Elements in Order: Understanding the Periodic Table

Introduction:

Understanding chemistry, a complex subject, can be made easier if we organize all the facts into a simple pattern. The periodic table of elements has helped with this for over 100 years. It’s a big achievement in the history of chemistry. The table helps us understand how the properties of elements and their compounds behave in a predictable pattern. In this lesson, we’ll learn more about the periodic table and how it helps us understand the behavior of different elements.

History of the Periodic Table: Contributions and Development:

The history of the periodic table is fascinating. Some notable contributors to the field were Al-Razi, Dobereiner, Newland, and Mendeleev. Al-Razi classified substances based on their physical and chemical properties. Dobereiner arranged the known elements into groups called Triads in 1829, each containing three elements with similar properties. In 1864, Newland classified 62 known elements in order of increasing atomic mass and noticed that every eighth element had properties in common with the first one, which he called the Law of Octaves.

👉 Enroll now in the MDCAT-2025 Online Course and let your preparation speak for itself.

In 1871, Dmitri Mendeleev, a Russian chemist, created a more comprehensive periodic table. He arranged the elements with similar chemical properties into eight vertical columns called Groups and horizontal rows called Periods based on their atomic masses. Mendeleev also left gaps for undiscovered elements and predicted their properties based on their positions in the table. For example, he predicted the properties of germanium, which had not yet been discovered at the time. In 1886, germanium was discovered and found to match Mendeleev’s predictions.

The essential features of the Periodic Table:

Groups and Periods:

The modern periodic table arranges all elements in ascending order of their atomic numbers. The periodic table has eight vertical columns called Groups, which contain elements with similar properties. Groups are numbered I to VIII using Roman numerals, and each group is divided into two subgroups: A and B. A subgroups contain representative elements, while B subgroups contain less typical elements called transition elements, which are arranged in the center of the table. The periodic table also has horizontal rows called Periods. There are seven periods in the table, numbered 1 to 7 using Arabic numerals.

• Period 1 contains only two elements: hydrogen and helium.
• Periods 2 and 3 are short periods that each contain eight representative elements belonging to the A subgroup. In these periods, every eighth element has properties resembling the first element. Lithium and beryllium, which are in the 2nd period, have properties that are similar to sodium and magnesium in the 3rd period, respectively. Boron and aluminum both have an oxidation state of +3, and fluorine in the 2nd period is similar in some ways to chlorine in the 3rd period.
• Periods 4 and 5 are long periods that each contain eighteen elements, with eight representative elements belonging to the A subgroup and ten transition elements belonging to the B subgroup. The repetition of properties among these elements occurs after 18 elements.
• Period 6 is also a long period with thirty-two elements, including eight representative elements, ten transition elements, and fourteen elements called Lanthanides, which have similar properties and are shown separately at the bottom of the periodic table.
• Period 7 is currently incomplete, containing only two normal elements (₈₇Fr and ₈₈Ra), ten transition elements, and fourteen inner transition elements called Actinides, which follow ₈₉Ac. The Actinides are also shown at the bottom of the table, under the Lanthanides. Due to their scarcity, inner transition elements are also known as rare earth elements.

Families in the Periodic Table:

When you study the periodic table, you may notice that certain rows of elements with similar properties have been assigned common names, such as transition elements, Lanthanides, Actinides, or Rare Earth elements. Some typical elements belonging to sub-group A have also been assigned family names because of their peculiar characteristics.

For example, elements of group IA are known as Alkali Metals because of their property to form strong alkalis with water.

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

Similarly, the elements of group IIA are known as Alkaline Earth Metals because of their presence in the Earth’s crust and alkaline character.

The Halogen family is another important family in the periodic table, with elements of group VIIA named “Halogens” because of their salt-forming properties. Group VIIIA, also known as Noble Gases, is the least reactive group of gases in the periodic table. These family names are useful for quickly recognizing an element in the periodic table.

