture as temp. and entropy also effect solubility.

Entropy (S): used to describe the level of disorder in a chemical system. The greater the entropy, the greater the dispersal of energy and therefore the greater disorder. 

Units for entropy: J K-1 mol-1

In general: Solids have the smallest entropies and gases have the largest entropies. During changes of state, entropy values change: Eg, melting and boiling increases entropy values.

Predicting entropy changes: At 0K, entropy values are 0. Above 0K, energy becomes disperesed so all substances have a positive entropy value. 

• If a system becomes more random, the energy will be spread out more - entropy change will be positive.
• If a system becomes less random, energy becomes more concentrated - entropy change will be negative. 

Change in number of gaseous molecules: 

• A reaction that produces a more moles of gas will have a positive entropy change because the presence of a gas increases the disorder of particles. 
• A reaction that produces a few number of moles of gas (Eg, Ammonia) will have a negative entropy change because particles become less disordered. 

All substances have a standard entropy value, found in data books. This is the entropy of one mole of a substance, under standard conditions. They are all positive values.

Calculating entropy changes:

Change in entropy = entropies of procducts - entropies of reactants 

Feasbility: describes whether a reaction is able to take place and is energetically feasable. 

Free energy: the overall energy change in a chemical reaction is called the free energy change. It is made up of two types of energy: 

• the enthalpy change with its surroundings (loss or gain of energy).
• the entropy change at the temperature of the reaction. 

The Gibbs eqation shows the relationship between these two types of energy:               

conditions for feasibility: It depends on the balance between the two types of energy. For a reaction to be feasible:  the free energy change (deltaG) < 0

NB. When working out  the Gibbs equation, the entropy change is usually given in JK-1mol-1 so it must be converted to KJK-1mol-1, by dividing by 1000. The balance betweent these two energies: 

Endothermic reactions can take place at room temperature?

Some ionic compounds dissolve in water at room temperature. This is dependant on entropy and free energy. NB. if delta G is < 0, then it will be feasible. 

Limitations of predictions made for feasibility: 

The free energy change is useful for predicting feasibility, but there are many equations that have a negative deltaG and still dont take place.

This is because many reactions require a very large activation energy in order to start the reaction, and therefore the reactions have a very slow rate. With a catalyst many of these reactions would then take place.

Therefore, the Free engergy change tells us the thermodynamic feasibility, but does not take into account the kinetics or rate of reaction.