Chapter 22 Flashcards
Ionic bonds
Electrostatic attraction between oppositely-charged ions in the lattice structure
Lattice Enthalpy
Measure of the strength of ionic bonding in a giant ionic lattice
Lattice Enthalpy DEFINITION
Enthalpy change that accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions
Equation for KCl lattice Enthalpy?
K+ (g) + Cl- (g) -> KCl (s)
Gaseous ions -> solid ionic compounds
It is an EXOTHERMIC change and will always be negative as it involves bond formation (which releases energy)
Depict lattice Enthalpy on a cycle?
It is a downwards arrow from gaseous ions to the ionic lattice
What is a born-Haber cycle?
Cannot be measured directly and must be calculated indirectly
Born-Haber route for KCl Route 1?
Elements in standard states -> gaseous atoms (endothermic - formation of gaseous atoms)
Gaseous atoms -> gaseous ions (endothermic - formation of gaseous ions)
Gaseous ions -> ionic lattice (LATTICE ENTHALPY - EXOTHERMIC)
Born Haber cycle Route 2?
Changing from standard state directly to the ionic lattice - this is the Enthalpy change of FORMATION - this is Exothermic (just one Enthalpy change)
Standard Enthalpy change of formation?
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 their standard states
Standard Enthalpy change of atomisation? Give example for Cl2
Enthalpy change that takes place for the formation of one mole of gaseous atoms from the element in its standard state under standard conditions
1/2Cl2 (g) -> Cl(g)
Nature of Enthalpy change of atomisation?
Always endothermic as bonds are broken to form gaseous atoms ; when the element is a gas in its standard state, the standard Enthalpy change of atomisation = bond Enthalpy of bond being broken
Cl2(g) -> 2Cl(g) = +242kJ/mol
1/2 Cl2(f) -> Cl(g) = 121kJ/mol
First ionisation energy?
Enthalpy change required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions
Na (g) -> Na+ (g) + e-
Nature of ionisation energies?
Endothermic because energy required to overcome the attraction between a negative electron and the positive nucleus
Electron affinity
Opposite of ionisation energy - measures the energy to GAIN electrons - first electron affinity is the 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
Cl(g) + e- -> Cl- (g)
Nature of electron affinity
Exothermic because the electron being added is attracted in towards the nucleus
Ionisation energy/electron affinity of NaCl?
Ionisation energy of Na to Na+ (endothermic) - release an electron
Electron affinity of Cl to Cl- (Exothermic) - gain that electron
Second electron affinity?
With Oxygen - 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
First electron affinity Exothermic
What happens when salt dissolves in water?
Water molecules are able to break up the giant ionic lattice structure and overcome the strong electrostatic attractions between oppositely charged ions - this is what happens when salt dissolves in water
Standard Enthalpy of solution
Enthalpy change that takes place 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
Equation for standard Enthalpy change of solution of NaCl
Na+Cl- (s) + aq -> Na+ (aq) + Cl- (aq)
How can Enthalpy change of solution be determined experimentally?
Weigh a sample of the ionic lattice and pour a set amount of water into the plastic cup in the beaker (measure temperature) ; quickly tip all KCl into the water and stir until fully dissolved and the temperature no longer changes
Maths in experimentally determining Enthalpy change of solution
q = mc(change in temperature)
Where mass = mass of water (g) + mass of KCl
c = 4.18 (SHC)
Change in temperature
Figure out moles of KCl and then divide kJ of energy by moles
What happens when a solid ionic compound dissolves in water?
Ionic lattice breaks up forming separate gaseous ions - opposite energy change from lattice energy which forms the ionic lattice from gaseous ions
Water molecules are attracted to and surround the ions - separate gaseous ions interact with polar water molecules to form hydrated aqueous ions… this energy change is called the Enthalpy change of hydration
Enthalpy change of hydration
Enthalpy change that accompanies the dissolving of gaseous ions in water to form one mole of aqueous ions
NaCl (s) -> Na+(g) + Cl-(g) is endothermic you are breaking up the lattice (part of this is Enthalpy change of hydration (Exothermic or endothermic)
Exact opposite (just sign change) is formation of lattice
What are three levels in Enthalpy change hydration cycles?
Gaseous ions
Aqueous ions
Ionic lattice
Gaseous ions to ionic lattice
Lattice Enthalpy
Gaseous ions to aqueous ions
Enthalpy change of hydration
Ionic lattice to aqueous ions
Enthalpy change of solution
Ionic compounds
High melting and boiling points
Soluble in polar solvents
Conduct electricity when molten or in aqueous solution
Factors affecting lattice Enthalpy?
Effect of ionic size and effect of ionic charge
Effect of ionic size on lattice Enthalpy?
Ionic radius increases
Attraction between ions decreases
Lattice energy less negative
Melting point decreases
Effect of ion charge?
Ionic charge increases (best ratio 2:2 like CaO for example)
Attraction between ions increases
Lattice energy becomes more negative
Melting point increases
Opposing effects with oppositely charged ions?
Na+ Mg2+ and Al3+
Increasing charge means more attraction
Decreasing size (more protons to same number of electrons)
S2- and Cl-
Increasing charge gives more attraction
Increasing size gives less attraction
Uses of metal oxides?
Protective coating for the insides of furnaces
What other factors needed for melting points?
Picking of ions
Hydration Enthalpy factors?
Ionic size - as you go down the group, ionic radius increases and attraction between ion and water molecules decreases so hydration energy is less negative
Ionic charge - as you go along the period, ionic charge increases so attraction with water molecules increases and hydration energy becomes more negative
Predicting solubility
If sum of hydration enthalpies is larger than the magnitude of lattice Enthalpy then the overall Enthalpy change will be Exothermic and the compound should dissolve
What else does solubility depend on?
Temperature and entropy
What is entropy?
A measure of disorder (S)
Greater the entropy?
Greater the dispersal of energy and greater the disorder - tendency for energy to spread out
Greater the energy is spread out per kelvin per mole
General entropy rules?
Solids have smallest entropies
Gases have greatest entropies
If a system changes to become more random?
Energy is spread out more and entropy change is positive
Vice versa
Solid to liquid to gas
Entropy increases
Melting and boiling
Increases randomness and thus increases disorder and increases entropy
Standard entropy?
Entropy of one mole of a substance under standard conditions (100kPa and 298 K) - they are ALWAYS POSITIVE and have units J/K/mol
Calculating entropy changes?
Sum of Entropy of products - sum of entropy of reactants
Feasibility?
A reaction can happen if products have a lower overall energy than the reactants - spontaneous is same as energetically feasible
Free energy
Overall change in energy during a chemical reaction made of
1) Enthalpy change (heat transfer between chemical system and surroundings)
2) Entropy change at the temperature of the reaction ; dispersal of energy within the chemical system itself
Gibbs equation
Enthalpy change with sureoundings - Temperature in Kelvin*Entropy of system
If gibbs<0
Reaction is feasible
Convert Celsius to Kelvin
+273
What to remember in Gibbs energy?
Entropy should be divided by 1000 to become kJ/K/mol
As temperature increases, entropy becomes more significant
Limitations of predictions made for feasibility
Free energy change is useful for predicting feasibility but many reactions have a negative Gibbs energy and do not seem to take place
BECAUSE IT DOES NOT TAKE INTO ACCOUNT ACTIVATION ENERGY - VERY SLOW RATE (NO ACCOUNT OF KINETICS OR RATE OF REACTION)