Blocks in the Periodic Table

Elements in the periodic table can be divided into four groups based on how their electrons are arranged. This is helpful in understanding how they bond with other elements. The s-block elements are the elements in groups IA and IIA, which have electrons in the s orbital. The p-block elements are in groups IIIA to VIIIA, except for helium, and have electrons in the p orbital. The transition elements have electrons in the d-orbital, so they are called d-block elements. Lanthanides and Actinides have electrons in the f-orbital, which is why they are called f-block elements. This classification can help us understand how elements react with each other and their properties, especially their valency or oxidation state.

Classification of Elements Based on Metallic Character

The periodic table can also be classified based on the metallic character of elements. Metals are generally located on the left side, bottom, and center of the table, while non-metals are typically found in the upper right corner. Some elements, such as the lower members of groups IIIA, IVA, and VA, possess properties of both metals and non-metals and are called metalloids. Non-metals are found in groups IVA to VIIIA at the top right corner of the table, while metalloids are located just below the “steps” such as Si, As, and Te. All elements, except for hydrogen, are metals. This classification is useful in understanding the properties of elements and their behavior in chemical reactions.

The Physical Properties and Trends in the Periodic Table

The modern periodic table arranges elements in order of their atomic numbers and groups and periods are based on similar properties. However, as the number of protons and electrons increases, the properties of the elements gradually change within a group or period. We will look at some patterns in physical properties.

Atomic Size

When moving down a group in the periodic table, from top to bottom, the number of electron shells in an element increases. As a result, the atomic radius of the element also increases.

Refer to the tables below to compare the number of electron shells and atomic radii of elements in the alkali metal and halogen groups.

Element	Symbol	Atomic Number	Atomic Radius (pm)	Ionic Radius (pm)	Trend
Lithium	Li	3	134	74	Smallest
Sodium	Na	11	154	102
Potassium	K	19	196	138
Rubidium	Rb	37	211	149
Cesium	Cs	55	225	170	largest

 

Table 1.1: Trends in alkali metals

Element	Symbol	Atomic Number	Atomic Radius (pm)	Ionic Radius (pm)	Trend
fluorine	F	9	71	131	Smallest
Chlorine	Cl	17	99	181
Bromine	Br	35	114	196
Iodine	I	53	133	220	largest

 

 

Table 1.2: Trends in halogens

Element	Li	Be	B	C	N	O	F	Ne
Electron Configuration	2,1	2,2	2,3	2,4	2,5	2,6	2,7	2,8
Nuclear Charge	3+	4+	5+	6+	7+	8+	9+	10+
Atomic Radius (pm)	134	90	82	77	75	73	71	69
General Trend	(largest)							(smallest)

Table 1.3: Trends in halogens

Atomic Radius:

Atoms are incredibly tiny, and we can’t see them with even the most powerful microscopes. We can’t measure their size directly, but we can measure the distance between two bonded atoms of any element, and half of that distance is considered to be the radius of the atom. In the periodic table, the atomic radius increases as we go down a group because the atomic number increases, and an extra shell of electrons is added in each period. However, the atomic radius decreases as we go across a period from left to right. This is because the positive charge in the nucleus increases, pulling the negatively charged electrons closer to the nucleus. This effect is especially noticeable in longer periods that involve “d” and “f” subshells and is called Lanthanide Contraction.

Ionic Radius:

When an atom loses one or more electrons, it becomes a positive ion. This process makes the atom smaller for two reasons. First, losing electrons usually means losing the outermost shell of the atom. Second, the proton-electron balance is thrown off, and the remaining electrons are pulled closer to the nucleus due to the stronger attraction of the nuclear charge. Therefore, a positive ion is always smaller than its neutral parent atom.

🎓 Enroll Now – MDCAT 2025 Course by Mahida Academy

For example, the radius of a sodium (Na) atom is 157pm, while the radius of a sodium ion (Na⁺) is 95pm. On the other hand, adding one or more electrons to a neutral atom to form a negative ion makes the ion larger. This is because the electrons added to the shell increase repulsion between electrons, causing the shell to expand. For example, the radius of a fluorine (F) atom is 72pm, while the radius of a fluoride ion (F^–) is 136pm. In a group of the periodic table, ions with the same charge increase in size from top to bottom. On the other hand, isoelectronic positive ions within a period decrease in size from left to right due to increasing nuclear charge. The same trend is observed for isoelectronic negative ions in a period.