CI and II only[1]
The attraction between ions is stronger when the charges are larger. Ca2+ and S2− carry charges of 2, K+ and Cl− only 1, so I and II both explain it.
III is false: K+ and Ca2+ have the same number of electrons, and Ca2+ has more protons, so Ca2+ is smaller.
Ca giant lattice of oppositely charged ions[1]
A metal is a lattice of positive ions only, held together by the sea of delocalised electrons. There are no negative ions.
The other statements are correct.
BBoth statements are correct but statement 2 does not explain statement 1.[1]
Both statements are true. Mobile electrons explain why metals conduct electricity.
Malleability comes from layers of ions sliding, and the high melting point from the strong metallic bonds, not from the electrons being mobile.
BWF4[1]
The sixth ionisation energy falls sharply from X to Y. For Y it is the last valence electron, so Y has 6 valence electrons. For X the sixth electron comes from an inner shell, so X has 5 valence electrons.
W is the element before X, with 4 valence electrons, and forms WF4.
Abent, 110°[1]
Diazene is H–N=N–H. Each N has a single bond to H, a double bond to N (one region) and one lone pair: 2 bond pairs and 1 lone pair.
Three regions round each N point to the corners of a triangle, and the lone pair squeezes the angle below 120°, so the shape is bent at about 110°.
BT-shaped, trigonal pyramidal[1]
I has 7 valence electrons: 3 bond pairs and 2 lone pairs. The 5 pairs form a trigonal bipyramid with both lone pairs in the middle, so IF3 is T-shaped.
P has 5 valence electrons: 3 bond pairs and 1 lone pair, so PF3 is trigonal pyramidal.
Crow C[1]
3 bond pairs and 1 lone pair give a trigonal pyramid with a bond angle of 107°, as in NH3.
2 bond pairs and 1 lone pair give a bent shape just under 120°. 2 bond pairs and 2 lone pairs give about 104.5°. 3 bond pairs and 2 lone pairs give a T-shape with angles just under 90°.
APOCl3 and CCl4[1]
POCl3 has four regions round P (a P=O double bond and three P–Cl bonds) and no lone pairs, so it is tetrahedral, like CCl4.
AlCl3 is trigonal planar but NCl3 is pyramidal. SO2 is bent but CO2 is linear. ClF3 is T-shaped but SF4 is see-saw.
DNH3 and H3O+[1]
NH3 and H3O+ each have 3 bond pairs and 1 lone pair round the central atom: trigonal pyramidal, 107°.
CH3+ has 3 bond pairs and no lone pair (120°). H2O and NH2− have 2 bond pairs and 2 lone pairs (about 104.5°).
B180°[1]
The anion is attracted to the electron-deficient (slightly positive) region of the iodine atom.
That region lies directly opposite the X–I bond, so X, I and A− lie in a straight line: 180°.
Crow C[1]
Trichloromethane, CHCl3, is tetrahedral round C, and the C–H and C–Cl dipoles do not cancel, so it is polar.
XeF4 is square planar, SeF6 is octahedral, and BCl3 is trigonal planar but non-polar because its dipoles cancel.
C3 only[1]
The number of electrons increases from CH3CH2Cl to CH3CH2I, so the instantaneous dipole–induced dipole attractions get stronger and the boiling point rises (3).
Electronegativity falls from Cl to I, so 1 and 2 are wrong. Statement 4 is true, but boiling does not break covalent bonds, so it does not explain the trend.
ACO2[1]
CO2 molecules are non-polar, so only instantaneous dipole–induced dipole forces act between them.
Cu has metallic bonds, H2O has hydrogen bonds as well, and SiO2 is held by covalent bonds throughout.
C3 and 4 only[1]
From ethane to butane the number of electrons rises, so the instantaneous dipole–induced dipole forces get stronger (4). Butane and 2-methylpropane have the same number of electrons, but the chain-shaped butane has more surface contact (3).
