enthalpy and entropy year 2

Question 1

lattice enthalpy

Answer 1

• a measure of the strength of ionic bonding in a giant ionic lattice
• enthalpy change that accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions
• involves ionic bond formation so is an exothermic change and the value for the enthalpy change will always be negative
• cannot be measured directly so calculated indirectly using born-haber cycles

Question 2

standard enthalpy change of formation

Answer 2

• enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions, with all reactants and products in standard states

Question 3

enthalpy change of atomisation

Answer 3

• enthalpy change that takes place for the formation of one mole of gaseous atoms from the element in its standard state under standard conditions
• always endothermic as bonds are broken to form gaseous atoms
• when the bond is already a gas in its standard state, the enthalpy change of atomisation is related to the bond enthalpy of the bond being broken
• as its one mole of gaseous atoms forming, when drawing born haber cycles, remember to multiply by 2 if 2 atoms are created eg if I2 to2I

Question 4

first ionisation energy

Answer 4

• enthalpy change required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions
• always endothermic as energy is required to overcome the attraction between negative electrons and positive nuclei

Question 5

first electron affinity

Answer 5

• enthalpy change that takes place when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions
• exothermic as electron being added is attracted to the nuclei

Question 6

successive electron affinities

Answer 6

• happens when an ion needs a 2- charge eg O2
• second electron affinities are endothermic. A second electron is being gained by a negative ion, which repels the electron away, so energy must be put in to force the negatively-charged electron onto the negative ion. This would be shown on a born haber cycle by a positive energy change

Question 7

standard enthalpy change of solution

Answer 7

• enthalpy change when one mole of a solute dissolves in a solvent.
• If the solvent is water, the ions from the ionic lattice finish up surrounded with water molecules as aqueous ions
• eg Na+Cl- + aq –> Na+ (aq) + Cl- (aq)
• can be endothermic or exothermic
• attraction exists in three dimensions, so ions and water molecules above and bellow the plane of the paper
• partial charges of atoms on the water molecules are attracted to oppositely charged ions
• can work out experimentally, by using q=mcΔT. The mass should be the combined mass of the water and the solid added, as it is the solution that is changing temperature, so the mass of solution should be used in the calculations.

Question 8

dissolving process

Answer 8

When solid compound dissolves in water:
- the ionic lattice breaks up forming separate gaseous ions. This is the opposite energy change of lattice enthalpy
- water molecules are attracted to, and surround, the ions, to form hydrated aqueous ions. This energy change is called enthalpy change of hydration

Question 9

enthalpy change of hydration

Answer 9

• enthalpy change that accompanies the dissolving of gaseous ions in water to form one mole of** aqueous ions**
  eg Na+ + aq –> Na+ (aq)
• can be shown/ calculated on an energy cycle diagram, with the gaseous ions on the top, which are then split into aqueous ions and then into an ionic lattice

Question 10

affect of ionic size on lattice enthalpy

Answer 10

• as the size of the ion increases, the lattice enthalpy becomes LESS exothermic
• the ionic radius increases, the attraction between ions decreases, and the melting point decreases
• will have a lower charge density, than others with the same charge, if the size of the ion is larger

Question 11

effect of ionic charge on lattice enthalpy

Answer 11

• as ionic charge increases, lattice enthalpy becomes more exothermic
• the attraction between ions increase, the melting point increases, so lattice enthalpy is more negative
• also means has a higher charge density if ionic charge is bigger, than an ion of the same size but with a smaller charge
• as you go across the metal ions in group three of the periodic table, ion size decreases and ion charge increases, which explains why Al3+ will lead to compounds with a much higher lattice enthalpy. The size of the ion decreases as it is more protons attracting the same amount of electrons
• as you go from right to left for the non metals in group three you have two opposing effects. Increasing charge gives more attraction, but an increase in size gives less attraction
• magnitude of lattice enthalpy gives a good indication of the melting point of an ionic compound.

Question 12

factors affecting hydration enthalpies

Answer 12

• also affected by ionic size and charge
• as the size of an ion increases, the attraction between ion and water molecule decreases as ionic radius is increasing so hydration energy is LESS negative
• as charge of the ion increases, attraction with water molecules increases and so hydration energy becomes MORE negative

Question 13

predicting solubility

Answer 13

• if the sum of hydration enthalpies is larger than the magnitude of the lattice enthalpy, the enthalpy change of solution will be exothermic and the compound SHOULD dissolve
• however, many compounds with endothermic enthalpy changes of solution are soluble. Reasons for solubility also depends on temperature and entropy

Question 14

entropy

Answer 14

• things tend towards a state of disorder, as natural tendency for energy to spread rather than be concentrated in one place
  -entropy( ΔS) is a measure of the dispersal of energy in a system, which is greater the more disordered a system is. OR it is the number of ways that particles can be arranged
• greater the entropy, the greater the dispersal of energy and the greater the disorder
• units are J/K/mol. The greater the entropy, the greater that energy is spread out per Kelvin per mole
  In general:
• solids have the smallest entropies
• liquids have the greater entropies
• gases have the greatest entropies
• At 0K, there would be no energy and all substances would have an entropy value of zero. Above 0K, energy becomes dispersed amongst the particles and all substances have positive entropy
• in an equation you can predict whether entropy increases or decreases by comparing the physical states and amount of gas molecules on either side of an equation. Reactions that produce gases result in an increase in entropy. If number of gas molecules on the product side of the equation decreases, ΔS is negative

Question 15

changes of state

Answer 15

• entropy increases during changes of state that gives a more random arrangement of particles: solid-> liquid -> gas
  eg entropy increases when water melts, and then increases again when water evaporates

Question 16

standard entropies

Answer 16

• every substance has a standard entropy, which can be found in data books. It is the entropy of one mole of a substance under standard conditions (100kPa and 298K)
• standard entropies are always positive
• entropy change of a reaction= ΣSΘproducts - ΣSΘreactants. (where Σ = sum of).
• when calculating ensure to multiply standard entropy values by the coefficient

Question 17

Gibbs free energy

Answer 17

Gibbs’ equation:
ΔG = ΔH − TΔS
- can take the form of y=mx+c, where a graph is plotted of ΔG against T
- where ΔG is the free energy change (kJ/mol)
- ΔH is the enthalpy change. Transfer of heat between chemical systems and surroundings (kJ/mol)
- TΔS is the temperature(K) change times the entropy (J/K/mol). It is the dispersal of energy within the chemical system itself
- can also use G values to work out ΔG, this is also products- reactants

Question 18

conditions for feasibility

Answer 18

• feasibility- used to describe whether a reaction is able to happen and is energetically feasible (spontaneous)
• the feasibility of a reaction depends on the balance between ΔH and TΔS in the Gibbs’ equation. For a reaction to be feasible, there must be a decrease in free energy
  ΔG< 0
• for many reactions at room temperature ΔH has a much larger magnitude than TΔS and so ΔG is largely dependant on ΔH. As temperature increases, the TΔS term becomes more important
• to work out the minimum temperature for a reaction to be feasible, set ΔG to zero, and plug in ΔH and ΔS and then solve to find T

Question 19

limitations of predictions made for feasibility

Answer 19

• the free energy change is useful for predicting feasibility, but many reactions have a negative ΔG and don’t seem to take place
• eg decomposition of hydrogen peroxide
• this reaction doesn’t occur spontaneously at 298K, as there is a very large activation energy, which results in a very slow rate. Reaction takes take place by the addition of an MnO2 catalyst, which allows the reaction to take place via an alternative route with a lower activation energy
• although ΔG indicates the thermodynamic feasibility, it doesnt take into account the kinetics or rate of reaction