Energy

• AS Revision:• Energy terms

• Enthalpy change of reaction.ΔHr

• Enthalpy change of formation ΔHf

• Enthalpy change of combustion ΔHc

• Standard conditions• Hess’ Law

Enthalpy Change of Reaction ΔHr

• Enthalpy, H, is the stored energy in a compound. We cannot measure this.

• Enthalpy changes, ΔH are transfers of energy in or out of the system at constant pressure.

• ΔHr Enthalpy change for a reaction under standard conditions when the molar quantities from the balanced equation are in their standard states.

• E.g.2H2(g) + O2(g) → 2H2O(l) ΔHr =-571.6kJ

Enthalpy Change of Formation ΔHf

• The energy change when 1 mole of a compound is made from its constituent elements in their standard states, under standard conditions.

• E.g.H2(g) + ½O2(g) → H2O(l) ΔHf =-285.8kJmol-1

• The same reaction but half the energy given out.

• This is per mole of water made, NOT for the balanced equation which is the case in ΔHr .

Enthalpy Change of Combustion ΔHc

• The energy change when 1 mole of a substance burns completely in oxygen under standard conditions.

• E.g.H2(g) + ½O2(g) → H2O(l) ΔHc =-285.8kJmol-1

• Same equation, same conditions, same energy change, different name.

• Standard Conditions:• Temp: 298K• Pressure:1 atmosphere (101kPa)• Solutions: 1 mol dm-3

Energy Cycles

• Since some energy changes cannot be measured directly, at AS Hess’ Law was used.

• “If a reaction can take place by more than one route and the initial and final conditions are the same, the total enthalpy change is the same for each route.”

• This is usually used when combustion or formation data is available.

• E.g. C2H4(g) + HBr(g) → C2H5Br(l)

Lattice Enthalpy

• At A2 Hess’ Law is applied to IONIC Compounds.

• This is called a Born-Haber Cycle.• It is usually used to calculate lattice

enthalpy which is a measure of ionic bond strength.

• Lattice Enthalpy is the enthalpy change that accompanies the formation of 1 mole of an ionic compound from its GASEOUS ions under standard conditions.

• E.g. K+(g) + Cl-(g) → KCl(s) ΔHLE = -711 kJmol-1

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Ionic lattice of sodium chloride

Lattice Enthalpy

• This term is ALWAYS negative as energy is ALWAYS given out as oppositely charged ions come together.

• The higher the magnitude, the greater the electrostatic force between the oppositely charged ions.

• Lattice enthalpies cannot be measured directly because it is impossible to form one mole of an ionic lattice from gaseous ions.

Other Key Terms

• To complete Born-Haber Cycles, other key terms are needed as listed on p.167.

• Copy these into the spaces provided on the sheet. Complete the questions on page 167.

• Born-Haber Cycles can also be done as scale diagrams using graph paper or even lined paper.

• Then do q.1,2 on p. 195.• Q. p. 167., Q. p.169• Q. p.171, q. p.173

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Completed Born–Haber cycle for sodium chloride

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Construction of the Born–Haber cycle for caesium chloride

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Born–Haber cycle for sodium oxide

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Born–Haber cycle for calcium chloride

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Born–Haber cycle for copper(II) oxide

Trends in Lattice Enthalpy

• Lattice enthalpy is the result of electrostatic attractions between ions of opposite charge.

• The value of lattice enthalpy is affected by:• i) size of ions• ii) charge on the ions• Why?• See handout for group 1 halides.• The lattice enthalpy becomes less

exothermic as the size of the ion increases.• True for both anions and cations.

Ionic radius

• As the radius of the ion increases and the charge density on the surface of the ions decreases, the attraction between the ions decreases.

• Small ions pack more closely together in a lattice than larger ions.

• The lattice energy is less exothermic if the attraction is lower because less energy is given out as the lattice forms.

• This leads to weaker ionic bonding.

