Chapter 22 - Enthalpy and Entropy Flashcards
what is a ‘real terms’ explanation of what lattice enthalpy is
how strongly a giant ionic lattice is held together
what is the definition for lattice enthalpy
Lattice Enthalpy is the energy released when one mole of ionic crystal is formed from its constituent ions in their gaseous state under standard conditions of 298K and 100Kpa
what sign does lattice enthalpy usually have and why
negative because the crystal is being formed so energy is released (exothermic) because bonds are being made
what effect does increasing the charge of the ions have on the magnitude of the lattice enthalpy
it increases the lattice enthalpy as there is a greater electrostatic attraction between the ions
it tends to be that more positive ions are also smaller so there is a greater charge density
what effect does decreasing the size of the ion have on the magnitude of the lattice enthalpy
it increases the lattice enthalpy as there is the ‘same charge’ in a smaller ion, thus there is a greater charge density so the electrostatic attraction increases
can we measure lattice enthalpy directly
no
what is usually step 1 of a born haber cycle
atomisation - the molecules or lattice of the constituent elements breaks down to leave gaseous atoms
define standard enthalpy of atomisation
the enthalpy of atomisation of an element is the enthalpy change when one mole of gaseous atoms is formed from the element in its standard state, 298K and 100Kpa
what does the sign tend to be for enthalpy of atomisation and why
it tends to be positive (endothermic) as it requires energy to break down the molecules or lattice of the element in its standard state
what is usually step 2 of a born haber cycle
ionisation- the atoms are ionised (an electron is removed)
some atoms of elements may need 2nd or 3rd ionisations
define enthalpy of ionisation
i.e. ionisation energy
first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms (to form one mole of 1+ ions and one mole of electrons)
what is often the 3rd step of a born haber cycle
electron affinity- the formation of 1- ions from the non-metal
what sign does ionisation energy usually have and why
it usually has a positive value (endothermic) because it requires energy to overcome the electrostatic attraction between the positive nucleus and the negative electron
define electron affinity
the electron affinity is the enthalpy change accompanying the gain of one mole of electrons by one mole of atoms in the gaseous phase
what signs do the different stages of electron affinity have and why
- the first electron affinity is usually negative (exothermic) because there is an electrostatic attraction between the positive nucleus and the negative electron
- any further electron affinities will have a positive sign (endothermic) because energy is required to overcome the repulsion between a negative ion and a negative electron
define standard enthalpy of formation
the enthalpy change that occurs when one mole of a compound is formed from its constituent elements in their standard states and under standard conditions of 298K and 100Kpa
What do we need remember when adding values to a born-haber cycle
always multiply by the relevant number of moles because ionisation energies and energies of atomisation are for 1 mole formed
what is the overall energy change associated with dissolving called and define it
The standard enthalpy change of solution
“The standard enthalpy change of solution (deltaHsol) is the enthalpy change that takes place when 1 mol of an ionic solid dissolves completely in water at 298K and 101KPa”
what two factors affect the size of the enthalpy change of solution
- the energy required to break down the ionic lattice
- the energy released due to the attractions of the ions with the Delta +ve hydrogens and Delta -ve oxygens
explain in terms of the factors and enthalpies what an endo/exothermic enthalpy change of solution actually means
- endothermic, more energy was required to break down the ionic lattice than is released by the attractions to the water molecules
lattice enthalpy > enthalpy of hydration - exothermic, more energy is released by the attractions of the ions to the water molecules than is required to break down the ionic lattice
enthalpy of hydration > lattice enthalpy
what is the equation to represent the dissolving of sodium chloride
NaCl(s) —> Na+ (aq) + Cl-(aq)
explain in full the process of dissolving
- the ionic lattice breaks up and forms separate gaseous ions - this is represented by the lattice dissociation enthalpy
- the water molecules attract to and surround the ions to form aqueous/hydrated ions - represented by enthalpy of hydration
define the enthalpy change of hydration
” The enthalpy change of hydration is the enthalpy change which occurs when one mole of gaseous ions of an element are completely hydrated by water at 298K and 101KPa”
write two equations to represent the hydration of sodium and chloride ions
Na+ (g) —> Na+ (aq)
Cl- (g) –> Cl- (aq)
how can we represent the process of dissolving with a diagram
use another Born-Haber cycle style thing where there’s lattice dissociation enthalpy going up and enthalpy change of solution going down from the start and enthalpy of hydration going from the top to the bottom
what is the method for determining the enthalpy change of solution for a substance
- Weigh out a known mass of a substance, calculate the number of moles
- measure out 25 cm^3 of distilled water using a measuring cylinder
- record the initial temperature of the water
- add the substance
- record the max/min temp.
