Define standard enthalpy change of solution .

Standard enthalpy change of solution (Δ H sol θ ) is the enthalpy change when 1 mole of an ionic substance dissolves in sufficient water to form a very dilute solution under standard conditions.

True or False? The enthalpy change of hydration is always exothermic.

True. Hydration enthalpies are always exothermic because energy is released when ion-dipole bonds form between the ions and polar water molecules.

Define standard enthalpy change of hydration .

Standard enthalpy change of hydration (Δ H hyd θ ) is the enthalpy change when 1 mole of a specified gaseous ion dissolves in sufficient water to form a very dilute solution under standard conditions.

Write an equation to represent the standard enthalpy change of solution of KCl.

KCl (s) + aq → K + (aq) + Cl - (aq)

True or False? The enthalpy change of solution can only be exothermic.

False. Δ H sol θ can be either exothermic or endothermic depending on the relative magnitudes of the lattice energy and the hydration enthalpy.

What type of bonds form between water molecules and ions during hydration, and which part of the water molecule is attracted to cations?

Ion-dipole bonds form during hydration. The δ - oxygen atom of water is attracted to positive ions ( cations ). The δ + hydrogen atoms are attracted to negative ions ( anions ).

When an ionic solid dissolves, the .......... of the lattice must be overcome. This energy is compensated for by the .......... released during hydration.

When an ionic solid dissolves, the electrostatic attraction of the lattice must be overcome. This energy is compensated for by the enthalpy of hydration released during hydration.

The lattice energy of KCl is -711 kJ mol -1 . To break the lattice, .......... kJ mol -1 is needed. As Δ H sol θ of KCl is +26 kJ mol -1 , the hydration enthalpy must provide .......... kJ mol -1 .

The lattice energy of KCl is -711 kJ mol -1 . To break the lattice, +711 kJ mol -1 is needed. As Δ H sol θ of KCl is +26 kJ mol -1 , the hydration enthalpy must provide -685 kJ mol -1 .

State the relationship between Δ H sol θ , Δ H latt θ and Δ H hyd θ in an energy cycle.

Δ H sol θ = -Δ H latt θ + Δ H hyd θ This can be rearranged to: Δ H hyd θ = Δ H sol θ + Δ H latt θ

True or False? In the energy cycle, the direct route from ionic solid to hydrated ions represents the enthalpy change of solution.

True. The direct route (ionic solid → ions in aqueous solution) represents Δ H sol θ . The indirect route goes via the gaseous ions and involves reverse lattice enthalpy and hydration enthalpies.

In the energy cycle for KCl: Direct route: KCl (s) → KCl (aq) represents .......... Indirect route: KCl (s) → K + (g) + Cl - (g) → KCl (aq) involves .......... and ..........

In the energy cycle for KCl: Direct route: KCl (s) → KCl (aq) represents Δ H sol θ Indirect route: KCl (s) → K + (g) + Cl - (g) → KCl (aq) involves -Δ H latt θ and Δ H hyd θ

Why is a negative sign placed in front of Δ H latt θ in the energy cycle equation?

Lattice energy (Δ H latt θ ) refers to lattice formation. In the energy cycle, the lattice is being broken (dissociated), which is the reverse process, so a negative sign is used. This is sometimes referred to as lattice dissociation enthalpy .

True or False? The total enthalpy of hydration in the energy cycle is the sum of the hydration enthalpies of each individual ion.

True. Each ion (cation and anion) has its own Δ H hyd θ . The total is found by adding the hydration enthalpies of all ions together.

An energy cycle has two routes from ionic lattice to hydrated ions. State what each route represents.

Route 1 (direct) : ionic solid → ions in aqueous solution, representing Δ H sol θ . Route 2 (indirect) : ionic solid → gaseous ions → ions in aqueous solution, representing -Δ H latt θ + Δ H hyd θ .

According to Hess's law, the enthalpy change of the .......... route equals the enthalpy change of the .......... route in an energy cycle.

According to Hess's law, the enthalpy change of the direct route equals the enthalpy change of the indirect route in an energy cycle.