Hydrocarbons have no H on F, O or N, so there is no hydrogen bonding, and boiling breaks no covalent bonds.
DNH3 forms few hydrogen bonds on average and has the smallest bond dipole.[1]
NH3 has only one lone pair, so it forms 1 hydrogen bond per molecule on average, and N–H is the least polar of the three bonds, so each hydrogen bond is the weakest.
HF and NH3 both form 1 per molecule. H–F is more polar than H–O. Water boils highest because it forms 2 per molecule, not because it dimerises.
CThere are, on average, more hydrogen bonds between water molecules.[1]
Water has 2 H atoms and 2 lone pairs, so it forms 2 hydrogen bonds per molecule on average. HF has only 1 H atom, so it forms 1.
H–F is more polar than H–O, so each hydrogen bond between HF molecules is actually stronger: the extent, not the strength, explains the result.
CII and III only[1]
H2O has far fewer electrons than H2Se, so its instantaneous dipole–induced dipole (temporary dipole) forces are weaker (II). O is more electronegative than S, so H2O is more polar and has stronger permanent dipole–permanent dipole forces than H2S (III).
I is wrong: H2S is the more polar molecule. H2Se boils higher because of its stronger instantaneous dipole–induced dipole forces from its larger number of electrons.
BIV and V[1]
Liquid–liquid extraction separates liquids that are immiscible, because they form two layers.
Hexane (IV) is non-polar and water (V) is polar, so they do not mix. The pairs in the other options mix: ethanol and hexane (A), acetone and hexane (C), and chloroform and ethanol (D).
Drow D[1]
Simple molecular substances are separate molecules held to one another by weak intermolecular forces.
A giant molecular lattice is atoms, not molecules, joined by covalent bonds. Each ion attracts all the oppositely charged ions round it. Only the valence electrons of a metal are delocalised.
Ba giant tetrahedral network of covalent bonds[1]
Hardness and a very high melting point mean many strong covalent bonds throughout the solid, as in diamond: each Si and C bonded to 4 others in a tetrahedral network.
Si and C are both non-metals, so they do not form ions. A layer structure would be soft, and a simple molecular structure would melt at a low temperature.
Cionic, simple molecular, giant molecular[1]
X melts high and conducts when molten: ionic. Y melts low and does not conduct: simple molecular.
Z melts very high but does not conduct when molten: giant molecular.
Crow C[1]
Sodium sulfate is a lattice of Na+ and SO42− ions (ionic bonding), and the S and O atoms inside each sulfate ion are joined by covalent bonds.
Ice is simple molecular. Iodine is non-polar, so it has instantaneous dipole–induced dipole forces, not permanent dipoles. Copper has metallic bonding.
Bbond length[1]
The atoms get larger down the group, so the distance between the two nuclei increases. Their larger, more diffuse orbitals overlap less effectively, so the bond energy falls: Cl–Cl 244, Br–Br 193 and I–I 151 kJ mol−1.
The number of electrons increases, so the instantaneous dipole–induced dipole forces get stronger and the halogens get less volatile.
DA dative bond can be formed from the head-on overlap of orbitals.[1]
Once formed, a dative bond is identical to any other covalent bond, so it can be a σ bond.
A is false: π bonds are weaker than σ bonds, as sideways overlap is less effective than head-on overlap. B is false: when the number of shared pairs increases, the bond length usually decreases. C is false: some non-polar bonds are stronger than polar bonds, because the number of shared pairs and how well the orbitals overlap matter more.
A1, 2 and 3[1]
ICl2− is linear (2 bond pairs, 3 lone pairs), planar, 180°. ClF3 is T-shaped (3 bond pairs, 2 lone pairs), planar, 90°. ClO3− is trigonal pyramidal (3 bond pairs, 1 lone pair), 107°.
1 is correct: the bond angles are ClF3 < ClO3− < ICl2−. 2 is correct: ClF3 and ClO3− have a net dipole moment, and only ICl2− is non-polar. 3 is correct: only ClO3− is not planar.