Periodicity

Group 5

Group 7

Group 15

Group 17

Soft solids

Hard solids

Liquids

Gases

$Li^{+}$

$Cl^{-}$

$Na^{+}$

$K^{+}$

Cerium

Plutonium

Lutetium

Barium

Elements are arranged by increasing atomic mass.

The group number corresponds to the number of electron shells.

The period number corresponds to the number of valence electrons.

Elements with the same number of valence electrons are in the same group.

They have the highest values in their respective periods.

They have the lowest values in their respective periods.

Electronegativity is generally not assigned to them as they do not readily form bonds.

Their electronegativity increases down the group.

Its highest energy occupied subshell is a d-subshell.

It has 10 valence electrons.

It is located in Group 2 of the periodic table.

It always forms ions with a $2+$ charge.

Atomic radius

Electronegativity

First ionization energy

Oxidizing power

$Br_2(aq) + 2NaCl(aq) \rightarrow Cl_2(aq) + 2NaBr(aq)$

$I_2(aq) + 2KBr(aq) \rightarrow Br_2(aq) + 2KI(aq)$

$Cl_2(aq) + 2KI(aq) \rightarrow I_2(aq) + 2KCl(aq)$

$I_2(aq) + 2NaCl(aq) \rightarrow Cl_2(aq) + 2NaI(aq)$

$O^{2-}$

$F^-$

$Na^+$

$Mg^{2+}$

Atomic radius

Melting point

Electronegativity

Ionic radius

Group 1

Group 2

Group 13

Group 14

Fluorine

Chlorine

Bromine

Iodine

It decreases

It increases

It remains constant

It increases then decreases

Silicon

Phosphorus

Sulfur

Chlorine

The $Na^{+}$ ion has a higher nuclear charge than the $Na$ atom.

The $Na^{+}$ ion is isoelectronic with the $Ar$ atom.

The ratio of protons to electrons is higher in the $Na^{+}$ ion than in the $Na$ atom.

There is increased electron-electron repulsion in the $Na^{+}$ ion compared to the $Na$ atom.

Argon is often measured as larger than Chlorine because its radius is typically reported as a Van der Waals radius.

Argon is significantly smaller than Chlorine due to the massive increase in nuclear charge.

Argon and Chlorine have identical radii because they are in the same period.

Argon is larger because it has an extra energy level compared to Chlorine.

Metallic, forms a basic oxide, high electrical conductivity.

Non-metallic, forms an acidic oxide $P_4O_{10}$, exists as $P_4$ molecules.

Metalloid, forms an amphoteric oxide, semi-conductor.

Non-metallic, forms a neutral oxide, low melting point.

$Na$

$Mg$

$Al$

$Si$

Boron has a higher nuclear charge than Beryllium.

Beryllium has a full $2s$ subshell, which is more stable than the single $2p$ electron in Boron.

The $2p$ electron in Boron is further from the nucleus and more shielded than the $2s$ electrons in Beryllium.

Beryllium is smaller than Boron.

Lithium

Calcium

Iron

Chlorine

Electron affinity

First ionization energy

Electronegativity

Lattice energy

Carbon

Phosphorus

Neon

Sulfur

Decreases

Increases

Stays the same

Fluctuates randomly

Strongly acidic

Weakly acidic

Weakly alkaline

Neutral

Cesium

Lithium

Fluorine

Helium

$SiO_{2}$ is a polar molecule while $P_{4}O_{10}$ is non-polar.

$SiO_{2}$ has a giant covalent structure while $P_{4}O_{10}$ is molecular covalent.

$SiO_{2}$ has stronger London dispersion forces.

The $Si-O$ bond is stronger than the $P-O$ bond.

The shielding effect of the $3d$ subshell decreases significantly when it is fully occupied.

The effective nuclear charge ($Z_{eff}$) decreases as electrons are added to the $4s$ subshell.

The electron-electron repulsion within the $3d$ subshell becomes significant as it reaches its maximum capacity, counteracting the increase in nuclear charge.

The atoms at the end of the transition series have a lower coordination number in their metallic lattice, leading to a larger atomic radius.

Basic $\rightarrow$ Acidic $\rightarrow$ Amphoteric

Basic $\rightarrow$ Amphoteric $\rightarrow$ Acidic

Acidic $\rightarrow$ Amphoteric $\rightarrow$ Basic

Amphoteric $\rightarrow$ Basic $\rightarrow$ Acidic

Variable oxidation states

Formation of colored complexes

Magnetic properties

Catalytic activity

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