Rearrange the energy cycle equation to show how Δ H latt θ can be calculated.

Δ H hyd θ = Δ H sol θ + Δ H latt θ Rearranging: Δ H latt θ = Δ H hyd θ - Δ H sol θ

True or False? An energy cycle can only be used to calculate the enthalpy of hydration.

False. An energy cycle can be used to calculate any unknown value among Δ H sol θ , Δ H latt θ and Δ H hyd θ , provided the other two are known.

State the equation relating Δ H hyd θ , Δ H latt θ and Δ H sol θ used in energy cycle calculations.

Δ H hyd θ = Δ H latt θ + Δ H sol θ

For MgCl 2 , the total Δ H hyd θ = Δ H hyd θ [Mg 2+ ] + .......... × Δ H hyd θ [Cl - ] because there are .......... chloride ions.

For MgCl 2 , the total Δ H hyd θ = Δ H hyd θ [Mg 2+ ] + 2 × Δ H hyd θ [Cl - ] because there are two chloride ions.

True or False? When calculating Δ H hyd θ [Cl - ] in KCl, only the hydration enthalpy of the cation is needed.

False. Both cation and anion hydration enthalpies contribute. The total Δ H hyd θ = Δ H hyd θ [K + ] + Δ H hyd θ [Cl - ], so the cation value is used to isolate the anion value.

Using the data below, calculate Δ H hyd θ [Cl - ] in KCl. Δ H latt θ [KCl] = -711, Δ H sol θ [KCl] = +26, Δ H hyd θ [K + ] = -322 (all kJ mol -1 )

Δ H hyd θ [K + ] + Δ H hyd θ [Cl - ] = Δ H latt θ + Δ H sol θ -322 + Δ H hyd θ [Cl - ] = (-711) + (+26) = -685 Δ H hyd θ [Cl - ] = -685 + 322 = -363 kJ mol -1

True or False? The equation Δ H hyd θ = Δ H latt θ + Δ H sol θ can be rearranged to find any unknown enthalpy in the cycle.

True. The equation can be rearranged to find Δ H sol θ , Δ H latt θ or individual ionic hydration enthalpies, provided the other values are given.

In an energy cycle calculation for MgCl 2 , why must Δ H hyd θ [Cl - ] be doubled?

MgCl 2 contains two Cl - ions per formula unit. When 1 mole of MgCl 2 dissolves, 2 moles of Cl - ions are hydrated, so the hydration enthalpy of Cl - must be multiplied by 2.

To find Δ H latt θ from an energy cycle: Δ H latt θ = Δ H hyd θ .......... Δ H sol θ

To find Δ H latt θ from an energy cycle: Δ H latt θ = Δ H hyd θ - Δ H sol θ

Using these data, calculate Δ H hyd θ [Mg 2+ ]. Δ H latt θ [MgCl 2 ] = -2592, Δ H sol θ [MgCl 2 ] = -55, Δ H hyd θ [Cl - ] = -363 (all kJ mol -1 )

Δ H hyd θ [Mg 2+ ] = Δ H latt θ + Δ H sol θ - 2Δ H hyd θ [Cl - ] = (-2592) + (-55) - (2 × -363) = -2647 + 726 = -1921 kJ mol -1

True or False? Hess's law states that the total enthalpy change depends on the route taken between reactants and products.

False. Hess's law states that the total enthalpy change is independent of the route taken. This allows energy cycle calculations to equate the direct and indirect routes.

State the two factors that affect the enthalpy change of hydration of an ion.

Ionic radius : smaller ions have higher charge density and more exothermic hydration enthalpies. Ionic charge : higher charge leads to greater charge density and more exothermic hydration enthalpies.

True or False? As ionic radius decreases, the enthalpy change of hydration becomes more exothermic.

True. Smaller ions have higher charge density, forming stronger ion-dipole attractions with water molecules. More energy is released during hydration.

Define ion-dipole attraction .

An ion-dipole attraction is the electrostatic force between an ion and the oppositely charged end of a polar water molecule, responsible for the release of energy during hydration.

