Question

1. The bond dissociation enthalpies for what series of group 17 binary compounds can be used as an indicator for group trends in sigma bond enthalpies .?

2. What are the dominant attractive (bonding) contributions that make the dimerization of   N_2(g) to a hypothetical T_d   N₄ molecule highly thermodynamically unfavorable   and that make T_d P_4(s) thermodynamically favorable compared to P_2(g) ?

3.The standard enthalpies of vaporization and atomization for the As₄ molecular solid standard state of elemental arsenic (As_(s)) are 35 kj and 301kj per mole of arsenic atoms respectively . What is the average effective As-As bond dissociation enthalpy for the standard state of As_(s) ?

Answer 1

1. The binary compounds whose bond dissociation enthalpies can be used as an indicator for group trends in sigma bond enthalpies are Halogen hydrides HX, rather than dihalogen compounds X₂ (X=F,Cl,Br,I) because:

In case of dihalogens, the high repulsion between lone pairs of F atoms in F₂ weakens the Bond and hence it becomes lower than other dihalogens, which greatly underestimates F's capability to form strong covalent bonds. In HF, as H is devoid of lone pairs, such deviation is not observed and hence the HX bond enthalpies are better indicator of group trends in sigma bond enthalpies.(HX has only single sigma bond)

2. The ability to form bonding decreases down the group because, the orbital overlap decrease due to Increasing atomic size and as the orbitals become more diffused. Dinitrogen has very strong triple bond composed of one sigma and two bonds, with high bond enthalpy of 226 kcal/mol. N-N single bond is much weaker at 40 kcal/mol, whereas in case of phosphorous the 6 P-P triple bonds in 2P₂ s has lower (2X117 kcal/mol=234 kcal/mol) bond enthalpy than 6 P-P single bond (6X47.8 kcal/mol=287 kcal/mol)in P_4.  (Enthalpies at standard conditions)

3. As_4(s)As_4(g)   

As_4(g)4As_(g)

   

Answer 2

1. The binary compounds whose bond dissociation enthalpies can be used as an indicator for group trends in sigma bond enthalpies are Halogen hydrides HX, rather than dihalogen compounds X₂ (X=F,Cl,Br,I) because:

In case of dihalogens, the high repulsion between lone pairs of F atoms in F₂ weakens the Bond and hence it becomes lower than other dihalogens, which greatly underestimates F's capability to form strong covalent bonds. In HF, as H is devoid of lone pairs, such deviation is not observed and hence the HX bond enthalpies are better indicator of group trends in sigma bond enthalpies.(HX has only single sigma bond)

2. The ability to form bonding decreases down the group because, the orbital overlap decrease due to Increasing atomic size and as the orbitals become more diffused. Dinitrogen has very strong triple bond composed of one sigma and two bonds, with high bond enthalpy of 226 kcal/mol. N-N single bond is much weaker at 40 kcal/mol, whereas in case of phosphorous the 6 P-P triple bonds in 2P₂ s has lower (2X117 kcal/mol=234 kcal/mol) bond enthalpy than 6 P-P single bond (6X47.8 kcal/mol=287 kcal/mol)in P_4.  (Enthalpies at standard conditions)

3. As_4(s)As_4(g)   

As_4(g)4As_(g)