- calculate energy change using q=mc(deltaT) (remember m is mass of water + substance)
- use q/n for enthalpy change
- improve using pipette not measuring cylinder and using a cooling curve
what are the two factors which affect lattice enthalpy and enthalpy of hydration and what is the overall effect of these factors
- ionic charge
- ionic size
- these combine to have the effect of charge density
explain how ionic size affects lattice enthalpy
- as ionic size increases, charge density decreases
- thus attraction between ions decreases
- so less energy is required to overcome the electrostatic forces between ions
- so melting point and lattice enthalpy both decrease
explain how ionic charge affects lattice enthalpy
- a greater ionic charge gives a higher charge density
- thus there is a greater attraction between ions
- thus more energy is required to overcome the electrostatic attraction between ions
- so melting point and lattice enthalpy increase
explain how ionic size affects enthalpy of hydration
- a greater ionic size gives a lower charge density
- thus there is a lower attraction to the polar water molecules
- thus the enthalpy of hydration is lower/ less energy given out/ less negative
explain how ionic charge affects enthalpy of hydration
- as ionic charge increases, charge density increases
- thus there is a greater electrostatic attraction to the polar water molecules
- so enthalpy of hydration increases/ more energy given out/ more negative
define entropy and give its units
” Entropy, S, is a measure of disorder in a system, it is the number of ways particles can be arranged and the number of ways energy can be shared between particles”
units = J K^-1 mol^-1
what generally happens to entropy as the state of the substance changes and why
- entropy tends to increase as you go from solid to liquid to gas
- because melting and boiling increase randomness of particles
- and energy is more spread
what is the effect of increasing temperature on entropy and why
- increasing temperature increases entropy
- because there are a greater number of ways for the energies in the particles to be distributed
what is the effect of gases and dissolving on entropy
- more moles of gas increase disorder of particles because the energy is more spread out so entropy increase
- dissolving a substance also increases entropy for the same reason
what are the units of enthalpy and entropy
enthalpy = KJmol^-1 entropy = J K^-1 mol^-1
do enthalpy and entropy have absolute 0 values
- enthalpy has no absolute 0 value, we just measure changes from an arbitrary 0 value
- entropy does have an absolute 0 at temp = 0K
how to calculate entropy of system
change in entropy of the system (deltaSsys) = (sum of entropy of products) - (sum of entropy of reactants)
how to calculate entropy of surroundings
entropy of surroundings (deltaSsurroundings) = - (deltaH)/T
calculating overall entropy change
overall entropy change = entropy change of system + entropy change of surroundings
(deltaStot) = (deltaSsys) + (deltaSsurround)
what must deltaS be for a reaction to be feasible
- it must be positive
what must happen for a reaction to be feasible overall
- the overall energy change must be negative
what is the Gibbs equation
deltaG = deltaH - TdeltaS
what are the units of the gibbs equation and what should we watch for
deltaG = Kjmol^-1 deltaH = Kj T = kelvin deltaS= kj K^-1 mol^-1
how to work out how conditions vs sign of enthalpy and entropy affect feasibility
draw graphs of deltaG against T for different combinations of +ve and -ve and where deltaG < 0, the reaction is feasible
how to find mimimum/max temp that a reaction is feasible at given the gibbs equation
the limit temperature occurs where deltaG = 0
thus deltaH - TdeltaS = 0
so T = deltaH/deltaS
what are the limitations of working with the gibbs equation
it gives no indication of activation energy or rate of reaction
when is a reaction feasible (energy) and why is this
when deltaG < 0
so energy is released
deltaG = -T(deltaStot)
so if deltaStot must be +ve and temperature is in Kelvin then DeltaG must be negative
