Periodic Properties and Their Trends: MDCAT Chemistry notes

Why the trends happen

Two factors control almost every periodic trend.

• Across a period, the nuclear charge increases while electrons enter the same shell. The shielding effect stays essentially constant, so the outer electrons are pulled in more tightly.
• Down a group, a new shell is added each time. Shielding by inner shells increases and the outer electron is farther away, so it is held more loosely. Rb shows more shielding than Li, Na or K.

Summary of trends

Atomic and ionic radius

• Alkali metals have the largest atoms in each period. Na > Mg > Al > P > S, and Ba is the largest of Be, Mg, Ca and Ba.
• Covalent radius is half the single-bond distance between two similar atoms. For H it is about 37 pm. A bond length is the sum of the two covalent radii. Example: C–Cl in $\mathrm{CH_3Cl}$ is 176.7 pm and $r_{\mathrm{Cl}}$ is 99.4 pm, so $r_{\mathrm{C}} = 77.3$ pm.
• A cation is smaller than its atom. An anion is larger, because the extra electron–electron repulsion expands the cloud. The fluoride ion is about 136 pm, much larger than the F atom.
• Ionic radius falls across a period because the nuclear charge increases.
• Larger atoms are more polarizable.

Ionization energy (IE)

IE falls down a group because the atomic radius and shielding increase, so Ba < Ca < Mg. It rises across a period because the nuclear charge increases.

• Mg > Al: Mg has a filled $3s^2$ subshell, and the 3p electron of Al is easier to remove.
• N > O: N has a stable half-filled $2p^3$. In O, two electrons are paired in one p orbital and repel each other.

A high IE means the element is less metallic, less likely to form cations and a weaker reducing agent.

Electron affinity and electronegativity

• Electron affinity is the energy released when an electron is added to a gaseous atom.
• Order in group VII: Cl > F > Br > I. Fluorine's small, crowded 2p shell repels the incoming electron.
• Electronegativity has no unit because it is a relative (Pauling) scale. F is the most electronegative element (4.0), because it is the smallest atom in its period with a high nuclear charge. HF therefore has the largest electronegativity difference of the hydrogen halides.

Other trends asked

• Melting points: in periods 2 and 3 the maximum is at C and Si, which are giant covalent networks. In group I, melting and boiling points fall down the group, so Cs is lowest. In group VIA they rise down the group, so O is lowest.
• Oxides of period 3: $\mathrm{Na_2O}$ and MgO are basic, $\mathrm{Al_2O_3}$ is amphoteric, and the oxides of P and S are acidic. Al quickly forms a protective oxide film.
• Hydration energy is highest for small, highly charged ions, so $\mathrm{Al^{3+}}$ has the most. Li⁺ has the highest in group I.
• Lithium is anomalous because of its small radius and high charge density.
• Ionic character rises with cation size and falls with cation charge, so CsCl is the most ionic and $\mathrm{BeCl_2}$ the least.
• Hydride boiling points rise down a group because of stronger dispersion forces, so period-4 hydrides boil higher than period-3 hydrides. The first members are exceptions where hydrogen bonding is present.

Common MDCAT traps

• Shielding is constant across a period. It does not increase.
• F has the highest electronegativity but not the highest electron affinity. Cl does.
• Remember the IE dips at Al (after Mg) and O (after N).
• Covalent radius is defined using two identical atoms.
• Aluminium oxide is amphoteric, not basic.

Quick revision

• Radius decreases across a period and increases down a group.
• F is the most electronegative element, and electronegativity has no unit.
• Metallic character is greatest at the bottom left of the table, so Rb is more metallic than K, Ca or Fe.
• The fluoride ion is about 136 pm.
• Low alkali-metal density comes from large atomic volume.