Both have a giant metallic lattice structure. Both K+ and Na+ have a charge of +1 and each atom contributes one electron to the sea of delocalised electrons, but K+ has a larger ionic radius, so K+ has a lower charge density.[1]
Hence, less energy is needed to overcome the weaker electrostatic attraction between the K+ cations and the sea of delocalised electrons.[1]
Both have a giant ionic lattice structure. F− has a smaller ionic radius than Cl− and the same charge, so the magnitude of the lattice energy of NaF is greater than that of NaCl.[1]
Hence, more energy is needed to overcome the stronger electrostatic attraction between Na+ and F− ions than between Na+ and Cl− ions.[1]
SiO2 has a giant molecular structure, CO2 has a simple molecular structure[1]
Hence, more energy is needed to overcome the stronger, extensive covalent bonds between Si and O atoms in SiO2 than the weaker instantaneous dipole–induced dipole interactions between CO2 molecules.[1]
It is linear, so the two bond dipoles cancel: no overall dipole moment.[1]
The hydrogen bonds in ice hold the molecules in an open structure, further apart than in the liquid, so ice is less dense.[1]
CThere are, on average, more hydrogen bonds between water molecules than there are between HF molecules.[1]
Each H2O can form two hydrogen bonds. HF tends to form fewer hydrogen bonds per molecule.
Bgiant covalent to giant metallic[1]
White tin must have a giant metallic structure as it conducts electricity in the solid and liquid state and it is malleable.
Grey tin cannot have a giant ionic structure as it does not conduct electricity in the liquid state, thus it has to be a giant covalent structure.
S is a Period 3 element and can expand its octet (i.e. have more than eight electrons) due to the presence of vacant and energetically accessible d orbitals, hence SF6 can exist.[1]
However, O is in Period 2 and does not have vacant and energetically accessible d orbitals to accommodate more than eight electrons, hence OF6 cannot be formed.[1]
NoteOrbitals that are energetically accessible have low energies and can be used for bonding.
A small amount of energy is required to overcome the weak instantaneous dipole–induced dipole interactions between Kr atoms, hence the low boiling point.[1]
A large amount of energy is required to overcome the strong electrostatic attraction between Rb+ cations and its sea of delocalised electrons in Rb, hence the high boiling point.[1]
Halogens have simple molecular structures held by weak instantaneous dipole–induced dipole interactions.
The number of electrons of the halogens increases down the group, thus the polarisability of the electron cloud increases.[1]
A larger amount of energy is required to overcome the stronger instantaneous dipole–induced dipole interactions between molecules down the group. Hence, boiling point increases down the group.[1]
Both HF and HCl have simple molecular structures.[1]
More energy is required to overcome the stronger intermolecular hydrogen bonds in HF than the weaker permanent dipole–permanent dipole interactions between HCl molecules during boiling. Thus, HF has a higher boiling point than HCl.[1]
Both NH3 and PH3 have simple molecular structures. NH3 molecules have stronger hydrogen bonding via the lone pair on N and the highly δ+ hydrogen of the N–H bond, while PH3 molecules have weaker permanent dipole–permanent dipole interactions.[1]
Thus, more energy is needed to overcome the stronger hydrogen bonds and NH3 has a higher boiling point.[1]
The bond angle in NH3 is smaller than that in BF3. There are three bond pairs and one lone pair around N in NH3, so NH3 has a trigonal pyramidal shape. Its bond angle is less than the tetrahedral angle of 109.5° because of the greater lone pair–bond pair repulsion compared with bond pair–bond pair repulsion.[1]
BF3 has a trigonal planar shape as it has three bond pairs and no lone pairs around B, so its bond angle is 120°.[1]
CH4 has a larger bond angle than NH3. CH4 has four bond pairs and no lone pairs around C, so the bond angle in CH4 is the ideal bond angle of 109.5°.[1]
However, NH3 has three bond pairs and one lone pair. Since the lone pair is closer to the nucleus of N, there is greater lone pair–bond pair repulsion than bond pair–bond pair repulsion, so the bond pairs are pushed closer together and the bond angle in NH3 is smaller.[1]
Both have a giant ionic lattice structure, held together by strong electrostatic attractions between oppositely charged ions.
|lattice energy| ∝ q+ × q−r+ + r−
Since Ca2+ and O2− have smaller cationic and anionic radii than Na+ and Cl−, and they also have charges of higher magnitude, the magnitude of the lattice energy of CaO is higher than that of NaCl.[1]
Hence, more energy is needed to overcome the stronger electrostatic attraction between Ca2+ and O2− ions than between Na+ and Cl− ions, and CaO has a much higher melting point.[1]
The Cl–C–Cl bond angle is larger than the H–C–H bond angle. Cl atoms are bigger and experience greater repulsion in their electron clouds.[1]
or The Cl–C–Cl bond angle is smaller than the H–C–H bond angle (actual data). Cl is more electronegative than H. Cl draws the C–Cl bond pair closer to itself. The C–Cl bond pair of electrons are thus further away from the nucleus of C and exert less repulsion than those in the C–H bond.
NoteEither answer scores if it is justified. Stating that the Cl–C–Cl bond angle is larger because Cl is more electronegative is a contradiction.
Trigonal pyramidal[1]
3 bond pairs and 1 lone pair[1]