Thermodynamics Part 1 Flashcards
What is the strength of the bonding in a lattice given by
It’s lattice enthalpy
What are the two ways to define lattice enthalpy
- enthalpy of lattice dissociation
- enthalpy of lattice formation
Enthalpy of lattice dissociation
The enthalpy change to separate one mile of solid ionic compound into its gaseous ions
Why are enthalpies of lattice dissociation always endothermic
Bonds are being broken
What is the enthalpy of lattice dissociation a measure of
The strength of the electrostatic forces of attraction between ions, massive of the strength of the ionic bonding
Enthalpy of lattice formation
The enthalpy change when one mole of solid ionic compound is formed from its gaseous ions I
Why are enthalpies of lattice formation always endothermic
Bonds are being formed
How will the values for dissociation and formation be similar
They’ll be the same but have the opposite sign
Standard enthalpy of formation
The enthalpy change that occurs when one mole of a compound is formed from its constituent elements with all reactants and products in their standard states
Equation to represent the standard enthalpy of formation potassium chloride
K(s) + 1/2Cl2(g) > KCl(s)
What is the standard enthalpy of formation always in ionic compounds and why
Expthermic becaude energy is released when the compound is formed
Standard enthalpy of atomisation
The enthalpy change when one mole of gaseous atoms it formed from the element in its standard state
Equation to represent atomisation of chlorine
1/2 Cl2(g) > Cl (g)
What is atomisation always and why
Endothermic because energy is needed to form gaseous atoms
Bond enthalpy
The energy required to break one mole of a given covalent bond in the molecules in the gaseous state
Equation to represent the bond enthalpy of a Cl-Cl bond
Cl2(g)> 2Cl(g)
What is the relationship between the value for the bond enthalpy of a Cl-Cl bond and the value for the atomisation of chlorine
Bond enthalpy= 2xatomisation energy
Why are bond enthalpies always endothermic
It takes energy to break bonds
First ionisation energy
The energy needed to remove one mole of electrons from one mole of atoms in the gaseous state
Equation to represent the first ionisation energy of lithium
Li(g)>Li+ + e-
Why are first IEs always endothermic
It takes energy to remove an electron
Second ionisation energy
The enthalpy change when one mole of gaseous 2+ ions are formed from one mole of gaseous 1+ ions
Equation to represent the second IE of calcium
Ca+(g) > Ca2+(g) + e-
Why is second IE always endothermic
It takes energy to remove an electron
Why is the value for second IE always larger than value for first IE
stronger attraction- same protons but fewer electrons
+ve ion attracting -ve electrons
First electron affinity
The enthalpy change when one mole of gaseous 1- ions are formed from one mole of gaseous atoms
First electron affinity of chlorine
Cl(g) + e- > Cl-(g)
Why is first electron affinity always expthermic
Energy is released when the nucleus of an atom attracts an electron
Second electron affinity
The enthalpy change when one mole of gaseous 2- ions are formed from one mole of gaseous 1- ions
Equation to represent the second electron affinity of oxygen
O-(g) + e- > O2-(g)
Why is second electron affinity always endothermic
energy is needed to overcome the repulsion between the electrons and anion (both negative)
The perfect ionic model
Ions are perfect spheres with only electrostatic forces of attractive (no covalent character)
When does the fact that experimental lattice enthalpies found using born haber cycles are usually different to the theoretical value provide evidence for
Some ionic compounds having covalent character
What is the relationship between the closeness of the experimental and theoretical value and the degree of ionic bonding
The closer the two values the greater the degree of ionic bonding
What must the energy produced when water surrounds the ions be the same as for a solid dissolves
The electrostatic forces of attraction between the ions
First step in solid dissolving
Ionic lattice breaks down into gaseous ions
MX(s) > M+(g) + X-(g)
Second step in solid dissolving
Hydration of gaseous ions
M+(g) + X-(g) > M+(aq) + X-(aq)
What is the resulting enthalpy change when a solid dissolves known as
The enthalpy change of solution
Enthalpy of solution
The enthalpy change when one mole of an ionic substance dissolves in enough solvent to form an infinitely dilute solution
Equation to represent the enthalpy of solution of sodium chloride
NaCl(s) > Na+(aq) + Cl-(aq)
Enthalpy of hydration
The enthalpy change when one mole of aqueous ions is formed from one mole of gaseous ions
Equation to represent the enthalpy of hydration of sodium ions
Na+(g) > Na+(aq)
Why are enthalpies of hydration always expthermic
Energy is released when the gaseous ions form bonds to water
What are lattice enthalpy values a measure of
The strength of attraction between oppositely charged ions
What affects how strong the forces of attraction between oppositely charged ions are
Charged density
What two things affect charge density
Size of ion
Size of charge
Why do lattice enthalpies become less expthermic NaCl> NaI
- Cl- smaller ion
- greater charge density
- greater attraction, stronger ionic bonding
- therefore enthalpy of lattice formation becomes more expthermic
Why do lattice enthalpies become more exothermic NaCl > MgCl2 > MgO
As charges increase on ions, stronger attraction between oppositely charged ions, stronger ionic bonding therefore more exothermic enthalpy of lattice formation
Why do hydration enthalpies become more exothermic I->Br->Cl-
Cl- smaller ion- greater charge density- greater attraction to water molecules
Why does hydration enthalpy become more exothermic Na+>Mg2+>Al3+
As charge on ions increases, attraction with water molecules increases massively- mucb more exothermic hydration enthalpy
