Which equation represents lattice enthalpy?

A. NaCl (g) → Na⁺ (g) + Cl⁻ (g)

B. NaCl (s) → Na⁺ (g) + Cl⁻ (g)

C. NaCl (s) → Na⁺ (aq) + Cl⁻ (aq)

D. NaCl (s) → Na⁺ (s) + Cl⁻ (s)

Which equation represents the standard enthalpy of atomization of bromine, Br₂?

A. $\frac{1}{2}$Br₂ (l) → Br (g)

B. Br₂ (l) → 2Br (g)

C. Br₂ (l) → 2Br (l)

D. $\frac{1}{2}$Br₂ (l) → Br (l)

Which equation represents hydration enthalpy?

A.  Na⁺(g) → Na⁺(aq)

B.  Na⁺(aq) → Na⁺(g)

C.  NaCl (s) → NaCl (aq)

D.  NaCl (aq) → NaCl (s)

Which value represents the lattice enthalpy, in kJ mol⁻¹, of strontium chloride, SrCl₂?

A.     – (–829) + 164 + 243 + 550 + 1064 – (–698)

B.     –829 + 164 + 243 + 550 + 1064 – 698

C.     – (–829) + 164 + 243 + 550 + 1064 – 698

D.     –829 + 164 + 243 + 550 + 1064 – (–698)

What is the enthalpy of solution of MgF₂(s) in kJ mol⁻¹?

Lattice enthalpy of MgF₂(s) = 2926 kJ mol⁻¹

Hydration enthalpy of Mg²⁺(g) = −1963 kJ mol⁻¹

Hydration enthalpy of F⁻(g) = −504 kJ mol⁻¹

A.     2926 − 1963 + 2(−504)

B.     2926 − 1963 − 504

C.     −2926 − (−1963) − (−504)

D.     −2926 − (−1963) − 2(−504)

Which equation represents the lattice enthalpy of magnesium sulfide?

A. MgS (s) → Mg (g) + S (g)

B. MgS (s) → Mg⁺ (g) + S^– (g)

C. MgS (s) → Mg²⁺ (g) + S^2– (g)

D. MgS (s) → Mg (s) + S (s)

Which equation represents enthalpy of hydration?

A.     Na(g) → Na⁺(aq) + e⁻

B.     Na⁺(g) → Na⁺(aq)

C.     NaCl(s) → Na⁺(g) + Cl⁻(g)

D.     NaCl(s) → Na⁺(aq) + Cl⁻(aq)

Which combination gives the standard hydration enthalpy of ${\mathrm{Na}}^{+} \left(\mathrm{g}\right)$?

A.  $4+359+790$

B.  $4+359-790$

C.  $-4-359+790$

D.  $4-359+790$

Which change is exothermic?

A.  $\frac{1}{2}$Cl₂ (g) → Cl (g)

B.   K (g) → K⁺ (g) + e⁻

C.   KCl (s) → K⁺ (g) + Cl⁻ (g)

D.   Cl (g) + e⁻ → Cl⁻ (g)

Determine the enthalpy of solution of copper(II) chloride, using data from sections 18 and 20 of the data booklet.

The enthalpy of hydration of the copper(II) ion is −2161 kJ mol⁻¹.

Deduce the change in enthalpy, ΔH, in kJ, when 56.00 g of ethanol is burned. Use section 13 in the data booklet.

What is the order of increasing (more exothermic) enthalpy of hydration?

Xⁿ⁺ (g) → Xⁿ⁺ (aq)

A.  Ca²⁺, Mg²⁺, K⁺, Na⁺

B.  Na⁺, K⁺, Mg²⁺, Ca²⁺

C.  K⁺, Na⁺, Ca²⁺, Mg²⁺

D.  Mg²⁺, Ca²⁺, Na⁺, K⁺

Justify why K₂O has a lower lattice enthalpy (absolute value) than Na₂O.

The standard enthalpy of formation of sodium oxide is −414 kJ mol⁻¹. Determine the lattice enthalpy of sodium oxide, in kJ mol⁻¹, using section 8 of the data booklet and your answers to (d)(i).

(If you did not get answers to (d)(i), use +850 kJ mol⁻¹ and +600 kJ mol⁻¹ respectively, but these are not the correct answers.)

Calculate values for the following changes using section 8 of the data booklet.

ΔH_atomisation (Na) = 107 kJ mol⁻¹
ΔH_atomisation (O) = 249 kJ mol⁻¹

$\frac{1}{2}$O₂(g) → O^2- (g):

Na (s) → Na⁺ (g):

Which represents electron affinity?

A.  Al²⁺(g) → Al³⁺(g) + e⁻

B.  C (g) + e⁻ → C⁻(g)

C.  Cl₂(g) → 2Cl (g)

D.  S (s) → S⁺(g) + e⁻

Which substance has the highest lattice enthalpy?

A.  $\text{KCl}$

B.  ${\text{CaCl}}_2$

C.  $\text{KF}$

D.  ${\text{CaF}}_{\text{2}}$

Consider the Born–Haber cycle for the formation of sodium oxide:

What is the lattice enthalpy, in kJ mol⁻¹, of sodium oxide?

A.  414 + 2(108) + 249 + 2(496) − 141 + 790

B.  414 + 2(108) + 249 + 2(496) + 141 + 790

C.  −414 + 2(108) + 249 + 2(496) − 141 + 790

D.  −414 − 2(108) − 249 − 2(496) + 141 − 790