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Trend in lattice enthalpies of the sodium halides

Charge on the ions• Lattice enthalpy becomes more exothermic as the charge

on the ion increases.• It is most exothermic for small, highly charged ions.• See this by comparing LiF with MgO.• LiF ΔHlatt = -1050 kJmol-1

• MgO ΔHlatt = -3923 kJ mol-1

• Shows that doubly charged ions attract each other more strongly than singly charged ions.

• Charge Density:• Assumes that charge is spread over the outer surface of the

ion and is charge per unit surface area:• Charge density = charge on ion• ionic radius• Surface area α radius 2

•

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Lattice enthalpy becomes more exothermic as the ionic charge increases

Magnesium oxide as a refractory

• The exceptionally high lattice enthalpy of magnesium oxide means that that the lattice is so strong that it takes a great deal of heat to decompose it. (mp 2853oC)

• It can be used to line furnaces – as a refractory.

• It is also used in ceramics, wire coatings and flame retardant particle boards

Exercise

• For each pair of compounds, suggest which will have the most exothermic lattice enthalpy:

• A) CaO and CaCl2• B) KCl and K2O

• C) BaI2 and SrI2• Place the following compounds in order of

increasingly exothermic lattice enthalpy.

• Li2O LiF MgO

• Explain why you have put them in this order.

Exercise

• For each pair of compounds, suggest which will have the most exothermic lattice enthalpy:

• A) CaO and CaCl2• B) KCl and K2O

• C) BaI2 and SrI2• Place the following compounds in order of

increasingly exothermic lattice enthalpy.

• LiF Li2O MgO

• Because LiF is singly charged ions, Li2O has one doubly charged on and MgO has two doubly charged ions.

Testing The Ionic Model

Lattice Enthalpies/kJmol-1

Compound Actual Value (from B-H cycle)

Theoretical Value

NaF -918 -912

NaCl -787 -770

NaBr -742 -735

NaI -705 -687

AgF -958 -920

AgCl -905 -833

AgBr -891 -816

AgI -889 -778

Theoretical Values compared with Actual Values

• By applying electrostatic calculations and knowing the arrangement of ions in the lattice scientists can work out a theoretical value for standard lattice enthalpies.

• The model is based on perfect ionic bonding, with perfectly spherical ions with even charge density.

• Alkali metal halides show close agreement so they are close to perfectly ionic.

• Silver halides are stronger, having covalency in their bonding so there is less agreement between actual and theoretical lattice enthalpies.

• Bonding is not purely ionic.

Enthalpy Change of Solution ΔHs

• The standard enthalpy change of solution is the enthalpy change that takes place when one mole of a compound is completely dissolved in water under standard conditions.

• This involves breaking the lattice so that the individual ions are free to move through the solution and then hydrating the ions.

• This happens because the δ+ hydrogen atoms in water molecules are attracted to anions and the δ- oxygen atoms are weakly attracted to the cations.

• As a result both anions and cations are surrounded by water molecules. http://www.chemit.co.uk/default.aspx?sitemapID=57

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Ionic lattice dissolves in polar solvents – each ion is surrounded by water molecules to become ‘hydrated’

Enthalpy of Hydration ΔHhyd

• The standard enthalpy change of hydration is the enthalpy change that takes place when one mole of isolated gaseous ions is dissolved in water forming one mole of hydrated ions under standard conditions.

• Each ion in the lattice has its own value for enthalpy of hydration.

• The overall value for the enthalpy change of solution depends on the relative values of the reverse lattice enthalpy (endothermic) and the exothermic hydration enthalpies.

• Do q. p.175

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Born–Haber cycle for calculating the enthalpy change of solution of KCl(s)

Effect of Ion Charge and Size on Hydration Enthalpy

• Ionic size• Small ions have a large charge density and

attract water molecules to them strongly, increasing the exothermic value of hydration enthalpy.

• Ionic Charge• More highly charged ions attract water

molecules to them more strongly, increasing the exothermic value of hydration enthalpy.

• Do q.p.177