Explain why the enthalpy of hydration of Mg 2+ is more exothermic than that of Ba 2+ .

Mg 2+ has a smaller ionic radius than Ba 2+ , giving it a higher charge density . This leads to stronger ion-dipole attractions with water molecules and more energy released during hydration.

True or False? Higher ionic charge leads to less exothermic hydration enthalpies.

False. Higher ionic charge leads to greater charge density, which strengthens ion-dipole attractions with water molecules and makes hydration enthalpies more exothermic.

The enthalpy of hydration of CaO is more exothermic than that of KCl because Ca 2+ and O 2- have .......... ionic charges and .......... ionic radii than K + and Cl - .

The enthalpy of hydration of CaO is more exothermic than that of KCl because Ca 2+ and O 2- have higher ionic charges and smaller ionic radii than K + and Cl - .

Which ion has the more exothermic enthalpy of hydration: Mg 2+ or Ca 2+ ? Give a reason.

Mg 2+ has the more exothermic enthalpy of hydration. It has a smaller ionic radius than Ca 2+ , giving it a higher charge density and stronger ion-dipole attractions with water molecules.

A smaller ionic radius leads to .......... charge density, .......... ion-dipole attractions and a .......... exothermic enthalpy of hydration.

A smaller ionic radius leads to higher charge density, stronger ion-dipole attractions and a more exothermic enthalpy of hydration.

Why are hydration enthalpies always exothermic?

Energy is released when ion-dipole bonds form between the ions and the polar water molecules during hydration. Bond formation always releases energy.

A LevelChemistryCambridge (CIE)Flashcards23. Chemical EnergeticsEnthalpies of Solution & Hydration

Enthalpies of Solution & Hydration (Cambridge (CIE) A Level Chemistry): Flashcards

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FrontEnthalpies of Solution & Hydration
Define standard enthalpy change of solution.

