Thermodynamics Flashcards
Hess’s law
Enthalpy change for a chemical reaction is the same regardless of route taken
Standard enthalpy of formation
Enthalpy change when one mole of a compound is formed from its elements under standard conditions
All reactants and products in their standard states
Standard Enthalpy of combustion
Enthalpy change when one mole of a compound is completely burned in oxygen under standard conditions
All reactants and products in their standard states
Standard Enthalpy of atomisation
Enthalpy change when one of gaseous atoms is formed from an element in its standard state
First ionisation energy
Enthalpy change when one mole of electron is removed from one mole of atoms in the gaseous state
First electron affinity
Standard Enthalpy change when one mole of gaseous atoms is converted into a mole of gaseous ions
Lattice formation Enthalpy
Enthalpy change when one mole of solid ionic compound is formed from gaseous ions
Lattice dissociation Enthalpy
Standard Enthalpy change when one mole of solid ionic compound dissociated into its gaseous ions
Standard Enthalpy of hydration
Enthalpy change when one molecule of gaseous ions is converted into one mole of aqueous ions
Standard Enthalpy of solution
Enthalpy change when one mole of solute dissolves in enough solvent to form a solution in which the ions are far enough apart to not interact with each other
Important born-hater cycle rules
Enthalpy change is positive - arrow points upwards
Enthalpy change is negative - arrow points downwards
State symbols in every stage of the cycle
Order of born haber cycle
Start with elements in their standard states
Atomise the metal
Atomise the non metal
Ionise the metal
Electron affinity for the non metal
Lattice formation of the whole compound
What are the 2 factors that determine how exothermic a lattice Enthalpy will be
Charge on the ions
Size of the ions
Charge on the ions
Higher charge on ion the stronger the ionic bond
Greater lattice Enthalpy
More energy given out in the formation of lattice
More energy needed to break apart the ionic lattice
Size of the ions
Smaller the ion the stronger the attraction so stronger the ionic binding
So greater the lattice Enthalpy
More negative lattice formation so
Stronger ionic bonds
In answers use CRAM
charge of the ions,radius of the ions,attraction between ions and more exothermic/endothermic
Actual and theoretical value
If both values are close then indicates that the compound shows almost purely ionic bonding
If actual value is higher than theoretical value it’s because compound has some covalent character
What is covalent character
+ve ion is quite small which quite a strong +ve charge
-ve ion is large with a diffuse electron cloud
+ve ion is strongly polarising
The electron cloud about the -ve ion becomes distorted
Some of the electron density is shared
Perfect ionic model
Purely ionic bonding with no covalent character
Ions act as point charges
Ions act as perfect spheres which cannot be distorted
Stages of Enthalpy of solution
1-lattice dissociation
2-hydration
Enthalpy of solution calculation
> H Sol = >H LD + >H Hyd
Why do chemical reactions take place
Reaction has to be feasible and spontaneous
What determines if a reaction is feasible
Temperature
Enthalpy - most feasible reactions are exothermic
Entropy
Entropy meaning
Entropy is a measure of disorder
Units are jk-1mol-1
Is disorder increases entropy is positive
Is disorder decreases entropy is negative
Entropy increases with change of state s-l-g
If mols of products are higher than mol of reactants increases entropy
Calculating entropy change
Entropy change = (sum of entropy of products) - (sum of entropy of reactants)
Gibbs free entropy change
Gibbs = Enthalpy change - temp x entropy change
Gibbs and graphs
Y= Gibbs
X = t
M =- entropy change
Y intercept = Enthalpy change
