Thermodynamics Flashcards
Enthalpy of formation
The enthalpy change when 1 mole of the compound is formed from its elements under standard conditions, all reactants and products being in their standard states.
Na(s) + 1/2Cl2(g) → NaCl(s)
Is enthalpy of formation endothermic or exothermic
Exothermic for most substances
Enthalpy of atomisation
The enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state
Na(s) → Na(g)
1/2O2 → O(g)
Is enthalpy of atomisation exothermic or endothermic
Endothermic
First ionisation energy/enthalpy
The enthalpy change needed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous ions with a charge of +1.
Mg(g) → Mg+(g) + e-
Bond dissociation enthalpy
The standard molar enthalpy change when one mole of a covalent bond is broken into two gaseous atoms
Cl2(g) → 2Cl(g)
CH4(g) → CH3(g) + H(g)
Is ionisation energy endothermic or exothermic
Endothermic
Is bond dissociation enthalpy endothermic or exothermic
Endothermic
What happens to the bond dissociation enthalpy for diatomic molecules
It is the same as x2 of the enthalpy of atomisation of the element
First electron affinity
The enthalpy change that occurs when 1 mole of gaseous atoms gain 1 mole of electrons to form 1 mole of gaseous ions with a -1 charge.
O(g) + e- → O-(g)
Is first electron affinity exothermic or endothermic
Exothermic for atoms that normally form negative ions / non-metals
Why is first electron affinity exothermic for many non-metals
The ion is more stable than the atom and there is an attraction between the nucleus and the electron.
Second electron affinity
The enthalpy change when one mole of gaseous 1- ions gains one electron per ion to produce gaseous 2- ions.
O-(g) + e- → O2-(g)
Is the second electron affinity exothermic or endothermic
Endothermic
Why is the second electron affinity endothermic
It takes energy to overcome the repulsive force between the negative ion and the electron
Enthalpy of lattice formation
The enthalpy change when 1 mole of an ionic compound is formed from its constituents in gaseous form.
Na+(g) + Cl-(g) → NaCl(s)
Is enthalpy of lattice formation exothermic or endothermic
Exothermic
Enthalpy of lattice dissociation
The enthalpy change when one mole of an ionic compounds dissociates into its constituent ions in gaseous form.
Is enthalpy of lattice dissociation exothermic or endothermic
Endothermic
Enthalpy of hydration
The enthalpy change when one mole of gaseous ions become aqueous ions.
X+(g) + aq → X+ (aq)
Mean bond enthalpy
The energy required to break a covalent bond in a molecule averaged over a range of compounds.
Why is the first ionisation energy of barium less negative/ less endothermic than that of calcium
The atomic radius of barium is bigger than calcium
There is more shielding of the outer electrons in barium
The nuclear attraction decreases, so less energy is needed to remove the outer electron in barium.
Describe the steps to draw a born haber cycle
- Write the lattice enthalpy equation.
- If there is a number in front of any ions, multiply any step involving that ion by the number
- Enthalpy of formation
- Enthalpy of atomisation of positive ion
- Ionisation energies
- Enthalpy of atomisation of negative ion
- Electron affinities
- Lattice enthalpy
How to calculate enthalpy of formation from born haber cycle
Enthalpy of formation = everything else added up
Why is the formula of magnesium chloride MgCl2 and not MgCl3
A large amount of energy is needed to remove the third electron from a magnesium atom as it is removed from a different stable shell which is closer to the nucleus and shielded by fewer inner shells.
The large amount of energy needed for the 3rd ionisation energy of magnesium is not compensated for by the lattice enthalpy. Therefore the enthalpy of formation of MgCl3 would be highly endothermic and so would not form as it is energetically less stable than the elements.
Which two factors affect the magnitude of lattice enthalpy
Radius of ion/size of ion
Charge of the ion
Why is the lattice enthalpy of CaO less negative than MgO
The charge on both positive ions is the same, +2, but Ca 2+ is bigger than Mg 2+.
Ca 2+ has a smaller charge density than Mg2+.
The electrostatic forces of attraction between Ca2+ and O2- is weaker, therefore the lattice enthalpy of CaO is less negative than MgO.
Polarising power
The ability of the positive ion to distort the electron cloud of the negative ion
What factors affect polarising power
Charge and size of ion.
A smaller, more charge positive ion is more polarising.
A larger negative ion is more easily polarised
Which ion has a greater polarising power, Na+ or Al3+
Al3+, as it is smaller than Na+ and has a greater charge.
Al3+ has a greater charge density than Na+ so is more polarising
What happens to covalent character as you go down a group
Increases
What are experimental lattice enthalpies
Values of lattice enthalpy calculated from Born-haber cycles
What are theoretical lattice enthalpies
Values of lattice enthalpy calculated on the basis of a purely ionic model existing
What does the ionic model assume
Perfectly spherical ions
Ions 100% ionic
Attractions purely electrostatic
Uniformly distributed charge
Is the experimental more or less negative than the theoretical value for lattice enthalpy and why
Experimental value is always more negative than theoretical value as covalent character strengthens the lattice
Why are the calculated and experimental values for lattice enthalpy in BaCl2 similar
The bonding in barium chloride is almost 100% ionic
Ba2+ and Cl- exist as separate spherical ions
The large Ba2+ ions have a low polarising power so cannot distort the electron cloud of Cl-
Explain why the theoretical enthalpy of lattice dissociation for silver fluoride is different from the experimental value that can be calculated using a Born–Haber cycle.
