Lattice Enthalpy

Question 1

Discuss whether a positive ion will have a smaller or larger radius than the atom.

Answer 1

The positive ion will have fewer electron shells and a higher proton to electron ratio than the atom, so a greater nuclear attraction on the outer electrons, so smaller radius.

Question 2

Discuss whether a negative ion will have a smaller or larger radius than the atom.

Answer 2

The negative ion will have more electron shells and a smaller proton to electron ratio so there is a weaker nuclear attraction on the outer electrons, hence a larger radius.

Question 3

Definition of an ionic bond

Answer 3

Electrostatic attraction between oppositely charged ions.

Question 4

Definition of lattice enthalpy

Answer 4

The enthalpy change when one mole of solid ionic lattice is formed from its gaseous ions (under standard conditions).

Question 5

Write the equation for the standard lattice enthalpy of NaCl(s).

Answer 5

Na+(g) + Cl-(g) …. NaCl(s)

Question 6

Write the eqaution for the standard lattice enthalpy of MgF2(s).

Answer 6

Mg^2+(g) + F^2-(g) …. MgF2(s)

Question 7

What type of enthalpy change is bond making? (3 points).

Answer 7

1. Bond making is exothermic.
2. As the ions are forming attractions with each other, they are making bonds so lattice enthalpy values are always negative.
3. The more negative the lattice enthalpy value is, the stronger the ionic bond.

Question 8

What factors affect the value of lattice enthalpy?

Answer 8

1. Charge on the ions- the greater the chanrge, the stronger the attraction for ions of opposite charge.
2. Size of ions.

Question 9

How will the charge and the size of ions change across a period for positive ions?

Answer 9

Proton to electron ratio increases. Nuclear attraction for outer electron(s) increases. Ionic radius decreases.

Question 10

How will the charge and the size of ions change across a period for negative ions? (3 points)

Answer 10

1. Proton to electron ratio increases.
2. Nuclear attraction for outer electron(s) decreases.
3. Ionic radius decreases.

Question 11

How will the charge and size of positive and negative ions change down a group?

Answer 11

Charge remains the same. Number of electron shells increases.Ionic radius increases.

Question 12

Describe the relationship between the size of ion and charge density.

Answer 12

The size of the ion dictates the charge density- how spread out the charge is over the ion.
If an ion is small, the charge is not spread out, so the charge density will be high.
The higher the charge density, the stronger the attraction for ions of opposite charge.

Question 13

EXAMPLE QU: Which compound will have the more exothermic lattice enthalpy, MgO or Na20?

Answer 13

Mg^2+ has a greater charge than Na+.
Mg^2+ has a smaller ionic radius than Na+.
Mg^2+ has a greater change density then Na+.
Mg^2+ is more strongly attracted to O^2-.
MgO has the more exothermic lattice enthalpy.

Question 14

Definition of standard enthalpy change of formation and write an equation for the standard formation of NaCl(s).

Answer 14

The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.
Na(s) + 1/2Cl2(g) …. NaCl(s)

Question 14

Definition of standard enthalpy change of formation and write an equation for the standard formation of NaCl(s).

Answer 14

The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.
Na(s) + 1/2Cl2(g) …. NaCl(s)

Question 15

Definition of standard enthalpy change of atomisation and write an equation for the standard atomisation of calcium and another equation for the standard atomisation of chlorine.

Answer 15

The enthalpy change when one mole of gaseous atoms are formed from an element in its standard states.
Ca(s) …. Ca(g)
1/2Cl2(g) …. Cl(g)

Question 16

Definition of standard enthalpy change of first ionisation energies and write an equation for the first ionisation energy of magnesium.

Answer 16

The enthalpy change when 1 mole of electrons is removed from one mole of gaseous atoms.
Mg(g) …. Mg+(g) + e-

Question 17

Definition of standard enthalpy change of second ionisation energies and write an equation to for the the second ionisation energy of magnesium.

Answer 17

The enthalpy change when 1 mole of electrons is removed from one mole of gaseous atoms.
Mg(g) …. Mg^2+(g) + e-

Question 18

Definition of standard enthalpy change for first electron affinity and write an equation for the standard first electron affinity of chlorine.

Answer 18

The enthalpy change when one mole of gaseous 1- ions are made from one mole of gaseous atoms.
Cl(g) + e- …. Cl-(g)

Question 19

Definition of standard enthalpy change for second electron affinity and write an equation for the seond electron affinity of oxygen.

Answer 19

The enthalpy change when one mole of gaseous 2- ions are made from one mole of gaseous 1- ions.
O-(g) + e- …. O^2-(g)

Question 20

Why are second electron affinities endothermic?

Answer 20

Energy needed to overcome the repulsion between the negative electron and negative ion as only opposites attract.

Question 21

Definition of Hess’s law

Answer 21

The enthalpy change of a reaction is independent of the route taken providing the initial and final conditions are the same.

Question 22

Definition of enthalpy change of hydration

Answer 22

The enthalpy change when one mole of aqueous ions are formed from their gaseous ions in their standard states under standard conditions. (The sign is always negative!)

Question 23

Describe and explain the two factors that affect how exothermic the enthalpy of hydration is.

Answer 23

1. Charge on the ion- The greater the charge on the ion, the greater the attraction for water molecules. Enthalpy of hydration become more negative/exothermic.
2. Ionic radius- As ionic readius becomes smaller, the attraction for water increases. Enthalpy of hydration become more negative/exothermic.

Question 24

Definition of enthalpy change of solution

Answer 24

The enthalpy change when one mole of a compound is completely dissolved in water under standard conditions.