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Define standard enthalpy change of solution.
Standard enthalpy change of solution (ΔH_sol^θ) is the enthalpy change when 1 mole of an ionic substance dissolves in sufficient water to form a very dilute solution under standard conditions.
True or False?
The enthalpy change of hydration is always exothermic.
True.
Hydration enthalpies are always exothermic because energy is released when ion-dipole bonds form between the ions and polar water molecules.
Define standard enthalpy change of hydration.
Standard enthalpy change of hydration (ΔH_hyd^θ) is the enthalpy change when 1 mole of a specified gaseous ion dissolves in sufficient water to form a very dilute solution under standard conditions.
Write an equation to represent the standard enthalpy change of solution of KCl.
KCl (s) + aq → K⁺ (aq) + Cl^- (aq)
True or False?
The enthalpy change of solution can only be exothermic.
False.
ΔH_sol^θ can be either exothermic or endothermic depending on the relative magnitudes of the lattice energy and the hydration enthalpy.
What type of bonds form between water molecules and ions during hydration, and which part of the water molecule is attracted to cations?
Ion-dipole bonds form during hydration. The δ^- oxygen atom of water is attracted to positive ions (cations). The δ⁺ hydrogen atoms are attracted to negative ions (anions).
When an ionic solid dissolves, the .......... of the lattice must be overcome. This energy is compensated for by the .......... released during hydration.
When an ionic solid dissolves, the electrostatic attraction of the lattice must be overcome. This energy is compensated for by the enthalpy of hydration released during hydration.
Write an equation that represents the standard enthalpy change of hydration of Mg²⁺ ions.
Mg²⁺ (g) + aq → Mg²⁺ (aq)
The lattice energy of KCl is -711 kJ mol^-1. To break the lattice, .......... kJ mol^-1 is needed. As ΔH_sol^θ of KCl is +26 kJ mol^-1, the hydration enthalpy must provide .......... kJ mol^-1.
The lattice energy of KCl is -711 kJ mol^-1. To break the lattice, +711 kJ mol^-1 is needed. As ΔH_sol^θ of KCl is +26 kJ mol^-1, the hydration enthalpy must provide -685 kJ mol^-1.
State the relationship between ΔH_sol^θ, ΔH_latt^θ and ΔH_hyd^θ in an energy cycle.
ΔH_sol^θ = -ΔH_latt^θ + ΔH_hyd^θ
This can be rearranged to:
ΔH_hyd^θ = ΔH_sol^θ + ΔH_latt^θ
True or False?
In the energy cycle, the direct route from ionic solid to hydrated ions represents the enthalpy change of solution.
True.
The direct route (ionic solid → ions in aqueous solution) represents ΔH_sol^θ. The indirect route goes via the gaseous ions and involves reverse lattice enthalpy and hydration enthalpies.
In the energy cycle for KCl:
Direct route: KCl (s) → KCl (aq) represents ..........
Indirect route: KCl (s) → K⁺ (g) + Cl^- (g) → KCl (aq) involves .......... and ..........
In the energy cycle for KCl:
Direct route: KCl (s) → KCl (aq) represents ΔH_sol^θ
Indirect route: KCl (s) → K⁺ (g) + Cl^- (g) → KCl (aq) involves -ΔH_latt^θ and ΔH_hyd^θ
Why is a negative sign placed in front of ΔH_latt^θ in the energy cycle equation?
Lattice energy (ΔH_latt^θ) refers to lattice formation. In the energy cycle, the lattice is being broken (dissociated), which is the reverse process, so a negative sign is used. This is sometimes referred to as lattice dissociation enthalpy.
True or False?
The total enthalpy of hydration in the energy cycle is the sum of the hydration enthalpies of each individual ion.
True.
Each ion (cation and anion) has its own ΔH_hyd^θ. The total is found by adding the hydration enthalpies of all ions together.
An energy cycle has two routes from ionic lattice to hydrated ions. State what each route represents.
Route 1 (direct): ionic solid → ions in aqueous solution, representing ΔH_sol^θ.
Route 2 (indirect): ionic solid → gaseous ions → ions in aqueous solution, representing -ΔH_latt^θ + ΔH_hyd^θ.
According to Hess's law, the enthalpy change of the .......... route equals the enthalpy change of the .......... route in an energy cycle.
According to Hess's law, the enthalpy change of the direct route equals the enthalpy change of the indirect route in an energy cycle.
Rearrange the energy cycle equation to show how ΔH_latt^θ can be calculated.
ΔH_hyd^θ = ΔH_sol^θ + ΔH_latt^θ
Rearranging: ΔH_latt^θ = ΔH_hyd^θ - ΔH_sol^θ
True or False?
An energy cycle can only be used to calculate the enthalpy of hydration.
False.
An energy cycle can be used to calculate any unknown value among ΔH_sol^θ, ΔH_latt^θ and ΔH_hyd^θ, provided the other two are known.
