Periodic Trends in Physical Properties

Periodic trends in physical properties

Down the group

i. The atomic size (atomic radius)

The atomic radius of an uncombined atom cannot be defined strictly because of the uncertain boundary of electron clouds. The distance between the nuclei of chemically or covalently combined atoms can be measured accurately by x-ray diffraction method.

Therefore, the atomic radius is defined as half the distance between the nuclei of two similar/identical atoms joined by a single covalent or metallic bond. There is a significant regular increase in atomic radii among elements down the group. This trend is due to the increase in the number of electrons, number of shells, and the increase in size as new shells are added. The following table shows the trend down Group IA.

Element	Atomic Radius (Å)
Li	1.52
Na	1.86
K	2.27
Rb	2.65
Cs	2.62

ii. Ionic radii

The ionic radius of an atom of an element is measured in the same way as the atomic radius. The ionic radii increase down the group for similar reasons as atomic radii. Cations are smaller than their parent atoms, while anions are larger because electrons are either removed or added. The ionic radii trend for Group VIIA elements is shown below.

Element	Ionic Radius (Å)	Ionic Type
Cl	0.99	Cl–
F	0.64	F–
Br	1.14	Br–
I	1.33	I–

iii. Ionization energy

Ionization energy refers to the energy required to remove an electron from a gaseous atom or ion. Ionization energy decreases down the group due to increasing atomic size, with the outer electrons becoming less tightly bound. The following trends show the variation of ionization energy down a group, where alkali metals exhibit the lowest ionization energies.

iv. Electron affinity (E.A)

Electron affinity is the energy change that occurs when an electron is added to a gaseous atom or ion. Halogens have the most negative electron affinities, which makes them readily form halide ions. There is a general decrease in electron affinity down the group, as shown in the trend for Group VIIA elements.

v. Electronegativity

Electronegativity measures the attraction an atom exerts on electron pairs in a covalent bond. Electronegativity decreases as one moves down a group due to the increase in atomic size and shielding effects from inner electrons. For example, Li is much more electronegative compared to other group members.

Across the period

i. Atomic radii (atomic size)

The atomic radii decrease steadily across each period due to the increase in effective nuclear charge, which pulls electrons closer to the nucleus. The atomic size of elements decreases from Li to F and then slightly increases for noble gases like Ne.

Element	Atomic Radius (Å)
Li	1.52
Be	1.12
B	0.85
C	0.77
F	0.64
Ne	0.38

ii. Ionization energy

Ionization energy increases across a period as the effective nuclear charge increases, making it harder to remove electrons. The trend shows that elements like Beryllium (Be) and Nitrogen (N) have high ionization energies due to their stable electron configurations.

iii. Ionic radii

Ionic radii exhibit periodic trends, with cations being smaller than their parent atoms and anions being larger. The trend follows the same pattern as atomic radii, with variations for transition metals and lanthanides.

iv. Electron affinity

Non-metals tend to have higher electron affinities than metals. As the effective nuclear charge increases, electron affinity becomes more exothermic, making elements more likely to gain electrons. This trend is observed across the period.

v. Electronegativity

Electronegativity increases across a period due to the increase in nuclear charge and the decrease in atomic size. Halogens have the highest electronegativity in each period.

vi. Melting point

Melting points vary across periods due to the different types of forces between atoms. In Period 3, for example, the melting point increases sharply from sodium (Na) to magnesium (Mg) due to the number of valence electrons contributing to the metallic bonding.