Grade 11 → Classification of elements and periodicity in properties ↓
Periodic trends in properties
The periodic table is a powerful tool in chemistry. It is arranged in such a way that the elements are listed in order of increasing atomic number, creating periodic patterns or trends. These trends provide valuable information about the properties of elements, helping scientists predict behavior in chemical reactions. Let's look at these trends in detail.
1. Atomic radius
The atomic radius is one of the most fundamental chemical properties. It refers to the size of the atom, more precisely, the distance from the nucleus to the outermost shell of electrons.
Trends over a period
As you move from left to right across a period, the atomic number increases. This means that more protons are added to the nucleus, and electrons are added to the same shell. The increased nuclear charge pulls the electrons closer to the nucleus, resulting in a smaller atomic radius.
Na
(sodium) is larger than that of Cl
(chlorine).
Downward trend in the group
The atomic radius increases as you move down a group. This is because each successive element has an additional electron shell, which more than accounts for the increased nuclear charge, increasing the size of the atom.
(K)
has a much larger radius than Lithium (Li)
.
2. Ionization energy
Ionization energy is the energy required to remove an electron from an isolated gaseous atom. It provides information about the strength of the electron bond of the atom.
Trends over a period
Ionization energy generally increases across a period. With a greater number of protons in the nucleus, the electrons are bound more tightly, making it harder to remove one.
Na
is less than that of Ar
due to increase in nuclear charge across the period.
Downward trend in the group
Ionization energy decreases as we go down the group. The outer electrons are farther from the nucleus and experience a shielding effect from the inner-shell electrons, making them easier to remove.
Li
with Cs
, caesium has a lower ionization energy because its outer electron is farther from the nucleus.
3. Electron affinity
Electron affinity measures how easily an atom can gain an electron. This property represents the energy change when an electron is added to a neutral atom.
Trends over a period
Generally, electron affinity becomes more negative (more energy is released) as you proceed across a period from left to right. Atoms more eagerly accept electrons to achieve a full valence shell.
(Cl)
has a more negative electron affinity than sodium (Na)
.
Downward trend in the group
Electron affinity becomes less negative as we go down the group. The additional electron shells reduce the positive attraction of the nucleus on the incoming electrons.
F
is greater than I
(iodine).
4. Electronegativity
Electronegativity indicates the tendency of an atom to attract shared electrons in a chemical bond. It is an important factor in determining the nature of the bond between atoms.
Trends over a period
Electronegativity increases as you move across a period. Atoms on the right side of the table have more protons, creating a greater attraction for electrons.
(O)
is more electronegative than carbon (C)
.
Downward trend in a group
Electronegativity decreases as we go down the group. Larger atomic radius and additional electron shells weaken the pull on shared electrons.
(F)
is more electronegative than iodine (I)
.
Conclusion
Periodic trends in atomic radius, ionization energy, electron affinity and electronegativities reveal the systematic and predictable nature of the behaviour of elements. Understanding these trends helps scientists make informed predictions about chemical reactivity and the nature of bonds. Such patterns are fundamental to guiding intensive explorations, experiments and theoretical developments in chemistry.