Lesson 1: Periodic Classification Evolution

Video Lesson

Competencies

At the end of this section, you should be able to:

• Describe Dobereiner’s Triads and their role in the development of the periodic table;
• Understand Newlands’ Law of Octaves and how it relates to the classification of elements;
• Explain Mendeleev’s approach to classifying elements and his periodic law;
• Describe periodicity;

Brainstorming Questions

• Why do you think Dobereiner’s Triads could not be generalized as more elements were discovered?
• Why were the inert gases and elements beyond calcium not accounted for in Newlands’ Law of Octaves?
• How does Mendeleev’s periodic table illustrate the concept of periodicity in elements, and what role did atomic mass play in this arrangement?

Key Term

1.1. Dobereiner’s Triads

In the 19th century, chemists were on a quest to understand the properties of elements and how they related to each other. At that time, they had a limited understanding of atomic structure. The concepts of electrons and protons were not yet discovered, so scientists based their ideas primarily on atomic masses—the weight of atoms relative to one another. Despite the incomplete knowledge, they began to organize elements into a table, which we now know as the periodic table.

What Were Dobereiner’s Triads?

In 1829, German chemist Johann Wolfgang Döbereiner made a significant contribution to chemistry by identifying a pattern among elements. He noticed that certain groups of three elements, which he called triads, had similar chemical and physical properties. What was remarkable about these triads was that the atomic weight of the middle element was approximately the average of the atomic weights of the other two elements in the triad.

Understanding Dobereiner’s Triads with Examples

To understand how Dobereiner’s Triads worked, let’s look at some examples. Döbereiner found that the elements in these triads had a consistent pattern:

• Lithium (Li), Sodium (Na), and Potassium (K)
  • Atomic weights: Li = 7, Na = 23, K = 39
  • Average of Li and K: (7 + 39) / 2 = 23 (which is approximately the atomic weight of Na)
• Calcium (Ca), Strontium (Sr), and Barium (Ba)
  • Atomic weights: Ca = 40, Sr = 87.6, Ba = 137
  • Average of Ca and Ba: (40 + 137) / 2 = 88.5 (which is close to the atomic weight of Sr)

Significance of Dobereiner’s Triads

Dobereiner’s Triads were significant because they provided early evidence that there was some order in the way elements behaved and interacted. The idea that elements could be grouped based on their atomic weights and properties suggested that there might be a deeper connection between atomic mass and chemical behavior. This was a stepping stone toward the development of the modern periodic table.

Limitations and Evolution

While Döbereiner’s Triads were an important discovery, they had limitations. As more elements were discovered, not all of them could be grouped into triads, and the simple average of atomic weights did not always predict chemical behavior accurately. For example, not all elements fit neatly into these groups, and as scientists learned more about atomic structure, they realized that atomic mass alone did not fully explain the properties of elements.

1.2. Newlands’ Law of Octaves

In the mid-19th century, English chemist John Alexander Reina Newlands made a significant attempt to classify elements based on their atomic weights. His innovative approach led to the formulation of what is known as Newlands’ Law of Octaves.

Newlands’ Law of Octaves states that when elements are arranged in increasing order of their atomic weight, the properties of every eighth element are similar to those of the first one. Newlands compared this pattern to the repetition of musical notes in an octave, where the eighth note repeats the properties of the first note.

 Limitations of Newlands’ Classification

While Newlands’ Law of Octaves was groundbreaking, it had some notable limitations:

1. Inert Gases: At the time of Newlands’ work, the noble gases (inert gases) had not been discovered. Therefore, his system could not account for these elements, which were later found to fit into the periodic table in a way that Newlands’ model did not accommodate.
2. Beyond Calcium: Newlands observed that the repeating pattern of every eighth element did not hold true beyond calcium (Ca). This limitation indicated that the periodicity he proposed was not universal for all known elements.

1.3. Mendeleev’s Classification of the Elements

In the 19th century, Dmitri Mendeleev, a Russian chemist, made a significant contribution to chemistry by creating a periodic table that systematically organized elements based on their properties and atomic masses. Mendeleev’s work built upon earlier attempts to classify elements, and his table laid the groundwork for the modern periodic table we use today.

Mendeleev’s Periodic Law Mendeleev’s periodic law states that the physical and chemical properties of elements are periodic functions of their atomic weights. Mendeleev’s periodic table was organized into seven horizontal rows (periods) and eight vertical columns (groups). Each period contained elements with increasing atomic weights, while elements in the same group shared similar properties.

Periodicity and Patterns

The periodic table reveals regular patterns in the properties of elements. Elements within the same group exhibit similar chemical and physical properties. For example, alkali metals (Group I) are highly reactive and have similar reactivity trends down the group.

Merits and Limitations of Mendeleev’s Table Merits:

1. Systematic Study: Mendeleev’s table allowed for a more systematic study of element properties.
2. Prediction of New Elements: Mendeleev predicted the existence and properties of undiscovered elements based on gaps in his table, such as germanium and gallium.

Limitations:

1. Inert Gases Missing: Mendeleev’s table did not include inert gases, which were discovered later and added to the modern periodic table.
2. Incorrect Positions: Some elements were placed in incorrect positions due to inconsistencies in atomic weights, such as potassium and Argon.

Figure: Mendeleev’s Periodic Table