State the equation relating ΔH_hyd^θ, ΔH_latt^θ and ΔH_sol^θ used in energy cycle calculations.
ΔH_hyd^θ = ΔH_latt^θ + ΔH_sol^θ
For MgCl₂, the total ΔH_hyd^θ = ΔH_hyd^θ[Mg²⁺] + .......... × ΔH_hyd^θ[Cl^-] because there are .......... chloride ions.
For MgCl₂, the total ΔH_hyd^θ = ΔH_hyd^θ[Mg²⁺] + 2 × ΔH_hyd^θ[Cl^-] because there are two chloride ions.
True or False?
When calculating ΔH_hyd^θ[Cl^-] in KCl, only the hydration enthalpy of the cation is needed.
False.
Both cation and anion hydration enthalpies contribute. The total ΔH_hyd^θ = ΔH_hyd^θ[K⁺] + ΔH_hyd^θ[Cl^-], so the cation value is used to isolate the anion value.
Using the data below, calculate ΔH_hyd^θ[Cl^-] in KCl.
ΔH_latt^θ[KCl] = -711, ΔH_sol^θ[KCl] = +26, ΔH_hyd^θ[K⁺] = -322 (all kJ mol^-1)
ΔH_hyd^θ[K⁺] + ΔH_hyd^θ[Cl^-] = ΔH_latt^θ + ΔH_sol^θ
-322 + ΔH_hyd^θ[Cl^-] = (-711) + (+26) = -685
ΔH_hyd^θ[Cl^-] = -685 + 322 = -363 kJ mol^-1
True or False?
The equation ΔH_hyd^θ = ΔH_latt^θ + ΔH_sol^θ can be rearranged to find any unknown enthalpy in the cycle.
True.
The equation can be rearranged to find ΔH_sol^θ, ΔH_latt^θ or individual ionic hydration enthalpies, provided the other values are given.
In an energy cycle calculation for MgCl₂, why must ΔH_hyd^θ[Cl^-] be doubled?
MgCl₂ contains two Cl^- ions per formula unit. When 1 mole of MgCl₂ dissolves, 2 moles of Cl^- ions are hydrated, so the hydration enthalpy of Cl^- must be multiplied by 2.
To find ΔH_latt^θ from an energy cycle:
ΔH_latt^θ = ΔH_hyd^θ .......... ΔH_sol^θ
To find ΔH_latt^θ from an energy cycle:
ΔH_latt^θ = ΔH_hyd^θ - ΔH_sol^θ
Using these data, calculate ΔH_hyd^θ[Mg²⁺].
ΔH_latt^θ[MgCl₂] = -2592, ΔH_sol^θ[MgCl₂] = -55, ΔH_hyd^θ[Cl^-] = -363 (all kJ mol^-1)
ΔH_hyd^θ[Mg²⁺] = ΔH_latt^θ + ΔH_sol^θ - 2ΔH_hyd^θ[Cl^-]
= (-2592) + (-55) - (2 × -363)
= -2647 + 726 = -1921 kJ mol^-1
True or False?
Hess's law states that the total enthalpy change depends on the route taken between reactants and products.
False.
Hess's law states that the total enthalpy change is independent of the route taken. This allows energy cycle calculations to equate the direct and indirect routes.
State the two factors that affect the enthalpy change of hydration of an ion.
1. Ionic radius: smaller ions have higher charge density and more exothermic hydration enthalpies.
2. Ionic charge: higher charge leads to greater charge density and more exothermic hydration enthalpies.
True or False?
As ionic radius decreases, the enthalpy change of hydration becomes more exothermic.
True.
Smaller ions have higher charge density, forming stronger ion-dipole attractions with water molecules. More energy is released during hydration.
Define ion-dipole attraction.
An ion-dipole attraction is the electrostatic force between an ion and the oppositely charged end of a polar water molecule, responsible for the release of energy during hydration.
Explain why the enthalpy of hydration of Mg²⁺ is more exothermic than that of Ba²⁺.
Mg²⁺ has a smaller ionic radius than Ba²⁺, giving it a higher charge density. This leads to stronger ion-dipole attractions with water molecules and more energy released during hydration.
True or False?
Higher ionic charge leads to less exothermic hydration enthalpies.
False.
Higher ionic charge leads to greater charge density, which strengthens ion-dipole attractions with water molecules and makes hydration enthalpies more exothermic.
The enthalpy of hydration of CaO is more exothermic than that of KCl because Ca²⁺ and O^2- have .......... ionic charges and .......... ionic radii than K⁺ and Cl^-.
The enthalpy of hydration of CaO is more exothermic than that of KCl because Ca²⁺ and O^2- have higher ionic charges and smaller ionic radii than K⁺ and Cl^-.
Which ion has the more exothermic enthalpy of hydration: Mg²⁺ or Ca²⁺? Give a reason.
Mg²⁺ has the more exothermic enthalpy of hydration. It has a smaller ionic radius than Ca²⁺, giving it a higher charge density and stronger ion-dipole attractions with water molecules.
A smaller ionic radius leads to .......... charge density, .......... ion-dipole attractions and a .......... exothermic enthalpy of hydration.
A smaller ionic radius leads to higher charge density, stronger ion-dipole attractions and a more exothermic enthalpy of hydration.
Why are hydration enthalpies always exothermic?
Energy is released when ion-dipole bonds form between the ions and the polar water molecules during hydration. Bond formation always releases energy.

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