Experimental lattice enthalpy value allows for / includes covalent interaction /
non–spherical ions / distorted ions / polarisation
OR AgF has covalent character
Theoretical lattice enthalpy value assumes only ionic interaction / point charges
/ no covalent / perfect spheres / perfectly ionic
Entropy
Measure of disorder.
A measure of the dispersal of energy in a system at a specific temperature.
Entropies of the 3 states of matter
Gases > Liquids > Solids
Why is the entropy at 0K zero
The particles have no energy and so are not moving. There is no disorder.
Examples of change which increase entropy
Increase temperature
Melt a solid / boil a liquid
Increase in number of gas molecules
Increase in the number of moles in the same physical state
Ionic solid dissolves in water
Large molecule breaks into smaller molecule
Explain the meaning of spontaneous in a thermodynamic context
delta G is negative
What happens to delta S if system is more disordered
delta S is more positive
What happens to delta S if the system is less disordered
delta S is negative
Why does ammonia have a higher entropy than water
Ammonia is a gas whereas water is a liquid in standard conditions.
The molecules in ammonia move randomly in all directions and are more disordered than the molecules in liquid water.
Why does propane have a higher entropy than argon even though they are both gases
Argon consists of single atoms. Propane has molecules made of several atoms. Larger molecules with more atoms have more ways to rotate and vibrate increasing the number of ways that the energy of molecules can be spread
Why does pentene have a higher entropy than cyclopentane. They both have the same molecular formula and are both gases.
The ring in cyclopentane restricts the freedom of the molecules to rotate about single bonds so cyclopentane is less disordered.
Delta S =
Sum of S of products - Sum of S of reactants
Units of delta S
J K^-1 mol^-1
When is a reaction feasible
When delta G is 0 or negative
delta G =
delta H - (T x delta S)
units of delta G
kJmol^-1
units of temperature in Gibbs free energy equation
K
Explain why this line obeys the mathematical equation for a straight line, y=mx + c with the aid of a thermodynamic equation
delta G = delta H - (T x delta S)
delta H is the y intercept
-delta S is the gradient
Explain the difference between the hydration enthalpies of the magnesium and sodium ions
Magnesium ion is smaller and more charged so has a greater charge density than the sodium ion.
Magnesium ion attracts water more strongly
A 5.00 g sample of potassium chloride was added to 50.0 g of water initially at 20.0 °C. The mixture was stirred and as the potassium chloride dissolved, the temperature of the solution
decreased. Describe the steps you would take to determine an accurate minimum temperature that is not influenced by heat from the surroundings.
Start a clock when KCl is added to water
Record the temperature every subsequent minute/ at regular intervals for about 5 minutes
Plot a graph of temperature vs time
Extrapolate back to time of mixing = 0 and determine the temperature
In terms of electrostatic forces, suggest why the electron affinity of fluorine is negative
There is an attraction between the nucleus / protons and (the added) electron(s)
Energy is released (when the electron is gained)
Suggest why the hydration of a chloride ion is exothermic
Water is polar / water has Hδ+
Chloride ion attracts (the H in) water molecules
The enthalpy of solution for potassium chloride is +17.2 kJ mol−1. Explain why the free-energy change for the dissolving of potassium chloride in water is negative, even though the enthalpy change is positive.
The entropy change is positive / entropy increases
Because 1 mol (solid) → 2 mol (aqueous ions) / no of particles increases
Therefore TΔS > ΔH
Explain why the evaporation of water is spontaneous even though this change is
endothermic. In your answer, refer to the change in the arrangement of water molecules and the entropy change.
The molecules become more disordered / random when water changes from a liquid
to a gas / evaporates.
Therefore the entropy change is positive / Entropy increases
TΔS>ΔH
ΔG<0
Explain why the hydration enthalpy of the fluoride ion is more negative than the hydration enthalpy of the chloride ion
Fluoride ions are smaller and so have greater charge density.
So negative charge attracts the delta positive H on water more strongly.
What does a negative value for enthalpy of solution mean
The ionic compound wil dissolve
The enthalpy change of hydration of Rb+ is -280. Explain how the enthalpy change of hydration of Na+ will differ from this value.
The enthalpy change of hydration of Na+ will be more negative than that of Rb+. Na+ is smaller than Rb+, so Na+ has a greater charge density than Rb+. There is a stronger force of attraction between Na+ and water molecules.
What does a born haber cycle look like for ionic compound which is soluble in water
The base has nothing on the left.
There is an arrow downwards to the base for enthalpy of solution.
From the top of the enthalpy of solution arrow there is the solid compound.
There is an arrow pointing down onto the solid compound from above which is lattice enthalpy.
On the top base from the LE arrow is the gaseous ions.
On the right of the top base is an arrow pointing down for enthalpy of hydration of 1 ion
Another arrow after that downwards for enthalpy of hydration of the other ion.
This goes all the way to the base.
What does a born haber cycle look like for ionic compound which is INsoluble in water
On the base is the solid ionic compound.
There is an arrow from above pointing down to the solid which is lattice enthalpy.
The top base has the gaseous ions.
From the right of the top base there are 2 arrows down from hydration of each ion.
After the second hydration arrow downwards, there is an arrow pointing up from the base for enthalpy of solution.
Why do some solids with a positive enthalpy of solution dissolve
∆G = ∆H - T∆S
∆H is positive because the reaction is endothermic
When the solid dissolves, the disorder increases so ∆S is positive
The magnitude of T∆S is greater than ∆H so ∆G is negative.
Since ∆G<0 the change is feasible and compound dissolves
