Chemistry: Thermodynamics Part 1
This flashcard set explains lattice enthalpy as a measure of the strength of bonding in an ionic lattice. It defines two types: lattice dissociation enthalpy, the energy required to separate one mole of a solid ionic compound into gaseous ions, and lattice formation enthalpy.
What is the strength of the bonding in a lattice given by
It’s lattice enthalpy
Key Terms
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
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| Term | Definition |
|---|---|
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 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 |
|
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 |
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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 |